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diff --git a/.gitattributes b/.gitattributes new file mode 100644 index 0000000..6833f05 --- /dev/null +++ b/.gitattributes @@ -0,0 +1,3 @@ +* text=auto +*.txt text +*.md text diff --git a/12787-0.txt b/12787-0.txt new file mode 100644 index 0000000..a4b2dee --- /dev/null +++ b/12787-0.txt @@ -0,0 +1,8166 @@ +*** START OF THE PROJECT GUTENBERG EBOOK 12787 *** + +[Transcriber's notes: In the chemical equations, superscripts are +indicated with a ^ and subscripts are indicated with a _. The affected +item is enclosed in curly brackets {}. Examples are H^{+} for hydrogen +ion and H_{2}O for water. Since the underscore is already being used +in this project, italics are designated by an exclamation point +before and after the italicized word or phrase.] + + + +AN INTRODUCTORY COURSE + +OF + +QUANTITATIVE + +CHEMICAL ANALYSIS + +WITH + +EXPLANATORY NOTES + + +BY + +HENRY P. TALBOT + +PROFESSOR OF INORGANIC CHEMISTRY AT THE MASSACHUSETTS INSTITUTE OF +TECHNOLOGY + +SIXTH EDITION, COMPLETELY REWRITTEN + + + + +PREFACE + + +This Introductory Course of Quantitative Analysis has been prepared +to meet the needs of students who are just entering upon the subject, +after a course of qualitative analysis. It is primarily intended to +enable the student to work successfully and intelligently without the +necessity for a larger measure of personal assistance and supervision +than can reasonably be given to each member of a large class. To this +end the directions are given in such detail that there is very little +opportunity for the student to go astray; but the manual is not, the +author believes, on this account less adapted for use with small +classes, where the instructor, by greater personal influence, can +stimulate independent thought on the part of the pupil. + +The method of presentation of the subject is that suggested by +Professor A.A. Noyes' excellent manual of Qualitative Analysis. For +each analysis the procedure is given in considerable detail, and +this is accompanied by explanatory notes, which are believed to be +sufficiently expanded to enable the student to understand fully the +underlying reason for each step prescribed. The use of the book +should, nevertheless, be supplemented by classroom instruction, mainly +of the character of recitations, and the student should be taught to +consult larger works. The general directions are intended to emphasize +those matters upon which the beginner in quantitative analysis must +bestow special care, and to offer helpful suggestions. The student +can hardly be expected to appreciate the force of all the statements +contained in these directions, or, indeed, to retain them all in +the memory after a single reading; but the instructor, by frequent +reference to special paragraphs, as suitable occasion presents itself, +can soon render them familiar to the student. + +The analyses selected for practice are those comprised in the first +course of quantitative analysis at the Massachusetts Institute of +Technology, and have been chosen, after an experience of years, +as affording the best preparation for more advanced work, and as +satisfactory types of gravimetric and volumetric methods. From the +latter point of view, they also seem to furnish the best insight into +quantitative analysis for those students who can devote but a limited +time to the subject, and who may never extend their study beyond the +field covered by this manual. The author has had opportunity to test +the efficiency of the course for use with such students, and has found +the results satisfactory. + +In place of the usual custom of selecting simple salts as material for +preliminary practice, it has been found advantageous to substitute, in +most instances, approximately pure samples of appropriate minerals or +industrial products. The difficulties are not greatly enhanced, while +the student gains in practical experience. + +The analytical procedures described in the following pages have been +selected chiefly with reference to their usefulness in teaching the +subject, and with the purpose of affording as wide a variety of +processes as is practicable within an introductory course of this +character. The scope of the manual precludes any extended attempt to +indicate alternative procedures, except through general references to +larger works on analytical chemistry. The author is indebted to the +standard works for many suggestions for which it is impracticable to +make specific acknowledgment; no considerable credit is claimed by him +for originality of procedure. + +For many years, as a matter of convenience, the classes for which this +text was originally prepared were divided, one part beginning with +gravimetric processes and the other with volumetric analyses. After a +careful review of the experience thus gained the conclusion has been +reached that volumetric analysis offers the better approach to the +subject. Accordingly the arrangement of the present (the sixth) +edition of this manual has been changed to introduce volumetric +procedures first. Teachers who are familiar with earlier editions +will, however, find that the order of presentation of the material +under the various divisions is nearly the same as that previously +followed, and those who may still prefer to begin the course of +instruction with gravimetric processes will, it is believed, be able +to follow that order without difficulty. + +Procedures for the determination of sulphur in insoluble sulphates, +for the determination of copper in copper ores by iodometric methods, +for the determination of iron by permanganate in hydrochloric acid +solutions, and for the standardization of potassium permanganate +solutions using sodium oxalate as a standard, and of thiosulphate +solutions using copper as a standard, have been added. The +determination of silica in silicates decomposable by acids, as a +separate procedure, has been omitted. + +The explanatory notes have been rearranged to bring them into closer +association with the procedures to which they relate. The number of +problems has been considerably increased. + +The author wishes to renew his expressions of appreciation of the +kindly reception accorded the earlier editions of this manual. He has +received helpful suggestions from so many of his colleagues within the +Institute, and friends elsewhere, that his sense of obligation must +be expressed to them collectively. He is under special obligations +to Professor L.F. Hamilton for assistance in the preparation of the +present edition. + +HENRY P. TALBOT + +!Massachusetts Institute of Technology, September, 1921!. + + + + +CONTENTS + + +PART I. INTRODUCTION + +SUBDIVISIONS OF ANALYTICAL CHEMISTRY + +GENERAL DIRECTIONS + Accuracy and Economy of Time; Notebooks; Reagents; Wash-bottles; + Transfer of Liquids + + +PART II. VOLUMETRIC ANALYSIS + +GENERAL DISCUSSION + Subdivisions; The Analytical Balance; Weights; Burettes; + Calibration of Measuring Devices +GENERAL DIRECTIONS + Standard and Normal Solutions + +!I. Neutralization Methods! + +ALKALIMETRY AND ACIDIMETRY + Preparation and Standardization of Solutions; Indicators +STANDARDIZATION OF HYDROCHLORIC ACID +DETERMINATION OF TOTAL ALKALINE STRENGTH OF SODA ASH +DETERMINATION OF ACID STRENGTH OF OXALIC ACID + +!II. Oxidation Processes! + +GENERAL DISCUSSION +BICHROMATE PROCESS FOR THE DETERMINATION OF IRON +DETERMINATION OF IRON IN LIMONITE BY THE BICHROMATE PROCESS +DETERMINATION OF CHROMIUM IN CHROME IRON ORE +PERMANGANATE PROCESS FOR THE DETERMINATION OF IRON +DETERMINATION OF IRON IN LIMONITE BY THE PERMANGANATE PROCESS +DETERMINATION OF IRON IN LIMONITE BY THE ZIMMERMANN-REINHARDT PROCESS +DETERMINATION OF THE OXIDIZING POWER OF PYROLUSITE +IODIMETRY +DETERMINATION OF COPPER IN ORES +DETERMINATION OF ANTIMONY IN STIBNITE +CHLORIMETRY +DETERMINATION OF AVAILABLE CHLORINE IN BLEACHING POWDER + +!III. Precipitation Methods! + +DETERMINATION OF SILVER BY THE THIOCYANATE PROCESS + + +PART III. GRAVIMETRIC ANALYSIS + +GENERAL DIRECTIONS + Precipitation; Funnels and Filters; Filtration and Washing of + Precipitates; Desiccators; Crucibles and their Preparation + for Use; Ignition of Precipitates +DETERMINATION OF CHLORINE IN SODIUM CHLORIDE +DETERMINATION OF IRON AND OF SULPHUR IN FERROUS AMMONIUM SULPHATE +DETERMINATION OF SULPHUR IN BARIUM SULPHATE +DETERMINATION OF PHOSPHORIC ANHYDRIDE IN APATITE +ANALYSIS OF LIMESTONE + Determination of Moisture; Insoluble Matter and Silica; Ferric + Oxide and Alumina; Calcium; Magnesium; Carbon Dioxide +ANALYSIS OF BRASS + Electrolytic Separations; Determination of Lead, Copper, Iron + and Zinc. +DETERMINATION OF SILICA IN SILICATES + +PART IV. STOICHIOMETRY + +SOLUTIONS OF TYPICAL PROBLEMS +PROBLEMS + +APPENDIX + +ELECTROLYTIC DISSOCIATION THEORY +FOLDING OF A FILTER PAPER +SAMPLE NOTEBOOK PAGES +STRENGTH OF REAGENTS +DENSITIES AND VOLUMES OF WATER +CORRECTIONS FOR CHANGE OF TEMPERATURE OF STANDARD SOLUTIONS +ATOMIC WEIGHTS +LOGARITHM TABLES + + + + +QUANTITATIVE CHEMICAL ANALYSIS + + + + +PART I + +INTRODUCTION + +SUBDIVISIONS OF ANALYTICAL CHEMISTRY + + +A complete chemical analysis of a body of unknown composition involves +the recognition of its component parts by the methods of !qualitative +analysis!, and the determination of the proportions in which these +components are present by the processes of !quantitative analysis!. +A preliminary qualitative examination is generally indispensable, if +intelligent and proper provisions are to be made for the separation of +the various constituents under such conditions as will insure accurate +quantitative estimations. + +It is assumed that the operations of qualitative analysis are familiar +to the student, who will find that the reactions made use of in +quantitative processes are frequently the same as those employed in +qualitative analyses with respect to both precipitation and systematic +separation from interfering substances; but it should be noted that +the conditions must now be regulated with greater care, and in such +a manner as to insure the most complete separation possible. For +example, in the qualitative detection of sulphates by precipitation +as barium sulphate from acid solution it is not necessary, in most +instances, to take into account the solubility of the sulphate +in hydrochloric acid, while in the quantitative determination of +sulphates by this reaction this solubility becomes an important +consideration. The operations of qualitative analysis are, therefore, +the more accurate the nearer they are made to conform to quantitative +conditions. + +The methods of quantitative analysis are subdivided, according +to their nature, into those of !gravimetric analysis, volumetric +analysis!, and !colorimetric analysis!. In !gravimetric! processes the +constituent to be determined is sometimes isolated in elementary +form, but more commonly in the form of some compound possessing a +well-established and definite composition, which can be readily and +completely separated, and weighed either directly or after ignition. +From the weight of this substance and its known composition, the +amount of the constituent in question is determined. + +In !volumetric! analysis, instead of the final weighing of a definite +body, a well-defined reaction is caused to take place, wherein the +reagent is added from an apparatus so designed that the volume of the +solution employed to complete the reaction can be accurately measured. +The strength of this solution (and hence its value for the reaction +in question) is accurately known, and the volume employed serves, +therefore, as a measure of the substance acted upon. An example will +make clear the distinction between these two types of analysis. +The percentage of chlorine in a sample of sodium chloride may be +determined by dissolving a weighed amount of the chloride in water +and precipitating the chloride ions as silver chloride, which is +then separated by filtration, ignited, and weighed (a !gravimetric! +process); or the sodium chloride may be dissolved in water, and a +solution of silver nitrate containing an accurately known amount of +the silver salt in each cubic centimeter may be cautiously added from +a measuring device called a burette until precipitation is complete, +when the amount of chlorine may be calculated from the number of cubic +centimeters of the silver nitrate solution involved in the reaction. +This is a !volumetric! process, and is equivalent to weighing without +the use of a balance. + +Volumetric methods are generally more rapid, require less apparatus, +and are frequently capable of greater accuracy than gravimetric +methods. They are particularly useful when many determinations of the +same sort are required. + +In !colorimetric! analyses the substance to be determined is converted +into some compound which imparts to its solutions a distinct color, +the intensity of which must vary in direct proportion to the amount of +the compound in the solution. Such solutions are compared with respect +to depth of color with standard solutions containing known amounts of +the colored compound, or of other similar color-producing substance +which has been found acceptable as a color standard. Colorimetric +methods are, in general, restricted to the determinations of very +small quantities, since only in dilute solutions are accurate +comparisons of color possible. + + + + +GENERAL DIRECTIONS + + +The following paragraphs should be read carefully and thoughtfully. A +prime essential for success as an analyst is attention to details and +the avoidance of all conditions which could destroy, or even lessen, +confidence in the analyses when completed. The suggestions here given +are the outcome of much experience, and their adoption will tend to +insure permanently work of a high grade, while neglect of them will +often lead to disappointment and loss of time. + + +ACCURACY AND ECONOMY OF TIME + +The fundamental conception of quantitative analysis implies a +necessity for all possible care in guarding against loss of material +or the introduction of foreign matter. The laboratory desk, and all +apparatus, should be scrupulously neat and clean at all times. A +sponge should always be ready at hand, and desk and filter-stands +should be kept dry and in good order. Funnels should never be allowed +to drip upon the base of the stand. Glassware should always be +wiped with a clean, lintless towel just before use. All filters and +solutions should be covered to protect them from dust, just as far as +is practicable, and every drop of solution or particle of precipitate +must be regarded as invaluable for the success of the analysis. + +An economical use of laboratory hours is best secured by acquiring +a thorough knowledge of the character of the work to be done before +undertaking it, and then by so arranging the work that no time shall +be wasted during the evaporation of liquids and like time-consuming +operations. To this end the student should read thoughtfully not only +the !entire! procedure, but the explanatory notes as well, before +any step is taken in the analysis. The explanatory notes furnish, in +general, the reasons for particular steps or precautions, but they +also occasionally contain details of manipulation not incorporated, +for various reasons, in the procedure. These notes follow the +procedures at frequent intervals, and the exact points to which they +apply are indicated by references. The student should realize that a +!failure to study the notes will inevitably lead to mistakes, loss of +time, and an inadequate understanding of the subject!. + +All analyses should be made in duplicate, and in general a close +agreement of results should be expected. It should, however, be +remembered that a close concordance of results in "check analyses" is +not conclusive evidence of the accuracy of those results, although the +probability of their accuracy is, of course, considerably enhanced. +The satisfaction in obtaining "check results" in such analyses must +never be allowed to interfere with the critical examination of the +procedure employed, nor must they ever be regarded as in any measure a +substitute for absolute truth and accuracy. + +In this connection it must also be emphasized that only the operator +himself can know the whole history of an analysis, and only he can +know whether his work is worthy of full confidence. No work should be +continued for a moment after such confidence is lost, but should +be resolutely discarded as soon as a cause for distrust is fully +established. The student should, however, determine to put forth his +best efforts in each analysis; it is well not to be too ready to +condone failures and to "begin again," as much time is lost in these +fruitless attempts. Nothing less than !absolute integrity! is or can +be demanded of a quantitative analyst, and any disregard of this +principle, however slight, is as fatal to success as lack of chemical +knowledge or inaptitude in manipulation can possibly be. + + +NOTEBOOKS + +Notebooks should contain, beside the record of observations, +descriptive notes. All records of weights should be placed upon the +right-hand page, while that on the left is reserved for the notes, +calculations of factors, or the amount of reagents required. + +The neat and systematic arrangement of the records of analyses is +of the first importance, and is an evidence of careful work and an +excellent credential. Of two notebooks in which the results may be, +in fact, of equal value as legal evidence, that one which is neatly +arranged will carry with it greater weight. + +All records should be dated, and all observations should be recorded +at once in the notebook. The making of records upon loose paper is a +practice to be deprecated, as is also that of copying original entries +into a second notebook. The student should accustom himself to orderly +entries at the time of observation. Several sample pages of systematic +records are to be found in the Appendix. These are based upon +experience; but other arrangements, if clear and orderly, may prove +equally serviceable. The student is advised to follow the sample pages +until he is in a position to plan out a system of his own. + + +REAGENTS + +The habit of carefully testing reagents, including distilled water, +cannot be too early acquired or too constantly practiced; for, in +spite of all reasonable precautionary measures, inferior chemicals +will occasionally find their way into the stock room, or errors will +be made in filling reagent bottles. The student should remember that +while there may be others who share the responsibility for the purity +of materials in the laboratory of an institution, the responsibility +will later be one which he must individually assume. + +The stoppers of reagent bottles should never be laid upon the desk, +unless upon a clean watch-glass or paper. The neck and mouth of all +such bottles should be kept scrupulously clean, and care taken that no +confusion of stoppers occurs. + + +WASH-BOTTLES + +Wash-bottles for distilled water should be made from flasks of about +750 cc. capacity and be provided with gracefully bent tubes, which +should not be too long. The jet should be connected with the tube +entering the wash-bottle by a short piece of rubber tubing in such +a way as to be flexible, and should deliver a stream about one +millimeter in diameter. The neck of the flask may be wound with cord, +or covered with wash-leather, for greater comfort when hot water is +used. It is well to provide several small wash-bottles for liquids +other than distilled water, which should invariably be clearly +labeled. + + +TRANSFER OF LIQUIDS + +Liquids should never be transferred from one vessel to another, nor to +a filter, without the aid of a stirring rod held firmly against the +side or lip of the vessel. When the vessel is provided with a lip it +is not usually necessary to use other means to prevent the loss of +liquid by running down the side; whenever loss seems imminent a !very +thin! layer of vaseline, applied with the finger to the edge of the +vessel, will prevent it. The stirring rod down which the liquid runs +should never be drawn upward in such a way as to allow the solution to +collect on the under side of the rim or lip of a vessel. + +The number of transfers of liquids from one vessel to another during +an analysis should be as small as possible to avoid the risk of slight +losses. Each vessel must, of course, be completely washed to insure +the transfer of all material; but it should be remembered that this +can be accomplished better by the use of successive small portions of +wash-water (perhaps 5-10 cc.), if each wash-water is allowed to drain +away for a few seconds, than by the addition of large amounts which +unnecessarily increase the volume of the solutions, causing loss of +time in subsequent filtrations or evaporations. + +All stirring rods employed in quantitative analyses should be rounded +at the ends by holding them in the flame of a burner until they begin +to soften. If this is not done, the rods will scratch the inner +surface of beakers, causing them to crack on subsequent heating. + + +EVAPORATION OF LIQUIDS + +The greatest care must be taken to prevent loss of solutions during +processes of evaporation, either from too violent ebullition, from +evaporation to dryness and spattering, or from the evolution of gas +during the heating. In general, evaporation upon the steam bath is to +be preferred to other methods on account of the impossibility of +loss by spattering. If the steam baths are well protected from dust, +solutions should be left without covers during evaporation; but +solutions which are boiled upon the hot plate, or from which gases are +escaping, should invariably be covered. In any case a watch-glass may +be supported above the vessel by means of a glass triangle, or other +similar device, and the danger of loss of material or contamination by +dust thus be avoided. It is obvious that evaporation is promoted by +the use of vessels which admit of the exposure of a broad surface to +the air. + +Liquids which contain suspended matter (precipitates) should always +be cautiously heated, since the presence of the solid matter is +frequently the occasion of violent "bumping," with consequent risk to +apparatus and analysis. + + + + +PART II + +VOLUMETRIC ANALYSIS + + +The processes of volumetric analysis are, in general, simpler than +those of gravimetric analysis and accordingly serve best as an +introduction to the practice of quantitative analysis. For their +execution there are required, first, an accurate balance with which +to weigh the material for analysis; second, graduated instruments in +which to measure the volume of the solutions employed; third, standard +solutions, that is, solutions the value of which is accurately known; +and fourth, indicators, which will furnish accurate evidence of the +point at which the desired reaction is completed. The nature of the +indicators employed will be explained in connection with the different +analyses. + +The process whereby a !standard solution! is brought into reaction is +called !titration!, and the point at which the reaction is exactly +completed is called the !end-point!. The !indicator! should show the +!end-point! of the !titration!. The volume of the standard solution +used then furnishes the measure of the substance to be determined as +truly as if that substance had been separated and weighed. + +The processes of volumetric analysis are easily classified, according +to their character, into: + +I. NEUTRALIZATION METHODS; such, for example, as those of acidimetry +and alkalimetry. + +II. OXIDATION PROCESSES; as exemplified in the determination of +ferrous iron by its oxidation with potassium bichromate. + +III. PRECIPITATION METHODS; of which the titration for silver with +potassium thiocyanate solution is an illustration. + +From a somewhat different standpoint the methods in each case may +be subdivided into (a) DIRECT METHODS, in which the substance to be +measured is directly determined by titration to an end-point with a +standard solution; and (b) INDIRECT METHODS, in which the substance +itself is not measured, but a quantity of reagent is added which is +known to be an excess with respect to a specific reaction, and the +unused excess determined by titration. Examples of the latter class +will be pointed out as they occur in the procedures. + + +MEASURING INSTRUMENTS + + +THE ANALYTICAL BALANCE + +For a complete discussion of the physical principles underlying the +construction and use of balances, and the various methods of weighing, +the student is referred to larger manuals of Quantitative Analysis, +such as those of Fresenius, or Treadwell-Hall, and particularly to +the admirable discussion of this topic in Morse's !Exercises in +Quantitative Chemistry!. + +The statements and rules of procedure which follow are sufficient +for the intelligent use of an analytical balance in connection with +processes prescribed in this introductory manual. It is, however, +imperative that the student should make himself familiar with these +essential features of the balance, and its use. He should fully +realize that the analytical balance is a delicate instrument which +will render excellent service under careful treatment, but such +treatment is an essential condition if its accuracy is to be depended +upon. He should also understand that no set of rules, however +complete, can do away with the necessity for a sense of personal +responsibility, since by carelessness he can render inaccurate not +only his own analyses, but those of all other students using the same +balance. + +Before making any weighings the student should seat himself before a +balance and observe the following details of construction: + +1. The balance case is mounted on three brass legs, which should +preferably rest in glass cups, backed with rubber to prevent slipping. +The front legs are adjustable as to height and are used to level the +balance case; the rear leg is of permanent length. + +2. The front of the case may be raised to give access to the balance. +In some makes doors are provided also at the ends of the balance case. + +3. The balance beam is mounted upon an upright in the center of the +case on the top of which is an inlaid agate plate. To the center of +the beam there is attached a steel or agate knife-edge on which the +beam oscillates when it rests on the agate plate. + +4. The balance beam, extending to the right and left, is graduated +along its upper edge, usually on both sides, and has at its +extremities two agate or steel knife-edges from which are suspended +stirrups. Each of these stirrups has an agate plate which, when the +balance is in action, rests upon the corresponding knife-edge of the +beam. The balance pans are suspended from the stirrups. + +5. A pointer is attached to the center of the beam, and as the beam +oscillates this pointer moves in front of a scale near the base of the +post. + +6. At the base of the post, usually in the rear, is a spirit-level. + +7. Within the upright is a mechanism, controlled by a knob at the +front of the balance case, which is so arranged as to raise the entire +beam slightly above the level at which the knife-edges are in contact +with the agate plates. When the balance is not in use the beam must +be supported by this device since, otherwise, the constant jarring +to which a balance is inevitably subjected, will soon dull the +knife-edges, and lessen the sensitiveness of the balance. + +8. A small weight, or bob, is attached to the pointer (or sometimes +to the beam) by which the center of gravity of the beam and its +attachments may be regulated. The center of gravity must lie very +slightly below the level of the agate plates to secure the desired +sensitiveness of the balance. This is provided for when the balance is +set up and very rarely requires alteration. The student should never +attempt to change this adjustment. + +9. Below the balance pans are two pan-arrests operated by a button +from the front of the case. These arrests exert a very slight upward +pressure upon the pans and minimize the displacement of the beam when +objects or weights are being placed upon the pans. + +10. A movable rod, operated from one end of the balance case, extends +over the balance beam and carries a small wire weight, called a rider. +By means of this rod the rider can be placed upon any desired division +of the scale on the balance beam. Each numbered division on the beam +corresponds to one milligram, and the use of the rider obviates the +placing of very small fractional weights on the balance pan. + +If a new rider is purchased, or an old one replaced, care must be +taken that its weight corresponds to the graduations on the beam of +the balance on which it is to be used. The weight of the rider in +milligrams must be equal to the number of large divisions (5, 6, 10, +or 12) between the central knife-edge and the knife-edge at the end of +the beam. It should be noted that on some balances the last division +bears no number. Each new rider should be tested against a 5 or +10-milligram weight. + +In some of the most recent forms of the balance a chain device +replaces the smaller weights and the use of the rider as just +described. + +Before using a balance, it is always best to test its adjustment. This +is absolutely necessary if the balance is used by several workers; it +is always a wise precaution under any conditions. For this purpose, +brush off the balance pans with a soft camel's hair brush. Then note +(1) whether the balance is level; (2) that the mechanism for raising +and lowering the beams works smoothly; (3) that the pan-arrests touch +the pans when the beam is lowered; and (4) that the needle swings +equal distances on either side of the zero-point when set in motion +without any load on the pans. If the latter condition is not +fulfilled, the balance should be adjusted by means of the adjusting +screw at the end of the beam unless the variation is not more than one +division on the scale; it is often better to make a proper allowance +for this small zero error than to disturb the balance by an attempt at +correction. Unless a student thoroughly understands the construction +of a balance he should never attempt to make adjustments, but should +apply to the instructor in charge. + +The object to be weighed should be placed on the left-hand balance pan +and the weights upon the right-hand pan. Every substance which +could attack the metal of the balance pan should be weighed upon a +watch-glass, and all objects must be dry and cold. A warm body gives +rise to air currents which vitiate the accuracy of the weighing. + +The weights should be applied in the order in which they occur in the +weight-box (not at haphazard), beginning with the largest weight which +is apparently required. After a weight has been placed upon the pan +the beam should be lowered upon its knife-edges, and, if necessary, +the pan-arrests depressed. The movement of the pointer will then +indicate whether the weight applied is too great or too small. When +the weight has been ascertained, by the successive addition of small +weights, to the nearest 5 or 10 milligrams, the weighing is completed +by the use of the rider. The correct weight is that which causes the +pointer to swing an equal number of divisions to the right and left +of the zero-point, when the pointer traverses not less than five +divisions on either side. + +The balance case should always be closed during the final weighing, +while the rider is being used, to protect the pans from the effect of +air currents. + +Before the final determination of an exact weight the beam should +always be lifted from the knife-edges and again lowered into place, +as it frequently happens that the scale pans are, in spite of the +pan-arrests, slightly twisted by the impact of the weights, the beam +being thereby virtually lengthened or shortened. Lifting the beam +restores the proper alignment. + +The beam should never be set in motion by lowering it forcibly upon +the knife-edges, nor by touching the pans, but rather by lifting the +rider (unless the balance be provided with some of the newer devices +for the purpose), and the swing should be arrested only when the +needle approaches zero on the scale, otherwise the knife-edges become +dull. For the same reason the beam should never be left upon its +knife-edges, nor should weights be removed from or placed on the +pans without supporting the beam, except in the case of the small +fractional weights. + +When the process of weighing has been completed, the weight should +be recorded in the notebook by first noting the vacant spaces in the +weight-box, and then checking the weight by again noting the weights +as they are removed from the pan. This practice will often detect and +avoid errors. It is obvious that the weights should always be returned +to their proper places in the box, and be handled only with pincers. + +It should be borne in mind that if the mechanism of a balance is +deranged or if any substance is spilled upon the pans or in the +balance case, the damage should be reported at once. In many instances +serious harm can be averted by prompt action when delay might ruin the +balance. + +Samples for analysis are commonly weighed in small tubes with cork +stoppers. Since the stoppers are likely to change in weight from +the varying amounts of moisture absorbed from the atmosphere, it is +necessary to confirm the recorded weight of a tube which has been +unused for some time before weighing out a new portion of substance +from it. + + +WEIGHTS + +The sets of weights commonly used in analytical chemistry range from +20 grams to 5 milligrams. The weights from 20 grams to 1 gram are +usually of brass, lacquered or gold plated. The fractional weights +are of German silver, gold, platinum or aluminium. The rider is of +platinum or aluminium wire. + +The sets of weights purchased from reputable dealers are usually +sufficiently accurate for analytical work. It is not necessary that +such a set should be strictly exact in comparison with the absolute +standard of weight, provided they are relatively correct among +themselves, and provided the same set of weights is used in all +weighings made during a given analysis. The analyst should assure +himself that the weights in a set previously unfamiliar to him are +relatively correct by a few simple tests. For example, he should make +sure that in his set two weights of the same denomination (i.e., two +10-gram weights, or the two 100-milligram weights) are actually equal +and interchangeable, or that the 500-milligram weight is equal to +the sum of the 200, 100, 100, 50, 20, 20 and 10-milligram weights +combined, and so on. If discrepancies of more than a few tenths of a +milligram (depending upon the total weight involved) are found, the +weights should be returned for correction. The rider should also be +compared with a 5 or 10-milligram weight. + +In an instructional laboratory appreciable errors should be reported +to the instructor in charge for his consideration. + +When the highest accuracy is desired, the weights may be calibrated +and corrections applied. A calibration procedure is described in a +paper by T.W. Richards, !J. Am. Chem. Soc.!, 22, 144, and in many +large text-books. + +Weights are inevitably subject to corrosion if not properly protected +at all times, and are liable to damage unless handled with great care. +It is obvious that anything which alters the weight of a single piece +in an analytical set will introduce an error in every weighing made +in which that piece is used. This source of error is often extremely +obscure and difficult to detect. The only safeguard against such +errors is to be found in scrupulous care in handling and protection +on the part of the analyst, and an equal insistence that if several +analysts use the same set of weights, each shall realize his +responsibility for the work of others as well as his own. + + +BURETTES + +A burette is made from a glass tube which is as uniformly cylindrical +as possible, and of such a bore that the divisions which are etched +upon its surface shall correspond closely to actual contents. + +The tube is contracted at one extremity, and terminates in either a +glass stopcock and delivery-tube, or in such a manner that a piece of +rubber tubing may be firmly attached, connecting a delivery-tube of +glass. The rubber tubing is closed by means of a glass bead. Burettes +of the latter type will be referred to as "plain burettes." + +The graduations are usually numbered in cubic centimeters, and the +latter are subdivided into tenths. + +One burette of each type is desirable for the analytical procedures +which follow. + + +PREPARATION OF A BURETTE FOR USE + +The inner surface of a burette must be thoroughly cleaned in order +that the liquid as drawn out may drain away completely, without +leaving drops upon the sides. This is best accomplished by treating +the inside of the burette with a warm solution of chromic acid in +concentrated sulphuric acid, applied as follows: If the burette is of +the "plain" type, first remove the rubber tip and force the lower +end of the burette into a medium-sized cork stopper. Nearly fill the +burette with the chromic acid solution, close the upper end with a +cork stopper and tip the burette backward and forward in such a way +as to bring the solution into contact with the entire inner surface. +Remove the stopper and pour the solution into a stock bottle to be +kept for further use, and rinse out the burette with water several +times. Unless the water then runs freely from the burette without +leaving drops adhering to the sides, the process must be repeated +(Note 1). + +If the burette has a glass stopcock, this should be removed after +the cleaning and wiped, and also the inside of the ground joint. The +surface of the stopcock should then be smeared with a thin coating of +vaseline and replaced. It should be attached to the burette by means +of a wire, or elastic band, to lessen the danger of breakage. + +Fill the burettes with distilled water, and allow the water to run out +through the stopcock or rubber tip until convinced that no air +bubbles are inclosed (Note 2). Fill the burette to a point above the +zero-point and draw off the water until the meniscus is just below +that mark. It is then ready for calibration. + +[Note 1: The inner surface of the burette must be absolutely clean if +the liquid is to run off freely. Chromic acid in sulphuric acid is +usually found to be the best cleansing agent, but the mixture must be +warm and concentrated. The solution can be prepared by pouring over a +few crystals of potassium bichromate a little water and then adding +concentrated sulphuric acid.] + +[Note 2: It is always necessary to insure the absence of air bubbles +in the tips or stopcocks. The treatment described above will usually +accomplish this, but, in the case of plain burettes it is sometimes +better to allow a little of the liquid to flow out of the tip while it +is bent upwards. Any air which may be entrapped then rises with the +liquid and escapes. + +If air bubbles escape during subsequent calibration or titration, an +error is introduced which vitiates the results.] + + +READING OF A BURETTE + +All liquids when placed in a burette form what is called a meniscus at +their upper surfaces. In the case of liquids such as water or +aqueous solutions this meniscus is concave, and when the liquids are +transparent accurate readings are best obtained by observing the +position on the graduated scales of the lowest point of the meniscus. +This can best be done as follows: Wrap around the burette a piece of +colored paper, the straight, smooth edges of which are held evenly +together with the colored side next to the burette (Note 1). Hold the +paper about two small divisions below the meniscus and raise or lower +the level of the eyes until the edge of the paper at the back of the +burette is just hidden from the eye by that in front (Note 2). Note +the position of the lowest point of the curve of the meniscus, +estimating the tenths of the small divisions, thus reading its +position to hundredths of a cubic centimeter. + +[Note 1: The ends of the colored paper used as an aid to accurate +readings may be fastened together by means of a gummed label. The +paper may then remain on the burette and be ready for immediate use by +sliding it up or down, as required.] + +[Note 2: To obtain an accurate reading the eye must be very nearly on +a level with the meniscus. This is secured by the use of the paper +as described. The student should observe by trial how a reading is +affected when the meniscus is viewed from above or below. + +The eye soon becomes accustomed to estimating the tenths of the +divisions. If the paper is held as directed, two divisions below the +meniscus, one whole division is visible to correct the judgment. It is +not well to attempt to bring the meniscus exactly to a division mark +on the burette. Such readings are usually less accurate than those in +which the tenths of a division are estimated.] + + +CALIBRATION OF GLASS MEASURING DEVICES + +If accuracy of results is to be attained, the correctness of all +measuring instruments must be tested. None of the apparatus offered +for sale can be implicitly relied upon except those more expensive +instruments which are accompanied by a certificate from the !National +Bureau of Standards! at Washington, or other equally authentic source. + +The bore of burettes is subject to accidental variations, and since +the graduations are applied by machine without regard to such +variations of bore, local errors result. + +The process of testing these instruments is called !calibration!. +It is usually accomplished by comparing the actual weight of water +contained in the instrument with its apparent volume. + +There is, unfortunately, no uniform standard of volume which has been +adopted for general use in all laboratories. It has been variously +proposed to consider the volume of 1000 grams of water at 4°, 15.5°, +16°, 17.5°, and even 20°C., as a liter for practical purposes, and to +consider the cubic centimeter to be one one-thousandth of that volume. +The true liter is the volume of 1000 grams of water at 4°C.; but +this is obviously a lower temperature than that commonly found in +laboratories, and involves the constant use of corrections if taken as +a laboratory standard. Many laboratories use 15.5°C. (60° F.) as the +working standard. It is plain that any temperature which is deemed +most convenient might be chosen for a particular laboratory, but it +cannot be too strongly emphasized that all measuring instruments, +including burettes, pipettes, and flasks, should be calibrated at that +temperature in order that the contents of each burette, pipette, +etc., shall be comparable with that of every other instrument, thus +permitting general interchange and substitution. For example, it is +obvious that if it is desired to remove exactly 50 cc. from a solution +which has been diluted to 500 cc. in a graduated flask, the 50 cc. +flask or pipette used to remove the fractional portion must give +a correct reading at the same temperature as the 500 cc. flask. +Similarly, a burette used for the titration of the 50 cc. of solution +removed should be calibrated under the same conditions as the +measuring flasks or pipettes employed with it. + +The student should also keep constantly in mind the fact that all +volumetric operations, to be exact, should be carried out as nearly at +a constant temperature as is practicable. The spot selected for +such work should therefore be subject to a minimum of temperature +variations, and should have as nearly the average temperature of +the laboratory as is possible. In all work, whether of calibration, +standardization, or analysis, the temperature of the liquids employed +must be taken into account, and if the temperature of these liquids +varies more than 3° or 4° from the standard temperature chosen for the +laboratory, corrections must be applied for errors due to expansion or +contraction, since volumes of a liquid measured at different times are +comparable only under like conditions as to temperature. Data to be +used for this purpose are given in the Appendix. Neglect of this +correction is frequently an avoidable source of error and annoyance in +otherwise excellent work. The temperature of all solutions at the time +of standardization should be recorded to facilitate the application of +temperature corrections, if such are necessary at any later time. + + +CALIBRATION OF THE BURETTES + +Two burettes, one at least of which should have a glass stopper, are +required throughout the volumetric work. Both burettes should be +calibrated by the student to whom they are assigned. + +PROCEDURE.--Weigh a 50 cc., flat-bottomed flask (preferably a +light-weight flask), which must be dry on the outside, to the nearest +centigram. Record the weight in the notebook. (See Appendix for +suggestions as to records.) Place the flask under the burette and draw +out into it about 10 cc. of water, removing any drop on the tip by +touching it against the inside of the neck of the flask. Do not +attempt to stop exactly at the 10 cc. mark, but do not vary more than +0.1 cc. from it. Note the time, and at the expiration of three minutes +(or longer) read the burette accurately, and record the reading in the +notebook (Note 1). Meanwhile weigh the flask and water to centigrams +and record its weight (Note 2). Draw off the liquid from 10 cc. to +about 20 cc. into the same flask without emptying it; weigh, and at +the expiration of three minutes take the reading, and so on throughout +the length of the burette. When it is completed, refill the burette +and check the first calibration. + +The differences in readings represent the apparent volumes, the +differences in weights the true volumes. For example, if an apparent +volume of 10.05 cc. is found to weigh 10.03 grams, it may be assumed +with sufficient accuracy that the error in that 10 cc. amounts to +-0.02 cc., or -0.002 for each cubic centimeter (Note 3). + +In the calculation of corrections the temperature of the water must be +taken into account, if this varies more than 4°C. from the laboratory +standard temperature, consulting the table of densities of water in +the Appendix. + +From the final data, plot the corrections to be applied so that they +may be easily read for each cubic centimeter throughout the burette. +The total correction at each 10 cc. may also be written on the burette +with a diamond, or etching ink, for permanence of record. + +[Note 1: A small quantity of liquid at first adheres to the side of +even a clean burette. This slowly unites with the main body of liquid, +but requires an appreciable time. Three minutes is a sufficient +interval, but not too long, and should be adopted in every instance +throughout the whole volumetric practice before final readings are +recorded.] + +[Note 2: A comparatively rough balance, capable of weighing to +centigrams, is sufficiently accurate for use in calibrations, for a +moment's reflection will show that it would be useless to weigh the +water with an accuracy greater than that of the readings taken on +the burette. The latter cannot exceed 0.01 cc. in accuracy, which +corresponds to 0.01 gram. + +The student should clearly understand that !all other weighings!, +except those for calibration, should be made accurately to 0.0001 +gram, unless special directions are given to the contrary. + +Corrections for temperature variations of less than 4°C. are +negligible, as they amount to less than 0.01 gram for each 10 grams of +water withdrawn.] + +[Note 3: Should the error discovered in any interval of 10 cc. on the +burette exceed 0.10 cc., it is advisable to weigh small portions (even +1 cc.) to locate the position of the variation of bore in the +tube rather than to distribute the correction uniformly over the +corresponding 10 cc. The latter is the usual course for small +corrections, and it is convenient to calculate the correction +corresponding to each cubic centimeter and to record it in the form +of a table or calibration card, or to plot a curve representing the +values. + +Burettes may also be calibrated by drawing off the liquid in +successive portions through a 5 cc. pipette which has been accurately +calibrated, as a substitute for weighing. If many burettes are to be +tested, this is a more rapid method.] + + +PIPETTES + +A !pipette! may consist of a narrow tube, in the middle of which is +blown a bulb of a capacity a little less than that which it is desired +to measure by the pipette; or it may be a miniature burette, without +the stopcock or rubber tip at the lower extremity. In either case, the +flow of liquid is regulated by the pressure of the finger on the top, +which governs the admission of the air. + +Pipettes are usually already graduated when purchased, but they +require calibration for accurate work. + + +CALIBRATION OF PIPETTES + +PROCEDURE.--Clean the pipette. Draw distilled water into it by sucking +at the upper end until the water is well above the graduation mark. +Quickly place the forefinger over the top of the tube, thus preventing +the entrance of air and holding the water in the pipette. Cautiously +admit a little air by releasing the pressure of the finger, and allow +the level of the water to fall until the lowest point of the meniscus +is level with the graduation. Hold the water at that point by pressure +of the finger and then allow the water to run out from the pipette +into a small tared, or weighed, beaker or flask. After a definite time +interval, usually two to three minutes, touch the end of the pipette +against the side of the beaker or flask to remove any liquid adhering +to it (Note 1). The increase in weight of the flask in grams +represents the volume of the water in cubic centimeters delivered by +the pipette. Calculate the necessary correction. + +[Note 1: A definite interval must be allowed for draining, and a +definite practice adopted with respect to the removal of the liquid +which collects at the end of the tube, if the pipette is designed to +deliver a specific volume when emptied. This liquid may be removed +at the end of a definite interval either by touching the side of the +vessel or by gently blowing out the last drops. Either practice, when +adopted, must be uniformly adhered to.] + + +FLASKS + +!Graduated or measuring flasks! are similar to the ordinary +flat-bottomed flasks, but are provided with long, narrow necks in +order that slight variations in the position of the meniscus with +respect to the graduation shall represent a minimum volume of liquid. +The flasks must be of such a capacity that, when filled with the +specified volume, the liquid rises well into the neck. + + +GRADUATION OF FLASKS + +It is a general custom to purchase the flasks ungraduated and to +graduate them for use under standard conditions selected for the +laboratory in question. They may be graduated for "contents" or +"delivery." When graduated for "contents" they contain a specified +volume when filled to the graduation at a specified temperature, and +require to be washed out in order to remove all of the solution from +the flask. Flasks graduated for "delivery" will deliver the specified +volume of a liquid without rinsing. A flask may, of course, be +graduated for both contents and delivery by placing two graduation +marks upon it. + +PROCEDURE.--To calibrate a flask for !contents!, proceed as follows: +Clean the flask, using a chromic acid solution, and dry it carefully +outside and inside. Tare it accurately; pour water into the flask +until the weight of the latter counterbalances weights on the opposite +pan which equal in grams the number of cubic centimeters of water +which the flask is to contain. Remove any excess of water with the aid +of filter paper (Note 1). Take the flask from the balance, stopper +it, place it in a bath at the desired temperature, usually 15.5° +or 17.5°C., and after an hour mark on the neck with a diamond the +location of the lowest point of the meniscus (Note 2). The mark may +be etched upon the flask by hydrofluoric acid, or by the use of an +etching ink now commonly sold on the market. + +To graduate a flask which is designed to !deliver! a specified volume, +proceed as follows: Clean the flask as usual and wipe all moisture +from the outside. Fill it with distilled water. Pour out the water +and allow the water to drain from the flask for three minutes. +Counterbalance the flask with weights to the nearest centigram. +Add weights corresponding in grams to the volume desired, and add +distilled water to counterbalance these weights. An excess of water, +or water adhering to the neck of the flask, may be removed by means of +a strip of clean filter paper. Stopper the flask, place it in a bath +at 15.5°C. or 17.5°C. and, after an hour, mark the location of the +lowest point of the meniscus, as described above. + +[Note 1: The allowable error in counterbalancing the water and +weights varies with the volume of the flask. It should not exceed one +ten-thousandth of the weight of water.] + +[Note 2: Other methods are employed which involve the use of +calibrated apparatus from which the desired volume of water may be run +into the dry flask and the position of the meniscus marked directly +upon it. For a description of a procedure which is most convenient +when many flasks are to be calibrated, the student is referred to the +!Am. Chem J.!, 16, 479.] + + + + +GENERAL DIRECTIONS FOR VOLUMETRIC ANALYSES + + +It cannot be too strongly emphasized that for the success of analyses +uniformity of practice must prevail throughout all volumetric work +with respect to those factors which can influence the accuracy of the +measurement of liquids. For example, whatever conditions are imposed +during the calibration of a burette, pipette, or flask (notably the +time allowed for draining), must also prevail whenever the flask or +burette is used. + +The student should also be constantly watchful to insure parallel +conditions during both standardization and analyst with respect to the +final volume of liquid in which a titration takes place. The value +of a standard solution is only accurate under the conditions which +prevailed when it was standardized. It is plain that the standard +solutions must be scrupulously protected from concentration or +dilution, after their value has been established. Accordingly, great +care must be taken to thoroughly rinse out all burettes, flasks, etc., +with the solutions which they are to contain, in order to remove all +traces of water or other liquid which could act as a diluent. It is +best to wash out a burette at least three times with small portions of +a solution, allowing each to run out through the tip before assuming +that the burette is in a condition to be filled and used. It is, of +course, possible to dry measuring instruments in a hot closet, but +this is tedious and unnecessary. + +To the same end, all solutions should be kept stoppered and away from +direct sunlight or heat. The bottles should be shaken before use to +collect any liquid which may have distilled from the solution and +condensed on the sides. + +The student is again reminded that variations in temperature of +volumetric solutions must be carefully noted, and care should always +be taken that no source of heat is sufficiently near the solutions to +raise the temperature during use. + +Much time may be saved by estimating the approximate volume of a +standard solution which will be required for a titration (if the data +are obtainable) before beginning the operation. It is then possible to +run in rapidly approximately the required amount, after which it is +only necessary to determine the end-point slowly and with accuracy. +In such cases, however, the knowledge of the approximate amount to be +required should never be allowed to influence the judgment regarding +the actual end-point. + + +STANDARD SOLUTIONS + +The strength or value of a solution for a specific reaction is +determined by a procedure called !Standardization!, in which the +solution is brought into reaction with a definite weight of a +substance of known purity. For example, a definite weight of pure +sodium carbonate may be dissolved in water, and the volume of a +solution of hydrochloric acid necessary to exactly neutralize the +carbonate accurately determined. From these data the strength or value +of the acid is known. It is then a !standard solution!. + + +NORMAL SOLUTIONS + +Standard solutions may be made of a purely empirical strength dictated +solely by convenience of manipulation, or the concentration may +be chosen with reference to a system which is applicable to all +solutions, and based upon chemical equivalents. Such solutions are +called !Normal Solutions! and contain such an amount of the reacting +substance per liter as is equivalent in its chemical action to one +gram of hydrogen, or eight grams of oxygen. Solutions containing one +half, one tenth, or one one-hundredth of this quantity per liter are +called, respectively, half-normal, tenth-normal, or hundredth-normal +solutions. + +Since normal solutions of various reagents are all referred to a +common standard, they have an advantage not possessed by empirical +solutions, namely, that they are exactly equivalent to each other. +Thus, a liter of a normal solution of an acid will exactly neutralize +a liter of a normal alkali solution, and a liter of a normal oxidizing +solution will exactly react with a liter of a normal reducing +solution, and so on. + +Beside the advantage of uniformity, the use of normal solutions +simplifies the calculations of the results of analyses. This is +particularly true if, in connection with the normal solution, the +weight of substance for analysis is chosen with reference to the +atomic or molecular weight of the constituent to be determined. (See +problem 26.) + +The preparation of an !exactly! normal, half-normal, or tenth-normal +solution requires considerable time and care. It is usually carried +out only when a large number of analyses are to be made, or when the +analyst has some other specific purpose in view. It is, however, a +comparatively easy matter to prepare standard solutions which differ +but slightly from the normal or half-normal solution, and these have +the advantage of practical equality; that is, two approximately +half-normal solutions are more convenient to work with than two which +are widely different in strength. It is, however, true that some of +the advantage which pertains to the use of normal solutions as regards +simplicity of calculations is lost when using these approximate +solutions. + +The application of these general statements will be made clear in +connection with the use of normal solutions in the various types of +volumetric processes which follow. + + + + +I. NEUTRALIZATION METHODS + +ALKALIMETRY AND ACIDIMETRY + + + + +GENERAL DISCUSSION + + +!Standard Acid Solutions! may be prepared from either hydrochloric, +sulphuric, or oxalic acid. Hydrochloric acid has the advantage of +forming soluble compounds with the alkaline earths, but its solutions +cannot be boiled without danger of loss of strength; sulphuric acid +solutions may be boiled without loss, but the acid forms insoluble +sulphates with three of the alkaline earths; oxalic acid can be +accurately weighed for the preparation of solutions, and its solutions +may be boiled without loss, but it forms insoluble oxalates with +three of the alkaline earths and cannot be used with certain of the +indicators. + +!Standard Alkali Solutions! may be prepared from sodium or potassium +hydroxide, sodium carbonate, barium hydroxide, or ammonia. Of sodium +and potassium hydroxide, it may be said that they can be used with all +indicators, and their solutions may be boiled, but they absorb carbon +dioxide readily and attack the glass of bottles, thereby losing +strength; sodium carbonate may be weighed directly if its purity is +assured, but the presence of carbonic acid from the carbonate is a +disadvantage with many indicators; barium hydroxide solutions may +be prepared which are entirely free from carbon dioxide, and such +solutions immediately show by precipitation any contamination from +absorption, but the hydroxide is not freely soluble in water; ammonia +does not absorb carbon dioxide as readily as the caustic alkalies, +but its solutions cannot be boiled nor can they be used with all +indicators. The choice of a solution must depend upon the nature of +the work in hand. + +A !normal acid solution! should contain in one liter that quantity of +the reagent which represents 1 gram of hydrogen replaceable by a base. +For example, the normal solution of hydrochloric acid (HCl) should +contain 36.46 grams of gaseous hydrogen chloride, since that amount +furnishes the requisite 1 gram of replaceable hydrogen. On the other +hand, the normal solution of sulphuric acid (H_{2}SO_{4}) should +contain only 49.03 grams, i.e., one half of its molecular weight in +grams. + +A !normal alkali solution! should contain sufficient alkali in a liter +to replace 1 gram of hydrogen in an acid. This quantity is represented +by the molecular weight in grams (40.01) of sodium hydroxide (NaOH), +while a sodium carbonate solution (Na_{2}CO_{3}) should contain but +one half the molecular weight in grams (i.e., 53.0 grams) in a liter +of normal solution. + +Half-normal or tenth-normal solutions are employed in most analyses +(except in the case of the less soluble barium hydroxide). Solutions +of the latter strength yield more accurate results when small +percentages of acid or alkali are to be determined. + + +INDICATORS + +It has already been pointed out that the purpose of an indicator is to +mark (usually by a change of color) the point at which just enough of +the titrating solution has been added to complete the chemical change +which it is intended to bring about. In the neutralization processes +which are employed in the measurement of alkalies (!alkalimetry!) +or acids (!acidimetry!) the end-point of the reaction should, in +principle, be that of complete neutrality. Expressed in terms of ionic +reactions, it should be the point at which the H^{+} ions from an +acid[Note 1] unite with a corresponding number of OH^{-} ions from a +base to form water molecules, as in the equation + +H^{+}, Cl^{-} + Na^{+}, OH^{-} --> Na^{+}, Cl^{-} + (H_{2}O). + +It is not usually possible to realize this condition of exact +neutrality, but it is possible to approach it with sufficient +exactness for analytical purposes, since substances are known which, +in solution, undergo a sharp change of color as soon as even a minute +excess of H^{+} or OH^{-} ions are present. Some, as will be seen, +react sharply in the presence of H^{+} ions, and others with OH^{-} +ions. These substances employed as indicators are usually organic +compounds of complex structure and are closely allied to the dyestuffs +in character. + +[Note 1: A knowledge on the part of the student of the ionic theory +as applied to aqueous solutions of electrolytes is assumed. A brief +outline of the more important applications of the theory is given in +the Appendix.] + + +BEHAVIOR OF ORGANIC INDICATORS + +The indicators in most common use for acid and alkali titrations are +methyl orange, litmus, and phenolphthalein. + +In the following discussion of the principles underlying the behavior +of the indicators as a class, methyl orange and phenolphthalein will +be taken as types. It has just been pointed out that indicators are +bodies of complicated structure. In the case of the two indicators +named, the changes which they undergo have been carefully studied by +Stieglitz (!J. Am. Chem. Soc.!, 25, 1112) and others, and it appears +that the changes involved are of two sorts: First, a rearrangement +of the atoms within the molecule, such as often occurs in organic +compounds; and, second, ionic changes. The intermolecular changes +cannot appropriately be discussed here, as they involve a somewhat +detailed knowledge of the classification and general behavior of +organic compounds; they will, therefore, be merely alluded to, and +only the ionic changes followed. + +Methyl orange is a representative of the group of indicators which, +in aqueous solutions, behave as weak bases. The yellow color which it +imparts to solutions is ascribed to the presence of the undissociated +base. If an acid, such as HCl, is added to such a solution, the acid +reacts with the indicator (neutralizes it) and a salt is formed, as +indicated by the equation: + +(M.o.)^{+}, OH^{-} + H^{+}, Cl^{-} --> (M.o.)^{+} Cl^{-} + (H_{2}O). + +This salt ionizes into (M.o.)^{+} (using this abbreviation for the +positive complex) and Cl^{-}; but simultaneously with this ionization +there appears to be an internal rearrangement of the atoms which +results in the production of a cation which may be designated as +(M'.o'.)^{+}, and it is this which imparts a characteristic red color +to the solution. As these changes occur in the presence of even a +very small excess of acid (that is, of H^{+} ions), it serves as the +desired index of their presence in the solution. If, now, an alkali, +such as NaOH, is added to this reddened solution, the reverse +series of changes takes place. As soon as the free acid present is +neutralized, the slightest excess of sodium hydroxide, acting as +a strong base, sets free the weak, little-dissociated base of the +indicator, and at the moment of its formation it reverts, because of +the rearrangement of the atoms, to the yellow form: + +OH^{-} + (M'.o'.)^{+} --> [M'.o'.OH] --> [M.o.OH]. + +Phenolphthalein, on the other hand, is a very weak, little-dissociated +acid, which is colorless in neutral aqueous solution or in the +presence of free H^{+} ions. When an alkali is added to such a +solution, even in slight excess, the anion of the salt which has +formed from the acid of the indicator undergoes a rearrangement of the +atoms, and a new ion, (Ph')^{+}, is formed, which imparts a pink color +to the solution: + +H^{+}, (Ph)^{-} + Na^{+}, OH^{-} --> (H_{2}O) + Na^{+}, (Ph)^{-} +--> Na^{+}, (Ph')^{-} + +The addition of the slightest excess of an acid to this solution, on +the other hand, occasions first the reversion to the colorless ion and +then the setting free of the undissociated acid of the indicator: + +H^{+}, (Ph')^{-} --> H^{+}, (Ph)^{-} --> (HPh). + +Of the common indicators methyl orange is the most sensitive toward +alkalies and phenolphthalein toward acids; the others occupy +intermediate positions. That methyl orange should be most sensitive +toward alkalies is evident from the following considerations: Methyl +orange is a weak base and, therefore, but little dissociated. It +should, then, be formed in the undissociated condition as soon as even +a slight excess of OH^{-} ions is present in the solution, and there +should be a prompt change from red to yellow as outlined above. On the +other hand, it should be an unsatisfactory indicator for use with weak +acids (acetic acid, for example) because the salts which it forms +with such acids are, like all salts of that type, hydrolyzed to a +considerable extent. This hydrolytic change is illustrated by the +equation: + +(M.o.)^{+} C_{2}H_{3}O_{2}^{-} + H^{+}, OH^{-} --> [M.o.OH] + H^{+}, +C_{2}H_{3}O_{2}^{-}. + +Comparison of this equation with that on page 30 will make it plain +that hydrolysis is just the reverse of neutralization and must, +accordingly, interfere with it. Salts of methyl orange with weak acids +are so far hydrolyzed that the end-point is uncertain, and methyl +orange cannot be used in the titration of such acids, while with +the very weak acids, such as carbonic acid or hydrogen sulphide +(hydrosulphuric acid), the salts formed with methyl orange are, in +effect, completely hydrolyzed (i.e., no neutralization occurs), and +methyl orange is accordingly scarcely affected by these acids. This +explains its usefulness, as referred to later, for the titration of +strong acids, such as hydrochloric acid, even in the presence of +carbonates or sulphides in solution. + +Phenolphthalein, on the other hand, should be, as it is, the best of +the common indicators for use with weak acids. For, since it is +itself a weak acid, it is very little dissociated, and its nearly +undissociated, colorless molecules are promptly formed as soon as +there is any free acid (that is, free H^{+} ions) in the solution. +This indicator cannot, however, be successfully used with weak bases, +even ammonium hydroxide; for, since it is weak acid, the salts +which it forms with weak alkalies are easily hydrolyzed, and as a +consequence of this hydrolysis the change of color is not sharp. +This indicator can, however, be successfully used with strong bases, +because the salts which it forms with such bases are much less +hydrolyzed and because the excess of OH^{-} ions from these bases also +diminishes the hydrolytic action of water. + +This indicator is affected by even so weak an acid as carbonic acid, +which must be removed by boiling the solution before titration. It is +the indicator most generally employed for the titration of organic +acids. + +In general, it may be stated that when a strong acid, such as +hydrochloric, sulphuric or nitric acid, is titrated against a strong +base, such as sodium hydroxide, potassium hydroxide, or barium +hydroxide, any of these indicators may be used, since very little +hydrolysis ensues. It has been noted above that the color change does +not occur exactly at theoretical neutrality, from which it follows +that no two indicators will show exactly the same end-point when acids +and alkalis are brought together. It is plain, therefore, that the +same indicator must be employed for both standardization and analysis, +and that, if this is done, accurate results are obtainable. + +The following table (Note 1) illustrates the variations in the volume +of an alkali solution (tenth-normal sodium hydroxide) required to +produce an alkaline end-point when run into 10 cc. of tenth-normal +sulphuric acid, diluted with 50 cc. of water, using five drops of each +of the different indicator solutions. + +==================================================================== + | | | | + INDICATOR | N/10 | N/10 |COLOR IN ACID|COLOR IN ALKA- + | H_{2}SO_{4}| NaOH |SOLUTION |LINE SOLUTION +_______________|____________|__________|_____________|______________ + | cc. | cc. | cc. | +Methyl orange | 10 | 9.90 | Red | Yellow +Lacmoid | 10 | 10.00 | Red | Blue +Litmus | 10 | 10.00 | Red | Blue +Rosalic acid | 10 | 10.07 | Yellow | Pink +Phenolphthalein| 10 | 10.10 | Colorless | Pink +==================================================================== + +It should also be stated that there are occasionally secondary +changes, other than those outlined above, which depend upon the +temperature and concentration of the solutions in which the indicators +are used. These changes may influence the sensitiveness of an +indicator. It is important, therefore, to take pains to use +approximately the same volume of solution when standardizing that is +likely to be employed in analysis; and when it is necessary, as is +often the case, to titrate the solution at boiling temperature, the +standardization should take place under the same conditions. It is +also obvious that since some acid or alkali is required to react with +the indicator itself, the amount of indicator used should be uniform +and not excessive. Usually a few drops of solution will suffice. + +The foregoing statements with respect to the behavior of indicators +present the subject in its simplest terms. Many substances other than +those named may be employed, and they have been carefully studied to +determine the exact concentration of H^{+} ions at which the color +change of each occurs. It is thus possible to select an indicator +for a particular purpose with considerable accuracy. As data of this +nature do not belong in an introductory manual, reference is made to +the following papers or books in which a more extended treatment of +the subject may be found: + +Washburn, E.W., Principles of Physical Chemistry (McGraw-Hill Book +Co.), (Second Edition, 1921), pp. 380-387. + +Prideaux, E.B.R., The Theory and Use of Indicators (Constable & Co., +Ltd.), (1917). + +Salm, E., A Study of Indicators, !Z. physik. Chem.!, 57 (1906), +471-501. + +Stieglitz, J., Theories of Indicators, !J. Am. Chem. Soc.!, 25 (1903), +1112-1127. + +Noyes, A.A., Quantitative Applications of the Theory of Indicators to +Volumetric Analysis, !J. Am. Chem. Soc.!, 32 (1911), 815-861. + +Bjerrum, N., General Discussion, !Z. Anal. Chem.!, 66 (1917), 13-28 +and 81-95. + +Ostwald, W., Colloid Chemistry of Indicators, !Z. Chem. Ind. +Kolloide!, 10 (1912), 132-146. + +[Note 1: Glaser, !Indikatoren der Acidimetrie und Alkalimetrie!. +Wiesbaden, 1901.] + + +PREPARATION OF INDICATOR SOLUTIONS + +A !methyl orange solution! for use as an indicator is commonly made by +dissolving 0.05-0.1 gram of the compound (also known as Orange III) in +a few cubic centimeters of alcohol and diluting with water to 100 cc. +A good grade of material should be secured. It can be successfully +used for the titration of hydrochloric, nitric, sulphuric, phosphoric, +and sulphurous acids, and is particularly useful in the determination +of bases, such as sodium, potassium, barium, calcium, and ammonium +hydroxides, and even many of the weak organic bases. It can also be +used for the determination, by titration with a standard solution of +a strong acid, of the salts of very weak acids, such as carbonates, +sulphides, arsenites, borates, and silicates, because the weak acids +which are liberated do not affect the indicator, and the reddening of +the solution does not take place until an excess of the strong acid +is added. It should be used in cold, not too dilute, solutions. Its +sensitiveness is lessened in the presence of considerable quantities +of the salts of the alkalies. + +A !phenolphthalein solution! is prepared by dissolving 1 gram of the +pure compound in 100 cc. of 95 per cent alcohol. This indicator is +particularly valuable in the determination of weak acids, especially +organic acids. It cannot be used with weak bases, even ammonia. It +is affected by carbonic acid, which must, therefore, be removed by +boiling when other acids are to be measured. It can be used in hot +solutions. Some care is necessary to keep the volume of the solutions +to be titrated approximately uniform in standardization and in +analysis, and this volume should not in general exceed 125-150 cc. for +the best results, since the compounds formed by the indicator undergo +changes in very dilute solution which lessen its sensitiveness. + +The preparation of a !solution of litmus! which is suitable for use +as an indicator involves the separation from the commercial litmus of +azolithmine, the true coloring principle. Soluble litmus tablets are +often obtainable, but the litmus as commonly supplied to the market is +mixed with calcium carbonate or sulphate and compressed into lumps. To +prepare a solution, these are powdered and treated two or three times +with alcohol, which dissolves out certain constituents which cause a +troublesome intermediate color if not removed. The alcohol is decanted +and drained off, after which the litmus is extracted with hot water +until exhausted. The solution is allowed to settle for some time, the +clear liquid siphoned off, concentrated to one-third its volume and +acetic acid added in slight excess. It is then concentrated to a +sirup, and a large excess of 95 per cent. alcohol added to it. This +precipitates the blue coloring matter, which is filtered off, washed +with alcohol, and finally dissolved in a small volume of water and +diluted until about three drops of the solution added to 50 cc. of +water just produce a distinct color. This solution must be kept in an +unstoppered bottle. It should be protected from dust by a loose plug +of absorbent cotton. If kept in a closed bottle it soon undergoes a +reduction and loses its color, which, however, is often restored by +exposure to the air. + +Litmus can be employed successfully with the strong acids and bases, +and also with ammonium hydroxide, although the salts of the latter +influence the indicator unfavorably if present in considerable +concentration. It may be employed with some of the stronger organic +acids, but the use of phenolphthalein is to be preferred. + + +PREPARATION OF STANDARD SOLUTIONS + +!Hydrochloric Acid and Sodium Hydroxide. Approximate Strength!, 0.5 N + + +PROCEDURE.--Measure out 40 cc. of concentrated, pure hydrochloric +acid into a clean liter bottle, and dilute with distilled water to an +approximate volume of 1000 cc. Shake the solution vigorously for a +full minute to insure uniformity. Be sure that the bottle is not too +full to permit of a thorough mixing, since lack of care at this point +will be the cause of much wasted time (Note 1). + +Weigh out, upon a rough balance, 23 grams of sodium hydroxide (Note +2). Dissolve the hydroxide in water in a beaker. Pour the solution +into a liter bottle and dilute, as above, to approximately 1000 cc. +This bottle should preferably have a rubber stopper, as the hydroxide +solution attacks the glass of the ground joint of a glass stopper, and +may cement the stopper to the bottle. Shake the solution as described +above. + +[Note 1: The original solutions are prepared of a strength greater +than 0.5 N, as they are more readily diluted than strengthened if +later adjustment is desired. + +Too much care cannot be taken to insure perfect uniformity of +solutions before standardization, and thoroughness in this respect +will, as stated, often avoid much waste of time. A solution once +thoroughly mixed remains uniform.] + +[Note 2: Commercial sodium hydroxide is usually impure and always +contains more or less carbonate; an allowance is therefore made for +this impurity by placing the weight taken at 23 grams per liter. If +the hydroxide is known to be pure, a lesser amount (say 21 grams) will +suffice.] + + +COMPARISON OF ACID AND ALKALI SOLUTIONS + +PROCEDURE.--Rinse a previously calibrated burette three times with the +hydrochloric acid solution, using 10 cc. each time, and allowing the +liquid to run out through the tip to displace all water and air +from that part of the burette. Then fill the burette with the acid +solution. Carry out the same procedure with a second burette, using +the sodium hydroxide solution. + +The acid solution may be placed in a plain or in a glass-stoppered +burette as may be more convenient, but the alkaline solution should +never be allowed to remain long in a glass-stoppered burette, as it +tends to cement the stopper to the burette, rendering it useless. It +is preferable to use a plain burette for this solution. + +When the burettes are ready for use and all air bubbles displaced from +the tip (see Note 2, page 17) note the exact position of the liquid in +each, and record the readings in the notebook. (Consult page 188.) Run +out from the burette into a beaker about 40 cc. of the acid and add +two drops of a solution of methyl orange; dilute the acid to about +80 cc. and run out alkali solution from the other burette, stirring +constantly, until the pink has given place to a yellow. Wash down the +sides of the beaker with a little distilled water if the solution has +spattered upon them, return the beaker to the acid burette, and add +acid to restore the pink; continue these alternations until the point +is accurately fixed at which a single drop of either solutions served +to produce a distinct change of color. Select as the final end-point +the appearance of the faintest pink tinge which can be recognized, or +the disappearance of this tinge, leaving a pure yellow; but always +titrate to the same point (Note 1). If the titration has occupied more +than the three minutes required for draining the sides of the burette, +the final reading may be taken immediately and recorded in the +notebook. + +Refill the burettes and repeat the titration. From the records of +calibration already obtained, correct the burette readings and make +corrections for temperature, if necessary. Obtain the ratio of the +sodium hydroxide solution to that of hydrochloric acid by dividing +the number of cubic centimeters of acid used by the number of cubic +centimeters of alkali required for neutralization. The check results +of the two titrations should not vary by more than two parts in one +thousand (Note 2). If the variation in results is greater than this, +refill the burettes and repeat the titration until satisfactory values +are obtained. Use a new page in the notebook for each titration. +Inaccurate values should not be erased or discarded. They should be +retained and marked "correct" or "incorrect," as indicated by the +final outcome of the titrations. This custom should be rigidly +followed in all analytical work. + +[Note 1: The end-point should be chosen exactly at the point of +change; any darker tint is unsatisfactory, since it is impossible to +carry shades of color in the memory and to duplicate them from day to +day.] + +[Note 2: While variation of two parts in one thousand in the values +obtained by an inexperienced analyst is not excessive, the idea must +be carefully avoided that this is a standard for accurate work to be +!generally applied!. In many cases, after experience is gained, the +allowable error is less than this proportion. In a few cases a +larger variation is permissible, but these are rare and can only +be recognized by an experienced analyst. It is essential that the +beginner should acquire at least the degree of accuracy indicated if +he is to become a successful analyst.] + + + + +STANDARDIZATION OF HYDROCHLORIC ACID + +SELECTION AND PREPARATION OF STANDARD + + +The selection of the best substance to be used as a standard for acid +solutions has been the subject of much controversy. The work of Lunge +(!Ztschr. angew. Chem.! (1904), 8, 231), Ferguson (!J. Soc. Chem. +Ind.! (1905), 24, 784), and others, seems to indicate that the best +standard is sodium carbonate prepared from sodium bicarbonate by +heating the latter at temperature between 270° and 300°C. The +bicarbonate is easily prepared in a pure state, and at the +temperatures named the decomposition takes place according to the +equation + +2HNaCO_{3} --> Na_{2}CO_{3} + H_{2}O + CO_{2} + +and without loss of any carbon dioxide from the sodium carbonate, such +as may occur at higher temperatures. The process is carried out as +described below. + +PROCEDURE.--Place in a porcelain crucible about 6 grams (roughly +weighed) of the purest sodium bicarbonate obtainable. Rest the +crucible upon a triangle of iron or copper wire so placed within a +large crucible that there is an open air space of about three eighths +of an inch between them. The larger crucible may be of iron, nickel or +porcelain, as may be most convenient. Insert the bulb of a thermometer +reading to 350°C. in the bicarbonate, supporting it with a clamp so +that the bulb does not rest on the bottom of the crucible. Heat +the outside crucible, using a rather small flame, and raise the +temperature of the bicarbonate fairly rapidly to 270°C. Then regulate +the heat in such a way that the temperature rises !slowly! to 300°C. +in the course of a half-hour. The bicarbonate should be frequently +stirred with a clean, dry, glass rod, and after stirring, should be +heaped up around the bulb of the thermometer in such a way as to cover +it. This will require attention during most of the heating, as the +temperature should not be permitted to rise above 310°C. for any +length of time. At the end of the half-hour remove the thermometer and +transfer the porcelain crucible, which now contains sodium carbonate, +to a desiccator. When it is cold, transfer the carbonate to a +stoppered weighing tube or weighing-bottle. + + +STANDARDIZATION + +PROCEDURE.--Clean carefully the outside of a weighing-tube, or +weighing-bottle, containing the pure sodium carbonate, taking care +to handle it as little as possible after wiping. Weigh the tube +accurately to 0.0001 gram, and record the weight in the notebook. Hold +the tube over the top of a beaker (200-300 cc.) and cautiously remove +the stopper, making sure that no particles fall from it or from the +tube elsewhere than in the beaker. Pour out from the tube a portion +of the carbonate, replace the stopper and determine approximately how +much has been removed. Continue this procedure until 1.00 to 1.10 +grams has been taken from the tube. Then weigh the tube accurately +and record the weight under the first weight in the notebook. +The difference in the two weights is the weight of the carbonate +transferred to the beaker. Proceed in the same way to transfer a +second portion of the carbonate from the tube to another beaker of +about the same size as the first. The beakers should be labeled and +plainly marked to correspond with the entries in the notebook. + +Pour over the carbonate in each beaker about 80 cc. of water, stir +until solution is complete, and add two drops of methyl orange +solution. Fill the burettes with the standard acid and alkali +solutions, noting the initial readings of the burettes and temperature +of the solutions. Run in acid from the burette, stirring and avoiding +loss by effervescence, until the solution has become pink. Wash down +the sides of the beaker with a !little! water from a wash-bottle, and +then run in alkali from the other burette until the pink is replaced +by yellow; then finish the titration as described on page 37. Note the +readings of the burettes after the proper interval, and record them in +the notebook. Repeat the procedure, using the second portion of sodium +carbonate. Apply the necessary calibration corrections to the volumes +of the solutions used, and correct for temperature if necessary. + +From the data obtained, calculate the volume of the hydrochloric +acid solution which is equivalent to the volume of sodium hydroxide +solution used in this titration. Subtract this volume from the volume +of hydrochloric acid. The difference represents the volume of acid +used to react with the sodium carbonate. Divide the weight of sodium +carbonate by this volume in cubic centimeters, thus obtaining the +weight of sodium carbonate equivalent to each cubic centimeter of the +acid. + +From this weight it is possible to calculate the corresponding weight +of HCl in each cubic centimeter of the acid, and in turn the relation +of the acid to the normal. + +If, however, it is recalled that normal solutions are equivalent to +each other, it will be seen that the same result may be more readily +reached by dividing the weight in grams of sodium carbonate per cubic +centimeter just found by titration by the weight which would be +contained in the same volume of a normal solution of sodium carbonate. +A normal solution of sodium carbonate contains 53.0 grams per liter, +or 0.0530 gram per cc. (see page 29). The relation of the acid +solution to the normal is, therefore, calculated by dividing the +weight of the carbonate to which each cubic centimeter of the acid is +equivalent by 0.0530. The standardization must be repeated until the +values obtained agree within, at most, two parts in one thousand. + +When the standard of the acid solution has been determined, calculate, +from the known ratio of the two solutions, the relation of the sodium +hydroxide solution to a normal solution (Notes 1 and 2). + +[Note 1: In the foregoing procedure the acid solution is standardized +and the alkali solution referred to this standard by calculation. It +is equally possible, if preferred, to standardize the alkali solution. +The standards in a common use for this purpose are purified +oxalic acid (H_{2}C_{2}O_{4}.2H_{2}O), potassium acid oxalate +(KHC_{2}O_{4}.H_{2}O or KHC_{2}O_{4}), potassium tetroxalate +(KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O), or potassium acid tartrate +(KHC_{4}O_{6}), with the use of a suitable indicator. The oxalic acid +and the oxalates should be specially prepared to insure purity, +the main difficulty lying in the preservation of the water of +crystallization. + +It should be noted that the acid oxalate and the acid tartrate each +contain one hydrogen atom replaceable by a base, while the tetroxalate +contains three such atoms and the oxalic acid two. Each of the two +salts first named behave, therefore, as monobasic acids, and the +tetroxalate as a tribasic acid.] + +[Note 2: It is also possible to standardize a hydrochloric acid +solution by precipitating the chloride ions as silver chloride and +weighing the precipitate, as prescribed under the analysis of sodium +chloride to be described later. Sulphuric acid solutions may be +standardized by precipitation of the sulphate ions as barium sulphate +and weighing the ignited precipitate, but the results are not above +criticism on account of the difficulty in obtaining large precipitates +of barium sulphate which are uncontaminated by inclosures or are not +reduced on ignition.] + + + + +DETERMINATION OF THE TOTAL ALKALINE STRENGTH OF SODA ASH + + +Soda ash is crude sodium carbonate. If made by the ammonia process it +may contain also sodium chloride, sulphate, and hydroxide; when made +by the Le Blanc process it may contain sodium sulphide, silicate, and +aluminate, and other impurities. Some of these, notably the hydroxide, +combine with acids and contribute to the total alkaline strength, +but it is customary to calculate this strength in terms of sodium +carbonate; i.e., as though no other alkali were present. + +PROCEDURE.--In order to secure a sample which shall represent the +average value of the ash, it is well to take at least 5 grams. As this +is too large a quantity for convenient titration, an aliquot portion +of the solution is measured off, representing one fifth of the entire +quantity. This is accomplished as follows: Weigh out on an analytical +balance two samples of soda ash of about 5 grams each into beakers +of about 500 cc. capacity. (The weighings need be made to centigrams +only.) Dissolve the ash in 75 cc. of water, warming gently, and filter +off the insoluble residue; wash the filter by filling it at least +three times with distilled water, and allowing it to drain, adding the +washings to the main filtrate. Cool the filtrate to approximately the +standard temperature of the laboratory, and transfer it to a 250 cc. +measuring flask, washing out the beaker thoroughly. Add distilled +water of laboratory temperature until the lowest point of the meniscus +is level with the graduation on the neck of the flask and remove any +drops of water that may be on the neck above the graduation by means +of a strip of filter paper; make the solution thoroughly uniform by +pouring it out into a dry beaker and back into the flask several +times. Measure off 50 cc. of the solution in a measuring flask, or +pipette, either of which before use should, unless they are dry on the +inside, be rinsed out with at least two small portions of the soda ash +solution to displace any water. + +If a flask is used, fill it to the graduation with the soda ash +solution and remove any liquid from the neck above the graduation with +filter paper. Empty it into a beaker, and wash out the small flask, +unless it is graduated for !delivery!, using small quantities of +water, which are added to the liquid in the beaker. A second 50 cc. +portion from the main solution should be measured off into a second +beaker. Dilute the solutions in each beaker to 100 cc., add two drops +of a solution of methyl orange (Note 1) and titrate for the alkali +with the standard hydrochloric acid solution, using the alkali +solution to complete the titration as already prescribed. + +From the volumes of acid and alkali employed, corrected for burette +errors and temperature changes, and the data derived from the +standardization, calculate the percentage of alkali present, assuming +it all to be present as sodium carbonate (Note 2). + +[Note 1: The hydrochloric acid sets free carbonic acid which is +unstable and breaks down into water and carbon dioxide, most of which +escapes from the solution. Carbonic acid is a weak acid and, as such, +does not yield a sufficient concentration of H^{+} ions to cause the +indicator to change to a pink (see page 32). + +The chemical changes involved may be summarized as follows: + +2H^{+}, 2Cl^{-} + 2Na^{+}, CO_{3}^{--} --> 2Na^{+}, 2Cl^{-} + +[H_{2}CO_{3}] --> H_{2}O + CO_{2}] + +[Note 2: A determination of the alkali present as hydroxide in soda +ash may be determined by precipitating the carbonate by the addition +of barium chloride, removing the barium carbonate by filtration, and +titrating the alkali in the filtrate. + +The caustic alkali may also be determined by first using +phenolphthalein as an indicator, which will show by its change from +pink to colorless the point at which the caustic alkali has been +neutralized and the carbonate has been converted to bicarbonate, and +then adding methyl orange and completing the titration. The amount of +acid necessary to change the methyl orange to pink is a measure of one +half of the carbonate present. The results of the double titration +furnish the data necessary for the determination of the caustic alkali +and of the carbonate in the sample.] + + + + +DETERMINATION OF THE ACID STRENGTH OF OXALIC ACID + + +PROCEDURE.--Weigh out two portions of the acid of about 1 gram +each. Dissolve these in 50 cc. of warm water. Add two drops of +phenolphthalein solution, and run in alkali from the burette until the +solution is pink; add acid from the other burette until the pink is +just destroyed, and then add 0.3 cc. (not more) in excess. Heat the +solution to boiling for three minutes. If the pink returns during the +boiling, discharge it with acid and again add 0.3 cc. in excess and +repeat the boiling (Note 1). If the color does not then reappear, add +alkali until it does, and a !drop or two! of acid in excess and boil +again for one minute (Note 2). If no color reappears during this time, +complete the titration in the hot solution. The end-point should be +the faintest visible shade of color (or its disappearance), as the +same difficulty would exist here as with methyl orange if an attempt +were made to match shades of pink. + +From the corrected volume of alkali required to react with the +oxalic acid, calculate the percentage of the crystallized acid +(H_{2}C_{2}O_{4}.2H_{2}O) in the sample (Note 3). + +[Note 1: All commercial caustic soda such as that from which the +standard solution was made contains some sodium carbonate. This reacts +with the oxalic acid, setting free carbonic acid, which, in turn, +forms sodium bicarbonate with the remaining carbonate: + +H_{2}CO_{3} + Na_{2}CO_{3} --> 2HNaCO_{3}. + +This compound does not hydrolyze sufficiently to furnish enough OH^{-} +ions to cause phenolphthalein to remain pink; hence, the color of +the indicator is discharged in cold solutions at the point at which +bicarbonate is formed. If, however, the solution is heated to boiling, +the bicarbonate loses carbon dioxide and water, and reverts to sodium +carbonate, which causes the indicator to become again pink: + +2HNaCO_{3} --> H_{2}O + CO_{2} + Na_{2}CO_{3}. + +By adding successive portions of hydrochloric acid and boiling, the +carbonate is ultimately all brought into reaction. + +The student should make sure that the difference in behavior of the +two indicators, methyl orange and phenolphthalein, is understood.] + +[Note 2: Hydrochloric acid is volatilized from aqueous solutions, +except such as are very dilute. If the directions in the procedure +are strictly followed, no loss of acid need be feared, but the amount +added in excess should not be greater than 0.3-0.4 cc.] + +[Note 3: Attention has already been called to the fact that the color +changes in the different indicators occur at varying concentrations +of H^{+} or OH^{-} ions. They do not indicate exact theoretical +neutrality, but a particular indicator always shows its color change +at a particular concentration of H^{+} or OH^{-} ions. The results +of titration with a given indicator are, therefore, comparable. As a +matter of fact, a small error is involved in the procedure as outlined +above. The comparison of the acid and alkali solutions was made, using +methyl orange as an indicator, while the titration of the oxalic acid +is made with the use of phenolphthalein. For our present purposes the +small error may be neglected but, if time permits, the student is +recommended to standardize the alkali solution against one of the +substances named in Note 1, page 41, and also to ascertain +the comparative value of the acid and alkali solutions, using +phenolphthalein as indicator throughout, and conducting the titrations +as described above. This will insure complete accuracy.] + + + + +II. OXIDATION PROCESSES + +GENERAL DISCUSSION + + +In the oxidation processes of volumetric analysis standard solutions +of oxidizing agents and of reducing agents take the place of the acid +and alkali solutions of the neutralization processes already studied. +Just as an acid solution was the principal reagent in alkalimetry, and +the alkali solution used only to make certain of the end-point, the +solution of the oxidizing agent is the principal reagent for the +titration of substances exerting a reducing action. It is, in general, +true that oxidizable substances are determined by !direct! titration, +while oxidizing substances are determined by !indirect! titration. + +The important oxidizing agents employed in volumetric solutions are +potassium bichromate, potassium permangenate, potassium ferricyanide, +iodine, ferric chloride, and sodium hypochlorite. + +The important reducing agents which are used in the form of standard +solutions are ferrous sulphate (or ferrous ammonium sulphate), oxalic +acid, sodium thiosulphate, stannous chloride, arsenious acid, and +potassium cyanide. Other reducing agents, as sulphurous acid, +sulphureted hydrogen, and zinc (nascent hydrogen), may take part in +the processes, but not as standard solutions. + +The most important combinations among the foregoing are: Potassium +bichromate and ferrous salts; potassium permanganate and ferrous +salts; potassium permanganate and oxalic acid, or its derivatives; +iodine and sodium thiosulphate; hypochlorites and arsenious acid. + + + + +BICHROMATE PROCESS FOR THE DETERMINATION OF IRON + + +Ferrous salts may be promptly and completely oxidized to ferric salts, +even in cold solution, by the addition of potassium bichromate, +provided sufficient acid is present to hold in solution the ferric and +chromic compounds which are formed. + +The acid may be either hydrochloric or sulphuric, but the former is +usually preferred, since it is by far the best solvent for iron and +its compounds. The reaction in the presence of hydrochloric acid is as +follows: + +6FeCl_{2} + K_{2}Cr_{2}O_{7} + 14HCl --> 6FeCl_{3} + 2CrCl_{3} + 2KCl ++ 7H_{2}O. + + +NORMAL SOLUTIONS OF OXIDIZING OR REDUCING AGENTS + +It will be recalled that the system of normal solutions is based upon +the equivalence of the reagents which they contain to 8 grams of +oxygen or 1 gram of hydrogen. A normal solution of an oxidizing agent +should, therefore, contain that amount per liter which is equivalent +in oxidizing power to 8 grams of oxygen; a normal reducing solution +must be equivalent in reducing power to 1 gram of hydrogen. In order +to determine what the amount per liter will be it is necessary to know +how the reagents enter into reaction. The two solutions to be employed +in the process under consideration are those of potassium bichromate +and ferrous sulphate. The reaction between them, in the presence of an +excess of sulphuric acid, may be expressed as follows: + +6FeSO_{4} + K_{2}Cr_{2}O_{7} + 7H_{2}SO_{4} --> 3Fe_{2}(SO_{4})_{3} + +K_{2}SO_{4} + Cr_{2}(SO_{4})_{3} + 7H_{2}O. + +If the compounds of iron and chromium, with which alone we are now +concerned, be written in such a way as to show the oxides of these +elements in each, they would appear as follows: On the left-hand side +of the equation 6(FeO.SO_{3}) and K_{2}O.2CrO_{3}; on the right-hand +side, 3(Fe_{2}O_{3}.3SO_{3}) and Cr_{2}O_{3}.3SO_{3}. A careful +inspection shows that there are three less oxygen atoms associated +with chromium atoms on the right-hand side of the equation than on the +left-hand, but there are three more oxygen atoms associated with iron +atoms on the right than on the left. In other words, a molecule of +potassium bichromate has given up three atoms of oxygen for oxidation +purposes; i.e., a molecular weight in grams of the bichromate (294.2) +will furnish 3 X 16 or 48 grams of oxygen for oxidation purposes. +As this 48 grams is six times 8 grams, the basis of the system, the +normal solution of potassium bichromate should contain per liter one +sixth of 294.2 grams or 49.03 grams. + +A further inspection of the dissected compounds above shows that six +molecules of FeO.SO_{3} were required to react with the three atoms of +oxygen from the bichromate. From the two equations + +3H_{2} + 3O --> 3H_{2}O +6(FeO.SO_{3}) + 3O --> 3(Fe_{2}O_{3}.3SO_{3}) + +it is plain that one molecule of ferrous sulphate is equivalent to one +atom of hydrogen in reducing power; therefore one molecular weight in +grams of ferrous sulphate (151.9) is equivalent to 1 gram of +hydrogen. Since the ferrous sulphate crystalline form has the formula +FeSO_{4}.7H_{2}O, a normal reducing solution of this crystalline salt +should contain 277.9 grams per liter. + + +PREPARATION OF SOLUTIONS + +!Approximate Strength 0.1 N! + +It is possible to purify commercial potassium bichromate by +recrystallization from hot water. It must then be dried and cautiously +heated to fusion to expel the last traces of moisture, but not +sufficiently high to expel any oxygen. The pure salt thus prepared, +may be weighed out directly, dissolved, and the solution diluted in a +graduated flask to a definite volume. In this case no standardization +is made, as the normal value can be calculated directly. It is, +however, more generally customary to standardize a solution of +the commercial salt by comparison with some substance of definite +composition, as described below. + +PROCEDURE.--Pulverize about 5 grams of potassium bichromate of good +quality. Dissolve the bichromate in distilled water, transfer the +solution to a liter bottle, and dilute to approximately 1000 cc. Shake +thoroughly until the solution is uniform. + +To prepare the solution of the reducing agent, pulverize about 28 +grams of ferrous sulphate (FeSO_{4}.7H_{2}O) or about 40 grams of +ferrous ammonium sulphate (FeSO_{4}.(NH_{4})_{2}SO_{4}.6H_{2}O) and +dissolve in distilled water containing 5 cc. of concentrated sulphuric +acid. Transfer the solution to a liter bottle, add 5 cc. concentrated +sulphuric acid, make up to about 1000 cc. and shake vigorously to +insure uniformity. + + +INDICATOR SOLUTION + +No indicator is known which, like methyl orange, can be used within +the solution, to show when the oxidation process is complete. Instead, +an outside indicator solution is employed to which drops of the +titrated solution are transferred for testing. The reagent used is +potassium ferricyanide, which produces a blue precipitate (or color) +with ferrous compounds as long as there are unoxidized ferrous ions in +the titrated solution. Drops of the indicator solution are placed upon +a glazed porcelain tile, or upon white cardboard which has been coated +with paraffin to render it waterproof, and drops of the titrated +solution are transferred to the indicator on the end of a stirring +rod. When the oxidation is nearly completed only very small amounts +of the ferrous compounds remain unoxidized and the reaction with the +indicator is no longer instantaneous. It is necessary to allow a brief +time to elapse before determining that no blue color is formed. Thirty +seconds is a sufficient interval, and should be adopted throughout the +analytical procedure. If left too long, the combined effect of light +and dust from the air will cause a reduction of the ferric compounds +already formed and a resultant blue will appear which misleads the +observer with respect to the true end-point. + +The indicator solution must be highly diluted, otherwise its own color +interferes with accurate observation. Prepare a fresh solution, as +needed each day, by dissolving a crystal of potassium ferricyanide +about the size of a pin's head in 25 cc. of distilled water. The salt +should be carefully tested with ferric chloride for the presence of +ferrocyanides, which give a blue color with ferric salts. + +In case of need, the ferricyanide can be purified by adding to its +solution a little bromine water and recrystallizing the compound. + + +COMPARISON OF OXIDIZING AND REDUCING SOLUTIONS + +PROCEDURE.--Fill one burette with each of the solutions, observing +the general procedure with respect to cleaning and rinsing already +prescribed. The bichromate solution is preferably to be placed in a +glass-stoppered burette. + +Run out from a burette into a beaker of about 300 cc. capacity nearly +40 cc. of the ferrous solution, add 15 cc. of dilute hydrochloric acid +(sp. gr. 1.12) and 150 cc. of water and run in the bichromate +solution from another burette. Since both solutions are approximately +tenth-normal, 35 cc. of the bichromate solution may be added without +testing. Test at that point by removing a very small drop of the +iron solution on the end of a stirring rod, mixing it with a drop of +indicator on the tile (Note 1). If a blue precipitate appears at once, +0.5 cc. of the bichromate solution may be added before testing again. +The stirring rod which has touched the indicator should be dipped in +distilled water before returning it to the iron solution. As soon as +the blue appears to be less intense, add the bichromate solution in +small portions, finally a single drop at a time, until the point is +reached at which no blue color appears after the lapse of thirty +seconds from the time of mixing solution and indicator. At the close +of the titration a large drop of the iron solution should be taken for +the test. To determine the end-point beyond any question, as soon as +the thirty seconds have elapsed remove another drop of the solution +of the same size as that last taken and mix it with the indicator, +placing it beside the last previous test. If this last previous test +shows a blue tint in comparison with the fresh mixture, the end-point +has not been reached; if no difference can be noted the reaction is +complete. Should the end-point be overstepped, a little more of the +ferrous solution may be added and the end-point definitely fixed. + +From the volumes of the solutions used, after applying corrections for +burette readings, and, if need be, for the temperature of solutions, +calculate the value of the ferrous solution in terms of the oxidizing +solution. + +[Note 1: The accuracy of the work may be much impaired by the removal +of unnecessarily large quantities of solution for the tests. At the +beginning of the titration, while much ferrous iron is still present, +the end of the stirring rod need only be moist with the solution; but +at the close of the titration drops of considerable size may properly +be taken for the final tests. The stirring rod should be washed to +prevent transfer of indicator to the main solution. This cautious +removal of solution does not seriously affect the accuracy of the +determination, as it will be noted that the volume of the titrated +solution is about 200 cc. and the portions removed are very +small. Moreover, if the procedure is followed as prescribed, the +concentration of unoxidized iron decreases very rapidly as the +titration is carried out so that when the final tests are made, though +large drops may be taken, the amount of ferrous iron is not sufficient +to produce any appreciable error in results. + +If the end-point is determined as prescribed, it can be as accurately +fixed as that of other methods; and if a ferrous solution is at +hand, the titration need consume hardly more time than that of the +permanganate process to be described later on.] + + +STANDARDIZATION OF POTASSIUM BICHROMATE SOLUTIONS + +!Selection of a Standard! + +A substance which will serve satisfactorily as a standard for +oxidizing solutions must possess certain specific properties: It must +be of accurately known composition and definite in its behavior as a +reducing agent, and it must be permanent against oxidation in the air, +at least for considerable periods. Such standards may take the form of +pure crystalline salts, such as ferrous ammonium sulphate, or may be +in the form of iron wire or an iron ore of known iron content. It is +not necessary that the standard should be of 100 per cent purity, +provided the content of the active reducing agent is known and no +interfering substances are present. + +The two substances most commonly used as standards for a bichromate +solution are ferrous ammonium sulphate and iron wire. A standard wire +is to be purchased in the market which answers the purpose well, and +its iron content may be determined for each lot purchased by a number +of gravimetric determinations. It may best be preserved in jars +containing calcium chloride, but this must not be allowed to come +into contact with the wire. It should, however, even then be examined +carefully for rust before use. + +If pure ferrous ammonium sulphate is used as the standard, clear +crystals only should be selected. It is perhaps even better to +determine by gravimetric methods once for all the iron content of a +large commercial sample which has been ground and well mixed. This +salt is permanent over long periods if kept in stoppered containers. + + +STANDARDIZATION + +PROCEDURE.--Weigh out two portions of iron wire of about 0.24-0.26 +gram each, examining the wire carefully for rust. It should be handled +and wiped with filter paper (not touched by the fingers), should +be weighed on a watch-glass, and be bent in such a way as not to +interfere with the movement of the balance. + +Place 30 cc. of hydrochloric acid (sp. gr. 1.12) in each of two 300 +cc. Erlenmeyer flasks, cover them with watch-glasses, and bring the +acid just to boiling. Remove them from the flame and drop in the +portions of wire, taking great care to avoid loss of liquid during +solution. Boil for two or three minutes, keeping the flasks covered +(Note 1), then wash the sides of the flasks and the watch-glass with +a little water and add stannous chloride solution to the hot liquid +!from a dropper! until the solution is colorless, but avoid more than +a drop or two in excess (Note 2). Dilute with 150 cc. of water and +cool !completely!. When cold, add rapidly about 30 cc. of mercuric +chloride solution. Allow the solutions to stand about three minutes +and then titrate without further delay (Note 3), add about 35 cc. of +the standard solution at once and finish the titration as prescribed +above, making use of the ferrous solution if the end-point should be +passed. + +From the corrected volumes of the bichromate solution required to +oxidize the iron actually know to be present in the wire, calculate +the relation of the standard solution to the normal. + +Repeat the standardization until the results are concordant within at +least two parts in one thousand. + + +[Note 1: The hydrochloric acid is added to the ferrous solution +to insure the presence of at least sufficient free acid for the +titration, as required by the equation on page 48. + +The solution of the wire in hot acid and the short boiling insure the +removal of compounds of hydrogen and carbon which are formed from the +small amount of carbon in the iron. These might be acted upon by the +bichromate if not expelled.] + +[Note 2: It is plain that all the iron must be reduced to the ferrous +condition before the titration begins, as some oxidation may have +occurred from the oxygen of the air during solution. It is also +evident that any excess of the agent used to reduce the iron must be +removed; otherwise it will react with the bichromate added later. + +The reagents available for the reduction of iron are stannous +chloride, sulphurous acid, sulphureted hydrogen, and zinc; of these +stannous chloride acts most readily, the completion of the reaction +is most easily noted, and the excess of the reagent is most readily +removed. The latter object is accomplished by oxidation to stannic +chloride by means of mercuric chloride added in excess, as the +mercuric salts have no effect upon ferrous iron or the bichromate. The +reactions involved are: + +2FeCl_{3} + SnCl_{2} --> 2FeCl_{2} + SnCl_{4} +SnCl_{2} + 2HgCl_{2} --> SnCl_{4} + 2HgCl + +The mercurous chloride is precipitated. + +It is essential that the solution should be cold and that the stannous +chloride should not be present in great excess, otherwise a secondary +reaction takes place, resulting in the reduction of the mercurous +chloride to metallic mercury: + +SnCl_{2} + 2HgCl --> SnCl_{4} + 2Hg. + +The occurrence of this secondary reaction is indicated by the +darkening of the precipitate; and, since potassium bichromate oxidizes +this mercury slowly, solutions in which it has been precipitated are +worthless as iron determinations.] + +[Note 3: The solution should be allowed to stand about three minutes +after the addition of mercuric chloride to permit the complete +deposition of mercurous chloride. It should then be titrated without +delay to avoid possible reoxidation of the iron by the oxygen of the +air.] + + + + +DETERMINATION OF IRON IN LIMONITE + + +PROCEDURE.--Grind the mineral (Note 1) to a fine powder. Weigh out +accurately two portions of about 0.5 gram (Note 2) into porcelain +crucibles; heat these crucibles to dull redness for ten minutes, +allow them to cool, and place them, with their contents, in beakers +containing 30 cc. of dilute hydrochloric acid (sp. gr. 1.12). Heat +at a temperature just below boiling until the undissolved residue is +white or until solvent action has ceased. If the residue is white, +or known to be free from iron, it may be neglected and need not be +removed by filtration. If a dark residue remains, collect it on a +filter, wash free from hydrochloric acid, and ignite the filter in a +platinum crucible (Note 3). Mix the ash with five times its weight of +sodium carbonate and heat to fusion; cool, and disintegrate the fused +mass with boiling water in the crucible. Unite this solution and +precipitate (if any) with the acid solution, taking care to avoid loss +by effervescence. Wash out the crucible, heat the acid solution +to boiling, add stannous chloride solution until it is colorless, +avoiding a large excess (Note 4); cool, and when !cold!, add 40 cc. of +mercuric chloride solution, dilute to 200 cc., and proceed with the +titration as already described. + +From the standardization data already obtained, and the known weight +of the sample, calculate the percentage of iron (Fe) in the limonite. + +[Note 1: Limonite is selected as a representative of iron ores in +general. It is a native, hydrated oxide of iron. It frequently occurs +in or near peat beds and contains more or less organic matter which, +if brought into solution, would be acted upon by the potassium +bichromate. This organic matter is destroyed by roasting. Since a high +temperature tends to lessen the solubility of ferric oxide, the heat +should not be raised above low redness.] + +[Note 2: It is sometimes advantageous to dissolve a large portion--say +5 grams--and to take one tenth of it for titration. The sample will +then represent more closely the average value of the ore.] + +[Note 3: A platinum crucible may be used for the roasting of the +limonite and must be used for the fusion of the residue. When used, it +must not be allowed to remain in the acid solution of ferric chloride +for any length of time, since the platinum is attacked and dissolved, +and the platinic chloride is later reduced by the stannous chloride, +and in the reduced condition reacts with the bichromate, thus +introducing an error. It should also be noted that copper and antimony +interfere with the determination of iron by the bichromate process.] + +[Note 4: The quantity of stannous chloride required for the reduction +of the iron in the limonite will be much larger than that added to the +solution of iron wire, in which the iron was mainly already in the +ferrous condition. It should, however, be added from a dropper to +avoid an unnecessary excess.] + + + + +DETERMINATION OF CHROMIUM IN CHROME IRON ORE + + +PROCEDURE.--Grind the chrome iron ore (Note 1) in an agate mortar +until no grit is perceptible under the pestle. Weigh out two portions +of 0.5 gram each into iron crucibles which have been scoured inside +until bright (Note 2). Weigh out on a watch-glass (Note 3), using the +rough balances, 5 grams of dry sodium peroxide for each portion, and +pour about three quarters of the peroxide upon the ore. Mix ore and +flux by thorough stirring with a dry glass rod. Then cover the mixture +with the remainder of the peroxide. Place the crucible on a triangle +and raise the temperature !slowly! to the melting point of the flux, +using a low flame, and holding the lamp in the hand (Note 4). Maintain +the fusion for five minutes, and stir constantly with a stout iron +wire, but do not raise the temperature above moderate redness (Notes 5 +and 6). + +Allow the crucible to cool until it can be comfortably handled (Note +7) and then place it in a 300 cc. beaker, and cover it with distilled +water (Note 8). The beaker must be carefully covered to avoid loss +during the disintegration of the fused mass. When the evolution of +gas ceases, rinse off and remove the crucible; then heat the solution +!while still alkaline! to boiling for fifteen minutes. Allow the +liquid to cool for a few minutes; then acidify with dilute sulphuric +acid (1:5), adding 10 cc. in excess of the amount necessary to +dissolve the ferric hydroxide (Note 9). Dilute to 200 cc., cool, add +from a burette an excess of a standard ferrous solution, and titrate +for the excess with a standard solution of potassium bichromate, using +the outside indicator (Note 10). + +From the corrected volumes of the two standard solutions, and their +relations to normal solutions, calculate the percentage of chromium in +the ore. + +[Note 1: Chrome iron ore is essentially a ferrous chromite, or +combination of FeO and Cr_{2}O_{3}. It must be reduced to a state of +fine subdivision to ensure a prompt reaction with the flux.] + +[Note 2: The scouring of the iron crucible is rendered much easier if +it is first heated to bright redness and plunged into cold water. In +this process oily matter is burned off and adhering scale is caused to +chip off when the hot crucible contracts rapidly in the cold water.] + +[Note 3: Sodium peroxide must be kept off of balance pans and should +not be weighed out on paper, as is the usual practice in the rough +weighing of chemicals. If paper to which the peroxide is adhering is +exposed to moist air it is likely to take fire as a result of +the absorption of moisture, and consequent evolution of heat and +liberation of oxygen.] + +[Note 4: The lamp should never be allowed to remain under the +crucible, as this will raise the temperature to a point at which the +crucible itself is rapidly attacked by the flux and burned through.] + +[Note 5: The sodium peroxide acts as both a flux and an oxidizing +agent. The chromic oxide is dissolved by the flux and oxidized to +chromic anhydride (CrO_{3}) which combines with the alkali to form +sodium chromate. The iron is oxidized to ferric oxide.] + +[Note 6: The sodium peroxide cannot be used in porcelain, platinum, or +silver crucibles. It attacks iron and nickel as well; but crucibles +made from these metals may be used if care is exercised to keep the +temperature as low as possible. Preference is here given to iron +crucibles, because the resulting ferric hydroxide is more readily +brought into solution than the nickelic oxide from a nickel crucible. +The peroxide must be dry, and must be protected from any admixture of +dust, paper, or of organic matter of any kind, otherwise explosions +may ensue.] + +[Note 7: When an iron crucible is employed it is desirable to allow +the fusion to become nearly cold before it is placed in water, +otherwise scales of magnetic iron oxide may separate from the +crucible, which by slowly dissolving in acid form ferrous sulphate, +which reduces the chromate.] + +[Note 8: Upon treatment with water the chromate passes into solution, +the ferric hydroxide remains undissolved, and the excess of peroxide +is decomposed with the evolution of oxygen. The subsequent boiling +insures the complete decomposition of the peroxide. Unless this is +complete, hydrogen peroxide is formed when the solution is acidified, +and this reacts with the bichromate, reducing it and introducing a +serious error.] + +[Note 9: The addition of the sulphuric acid converts the sodium +chromate to bichromate, which behaves exactly like potassium +bichromate in acid solution.] + +[Note 10: If a standard solution of a ferrous salt is not at hand, a +weight of iron wire somewhat in excess of the amount which would be +required if the chromite were pure FeO.Cr_{2}O_{3} may be weighed out +and dissolved in sulphuric acid; after reduction of all the iron by +stannous chloride and the addition of mercuric chloride, this solution +may be poured into the chromate solution and the excess of iron +determined by titration with standard bichromate solution.] + + + + +PERMANGANATE PROCESS FOR THE DETERMINATION OF IRON + + +Potassium permanganate oxidizes ferrous salts in cold, acid solution +promptly and completely to the ferric condition, while in hot acid +solution it also enters into a definite reaction with oxalic acid, by +which the latter is oxidized to carbon dioxide and water. + +The reactions involved are these: + +10FeSO_{4} + 2KMnO_{4} + 8H_{2}S_{4} --> 5Fe_{2}(SO_{4})_{3} + +K_{2}SO_{4} + 2MnSO_{4} + 8H_{2}O + +5C_{2}H_{2}O_{4}(2H_{2}O) + 2KMnO_{4} +3H_{2}SO_{4} --> K_{2}SO_{4} + +2MnSO_{4} + 10CO_{2} + 1 H_{2}O. + +These are the fundamental reactions upon which the extensive use of +potassium permanganate depends; but besides iron and oxalic acid the +permanganate enters into reaction with antimony, tin, copper, mercury, +and manganese (the latter only in neutral solution), by which these +metals are changed from a lower to a higher state of oxidation; and it +also reacts with sulphurous acid, sulphureted hydrogen, nitrous acid, +ferrocyanides, and most soluble organic bodies. It should be noted, +however, that very few of these organic compounds react quantitatively +with the permanganate, as is the case with oxalic acid and the +oxalates. + +Potassium permanganate is acted upon by hydrochloric acid; the action +is rapid in hot or concentrated solution (particularly in the presence +of iron salts, which appear to act as catalyzers, increasing the +velocity of the reaction), but slow in cold, dilute solutions. +However, the greater solubility of iron compounds in hydrochloric acid +makes it desirable to use this acid as a solvent, and experiments made +with this end in view have shown that in cold, dilute hydrochloric +acid solution, to which considerable quantities of manganous sulphate +and an excess of phosphoric acid have been added, it is possible to +obtain satisfactory results. + +It is also possible to replace the hydrochloric acid by evaporating +the solutions with an excess of sulphuric acid until the latter fumes. +This procedure is somewhat more time-consuming, but the end-point of +the permanganate titration is more permanent. Both procedures are +described below. + +Potassium permanganate has an intense coloring power, and since the +solution resulting from the oxidation of the iron and the reduction of +the permanganate is colorless, the latter becomes its own indicator. +The slightest excess is indicated with great accuracy by the pink +color of the solution. + + +PREPARATION OF A STANDARD SOLUTION + +!Approximate Strength 0.1 N! + +A study of the reactions given above which represent the oxidation of +ferrous compounds by potassium permanganate, shows that there are 2 +molecules of KMnO_{4} and 10 molecules of FeSO_{4} on the +left-hand side, and 2 molecules of MnSO_{4} and 5 molecules of +Fe_{2}(SO_{4})_{5} on the right-hand side. Considering only these +compounds, and writing the formulas in such a way as to show the +oxides of the elements in each, the equation becomes: + +K_{2}O.Mn_{2}O_{7} + 10(FeO.SO_{3}) --> K_{2}O.SO_{3} + 2(MnO.SO_{3}) ++ 5(Fe_{2}O_{3}.3SO_{3}). + +From this it appears that two molecules of KMnO_{4} (or 316.0 grams) +have given up five atoms (or 80 grams) of oxygen to oxidize the +ferrous compound. Since 8 grams of oxygen is the basis of normal +oxidizing solutions and 80 grams of oxygen are supplied by 316.0 grams +of KMnO_{4}, the normal solution of the permanganate should contain, +per liter, 316.0/10 grams, or 31.60 grams (Note 1). + +The preparation of an approximately tenth-normal solution of the +reagent may be carried out as follows: + +PROCEDURE.--Dissolve about 3.25 grams of potassium permanganate +crystals in approximately 1000 cc. of distilled water in a large +beaker, or casserole. Heat slowly and when the crystals have +dissolved, boil the solution for 10-15 minutes. Cover the solution +with a watch-glass; allow it to stand until cool, or preferably over +night. Filter the solution through a layer of asbestos. Transfer the +filtrate to a liter bottle and mix thoroughly (Note 2). + +[Note 1: The reactions given on page 61 are those which take place in +the presence of an excess of acid. In neutral solutions the reduction +of the permanganate is less complete, and, under these conditions, +two gram-molecular weights of KMnO_{4} will furnish only 48 grams +of oxygen. A normal solution for use under these conditions should, +therefore, contain 316.0/6 grams, or 52.66 grams.] + +[Note 2: Potassium permanganate solutions are not usually stable for +long periods, and change more rapidly when first prepared than after +standing some days. This change is probably caused by interaction +with the organic matter contained in all distilled water, except that +redistilled from an alkaline permanganate solution. The solutions +should be protected from light and heat as far as possible, since both +induce decomposition with a deposition of manganese dioxide, and it +has been shown that decomposition proceeds with considerable rapidity, +with the evolution of oxygen, after the dioxide has begun to form. As +commercial samples of the permanganate are likely to be contaminated +by the dioxide, it is advisable to boil and filter solutions through +asbestos before standardization, as prescribed above. Such solutions +are relatively stable.] + + +COMPARISON OF PERMANGANATE AND FERROUS SOLUTIONS + +PROCEDURE.--Fill a glass-stoppered burette with the permanganate +solution, observing the usual precautions, and fill a second burette +with the ferrous sulphate solution prepared for use with the potassium +bichromate. The permanganate solution cannot be used in burettes with +rubber tips, as a reduction takes place upon contact with the rubber. +The solution has so deep a color that the lower line of the meniscus +cannot be detected; readings must therefore be made from the upper +edge. Run out into a beaker about 40 cc. of the ferrous solution, +dilute to about 100 cc., add 10 cc. of dilute sulphuric acid, and run +in the permanganate solution to a slight permanent pink. Repeat, until +the ratio of the two solutions is satisfactorily established. + + +STANDARDIZATION OF A POTASSIUM PERMANGANATE SOLUTION + +!Selection of a Standard! + +Commercial potassium permanganate is rarely sufficiently pure to admit +of its direct weighing as a standard. On this account, and because +of the uncertainties as to the permanence of its solutions, it is +advisable to standardize them against substances of known value. Those +in most common use are iron wire, ferrous ammonium sulphate, sodium +oxalate, oxalic acid, and some other derivatives of oxalic acid. +With the exception of sodium oxalate, these all contain water of +crystallization which may be lost on standing. They should, therefore, +be freshly prepared, and with great care. At present, sodium oxalate +is considered to be one of the most satisfactory standards. + + +!Method A! + + +!Iron Standards! + +The standardization processes employed when iron or its compounds are +selected as standards differ from those applicable in connection with +oxalate standards. The procedure which immediately follows is that in +use with iron standards. + +As in the case of the bichromate process, it is necessary to reduce +the iron completely to the ferrous condition before titration. The +reducing agents available are zinc, sulphurous acid, or sulphureted +hydrogen. Stannous chloride may also be used when the titration is +made in the presence of hydrochloric acid. Since the excess of both +the gaseous reducing agents can only be expelled by boiling, with +consequent uncertainty regarding both the removal of the excess and +the reoxidation of the iron, zinc or stannous chlorides are the most +satisfactory agents. For prompt and complete reduction it is essential +that the iron solution should be brought into ultimate contact with +the zinc. This is brought about by the use of a modified Jones +reductor, as shown in Figure 1. This reductor is a standard apparatus +and is used in other quantitative processes. + +[Illustration: Fig. 1] + +The tube A has an inside diameter of 18 mm. and is 300 mm. long; the +small tube has an inside diameter of 6 mm. and extends 100 mm. below +the stopcock. At the base of the tube A are placed some pieces of +broken glass or porcelain, covered by a plug of glass wool about 8 mm. +thick, and upon this is placed a thin layer of asbestos, such as is +used for Gooch filters, 1 mm. thick. The tube is then filled with the +amalgamated zinc (Note 1) to within 50 mm. of the top, and on the zinc +is placed a plug of glass wool. If the top of the tube is not already +shaped like the mouth of a thistle-tube (B), a 60 mm. funnel is fitted +into the tube with a rubber stopper and the reductor is connected +with a suction bottle, F. The bottle D is a safety bottle to +prevent contamination of the solution by water from the pump. After +preparation for use, or when left standing, the tube A should be +filled with water, to prevent clogging of the zinc. + +[Note 1: The use of fine zinc in the reductor is not necessary and +tends to clog the tube. Particles which will pass a 10-mesh sieve, but +are retained by one of 20 meshes to the inch, are most satisfactory. +The zinc can be amalgamated by stirring or shaking it in a mixture of +25 cc. of normal mercuric chloride solution, 25 cc. of hydrochloric +acid (sp. gr. 1.12) and 250 cc. of water for two minutes. The solution +should then be poured off and the zinc thoroughly washed. It is then +ready for bottling and preservation under water. A small quantity of +glass wool is placed in the neck of the funnel to hold back foreign +material when the reductor is in use.] + + +STANDARDIZATION + +PROCEDURE.--Weigh out into Erlenmeyer flasks two portions of iron wire +of about 0.25 gram each. Dissolve these in hot dilute sulphuric acid +(5 cc. of concentrated acid and 100 cc. of water), using a covered +flask to avoid loss by spattering. Boil the solution for two or +three minutes after the iron has dissolved to remove any volatile +hydrocarbons. Meanwhile prepare the reductor for use as follows: +Connect the vacuum bottle with the suction pump and pour into the +funnel at the top warm, dilute sulphuric acid, prepared by adding 5 +cc. of concentrated sulphuric acid to 100 cc. of distilled water. See +that the stopcock (C) is open far enough to allow the acid to run +through slowly. Continue to pour in acid until 200 cc. have passed +through, then close the stopcock !while a small quantity of liquid +is still left in the funnel!. Discard the filtrate, and again +pass through 100 cc. of the warm, dilute acid. Test this with the +permanganate solution. A single drop should color it permanently; if +it does not, repeat the washing, until assured that the zinc is not +contaminated with appreciable quantities of reducing substances. Be +sure that no air enters the reductor (Note 1). + +Pour the iron solution while hot (but not boiling) through the +reductor at a rate not exceeding 50 cc. per minute (Notes 2 and 3). +Wash out the beaker with dilute sulphuric acid, and follow the iron +solution without interruption with 175 cc. of the warm acid and +finally with 75 cc. of distilled water, leaving the funnel partially +filled. Remove the filter bottle and cool the solution quickly under +the water tap (Note 4), avoiding unnecessary exposure to the oxygen of +the air. Add 10 cc. of dilute sulphuric acid and titrate to a faint +pink with the permanganate solution, adding it directly to the +contents of the vacuum flask. Should the end-point be overstepped, the +ferrous sulphate solution may be added. + +From the volume of the solution required to oxidize the iron in +the wire, calculate the relation to the normal of the permanganate +solution. The duplicate results should be concordant within two parts +in one thousand. + +[Note 1: The funnel of the reductor must never be allowed to empty. +If it is left partially filled with water the reductor is ready for +subsequent use after a very little washing; but a preliminary test is +always necessary to safeguard against error. + +If more than a small drop of permanganate solution is required to +color 100 cc. of the dilute acid after the reductor is well washed, an +allowance must be made for the iron in the zinc. !Great care! must be +used to prevent the access of air to the reductor after it has been +washed out ready for use. If air enters, hydrogen peroxide forms, +which reacts with the permanganate, and the results are worthless.] + +[Note 2: The iron is reduced to the ferrous condition by contact with +the zinc. The active agent may be considered to be !nascent! hydrogen, +and it must be borne in mind that the visible bubbles are produced by +molecular hydrogen, which is without appreciable effect upon ferric +iron. + +The rate at which the iron solution passes through the zinc should not +exceed that prescribed, but the rate may be increased somewhat when +the wash-water is added. It is well to allow the iron solution to run +nearly, but not entirely, out of the funnel before the wash-water +is added. If it is necessary to interrupt the process, the complete +emptying of the funnel can always be avoided by closing the stopcock. + +It is also possible to reduce the iron by treatment with zinc in a +flask from which air is excluded. The zinc must be present in excess +of the quantity necessary to reduce the iron and is finally completely +dissolved. This method is, however, less convenient and more tedious +than the use of the reductor.] + +[Note 3: The dilute sulphuric acid for washing must be warmed ready +for use before the reduction of the iron begins, and it is of the +first importance that the volume of acid and of wash-water should +be measured, and the volume used should always be the same in the +standardizations and all subsequent analyses.] + +[Note 4: The end-point is more permanent in cold than hot solutions, +possibly because of a slight action of the permanganate upon the +manganous sulphate formed during titration. If the solution turns +brown, it is an evidence of insufficient acid, and more should be +immediately added. The results are likely to be less accurate in this +case, however, as a consequence of secondary reactions between the +ferrous iron and the manganese dioxide thrown down. It is wiser to +discard such results and repeat the process.] + +[Note 5: The potassium permanganate may, of course, be diluted and +brought to an exactly 0.1 N solution from the data here obtained. The +percentage of iron in the iron wire must be taken into account in all +calculations.] + + +!Method B! + +!Oxalate Standards! + +PROCEDURE.--Weigh out two portions of pure sodium oxalate of 0.25-0.3 +gram each into beakers of about 600 cc. capacity. Add about 400 cc. of +boiling water and 20 cc. of manganous sulphate solution (Note 1). +When the solution of the oxalate is complete, heat the liquid, if +necessary, until near its boiling point (70-90°C.) and run in the +standard permanganate solution drop by drop from a burette, stirring +constantly until an end-point is reached (Note 2). Make a blank test +with 20 cc. of manganous sulphate solution and a volume of distilled +water equal to that of the titrated solution to determine the volume +of the permanganate solution required to produce a very slight pink. +Deduct this volume from the amount of permanganate solution used in +the titration. + +From the data obtained, calculate the relation of the permanganate +solution to the normal. The reaction involved is: + +5Na_{2}C_{2}O_{4} + 2KMnO_{4} + 8H_{2}SO_{4} --> 5Na_{2}SO_{4} + +K_{2}SO_{4} + 2MnSO_{4} + 10CO_{2} + 8H_{2}O + +[Note 1: The manganous sulphate titrating solution is made by +dissolving 20 grams of MnSO_{4} in 200 cubic centimeters of water and +adding 40 cc. of concentrated sulphuric acid (sp. gr. 1.84) and 40 cc. +or phosphoric acid (85%).] + +[Note 2: The reaction between oxalates and permanganates takes place +quantitatively only in hot acid solutions. The temperatures must not +fall below 70°C.] + + + + +DETERMINATION OF IRON IN LIMONITE + + +!Method A! + +The procedures, as here prescribed, are applicable to iron ores in +general, provided these ores contain no constituents which are reduced +by zinc or stannous chloride and reoxidized by permanganates. Many +iron ores contain titanium, and this element among others does +interfere with the determination of iron by the process described. +If, however, the solutions of such ores are treated with sulphureted +hydrogen or sulphurous acid, instead of zinc or stannous chloride to +reduce the iron, and the excess reducing agent removed by boiling, an +accurate determination of the iron can be made. + +PROCEDURE.--Grind the mineral to a fine powder. Weigh out two portions +of about 0.5 gram each into small porcelain crucibles. Roast the ore +at dull redness for ten minutes (Note 1), allow the crucibles to cool, +and place them and their contents in casseroles containing 30 cc. of +dilute hydrochloric acid (sp. gr. 1.12). + +Proceed with the solution of the ore, and the treatment of the +residue, if necessary, exactly as described for the bichromate process +on page 56. When solution is complete, add 6 cc. of concentrated +sulphuric acid to each casserole, and evaporate on the steam bath +until the solution is nearly colorless (Note 2). Cover the casseroles +and heat over the flame of the burner, holding the casserole in +the hand and rotating it slowly to hasten evaporation and prevent +spattering, until the heavy white fumes of sulphuric anhydride are +freely evolved (Note 3). Cool the casseroles, add 100 cc. of water +(measured), and boil gently until the ferric sulphate is dissolved; +pour the warm solution through the reductor which has been previously +washed; proceed as described under standardization, taking pains +to use the same volume and strength of acid and the same volume of +wash-water as there prescribed, and titrate with the permanganate +solution in the reductor flask, using the ferrous sulphate solution if +the end-point should be overstepped. + +From the corrected volume of permanganate solution used, calculate the +percentage of iron (Fe) in the limonite. + +[Note 1: The preliminary roasting is usually necessary because, even +though the sulphuric acid would subsequently char the carbonaceous +matter, certain nitrogenous bodies are not thereby rendered insoluble +in the acid, and would be oxidized by the permanganate.] + +[Note 2: The temperature of the steam bath is not sufficient to +volatilize sulphuric acid. Solutions may, therefore, be left to +evaporate overnight without danger of evaporation to dryness.] + +[Note 3: The hydrochloric acid, both free and combined, is displaced +by the less volatile sulphuric acid at its boiling point. Ferric +sulphate separates at this point, since there is no water to hold +it in solution and care is required to prevent bumping. The ferric +sulphate usually has a silky appearance and is easily distinguished +from the flocculent silica which often remains undissolved.] + + +!Zimmermann-Reinhardt Procedure! + + +!Method (B)! + +PROCEDURE.--Grind the mineral to a fine powder. Weigh out two portions +of about 0.5 gram each into small porcelain crucibles. Proceed with +the solution of the ore, treat the residue, if necessary, and reduce +the iron by the addition of stannous chloride, followed by mercuric +chloride, as described for the bichromate process on page 56. Dilute +the solution to about 400 cc. with cold water, add 10 cc. of the +manganous sulphate titrating solution (Note 1, page 68) and titrate +with the standard potassium permanganate solution to a faint pink +(Note 1). + +From the standardization data already obtained calculate the +percentage of iron (Fe) in the limonite. + +[Note 1: It has already been noted that hydrochloric acid reacts +slowly in cold solutions with potassium permanganate. It is, however, +possible to obtain a satisfactory, although somewhat fugitive +end-point in the presence of manganous sulphate and phosphoric acid. +The explanation of the part played by these reagents is somewhat +obscure as yet. It is possible that an intermediate manganic compound +is formed which reacts rapidly with the ferrous compounds--thus in +effect catalyzing the oxidizing process. + +While an excess of hydrochloric acid is necessary for the successful +reduction of the iron by stannous chloride, too large an amount +should be avoided in order to lessen the chance of reduction of the +permanganate by the acid during titration.] + + + + +DETERMINATION OF THE OXIDIZING POWER OF PYROLUSITE + +INDIRECT OXIDATION + + +Pyrolusite, when pure, consists of manganese dioxide. Its value as an +oxidizing agent, and for the production of chlorine, depends upon the +percentage of MnO_{2} in the sample. This percentage is determined +by an indirect method, in which the manganese dioxide is reduced and +dissolved by an excess of ferrous sulphate or oxalic acid in the +presence of sulphuric acid, and the unused excess determined by +titration with standard permanganate solution. + +PROCEDURE.--Grind the mineral in an agate mortar until no grit +whatever can be detected under the pestle (Note 1). Transfer it to a +stoppered weighing-tube, and weigh out two portions of about 0.5 gram +into beakers (400-500 cc.) Read Note 2, and then calculate in each +case the weight of oxalic acid (H_{2}C_{2}O_{4}.2H_{2}O) required to +react with the weights of pyrolusite taken. The reaction involved is + +MnO_{2} + H_{2}C_{2}O_{4}(2H_{2}O) + H_{2}SO_{4} --> MnSO_{4} + +2CO_{2} + 4H_{2}O. + +Weigh out about 0.2 gram in excess of this quantity of !pure! oxalic +acid into the corresponding beakers, weighing the acid accurately and +recording the weight in the notebook. Pour into each beaker 25 cc. of +water and 50 cc. of dilute sulphuric acid (1:5), cover and warm the +beaker and its contents gently until the evolution of carbon dioxide +ceases (Note 3). If a residue remains which is sufficiently colored to +obscure the end-reaction of the permanganate, it must be removed by +filtration. + +Finally, dilute the solution to 200-300 cc., heat the solution to a +temperature just below boiling, add 15 cc. of a manganese sulphate +solution and while hot, titrate for the excess of the oxalic acid with +standard permanganate solution (Notes 4 and 5). + +From the corrected volume of the solution required, calculate the +amount of oxalic acid undecomposed by the pyrolusite; subtract this +from the total quantity of acid used, and calculate the weight of +manganese dioxide which would react with the balance of the acid, and +from this the percentage in the sample. + +[Note 1: The success of the analysis is largely dependent upon the +fineness of the powdered mineral. If properly ground, solution should +be complete in fifteen minutes or less.] + +[Note 2: A moderate excess of oxalic acid above that required to react +with the pyrolusite is necessary to promote solution; otherwise the +residual quantity of oxalic acid would be so small that the last +particles of the mineral would scarcely dissolve. It is also desirable +that a sufficient excess of the acid should be present to react with a +considerable volume of the permanganate solution during the titration, +thus increasing the accuracy of the process. On the other hand, the +excess of oxalic acid should not be so large as to react with more of +the permanganate solution than is contained in a 50 cc. burette. If +the pyrolusite under examination is known to be of high grade, say 80 +per cent pure, or above the calculation of the oxalic acid needed may +be based upon an assumption that the mineral is all MnO_{2}. If the +quality of the mineral is unknown, it is better to weigh out three +portions instead of two and to add to one of these the amount of +oxalic prescribed, assuming complete purity of the mineral. Then run +in the permanganate solution from a pipette or burette to determine +roughly the amount required. If the volume exceeds the contents of a +burette, the amount of oxalic acid added to the other two portions is +reduced accordingly.] + +[Note 3: Care should be taken that the sides of the beaker are not +overheated, as oxalic acid would be decomposed by heat alone if +crystallization should occur on the sides of the vessel. Strong +sulphuric acid also decomposes the oxalic acid. The dilute acid +should, therefore, be prepared before it is poured into the beaker.] + +[Note 4: Ferrous ammonium sulphate, ferrous sulphate, or iron wire +may be substituted for the oxalic acid. The reaction is then the +following: + +2 FeSO_{4} + MnO_{2} + 2H_{2}SO_{4} --> Fe_{2}(SO_{4})_{3} + 2H_{2}O + +The excess of ferrous iron may also be determined by titration with +potassium bichromate, if desired. Care is required to prevent the +oxidation of the iron by the air, if ferrous salts are employed.] + +[Note 5: The oxidizing power of pyrolusite may be determined by other +volumetric processes, one of which is outlined in the following +reactions: + +MnO_{2} + 4HCl --> MnCl_{2} + Cl_{2} + 2H_{2}O +Cl_{2} + 2KI --> I_{2} + 2KCl +I_{2} + 2Na_{2}S_{2}O_{3} --> Na_{2}S_{4}O_{6} + 2NaI. + +The chlorine generated by the pyrolusite is passed into a solution of +potassium iodide. The liberated iodine is then determined by titration +with sodium thiosulphate, as described on page 78. This is a direct +process, although it involves three steps.] + + + + +IODIMETRY + + +The titration of iodine against sodium thiosulphate, with starch as an +indicator, may perhaps be regarded as the most accurate of volumetric +processes. The thiosulphate solution may be used in both acid and +neutral solutions to measure free iodine and the latter may, in turn, +serve as a measure of any substance capable of liberating iodine from +potassium iodide under suitable conditions for titration, as, for +example, in the process outlined in Note 5 on page 74. + +The fundamental reaction upon which iodometric processes are based is +the following: + +I_{2} + 2 Na_{2}S_{2}O_{3} --> 2 NaI + Na_{2}S_{4}O_{6}. + +This reaction between iodine and sodium thiosulphate, resulting in +the formation of the compound Na_{2}S_{4}O_{6}, called sodium +tetrathionate, is quantitatively exact, and differs in that +respect from the action of chlorine or bromine, which oxidize the +thiosulphate, but not quantitatively. + +NORMAL SOLUTIONS OF IODINE AND SODIUM THIOSULPHATE + +If the formulas of sodium thiosulphate and sodium tetrathionate are +written in a manner to show the atoms of oxygen associated +with sulphur atoms in each, thus, 2(Na_{2}).S_{2}O_{2} and +Na_{2}O.S_{4}O_{5}, it is plain that in the tetrathionate there are +five atoms of oxygen associated with sulphur, instead of the four +in the two molecules of the thiosulphate taken together. Although, +therefore, the iodine contains no oxygen, the two atoms of iodine +have, in effect, brought about the addition of one oxygen atoms to the +sulphur atoms. That is the same thing as saying that 253.84 grams of +iodine (I_{2}) are equivalent to 16 grams of oxygen; hence, since 8 +grams of oxygen is the basis of normal solutions, 253.84/2 or 126.97 +grams of iodine should be contained in one liter of normal iodine +solution. By a similar course of reasoning the conclusion is reached +that the normal solution of sodium thiosulphate should contain, +per liter, its molecular weight in grams. As the thiosulphate in +crystalline form has the formula Na_{2}S_{2}O_{3}.5H_{2}O, this weight +is 248.12 grams. Tenth-normal or hundredth-normal solutions are +generally used. + + +PREPARATION OF STANDARD SOLUTIONS + +!Approximate Strength, 0.1 N! + +PROCEDURE.--Weigh out on the rough balances 13 grams of commercial +iodine. Place it in a mortar with 18 grams of potassium iodide and +triturate with small portions of water until all is dissolved. Dilute +the solution to 1000 cc. and transfer to a liter bottle and mix +thoroughly (Note 1).[1] + +[Footnote 1: It will be found more economical to have a considerable +quantity of the solution prepared by a laboratory attendant, and to +have all unused solutions returned to the common stock.] + +Weigh out 25 grams of sodium thiosulphate, dissolve it in water which +has been previously boiled and cooled, and dilute to 1000 cc., also +with boiled water. Transfer the solution to a liter bottle and mix +thoroughly (Note 2). + +[Note 1: Iodine solutions react with water to form hydriodic acid +under the influence of the sunlight, and even at low room temperatures +the iodine tends to volatilize from solution. They should, therefore, +be protected from light and heat. Iodine solutions are not stable for +long periods under the best of conditions. They cannot be used in +burettes with rubber tips, since they attack the rubber.] + +[Note 2: Sodium thiosulphate (Na_{2}S_{2}O_{3}.5H_{2}O) is +rarely wholly pure as sold commercially, but may be purified by +recrystallization. The carbon dioxide absorbed from the air by +distilled water decomposes the salt, with the separation of sulphur. +Boiled water which has been cooled out of contact with the air should +be used in preparing solutions.] + + +INDICATOR SOLUTION + +The starch solution for use as an indicator must be freshly prepared. +A soluble starch is obtainable which serves well, and a solution of +0.5 gram of this starch in 25 cc. of boiling water is sufficient. The +solution should be filtered while hot and is ready for use when cold. + +If soluble starch is not at hand, potato starch may be used. Mix about +1 gram with 5 cc. of cold water to a smooth paste, pour 150 cc. of +!boiling! water over it, warm for a moment on the hot plate, and put +it aside to settle. Decant the supernatant liquid through a filter +and use the clear filtrate; 5 cc. of this solution are needed for a +titration. + +The solution of potato starch is less stable than the soluble starch. +The solid particles of the starch, if not removed by filtration, +become so colored by the iodine that they are not readily decolorized +by the thiosulphate (Note 1). + +[Note 1: The blue color which results when free iodine and starch +are brought together is probably not due to the formation of a true +chemical compound. It is regarded as a "solid solution" of iodine in +starch. Although it is unstable, and easily destroyed by heat, it +serves as an indicator for the presence of free iodine of remarkable +sensitiveness, and makes the iodometric processes the most +satisfactory of any in the field of volumetric analysis.] + + +COMPARISON OF IODINE AND THIOSULPHATE SOLUTIONS + +PROCEDURE.--Place the solutions in burettes (the iodine in a +glass-stoppered burette), observing the usual precautions. Run out 40 +cc. of the thiosulphate solution into a beaker, dilute with 150 cc. of +water, add 1 cc. to 2 cc. of the soluble starch solution, and titrate +with the iodine to the appearance of the blue of the iodo-starch. +Repeat until the ratio of the two solutions is established, +remembering all necessary corrections for burettes and for temperature +changes. + + +STANDARDIZATION OF SOLUTIONS + +Commercial iodine is usually not sufficiently pure to permit of its +use as a standard for thiosulphate solutions or the direct preparation +of a standard solution of iodine. It is likely to contain, beside +moisture, some iodine chloride, if chlorine was used to liberate the +iodine when it was prepared. It may be purified by sublimation after +mixing it with a little potassium iodide, which reacts with the iodine +chloride, forming potassium chloride and setting free the iodine. The +sublimed iodine is then dried by placing it in a closed container over +concentrated sulphuric acid. It may then be weighed in a stoppered +weighing-tube and dissolved in a solution of potassium iodide in a +stoppered flask to prevent loss of iodine by volatilization. About 18 +grams of the iodide and twelve grams of iodine per liter are required +for an approximately tenth-normal solution. + +An iodine solution made from commercial iodine may also be +standardized against arsenious oxide (As_{4}O_{6}). This substance +also usually requires purification by sublimation before use. + +The substances usually employed for the standardization of a +thiosulphate solution are potassium bromate and metallic copper. The +former is obtainable in pure condition or may be easily purified by +re-crystallization. Copper wire of high grade is sufficiently pure +to serve as a standard. Both potassium bromate and cupric salts in +solution will liberate iodine from an iodide, which is then titrated +with the thiosulphate solution. + +The reactions involved are the following: + +(a) KBrO_{3} + 6KI + 3H_{2}SO_{4} --> KBr + 3I_{2} + 3K_{2}SO_{4} + 3H_{2}O, + +(b) 3Cu + 8HNO_{3} --> 3Cu(NO_{3})_{2} + 2NO + 4H_{2}O, + 2Cu(NO_{3})_{2} + 4KI --> 2CuI + 4KNO_{3} + I_{2}. + +Two methods for the direct standardization of the sodium thiosulphate +solution are here described, and one for the direct standardization of +the iodine solution. + + +!Method A! + +PROCEDURE.--Weigh out into 500 cc. beakers two portions of about +0.150-0.175 gram of potassium bromate. Dissolve each of these in 50 +cc. of water, and add 10 cc. of a potassium iodide solution containing +3 grams of the salt in that volume (Note 1). Add to the mixture 10 cc. +of dilute sulphuric acid (1 volume of sulphuric acid with 5 volumes of +water), allow the solution to stand for three minutes, and dilute to +150 cc. (Note 2). Run in thiosulphate solution from a burette until +the color of the liberated iodine is nearly destroyed, and then add 1 +cc. or 2 cc. of starch solution, titrate to the disappearance of the +iodo-starch blue, and finally add iodine solution until the color +is just restored. Make a blank test for the amount of thiosulphate +solution required to react with the iodine liberated by the iodate +which is generally present in the potassium iodide solution, and +deduct this from the total volume used in the titration. + +From the data obtained, calculate the relation of the thiosulphate +solution to a normal solution, and subsequently calculate the similar +value for the iodine solution. + +[Note 1:--Potassium iodide usually contains small amounts of potassium +iodate as impurity which, when the iodide is brought into an acid +solution, liberates iodine, just as does the potassium bromate used as +a standard. It is necessary to determine the amount of thiosulphate +which reacts with the iodine thus liberated by making a "blank test" +with the iodide and acid alone. As the iodate is not always uniformly +distributed throughout the iodide, it is better to make up a +sufficient volume of a solution of the iodide for the purposes of the +work in hand, and to make the blank test by using the same volume of +the iodide solution as is added in the standardizing process. The +iodide solution should contain about 3 grams of the salt in 10 cc.] + +[Note 2: The color of the iodo-starch is somewhat less satisfactory in +concentrated solutions of the alkali salts, notably the iodides. The +dilution prescribed obviates this difficulty.] + + +!Method B! + +PROCEDURE.--Weigh out two portions of 0.25-0.27 gram of clean copper +wire into 250 cc. Erlenmeyer flasks (Note 1). Add to each 5 cc. of +concentrated nitric acid (sp. gr. 1.42) and 25 cc. of water, cover, +and warm until solution is complete. Add 5 cc. of bromine water and +boil until the excess of bromine is expelled. Cool, and add strong +ammonia (sp. gr. 0.90) drop by drop until a deep blue color indicates +the presence of an excess. Boil the solution until the deep blue is +replaced by a light bluish green, or a brown stain appears on the +sides of the flask (Note 2). Add 10 cc. of strong acetic acid (sp. +gr. 1.04), cool under the water tap, and add a solution of potassium +iodide (Note 3) containing about 3 grams of the salt, and titrate +with thiosulphate solution until the color of the liberated iodine +is nearly destroyed. Then add 1-2 cc. of freshly prepared starch +solution, and add thiosulphate solution, drop by drop, until the blue +color is discharged. + +From the data obtained, including the "blank test" of the iodide, +calculate the relation of the thiosulphate solution to the normal. + +[Note 1: While copper wire of commerce is not absolutely pure, the +requirements for its use as a conductor of electricity are such that +the impurities constitute only a few hundredths of one per cent and +are negligible for analytical purposes.] + +[Note 2: Ammonia neutralizes the free nitric acid. It should be added +in slight excess only, since the excess must be removed by boiling, +which is tedious. If too much ammonia is present when acetic acid is +added, the resulting ammonium acetate is hydrolyzed, and the ammonium +hydroxide reacts with the iodine set free.] + +[Note 3: A considerable excess of potassium iodide is necessary for +the prompt liberation of iodine. While a large excess will do no harm, +the cost of this reagent is so great that waste should be avoided.] + + +!Method C! + +PROCEDURE.--Weigh out into 500 cc. beakers two portions of 0.175-0.200 +gram each of pure arsenious oxide. Dissolve each of these in 10 cc. of +sodium hydroxide solution, with stirring. Dilute the solutions to 150 +cc. and add dilute hydrochloric acid until the solutions contain a few +drops in excess, and finally add to each a concentrated solution of +5 grams of pure sodium bicarbonate (NaHCO_{3}) in water. Cover the +beakers before adding the bicarbonate, to avoid loss. Add the starch +solution and titrate with the iodine to the appearance of the blue of +the iodo-starch, taking care not to pass the end-point by more than a +few drops (Note 1). + +From the corrected volume of the iodine solution used to oxidize the +arsenious oxide, calculate its relation to the normal. From the +ratio between the solutions, calculate the similar value for the +thiosulphate solution. + +[Note 1: Arsenious oxide dissolves more readily in caustic alkali than +in a bicarbonate solution, but the presence of caustic alkali during +the titration is not admissible. It is therefore destroyed by the +addition of acid, and the solution is then made neutral with the +solution of bicarbonate, part of which reacts with the acid, the +excess remaining in solution. + +The reaction during titration is the following: + +Na_{3}AsO_{3} + I_{2} + 2NaHCO_{3} --> Na_{3}AsO_{4} + 2NaI + 2CO_{2} ++ H_{2}O + +As the reaction between sodium thiosulphate and iodine is not always +free from secondary reactions in the presence of even the weakly +alkaline bicarbonate, it is best to avoid the addition of any +considerable excess of iodine. Should the end-point be passed by a few +drops, the thiosulphate may be used to correct it.] + + + + +DETERMINATION OF COPPER IN ORES + + +Copper ores vary widely in composition from the nearly pure copper +minerals, such as malachite and copper sulphide, to very low grade +materials which contain such impurities as silica, lead, iron, silver, +sulphur, arsenic, and antimony. In nearly all varieties there will be +found a siliceous residue insoluble in acids. The method here given, +which is a modification of that described by A.H. Low (!J. Am. Chem. +Soc.! (1902), 24, 1082), provides for the extraction of the copper +from commonly occurring ores, and for the presence of their common +impurities. For practice analyses it is advisable to select an ore of +a fair degree of purity. + +PROCEDURE.-- Weigh out two portions of about 0.5 gram each of the +ore (which should be ground until no grit is detected) into 250 cc. +Erlenmeyer flasks or small beakers. Add 10 cc. of concentrated nitric +acid (sp. gr. 1.42) and heat very gently until the ore is decomposed +and the acid evaporated nearly to dryness (Note 1). Add 5 cc. of +concentrated hydrochloric acid (sp. gr. 1.2) and warm gently. Then +add about 7 cc. of concentrated sulphuric acid (sp. gr. 1.84) and +evaporate over a free flame until the sulphuric acid fumes freely +(Note 2). It has then displaced nitric and hydrochloric acid from +their compounds. + +Cool the flask or beaker, add 25 cc. of water, heat the solution +to boiling, and boil for two minutes. Filter to remove insoluble +sulphates, silica and any silver that may have been precipitated as +silver chloride, and receive the filtrate in a small beaker, washing +the precipitate and filter paper with warm water until the filtrate +and washings amount to 75 cc. Bend a strip of aluminium foil (5 cm. x +12 cm.) into triangular form and place it on edge in the beaker. Cover +the beaker and boil the solution (being careful to avoid loss of +liquid by spattering) for ten minutes, but do not evaporate to small +volume. + +Wash the cover glass and sides of the beaker. The copper should now be +in the form of a precipitate at the bottom of the beaker or adhering +loosely to the aluminium sheet. Remove the sheet, wash it carefully +with hydrogen sulphide water and place it in a small beaker. Decant +the solution through a filter, wash the precipitated copper twice by +decantation with hydrogen sulphide water, and finally transfer the +copper to the filter paper, where it is again washed thoroughly, being +careful at all times to keep the precipitated copper covered with the +wash water. Remove and discard the filtrate and place an Erlenmeyer +flask under the funnel. Pour 15 cc. of dilute nitric acid (sp. gr. +1.20) over the aluminium foil in the beaker, thus dissolving any +adhering copper. Wash the foil with hot water and remove it. Warm this +nitric acid solution and pour it slowly through the filter paper, +thereby dissolving the copper on the paper, receiving the acid +solution in the Erlenmeyer flask. Before washing the paper, pour 5 cc. +of saturated bromine water (Note 3) through it and finally wash the +paper carefully with hot water and transfer any particles of copper +which may be left on it to the Erlenmeyer flask. Boil to expel the +bromine. Add concentrated ammonia drop by drop until the appearance of +a deep blue coloration indicates an excess. Boil until the deep blue +is displaced by a light bluish green coloration, or until brown stains +form on the sides of the flask. Add 10 cc. of strong acetic acid (Note +4) and cool under the water tap. Add a solution containing about 3 +grams of potassium iodide, as in the standardization, and titrate with +thiosulphate solution until the yellow of the liberated iodine is +nearly discharged. Add 1-2 cc. of freshly prepared starch solution and +titrate to the disappearance of the blue color. + +From the data obtained, calculate the percentage of copper (Cu) in the +ore. + +[Note 1: Nitric acid, because of its oxidizing power, is used as a +solvent for the sulphide ores. As a strong acid it will also dissolve +the copper from carbonate ores. The hydrochloric acid is added to +dissolve oxides of iron and to precipitate silver and lead. The +sulphuric acid displaces the other acids, leaving a solution +containing sulphates only. It also, by its dehydrating action, renders +silica from silicates insoluble.] + +[Note 2: Unless proper precautions are taken to insure the correct +concentrations of acid the copper will not precipitate quantitatively +on the aluminium foil; hence care must be taken to follow directions +carefully at this point. Lead and silver have been almost completely +removed as sulphate and chloride respectively, or they too would +be precipitated on the aluminium. Bismuth, though precipitated on +aluminium, has no effect on the analysis. Arsenic and antimony +precipitate on aluminium and would interfere with the titration if +allowed to remain in the lower state of oxidation.] + +[Note 3: Bromine is added to oxidize arsenious and antimonious +compounds from the original sample, and to oxidize nitrous acid formed +by the action of nitric acid on copper and copper sulphide.] + +[Note 4: This reaction can be carried out in the presence of sulphuric +and hydrochloric acids as well as acetic acid, but in the presence +of these strong acids arsenic and antimonic acids may react with the +hydriodic acid produced with the liberation of free iodine, thereby +reversing the process and introducing an error.] + + + + +DETERMINATION OF ANTIMONY IN STIBNITE + + +Stibnite is native antimony sulphide. Nearly pure samples of this +mineral are easily obtainable and should be used for practice, since +many impurities, notably iron, seriously interfere with the accurate +determination of the antimony by iodometric methods. It is, moreover, +essential that the directions with respect to amounts of reagents +employed and concentration of solutions should be followed closely. + +PROCEDURE.--Grind the mineral with great care, and weigh out two +portions of 0.35-0.40 gram into small, dry beakers (100 cc.). +Cover the beakers and pour over the stibnite 5 cc. of concentrated +hydrochloric acid (sp. gr. 1.20) and warm gently on the water bath +(Note 1). When the residue is white, add to each beaker 2 grams of +powdered tartaric acid (Note 2). Warm the solution on the water bath +for ten minutes longer, dilute the solution very cautiously by adding +water in portions of 5 cc., stopping if the solution turns red. It +is possible that no coloration will appear, in which case cautiously +continue the dilution to 125 cc. If a red precipitate or coloration +does appear, warm the solution until it is colorless, and again dilute +cautiously to a total volume of 125 cc. and boil for a minute (Note +3). + +If a white precipitate of the oxychloride separates during dilution +(which should not occur if the directions are followed), it is best to +discard the determination and to start anew. + +Carefully neutralize most of the acid with ammonium hydroxide solution +(sp. gr. 0.96), but leave it distinctly acid (Note 4). Dissolve 3 +grams of sodium bicarbonate in 200 cc. of water in a 500 cc. beaker, +and pour the cold solution of the antimony chloride into this, +avoiding loss by effervescence. Make sure that the solution contains +an excess of the bicarbonate, and then add 1 cc. or 2 cc. of starch +solution and titrate with iodine solution to the appearance of the +blue, avoiding excess (Notes 5 and 6). + +From the corrected volume of the iodine solution required to oxidize +the antimony, calculate the percentage of antimony (Sb) in the +stibnite. + +[Note 1: Antimony chloride is volatile with steam from its +concentrated solutions; hence these solutions must not be boiled until +they have been diluted.] + +[Note 2: Antimony salts, such as the chloride, are readily hydrolyzed, +and compounds such as SbOCl are formed which are often relatively +insoluble; but in the presence of tartaric acid compounds with complex +ions are formed, and these are soluble. An excess of hydrochloric acid +also prevents precipitation of the oxychloride because the H^{+} ions +from the acid lessen the dissociation of the water and thus prevent +any considerable hydrolysis.] + +[Note 3: The action of hydrochloric acid upon the sulphide sets free +sulphureted hydrogen, a part of which is held in solution by the acid. +This is usually expelled by the heating upon the water bath; but if it +is not wholly driven out, a point is reached during dilution at which +the antimony sulphide, being no longer held in solution by the acid, +separates. If the dilution is immediately stopped and the solution +warmed, this sulphide is again brought into solution and at the same +time more of the sulphureted hydrogen is expelled. This procedure must +be continued until the sulphureted hydrogen is all removed, since it +reacts with iodine. If no precipitation of the sulphide occurs, it +is an indication that the sulphureted hydrogen was all expelled on +solution of the stibnite.] + +[Note 4: Ammonium hydroxide is added to neutralize most of the acid, +thus lessening the amount of sodium bicarbonate to be added. The +ammonia should not neutralize all of the acid.] + +[Note 5: The reaction which takes place during titration may be +expressed thus: + +Na_{3}SbO_{3} + 2NaHCO_{3} + I_{2} --> Na_{3}SbO_{4} + 2NaI + H_{2}O + +2CO_{2}.] + +[Note 6: If the end-point is not permanent, that is, if the blue of +the iodo-starch is discharged after standing a few moments, the cause +may be an insufficient quantity of sodium bicarbonate, leaving the +solution slightly acid, or a very slight precipitation of an antimony +compound which is slowly acted upon by the iodine when the latter is +momentarily present in excess. In either case it is better to discard +the analysis and to repeat the process, using greater care in the +amounts of reagents employed.] + + + + +CHLORIMETRY + + +The processes included under the term !chlorimetry! comprise +those employed to determine chlorine, hypochlorites, bromine, and +hypobromites. The reagent employed is sodium arsenite in the presence +of sodium bicarbonate. The reaction in the case of the hypochlorites +is + +NaClO + Na_{3}AsO_{3} --> Na_{3}AsO_{4} + NaCl. + +The sodium arsenite may be prepared from pure arsenious oxide, +as described below, and is stable for considerable periods; but +commercial oxide requires resublimation to remove arsenic sulphide, +which may be present in small quantity. To prepare the solution, +dissolve about 5 grams of the powdered oxide, accurately weighed, +in 10 cc. of a concentrated sodium hydroxide solution, dilute the +solution to 300 cc., and make it faintly acid with dilute hydrochloric +acid. Add 30 grams of sodium bicarbonate dissolved in a little water, +and dilute the solution to exactly 1000 cc. in a measuring flask. +Transfer the solution to a dry liter bottle and mix thoroughly. + +It is possible to dissolve the arsenious oxide directly in a solution +of sodium bicarbonate, with gentle warming, but solution in sodium +hydroxide takes place much more rapidly, and the excess of the +hydroxide is readily neutralized by hydrochloric acid, with subsequent +addition of the bicarbonate to maintain neutrality during the +titration. + +The indicator required for this process is made by dipping strips of +filter paper in a starch solution prepared as described on page 76, +to which 1 gram of potassium iodide has been added. These strips are +allowed to drain and spread upon a watch-glass until dry. When touched +by a drop of the solution the paper turns blue until the hypochlorite +has all been reduced and an excess of the arsenite has been added. + + + + +DETERMINATION OF THE AVAILABLE CHLORINE IN BLEACHING POWDER + + +Bleaching powder consists mainly of a calcium compound which is a +derivative of both hydrochloric and hypochlorous acids. Its formula is +CaClOCl. Its use as a bleaching or disinfecting agent, or as a source +of chlorine, depends upon the amount of hypochlorous acid which it +yields when treated with a stronger acid. It is customary to express +the value of bleaching powder in terms of "available chlorine," by +which is meant the chlorine present as hypochlorite, but not the +chlorine present as chloride. + +PROCEDURE.--Weigh out from a stoppered test tube into a porcelain +mortar about 3.5 grams of bleaching powder (Note 1). Triturate the +powder in the mortar with successive portions of water until it is +well ground and wash the contents into a 500 cc. measuring flask +(Note 2). Fill the flask to the mark with water and shake thoroughly. +Measure off 25 cc. of this semi-solution in a measuring flask, or +pipette, observing the precaution that the liquid removed shall +contain approximately its proportion of suspended matter. + +Empty the flask or pipette into a beaker and wash it out. Run in the +arsenite solution from a burette until no further reaction takes place +on the starch-iodide paper when touched by a drop of the solution of +bleaching powder. Repeat the titration, using a second 25 cc. portion. + +From the volume of solution required to react with the bleaching +powder, calculate the percentage of available chlorine in the latter, +assuming the titration reaction to be that between chlorine and +arsenious oxide: + +As_{4}O_{6} + 4Cl_{2} + 4H_{2}O --> 2As_{2}O_{5} + 8HCl + +Note that only one twentieth of the original weight of bleaching +powder enters into the reaction. + +[Note 1: The powder must be triturated until it is fine, otherwise the +lumps will inclose calcium hypochlorite, which will fail to react with +the arsenious acid. The clear supernatant liquid gives percentages +which are below, and the sediment percentages which are above, the +average. The liquid measured off should, therefore, carry with it its +proper proportion of the sediment, so far as that can be brought about +by shaking the solution just before removal of the aliquot part for +titration.] + +[Note 2: Bleaching powder is easily acted upon by the carbonic acid in +the air, which liberates the weak hypochlorous acid. This, of course, +results in a loss of available chlorine. The original material for +analysis should be kept in a closed container and protected form the +air as far as possible. It is difficult to obtain analytical samples +which are accurately representative of a large quantity of the +bleaching powder. The procedure, as outlined, will yield results which +are sufficiently exact for technical purposes.] + + + + +III. PRECIPITATION METHODS + + + + +DETERMINATION OF SILVER BY THE THIOCYANATE PROCESS + + +The addition of a solution of potassium or ammonium thiocyanate to one +of silver in nitric acid causes a deposition of silver thiocyanate as +a white, curdy precipitate. If ferric nitrate is also present, the +slightest excess of the thiocyanate over that required to combine with +the silver is indicated by the deep red which is characteristic of the +thiocyanate test for iron. + +The reactions involved are: + +AgNO_{3} + KSCN --> AgSCN + KNO_{3}, +3KSCN + Fe(NO_{3})_{3} --> Fe(SCN)_{3} + 3KNO_{3}. + +The ferric thiocyanate differs from the great majority of salts in +that it is but very little dissociated in aqueous solutions, and the +characteristic color appears to be occasioned by the formation of the +un-ionized ferric salt. + +The normal solution of potassium thiocyanate should contain an amount +of the salt per liter of solution which would yield sufficient +(CNS)^{-} to combine with one gram of hydrogen to form HCNS, i.e., +a gram-molecular weight of the salt or 97.17 grams. If the ammonium +thiocyanate is used, the amount is 76.08 grams. To prepare the +solution for this determination, which should be approximately 0.05 +N, dissolve about 5 grams of potassium thiocyanate, or 4 grams of +ammonium thiocyanate, in a small amount of water; dilute this solution +to 1000 cc. in a liter bottle and mix as usual. + +Prepare 20 cc. of a saturated solution of ferric alum and add 5 cc. of +dilute nitric acid (sp. gr. 1.20). About 5 cc. of this solution should +be used as an indicator. + + +STANDARDIZATION + +PROCEDURE.--Crush a small quantity of silver nitrate crystals in a +mortar (Note 1). Transfer them to a watch-glass and dry them for an +hour at 110°C., protecting them from dust or other organic matter +(Note 2). Weigh out two portions of about 0.5 gram each and dissolve +them in 50 cc. of water. Add 10 cc. of dilute nitric acid which has +been recently boiled to expel the lower oxides of nitrogen, if any, +and then add 5 cc. of the indicator solution. Run in the thiocyanate +solution from a burette, with constant stirring, allowing the +precipitate to settle occasionally to obtain an exact recognition +of the end-point, until a faint red tinge can be detected in the +solution. + +From the data obtained, calculate the relation of the thiocyanate +solution to the normal. + +[Note 1: The thiocyanate cannot be accurately weighed; its solutions +must, therefore, be standardized against silver nitrate (or pure +silver), either in the form of a standard solution or in small, +weighed portions.] + +[Note 2: The crystals of silver nitrate sometimes inclose water which +is expelled on drying. If the nitrate has come into contact with +organic bodies it suffers a reduction and blackens during the heating. + +It is plain that a standard solution of silver nitrate (made by +weighing out the crystals) is convenient or necessary if many +titrations of this nature are to be made. In the absence of such a +solution the liability of passing the end-point is lessened by setting +aside a small fraction of the silver solution, to be added near the +close of the titration.] + + +DETERMINATION OF SILVER IN COIN + +PROCEDURE.-- Weigh out two portions of the coin of about 0.5 gram +each. Dissolve them in 15 cc. of dilute nitric acid (sp. gr. 1.2) and +boil until all the nitrous compounds are expelled (Note 1). Cool the +solution, dilute to 50 cc., and add 5 cc. of the indicator solution, +and titrate with the thiocyanate to the appearance of the faint red +coloration (Note 2). + +From the corrected volume of the thiocyanate solution required, +calculate the percentage of silver in the coin. + +[Note 1: The reaction with silver may be carried out in nitric acid +solutions and in the presence of copper, if the latter does not exceed +70 per cent. Above that percentage it is necessary to add silver in +known quantity to the solution. The liquid must be cold at the time of +titration and entirely free from nitrous compounds, as these sometimes +cause a reddening of the indicator solution. All utensils, distilled +water, the nitric acid and the beakers must be free from chlorides, +as the presence of these will cause precipitation of silver chloride, +thereby introducing an error.] + +[Note 2: The solution containing the silver precipitate, as well as +those from the standardization, should be placed in the receptacle for +"silver residues" as a matter of economy.] + + + + +PART III + +GRAVIMETRIC ANALYSIS + + + + +GENERAL DIRECTIONS + + +Gravimetric analyses involve the following principal steps: first, the +weighing of the sample; second, the solution of the sample; third, the +separation of some substance from solution containing, or bearing a +definite relation to, the constituent to be measured, under conditions +which render this separation as complete as possible; and finally, +the segregation of that substance, commonly by filtration, and the +determination of its weight, or that of some stable product formed +from it on ignition. For example, the gravimetric determination of +aluminium is accomplished by solution of the sample, by precipitation +in the form of hydroxide, collection of the hydroxide upon a filter, +complete removal by washing of all foreign soluble matter, and the +burning of the filter and ignition of the precipitate to aluminium +oxide, in which condition it is weighed. + +Among the operations which are common to nearly all gravimetric +analyses are precipitation, washing of precipitates, ignition of +precipitates, and the use of desiccators. In order to avoid burdensome +repetitions in the descriptions of the various gravimetric procedures +which follow, certain general instructions are introduced at this +point. These instructions must, therefore, be considered to be as much +a part of all subsequent procedures as the description of apparatus, +reagents, or manipulations. + +The analytical balance, the fundamentally important instrument in +gravimetric analysis, has already been described on pages 11 to 15. + + +PRECIPITATION + +For successful quantitative precipitations those substances are +selected which are least soluble under conditions which can be easily +established, and which separate from solution in such a state that +they can be filtered readily and washed free from admixed material. +In general, the substances selected are the same as those already +familiar to the student of Qualitative Analysis. + +When possible, substances are selected which separate in crystalline +form, since such substances are less likely to clog the pores of +filter paper and can be most quickly washed. In order to increase the +size of the crystals, which further promotes filtration and washing, +it is often desirable to allow a precipitate to remain for some time +in contact with the solution from which it has separated. The solution +is often kept warm during this period of "digestion." The small +crystals gradually disappear and the larger crystals increase in size, +probably as the result of the force known as surface tension, which +tends to reduce the surface of a given mass of material to a minimum, +combined with a very slightly greater solubility of small crystals as +compared with the larger ones. + +Amorphous substances, such as ferric hydroxide, aluminium hydroxide, +or silicic acid, separate in a gelatinous form and are relatively +difficult to filter and wash. Substances of this class also exhibit +a tendency to form, with pure water, what are known as colloidal +solutions. To prevent this as far as possible, they are washed with +solutions of volatile salts, as will be described in some of the +following procedures. + +In all precipitations the reagent should be added slowly, with +constant stirring, and should be hot when circumstances permit. +The slow addition is less likely to occasion contamination of the +precipitate by the inclosure of other substances which may be in the +solution, or of the reagent itself. + + +FUNNELS AND FILTERS + +Filtration in analytical processes is most commonly effected through +paper filters. In special cases these may be advantageously replaced +by an asbestos filter in a perforated porcelain or platinum crucible, +commonly known, from its originator, as a "Gooch filter." The +operation and use of a filter of this type is described on page 103. +Porous crucibles of a material known as alundum may also be employed +to advantage in special cases. + +The glass funnels selected for use with paper filters should have an +angle as near 60° as possible, and a narrow stem about six inches in +length. The filters employed should be washed filters, i.e., those +which have been treated with hydrochloric and hydrofluoric acids, and +which on incineration leave a very small and definitely known weight +of ash, generally about .00003 gram. Such filters are readily +obtainable on the market. + +The filter should be carefully folded to fit the funnel according to +either of the two well-established methods described in the Appendix. +It should always be placed so that the upper edge of the paper +is about one fourth inch below the top of the funnel. Under no +circumstances should the filter extend above the edge of the funnel, +as it is then utterly impossible to effect complete washing. + +To test the efficiency of the filter, fill it with distilled water. +This water should soon fill the stem completely, forming a continuous +column of liquid which, by its hydrostatic pressure, produces a gentle +suction, thus materially promoting the rapidity of filtration. Unless +the filter allows free passage of water under these conditions, it is +likely to give much trouble when a precipitate is placed upon it. + +The use of a suction pump to promote filtration is rarely altogether +advantageous in quantitative analysis, if paper filters are employed. +The tendency of the filter to break, unless the point of the filter +paper is supported by a perforated porcelain cone or a small "hardened +filter" of parchment, and the tendency of the precipitates to pass +through the pores of the filter, more than compensate for the possible +gain in time. On the other hand, filtration by suction may be useful +in the case of precipitates which do not require ignition before +weighing, or in the case of precipitates which are to be discarded +without weighing. This is best accomplished with the aid of the +special apparatus called a Gooch filter referred to above. + + +FILTRATION AND WASHING OF PRECIPITATES + +Solutions should be filtered while hot, as far as possible, since +the passage of a liquid through the pores of a filter is retarded by +friction, and this, for water at 100°C., is less than one sixth of the +resistance at 0°C. + +When the filtrate is received in a beaker, the stem of the funnel +should touch the side of the receiving vessel to avoid loss by +spattering. Neglect of this precaution is a frequent source of error. + +The vessels which contain the initial filtrate should !always! be +replaced by clean ones, properly labeled, before the washing of a +precipitate begins. In many instances a finely divided precipitate +which shows no tendency to pass through the filter at first, while the +solution is relatively dense, appears at once in the washings. Under +such conditions the advantages accruing from the removal of the first +filtrate are obvious, both as regards the diminished volume requiring +refiltration, and also the smaller number of washings subsequently +required. + +Much time may often be saved by washing precipitates by decantation, +i.e., by pouring over them, while still in the original vessel, +considerable volumes of wash-water and allowing them to settle. The +supernatant, clear wash-water is then decanted through the filter, +so far as practicable without disturbing the precipitate, and a new +portion of wash-water is added. This procedure can be employed to +special advantage with gelatinous precipitates, which fill up the +pores of the filter paper. As the medium from which the precipitate +is to settle becomes less dense it subsides less readily, and it +ultimately becomes necessary to transfer it to the filter and complete +the washing there. + +A precipitate should never completely fill a filter. The wash-water +should be applied at the top of the filter, above the precipitate. +It may be shown mathematically that the washing is most !rapidly! +accomplished by filling the filter well to the top with wash-water +each time, and allowing it to drain completely after each addition; +but that when a precipitate is to be washed with the !least possible +volume! of liquid the latter should be applied in repeated !small! +quantities. + +Gelatinous precipitates should not be allowed to dry before complete +removal of foreign matter is effected. They are likely to shrink and +crack, and subsequent additions of wash-water pass through these +channels only. + +All filtrates and wash-waters without exception must be properly +tested. !This lies at the foundation of accurate work!, and the +student should clearly understand that it is only by the invariable +application of this rule that assurance of ultimate reliability can +be secured. Every original filtrate must be tested to prove complete +precipitation of the compound to be separated, and the wash-waters +must also be tested to assure complete removal of foreign material. In +testing the latter, the amount first taken should be but a few +drops if the filtrate contains material which is to be subsequently +determined. When, however, the washing of the filter and precipitate +is nearly completed the amount should be increased, and for the final +test not less than 3 cc. should be used. + +It is impossible to trust to one's judgment with regard to the washing +of precipitates; the washings from !each precipitate! of a series +simultaneously treated must be tested, since the rate of washing will +often differ materially under apparently similar conditions, !No +exception can ever be made to this rule!. + +The habit of placing a clean common filter paper under the receiving +beaker during filtration is one to be commended. On this paper a +record of the number of washings can very well be made as the portions +of wash-water are added. + +It is an excellent practice, when possible, to retain filtrates and +precipitates until the completion of an analysis, in order that, in +case of question, they may be examined to discover sources of error. + +For the complete removal of precipitates from containing vessels, it +is often necessary to rub the sides of these vessels to loosen the +adhering particles. This can best be done by slipping over the end of +a stirring rod a soft rubber device sometimes called a "policeman." + + +DESICCATORS + +Desiccators should be filled with fused, anhydrous calcium chloride, +over which is placed a clay triangle, or an iron triangle covered with +silica tubes, to support the crucible or other utensils. The cover of +the desiccator should be made air-tight by the use of a thin coating +of vaseline. + +Pumice moistened with concentrated sulphuric acid may be used in place +of the calcium chloride, and is essential in special cases; but for +most purposes the calcium chloride, if renewed occasionally and not +allowed to cake together, is practically efficient and does not slop +about when the desiccator is moved. + +Desiccators should never remain uncovered for any length of time. The +dehydrating agents rapidly lose their efficiency on exposure to the +air. + + +CRUCIBLES + +It is often necessary in quantitative analysis to employ fluxes to +bring into solution substances which are not dissolved by acids. The +fluxes in most common use are sodium carbonate and sodium or potassium +acid sulphate. In gravimetric analysis it is usually necessary to +ignite the separated substance after filtration and washing, in order +to remove moisture, or to convert it through physical or chemical +changes into some definite and stable form for weighing. Crucibles +to be used in fusion processes must be made of materials which will +withstand the action of the fluxes employed, and crucibles to be used +for ignitions must be made of material which will not undergo any +permanent change during the ignition, since the initial weight of the +crucible must be deducted from the final weight of the crucible and +product to obtain the weight of the ignited substance. The three +materials which satisfy these conditions, in general, are platinum, +porcelain, and silica. + +Platinum crucibles have the advantage that they can be employed at +high temperatures, but, on the other hand, these crucibles can never +be used when there is a possibility of the reduction to the metallic +state of metals like lead, copper, silver, or gold, which would alloy +with and ruin the crucible. When platinum crucibles are used with +compounds of arsenic or phosphorus, special precautions are necessary +to prevent damage. This statement applies to both fusions and +ignitions. + +Fusions with sodium carbonate can be made only in platinum, since +porcelain or silica crucibles are attacked by this reagent. Acid +sulphate fusions, which require comparatively low temperatures, can +sometimes be made in platinum, although platinum is slightly attacked +by the flux. Porcelain or silica crucibles may be used with acid +fluxes. + +Silica crucibles are less likely to crack on heating than porcelain +crucibles on account of their smaller coefficient of expansion. +Ignition of substances not requiring too high a temperature may be +made in porcelain or silica crucibles. + +Iron, nickel or silver crucibles are used in special cases. + +In general, platinum crucibles should be used whenever such use is +practicable, and this is the custom in private, research or commercial +laboratories. Platinum has, however, become so valuable that it is +liable to theft unless constantly under the protection of the user. As +constant protection is often difficult in instructional laboratories, +it is advisable, in order to avoid serious monetary losses, to use +porcelain or silica crucibles whenever these will give satisfactory +service. When platinum utensils are used the danger of theft should +always be kept in mind. + + +PREPARATION OF CRUCIBLES FOR USE + +All crucibles, of whatever material, must always be cleaned, ignited +and allowed to cool in a desiccator before weighing, since all bodies +exposed to the air condense on their surfaces a layer of moisture +which increases their weight. The amount and weight of this moisture +varies with the humidity of the atmosphere, and the latter may change +from hour to hour. The air in the desiccator (see above) is kept at +a constant and low humidity by the drying agent which it contains. +Bodies which remain in a desiccator for a sufficient time (usually +20-30 minutes) retain, therefore, on their surfaces a constant weight +of moisture which is the same day after day, thus insuring constant +conditions. + +Hot objects, such as ignited crucibles, should be allowed to cool in +the air until, when held near the skin, but little heat is noticeable. +If this precaution is not taken, the air within the desiccator is +strongly heated and expands before the desiccator is covered. As the +temperature falls, the air contracts, causing a reduction of air +pressure within the covered vessel. When the cover is removed (which +is often rendered difficult) the inrush of air from the outside may +sweep light particles out of a crucible, thus ruining an entire +analysis. + +Constant heating of platinum causes a slight crystallization of the +surface which, if not removed, penetrates into the crucible. Gentle +polishing of the surface destroys the crystalline structure and +prevents further damage. If sea sand is used for this purpose, great +care is necessary to keep it from the desk, since beakers are easily +scratched by it, and subsequently crack on heating. + +Platinum crucibles stained in use may often be cleaned by the fusion +in them of potassium or sodium acid sulphate, or by heating with +ammonium chloride. If the former is used, care should be taken not +to heat so strongly as to expel all of the sulphuric acid, since the +normal sulphates sometimes expand so rapidly on cooling as to split +the crucible. The fused material should be poured out, while hot, on +to a !dry! tile or iron surface. + + +IGNITION OF PRECIPITATES + +Most precipitates may, if proper precautions are taken, be ignited +without previous drying. If, however, such precipitates can be dried +without loss of time to the analyst (as, for example, over night), it +is well to submit them to this process. It should, nevertheless, be +remembered that a partially dried precipitate often requires more care +during ignition than a thoroughly moist one. + +The details of the ignition of precipitates vary so much with the +character of the precipitate, its moisture content, and temperature to +which it is to be heated, that these details will be given under the +various procedures which follow. + + + + +DETERMINATION OF CHLORINE IN SODIUM CHLORIDE + + +!Method A. With the Use of a Gooch Filter! + +PROCEDURE.--Carefully clean a weighing-tube containing the sodium +chloride, handling it as little as possible with the moist fingers, +and weigh it accurately to 0.0001 gram, recording the weight at once +in the notebook (see Appendix). Hold the tube over the top of a beaker +(200-300 cc.), and cautiously remove the stopper, noting carefully +that no particles fall from it, or from the tube, elsewhere than into +the beaker. Pour out a small portion of the chloride, replace the +stopper, and determine by approximate weighing how much has been +removed. Continue this procedure until 0.25-0.30 gram has been taken +from the tube, then weigh accurately and record the weight beneath the +first in the notebook. The difference of the two weights represents +the weight of the chloride taken for analysis. Again weigh a second +portion of 0.25-0.30 gram into a second beaker of the same size as the +first. The beakers should be plainly marked to correspond with the +entries in the notebook. Dissolve each portion of the chloride in 150 +cc. of distilled water and add about ten drops of dilute nitric acid +(sp. gr. 1.20) (Note 2). Calculate the volume of silver nitrate +solution required to effect complete precipitation in each case, +and add slowly about 5 cc. in excess of that amount, with constant +stirring. Heat the solutions cautiously to boiling, stirring +occasionally, and continue the heating and stirring until the +precipitates settle promptly, leaving a nearly clear supernatant +liquid (Note 3). This heating should not take place in direct sunlight +(Note 4). The beaker should be covered with a watch-glass, and both +boiling and stirring so regulated as to preclude any possibility of +loss of material. Add to the clear liquid one or two drops of silver +nitrate solution, to make sure that an excess of the reagent is +present. If a precipitate, or cloudiness, appears as the drops fall +into the solution, heat again, and stir until the whole precipitate +has coagulated. The solution is then ready for filtration. + +Prepare a Gooch filter as follows: Fold over the top of a Gooch funnel +(Fig. 2) a piece of rubber-band tubing, such as is known as "bill-tie" +tubing, and fit into the mouth of the funnel a perforated porcelain +crucible (Gooch crucible), making sure that when the crucible is +gently forced into the mouth of the funnel an airtight joint results. +(A small 1 or 1-1/4-inch glass funnel may be used, in which case the +rubber tubing is stretched over the top of the funnel and then drawn +up over the side of the crucible until an air-tight joint is secured.) + +[ILLUSTRATION: FIG. 2] + +Fit the funnel into the stopper of a filter bottle, and connect the +filter bottle with the suction pump. Suspend some finely divided +asbestos, which has been washed with acid, in 20 to 30 cc. of water +(Note 1); allow this to settle, pour off the very fine particles, and +then pour some of the mixture cautiously into the crucible until an +even felt of asbestos, not over 1/32 inch in thickness, is formed. A +gentle suction must be applied while preparing this felt. Wash the +felt thoroughly by passing through it distilled water until all fine +or loose particles are removed, increasing the suction at the last +until no more water can be drawn out of it; place on top of the felt +the small, perforated porcelain disc and hold it in place by pouring a +very thin layer of asbestos over it, washing the whole carefully; +then place the crucible in a small beaker, and place both in a drying +closet at 100-110°C. for thirty to forty minutes. Cool the crucible +in a desiccator, and weigh. Heat again for twenty to thirty minutes, +cool, and again weigh, repeating this until the weight is constant +within 0.0003 gram. The filter is then ready for use. + +Place the crucible in the funnel, and apply a gentle suction, !after +which! the solution to be filtered may be poured in without disturbing +the asbestos felt. When pouring liquid onto a Gooch filter hold the +stirring-rod at first well down in the crucible, so that the liquid +does not fall with any force upon the asbestos, and afterward keep the +crucible will filled with the solution. + +Pour the liquid above the silver chloride slowly onto the filter, +leaving the precipitate in the beaker as far as possible. Wash the +precipitate twice by decantation with warm water; then transfer it +to the filter with the aid of a stirring-rod with a rubber tip and a +stream from the wash-bottle. + +Examine the first portions of the filtrate which pass through the +filter with great care for asbestos fibers, which are most likely to +be lost at this point. Refilter the liquid if any fibers are visible. +Finally, wash the precipitate thoroughly with warm water until free +from soluble silver salts. To test the washings, disconnect the +suction at the flask and remove the funnel or filter tube from the +suction flask. Hold the end of the tube over the mouth of a small test +tube and add from a wash-bottle 2-3 cc. of water. Allow the water to +drip through into the test tube and add a drop of dilute hydrochloric +acid. No precipitate or cloud should form in the wash-water (Note 16). +Dry the filter and contents at 100-110°C. until the weight is constant +within 0.0003 gram, as described for the preparation of the filter. +Deduct the weight of the dry crucible from the final weight, and from +the weight of silver chloride thus obtained calculate the percentage +of chlorine in the sample of sodium chloride. + +[Note 1: The washed asbestos for this type of filter is prepared by +digesting in concentrated hydrochloric acid, long-fibered asbestos +which has been cut in pieces of about 0.5 cm. in length. After +digestion, the asbestos is filtered off on a filter plate and washed +with hot, distilled water until free from chlorides. A small portion +of the asbestos is shaken with water, forming a thin suspension, which +is bottled and kept for use.] + +[Note 2: The nitric acid is added before precipitation to lessen the +tendency of the silver chloride to carry down with it other substances +which might be precipitated from a neutral solution. A large excess of +the acid would exert a slight solvent action upon the chloride.] + +[Note 3: The solution should not be boiled after the addition of the +nitric acid before the presence of an excess of silver nitrate is +assured, since a slight interaction between the nitric acid and the +sodium chloride is possible, by which a loss of chlorine, either as +such or as hydrochloric acid, might ensue. The presence of an excess +of the precipitant can usually be recognized at the time of its +addition, by the increased readiness with which the precipitate +coagulates and settles.] + +[Note 4: The precipitate should not be exposed to strong sunlight, +since under those conditions a reduction of the silver chloride ensues +which is accompanied by a loss of chlorine. The superficial alteration +which the chloride undergoes in diffused daylight is not sufficient +to materially affect the accuracy of the determination. It should be +noted, however, that a slight error does result from the effect of +light upon the silver chloride precipitate and in cases in which the +greatest obtainable accuracy is required, the procedure described +under "Method B" should be followed, in which this slight reduction of +the silver chloride is corrected by subsequent treatment with nitric +and hydrochloric acids.] + +[Note 5: The asbestos used in the Gooch filter should be of the finest +quality and capable of division into minute fibrous particles. A +coarse felt is not satisfactory.] + +[Note 6: The precipitate must be washed with warm water until it is +absolutely free from silver and sodium nitrates. It may be assumed +that the sodium salt is completely removed when the wash-water shows +no evidence of silver. It must be borne in mind that silver chloride +is somewhat soluble in hydrochloric acid, and only a single drop +should be added. The washing should be continued until no cloudiness +whatever can be detected in 3 cc. of the washings. + +Silver chloride is but slightly soluble in water. The solubility +varies with its physical condition within small limits, and is +about 0.0018 gram per liter at 18°C. for the curdy variety usually +precipitated. The chloride is also somewhat soluble in solutions of +many chlorides, in solutions of silver nitrate, and in concentrated +nitric acid. + +As a matter of economy, the filtrate, which contains whatever silver +nitrate was added in excess, may be set aside. The silver can be +precipitated as chloride and later converted into silver nitrate.] + +[Note 7: The use of the Gooch filter commends itself strongly when a +considerable number of halogen determinations are to be made, since +successive portions of the silver halides may be filtered on the same +filter, without the removal of the preceding portions, until the +crucible is about two thirds filled. If the felt is properly prepared, +filtration and washing are rapidly accomplished on this filter, and +this, combined with the possibility of collecting several precipitates +on the same filter, is a strong argument in favor of its use with any +but gelatinous precipitates.] + + +!Method B. With the Use of a Paper Filter! + +PROCEDURE.--Weigh out two portions of sodium chloride of about +0.25-0.3 gram each and proceed with the precipitation of the silver +chloride as described under Method A above. When the chloride is ready +for filtration prepare two 9 cm. washed paper filters (see Appendix). +Pour the liquid above the precipitates through the filters, wash twice +by decantation and transfer the precipitates to the filters, finally +washing them until free from silver solution as described. The funnel +should then be covered with a moistened filter paper by stretching it +over the top and edges, to which it will adhere on drying. It should +be properly labeled with the student's name and desk number, and then +placed in a drying closet, at a temperature of about 100-110°C., until +completely dry. + +The perfectly dry filter is then opened over a circular piece of +clean, smooth, glazed paper about six inches in diameter, placed upon +a larger piece about twelve inches in diameter. The precipitate is +removed from the filter as completely as possible by rubbing the sides +gently together, or by scraping them cautiously with a feather which +has been cut close to the quill and is slightly stiff (Note 1). In +either case, care must be taken not to rub off any considerable +quantity of the paper, nor to lose silver chloride in the form of +dust. Cover the precipitate on the glazed paper with a watch-glass to +prevent loss of fine particles and to protect it from dust from the +air. Fold the filter paper carefully, roll it into a small cone, and +wind loosely around !the top! a piece of small platinum wire (Note 2). +Hold the filter by the wire over a small porcelain crucible (which has +been cleaned, ignited, cooled in a desiccator, and weighed), ignite +it, and allow the ash to fall into the crucible. Place the crucible +upon a clean clay triangle, on its side, and ignite, with a low +flame well at its base, until all the carbon of the filter has been +consumed. Allow the crucible to cool, add two drops of concentrated +nitric acid and one drop of concentrated hydrochloric acid, and heat +!very cautiously!, to avoid spattering, until the acids have been +expelled; then transfer the main portion of the precipitate from the +glazed paper to the cooled crucible, placing the latter on the larger +piece of glazed paper and brushing the precipitate from the +smaller piece into it, sweeping off all particles belonging to the +determination. + +Moisten the precipitate with two drops of concentrated nitric acid and +one drop of concentrated hydrochloric acid, and again heat with great +caution until the acids are expelled and the precipitate is white, +when the temperature is slowly raised until the silver chloride just +begins to fuse at the edges (Note 3). The crucible is then cooled in a +desiccator and weighed, after which the heating (without the addition +of acids) is repeated, and it is again weighed. This must be continued +until the weight is constant within 0.0003 gram in two consecutive +weighings. Deduct the weight of the crucible, and calculate the +percentage of chlorine in the sample of sodium chloride taken for +analysis. + +[Note 1: The separation of the silver chloride from the filter is +essential, since the burning carbon of the paper would reduce a +considerable quantity of the precipitate to metallic silver, and its +complete reconversion to the chloride within the crucible, by means of +acids, would be accompanied by some difficulty. The small amount of +silver reduced from the chloride adhering to the filter paper after +separating the bulk of the precipitate, and igniting the paper +as prescribed, can be dissolved in nitric acid, and completely +reconverted to chloride by hydrochloric acid. The subsequent addition +of the two acids to the main portion of the precipitate restores the +chlorine to any chloride which may have been partially reduced by the +sunlight. The excess of the acids is volatilized by heating.] + +[Note 2: The platinum wire is wrapped around the top of the filter +during its incineration to avoid contact with any reduced silver from +the reduction of the precipitate. If the wire were placed nearer the +apex, such contact could hardly be avoided.] + +[Note 3: Silver chloride should not be heated to complete fusion, +since a slight loss by volatilization is possible at high +temperatures. The temperature of fusion is not always sufficient +to destroy filter shreds; hence these should not be allowed to +contaminate the precipitate.] + + + + +DETERMINATION OF IRON AND OF SULPHUR IN FERROUS AMMONIUM SULPHATE, + +FESO_{4}.(NH_{4})_{2}SO_{4}.6H_{2}O + + +DETERMINATION OF IRON + +PROCEDURE.--Weigh out into beakers (200-250 cc.) two portions of the +sample (Note 1) of about 1 gram each and dissolve these in 50 cc. of +water, to which 1 cc. of dilute hydrochloric acid (sp. gr. 1.12) has +been added (Note 2). Heat the solution to boiling, and while at the +boiling point add concentrated nitric acid (sp. gr. 1.42), !drop by +drop! (noting the volume used), until the brown coloration, which +appears after the addition of a part of the nitric acid, gives place +to a yellow or red (Note 3). Avoid a large excess of nitric acid, but +be sure that the action is complete. Pour this solution cautiously +into about 200 cc. of water, containing a slight excess of ammonia. +Calculate for this purpose the amount of aqueous ammonia required to +neutralize the hydrochloric and nitric acids added (see Appendix for +data), and also to precipitate the iron as ferric hydroxide from the +weight of the ferrous ammonium sulphate taken for analysis, assuming +it to be pure (Note 4). The volume thus calculated will be in excess +of that actually required for precipitation, since the acids are in +part consumed in the oxidation process, or are volatilized. Heat the +solution to boiling, and allow the precipitated ferric hydroxide to +settle. Decant the clear liquid through a washed filter (9 cm.), +keeping as much of the precipitate in the beaker as possible. Wash +twice by decantation with 100 cc. of hot water. Reserve the filtrate. +Dissolve the iron from the filter with hot, dilute hydrochloric acid +(sp. gr. 1.12), adding it in small portions, using as little as +possible and noting the volume used. Collect the solution in the +beaker in which precipitation took place. Add 1 cc. of nitric acid +(sp. gr. 1.42), boil for a few moments, and again pour into a +calculated excess of ammonia. + +Wash the precipitate twice by decantation, and finally transfer it to +the original filter. Wash continuously with hot water until finally +3 cc. of the washings, acidified with nitric acid (Note 5), show +no evidences of the presence of chlorides when tested with silver +nitrate. The filtrate and washings are combined with those from the +first precipitation and treated for the determination of sulphur, as +prescribed on page 112. + +[Note 1: If a selection of pure material for analysis is to be made, +crystals which are cloudy are to be avoided on account of loss of +water of crystallization; and also those which are red, indicating +the presence of ferric iron. If, on the other hand, the value of an +average sample of material is desired, it is preferable to grind the +whole together, mix thoroughly, and take a sample from the mixture for +analysis.] + +[Note 2: When aqueous solutions of ferrous compounds are heated in the +air, oxidation of the Fe^{++} ions to Fe^{+++} ions readily occurs in +the absence of free acid. The H^{+} and OH^{-} ions from water are +involved in the oxidation process and the result is, in effect, the +formation of some ferric hydroxide which tends to separate. Moreover, +at the boiling temperature, the ferric sulphate produced by the +oxidation hydrolyzes in part with the formation of a basic ferric +sulphate, which also tends to separate from solution. The addition of +the hydrochloric acid prevents the formation of ferric hydroxide, and +so far reduces the ionization of the water that the hydrolysis of the +ferric sulphate is also prevented, and no precipitation occurs on +heating.] + +[Note 3: The nitric acid, after attaining a moderate strength, +oxidizes the Fe^{++} ions to Fe^{+++} ions with the formation of an +intermediate nitroso-compound similar in character to that formed in +the "ring-test" for nitrates. The nitric oxide is driven out by heat, +and the solution then shows by its color the presence of ferric +compounds. A drop of the oxidized solution should be tested on +a watch-glass with potassium ferricyanide, to insure a complete +oxidation. This oxidation of the iron is necessary, since Fe^{++} ions +are not completely precipitated by ammonia. + +The ionic changes which are involved in this oxidation are perhaps +most simply expressed by the equation + +3Fe^{++} + NO_{3}^{-}+ 4H^{+} --> 3Fe^{+++} + 2H_{2}O + NO, + +the H^{+} ions coming from the acid in the solution, in this case +either the nitric or the hydrochloric acid. The full equation on which +this is based may be written thus: + +6FeSO_{4} + 2HNO_{3} + 6HCl --> 2Fe_{2}(SO_{4})_{3} + 2FeCl_{3} + 2NO ++ 4H_{2}O, + +assuming that only enough nitric acid is added to complete the +oxidation.] + +[Note 4: The ferric hydroxide precipitate tends to carry down some +sulphuric acid in the form of basic ferric sulphate. This tendency is +lessened if the solution of the iron is added to an excess of OH^{-} +ions from the ammonium hydroxide, since under these conditions +immediate and complete precipitation of the ferric hydroxide ensues. +A gradual neutralization with ammonia would result in the local +formation of a neutral solution within the liquid, and subsequent +deposition of a basic sulphate as a consequence of a local deficiency +of OH^{-} ions from the NH_{4}OH and a partial hydrolysis of the +ferric salt. Even with this precaution the entire absence of sulphates +from the first iron precipitate is not assured. It is, therefore, +redissolved and again thrown down by ammonia. The organic matter of +the filter paper may occasion a partial reduction of the iron during +solution, with consequent possibility of incomplete subsequent +precipitation with ammonia. The nitric acid is added to reoxidize this +iron. + +To avoid errors arising from the solvent action of ammoniacal +liquids upon glass, the iron precipitate should be filtered without +unnecessary delay.] + +[Note 5: The washings from the ferric hydroxide are acidified with +nitric acid, before testing with silver nitrate, to destroy the +ammonia which is a solvent of silver chloride. + +The use of suction to promote filtration and washing is permissible, +though not prescribed. The precipitate should not be allowed to dry +during the washing.] + + +!Ignition of the Iron Precipitate! + +Heat a platinum or porcelain crucible, cool it in a desiccator and +weigh, repeating until a constant weight is obtained. + +Fold the top of the filter paper over the moist precipitate of ferric +hydroxide and transfer it cautiously to the crucible. Wipe the inside +of the funnel with a small fragment of washed filter paper, if +necessary, and place the paper in the crucible. + +Incline the crucible on its side, on a triangle supported on a +ring-stand, and stand the cover on edge at the mouth of the crucible. +Place a burner below the front edge of the crucible, using a low flame +and protecting it from drafts of air by means of a chimney. The heat +from the burner is thus reflected into the crucible and dries +the precipitate without danger of loss as the result of a sudden +generation of steam within the mass of ferric hydroxide. As the drying +progresses the burner may be gradually moved toward the base of the +crucible and the flame increased until the paper of the filter begins +to char and finally to smoke, as the volatile matter is expelled. This +is known as "smoking off" a filter, and the temperature should not be +raised sufficiently high during this process to cause the paper to +ignite, as the air currents produced by the flame of the blazing paper +may carry away particles of the precipitate. + +When the paper is fully charred, move the burner to the base of the +crucible and raise the temperature to the full heat of the burner for +fifteen minutes, with the crucible still inclined on its side, but +without the cover (Note 1). Finally set the crucible upright in the +triangle, cover it, and heat at the full temperature of a blast lamp +or other high temperature burner. Cool and weigh in the usual manner +(Note 2). Repeat the strong heating until the weight is constant +within 0.0003 gram. + +From the weight of ferric oxide (Fe_{2}O_{3}) calculate the percentage +of iron (Fe) in the sample (Note 3). + +[Note 1: These directions for the ignition of the precipitate must be +closely followed. A ready access of atmospheric oxygen is of special +importance to insure the reoxidation to ferric oxide of any iron which +may be reduced to magnetic oxide (Fe_{3}O_{4}) during the combustion +of the filter. The final heating over the blast lamp is essential +for the complete expulsion of the last traces of water from the +hydroxide.] + +[Note 2: Ignited ferric oxide is somewhat hygroscopic. On this account +the weighings must be promptly completed after removal from the +desiccator. In all weighings after the first it is well to place the +weights upon the balance-pan before removing the crucible from the +desiccator. It is then only necessary to move the rider to obtain the +weight.] + +[Note 3: The gravimetric determination of aluminium or chromium is +comparable with that of iron just described, with the additional +precaution that the solution must be boiled until it contains but a +very slight excess of ammonia, since the hydroxides of aluminium and +chromium are more soluble than ferric hydroxide. + +The most important properties of these hydroxides, from a quantitative +standpoint, other than those mentioned, are the following: All are +precipitable by the hydroxides of sodium and potassium, but always +inclose some of the precipitant, and should be reprecipitated with +ammonium hydroxide before ignition to oxides. Chromium and aluminium +hydroxides dissolve in an excess of the caustic alkalies and form +anions, probably of the formula AlO_2^{-} and CrO_{2}^{-}. Chromium +hydroxide is reprecipitated from this solution on boiling. When first +precipitated the hydroxides are all readily soluble in acids, but +aluminium hydroxide dissolves with considerable difficulty after +standing or boiling for some time. The precipitation of the hydroxides +is promoted by the presence of ammonium chloride, but is partially +or entirely prevented by the presence of tartaric or citric acids, +glycerine, sugars, and some other forms of soluble organic matter. +The hydroxides yield on ignition an oxide suitable for weighing +(Al_{2}O_{3}, Cr_{2}O_{3}, Fe_{2}O_{3}).] + + + + +DETERMINATION OF SULPHUR + + +PROCEDURE.--Add to the combined filtrates from the ferric hydroxide +about 0.6 gram of anhydrous sodium carbonate; cover the beaker, and +then add dilute hydrochloric acid (sp. gr. 1.12) in moderate excess +and evaporate to dryness on the water bath. Add 10 cc. of concentrated +hydrochloric acid (sp. gr. 1.20) to the residue, and again evaporate +to dryness on the bath. Dissolve the residue in water, filter if not +clear, transfer to a 700 cc. beaker, dilute to about 400 cc., and +cautiously add hydrochloric acid until the solution shows a distinctly +acid reaction (Note 1). Heat the solution to boiling, and add !very +slowly! and with constant stirring, 20 cc. in excess of the calculated +amount of a hot barium chloride solution, containing about 20 grams +BaCl_{2}.2H_{2}O per liter (Notes 2 and 3). Continue the boiling for +about two minutes, allow the precipitate to settle, and decant the +liquid at the end of half an hour (Note 4). Replace the beaker +containing the original filtrate by a clean beaker, wash the +precipitated sulphate by decantation with hot water, and subsequently +upon the filter until it is freed from chlorides, testing the washings +as described in the determination of iron. The filter is then +transferred to a platinum or porcelain crucible and ignited, as +described above, until the weight is constant (Note 5). From the +weight of barium sulphate (BaSO_{4}) obtained, calculate the +percentage of sulphur (S) in the sample. + +[Note 1: Barium sulphate is slightly soluble in hydrochloric acid, +even dilute, probably as a result of the reduction in the degree of +dissociation of sulphuric acid in the presence of the H^{+} ions of +the hydrochloric acid, and possibly because of the formation of a +complex anion made up of barium and chlorine; hence only the smallest +excess should be added over the amount required to acidify the +solution.] + +[Note 2: The ionic changes involved in the precipitation of barium +sulphate are very simple: + +Ba^{++} + SO_{4}^{--} --> [BaSO_{4}] + +This case affords one of the best illustrations of the effect of an +excess of a precipitant in decreasing the solubility of a precipitate. +If the conditions are considered which exist at the moment when just +enough of the Ba^{++} ions have been added to correspond to the +SO_{4}^{--} ions in the solution, it will be seen that nearly all of +the barium sulphate has been precipitated, and that the small amount +which then remains in the solution which is in contact with the +precipitate must represent a saturated solution for the existing +temperature, and that this solution is comparable with a solution of +sugar to which more sugar has been added than will dissolve. It +should be borne in mind that the quantity of barium sulphate in +this !saturated solution is a constant quantity! for the existing +conditions. The dissolved barium sulphate, like any electrolyte, is +dissociated, and the equilibrium conditions may be expressed thus: + +(!Conc'n Ba^{++} x Conc'n SO_{4}^{--})/(Conc'n BaSO_{4}) = Const.!, + +and since !Conc'n BaSO_{4}! for the saturated solution has a constant +value (which is very small), it may be eliminated, when the expression +becomes !Conc'n Ba^{++} x Conc'n SO_{4}^{--} = Const.!, which is +the "solubility product" of BaSO_{4}. If, now, an excess of the +precipitant, a soluble barium salt, is added in the form of a +relatively concentrated solution (the slight change of volume of a few +cubic centimeters may be disregarded for the present discussion) +the concentration of the Ba^{++} ions is much increased, and as a +consequence the !Conc'n SO_{4}! must decrease in proportion if the +value of the expression is to remain constant, which is a requisite +condition if the law of mass action upon which our argument depends +holds true. In other words, SO_{4}^{--} ions must combine with some +of the added Ba^{++} ions to form [BaSO_{4}]; but it will be recalled +that the solution is already saturated with BaSO_{4}, and this freshly +formed quantity must, therefore, separate and add itself to the +precipitate. This is exactly what is desired in order to insure +more complete precipitation and greater accuracy, and leads to the +conclusion that the larger the excess of the precipitant added the +more successful the analysis; but a practical limit is placed upon +the quantity of the precipitant which may be properly added by other +conditions, as stated in the following note.] + +[Note 3: Barium sulphate, in a larger measure than most compounds, +tends to carry down other substances which are present in the solution +from which it separates, even when these other substances are +relatively soluble, and including the barium chloride used as the +precipitant. This is also notably true in the case of nitrates and +chlorates of the alkalies, and of ferric compounds; and, since in this +analysis ammonium nitrate has resulted from the neutralization of the +excess of the nitric acid added to oxidize the iron, it is essential +that this should be destroyed by repeated evaporation with a +relatively large quantity of hydrochloric acid. During evaporation a +mutual decomposition of the two acids takes place, and the nitric acid +is finally decomposed and expelled by the excess of hydrochloric acid. + +Iron is usually found in the precipitate of barium sulphate when +thrown down from hot solutions in the presence of ferric salts. This, +according to Kuster and Thiel (!Zeit. anorg. Chem.!, 22, 424), is due +to the formation of a complex ion (Fe(SO_{4})_{2}) which precipitates +with the Ba^{++} ion, while Richards (!Zeit. anorg. Chem.!, 23, 383) +ascribes it to hydrolytic action, which causes the formation of a +basic ferric complex which is occluded in the barium precipitate. +Whatever the character of the compound may be, it has been shown that +it loses sulphuric anhydride upon ignition, causing low results, even +though the precipitate contains iron. + +The contamination of the barium sulphate by iron is much less in the +presence of ferrous than ferric salts. If, therefore, the sulphur +alone were to be determined in the ferrous ammonium sulphate, the +precipitation by barium might be made directly from an aqueous +solution of the salt, which had been made slightly acid with +hydrochloric acid.] + +[Note 4: The precipitation of the barium sulphate is probably complete +at the end of a half-hour, and the solution may safely be filtered at +the expiration of that time if it is desired to hasten the analysis. + +As already noted, many precipitates of the general character of this +sulphate tend to grow more coarsely granular if digested for some time +with the liquid from which they have separated. It is therefore well +to allow the precipitate to stand in a warm place for several hours, +if practicable, to promote ease of filtration. The filtrate and +washings should always be carefully examined for minute quantities of +the sulphate which may pass through the pores of the filter. This is +best accomplished by imparting to the filtrate a gentle rotary motion, +when the sulphate, if present, will collect at the center of the +bottom of the beaker.] + +[Note 5: A reduction of barium sulphate to the sulphide may very +readily be caused by the reducing action of the burning carbon of the +filter, and much care should be taken to prevent any considerable +reduction from this cause. Subsequent ignition, with ready access +of air, reconverts the sulphide to sulphate unless a considerable +reduction has occurred. In the latter case it is expedient to add one +or two drops of sulphuric acid and to heat cautiously until the excess +of acid is expelled.] + +[Note 6: Barium sulphate requires about 400,000 parts of water for +its solution. It is not decomposed at a red heat but suffers loss, +probably of sulphur trioxide, at a temperature above 900°C.] + + + + +DETERMINATION OF SULPHUR IN BARIUM SULPHATE + + +PROCEDURE.--Weigh out, into platinum crucibles, two portions of about +0.5 gram of the sulphate. Mix each in the crucible with five to six +times its weight of anhydrous sodium carbonate. This can best be done +by placing the crucible on a piece of glazed paper and stirring the +mixture with a clean, dry stirring-rod, which may finally be wiped off +with a small fragment of filter paper, the latter being placed in the +crucible. Cover the crucible and heat until a quiet, liquid fusion +ensues. Remove the burner, and tip the crucible until the fused mass +flows nearly to its mouth. Hold it in that position until the mass has +solidified. When cold, the material may usually be detached in a lump +by tapping the crucible or gently pressing it near its upper edge. If +it still adheres, a cubic centimeter or so of water may be placed in +the cold crucible and cautiously brought to boiling, when the cake +will become loosened and may be removed and placed in about 250 cc. +of hot, distilled water to dissolve. Clean the crucible completely, +rubbing the sides with a rubber-covered stirring-rod, if need be. + +When the fused mass has completely disintegrated and nothing further +will dissolve, decant the solution from the residue of barium +carbonate (Note 1). Pour over the residue 20 cc. of a solution of +sodium carbonate and 10 cc. of water and heat to gentle boiling for +about three minutes (Note 2). Filter off the carbonate and wash it +with hot water, testing the slightly acidified washings for sulphate +and preserving any precipitates which appear in these tests. Acidify +the filtrate with hydrochloric acid until just acid, bring to boiling, +and slowly add hot barium chloride solution, as in the preceding +determination. Add also any tests from the washings in which +precipitates have appeared. Filter, wash, ignite, and weigh. + +From the weight of barium sulphate, calculate the percentage of +sulphur (S) in the sample. + +[Note 1: This alkaline fusion is much employed to disintegrate +substances ordinarily insoluble in acids into two components, one +of which is water soluble and the other acid soluble. The reaction +involved is: + +BaSO_{4} + Na_{2}CO_{3}, --> BaCO_{3}, + Na_{2}SO_{4}. + +As the sodium sulphate is soluble in water, and the barium carbonate +insoluble, a separation between them is possible and the sulphur can +be determined in the water-soluble portion. + +It should be noted that this method can be applied to the purification +of a precipitate of barium sulphate if contaminated by most of the +substances mentioned in Note 3 on page 114. The impurities pass into +the water solution together with the sodium sulphate, but, being +present in such minute amounts, do not again precipitate with the +barium sulphate.] + +[Note 2: The barium carbonate is boiled with sodium carbonate solution +before filtration because the reaction above is reversible; and it is +only by keeping the sodium carbonate present in excess until nearly +all of the sodium sulphate solution has been removed by filtration +that the reversion of some of the barium carbonate to barium sulphate +is prevented. This is an application of the principle of mass action, +in which the concentration of the reagent (the carbonate ion) is +kept as high as practicable and that of the sulphate ion as low as +possible, in order to force the reaction in the desired direction (see +Appendix).] + + + + +DETERMINATION OF PHOSPHORIC ANHYDRIDE IN APATITE + + +The mineral apatite is composed of calcium phosphate, associated with +calcium chloride, or fluoride. Specimens are easily obtainable which +are nearly pure and leave on treatment with acid only a slight +siliceous residue. + +For the purpose of gravimetric determination, phosphoric acid is +usually precipitated from ammoniacal solutions in the form of +magnesium ammonium phosphate which, on ignition, is converted into +magnesium pyrophosphate. Since the calcium phosphate of the apatite +is also insoluble in ammoniacal solutions, this procedure cannot be +applied directly. The separation of the phosphoric acid from the +calcium must first be accomplished by precipitation in the form of +ammonium phosphomolybdate in nitric acid solution, using ammonium +molybdate as the precipitant. The "yellow precipitate," as it is often +called, is not always of a definite composition, and therefore not +suitable for direct weighing, but may be dissolved in ammonia, and the +phosphoric acid thrown out as magnesium ammonium phosphate from the +solution. + +Of the substances likely to occur in apatite, silicic acid alone +interferes with the precipitation of the phosphoric acid in nitric +acid solution. + + +PRECIPITATION OF AMMONIUM PHOSPHOMOLYBDATE + +PROCEDURE.--Grind the mineral in an agate mortar until no grit is +perceptible. Transfer the substance to a weighing-tube, and weigh out +two portions, not exceeding 0.20 gram each (Note 1) into two beakers +of about 200 cc. capacity. Pour over them 20 cc. of dilute nitric acid +(sp. gr. 1.2) and warm gently until solvent action has apparently +ceased. Evaporate the solution cautiously to dryness, heat the residue +for about an hour at 100-110°C., and treat it again with nitric acid +as described above; separate the residue of silica by filtration on +a small filter (7 cm.) and wash with warm water, using as little as +possible (Note 2). Receive the filtrate in a beaker (200-500 cc.). +Test the washings with ammonia for calcium phosphate, but add all such +tests in which a precipitate appears to the original nitrate (Note 3). +The filtrate and washings must be kept as small as possible and should +not exceed 100 cc. in volume. Add aqueous ammonia (sp. gr. 0.96) until +the precipitate of calcium phosphate first produced just fails to +redissolve, and then add a few drops of nitric acid until this is +again brought into solution (Note 4). Warm the solution until it +cannot be comfortably held in the hand (about 60°C.) and, after +removal of the burner, add 75 cc. of ammonium molybdate solution which +has been !gently! warmed, but which must be perfectly clear. Allow +the mixture to stand at a temperature of about 50 or 60°C. for twelve +hours (Notes 5 and 6). Filter off the yellow precipitate on a 9 cm. +filter, and wash by decantation with a solution of ammonium nitrate +made acid with nitric acid.[1] Allow the precipitate to remain in the +beaker as far as possible. Test the washings for calcium with ammonia +and ammonium oxalate (Note 3). + +[Footnote 1: This solution is prepared as follows: Mix 100 cc. of +ammonia solution (sp. gr. 0.96) with 325 cc. of nitric acid (sp. gr. +1.2) and dilute with 100 cc. of water.] + +Add 10 cc. of molybdate solution to the nitrate, and leave it for +a few hours. It should then be carefully examined for a !yellow! +precipitate; a white precipitate may be neglected. + +[Note 1: Magnesium ammonium phosphate, as noted below, is slightly +soluble under the conditions of operation. Consequently the +unavoidable errors of analysis are greater in this determination than +in those which have preceded it, and some divergence may be expected +in duplicate analyses. It is obvious that the larger the amount of +substance taken for analysis the less will be the relative loss or +gain due to unavoidable experimental errors; but, in this instance, a +check is placed upon the amount of material which may be taken both by +the bulk of the resulting precipitate of ammonium phosphomolybdate +and by the excessive amount of ammonium molybdate required to effect +complete separation of the phosphoric acid, since a liberal excess +above the theoretical quantity is demanded. Molybdic acid is one of +the more expensive reagents.] + +[Note 2: Soluble silicic acid would, if present, partially separate +with the phosphomolybdate, although not in combination with +molybdenum. Its previous removal by dehydration is therefore +necessary.] + +[Note 3: When washing the siliceous residue the filtrate may be tested +for calcium by adding ammonia, since that reagent neutralizes the +acid which holds the calcium phosphate in solution and causes +precipitation; but after the removal of the phosphoric acid in +combination with the molybdenum, the addition of an oxalate is +required to show the presence of calcium.] + +[Note 4: An excess of nitric acid exerts a slight solvent +action, while ammonium nitrate lessens the solubility; hence the +neutralization of the former by ammonia.] + +[Note 5: The precipitation of the phosphomolybdate takes place more +promptly in warm than in cold solutions, but the temperature should +not exceed 60°C. during precipitation; a higher temperature tends to +separate molybdic acid from the solution. This acid is nearly white, +and its deposition in the filtrate on long standing should not be +mistaken for a second precipitation of the yellow precipitate. The +addition of 75 cc. of ammonium molybdate solution insures the presence +of a liberal excess of the reagent, but the filtrate should be tested +as in all quantitative procedures. + +The precipitation is probably complete in many cases in less than +twelve hours; but it is better, when practicable, to allow the +solution to stand for this length of time. Vigorous shaking or +stirring promotes the separation of the precipitate.] + +[Note 6: The composition of the "yellow precipitate" undoubtedly +varies slightly with varying conditions at the time of its formation. +Its composition may probably fairly be represented by the formula, +(NH_{4})_{3}PO_{4}.12MoO_{3}.H_{2}O, when precipitated under the +conditions prescribed in the procedure. Whatever other variations may +occur in its composition, the ratio of 12 MoO_{3}:1 P seems to +hold, and this fact is utilized in volumetric processes for the +determination of phosphorus, in which the molybdenum is reduced to +a lower oxide and reoxidized by a standard solution of potassium +permanganate. In principle, the procedure is comparable with that +described for the determination of iron by permanganate.] + + +PRECIPITATION OF MAGNESIUM AMMONIUM PHOSPHATE + +PROCEDURE.--Dissolve the precipitate of phosphomolybdate upon the +filter by pouring through it dilute aqueous ammonia (one volume of +dilute ammonia (sp. gr. 0.96) and three volumes of water, which +should be carefully measured), and receive the solution in the beaker +containing the bulk of the precipitate. The total volume of nitrate +and washings should not much exceed 100 cc. Acidify the solution with +dilute hydrochloric acid, and heat it nearly to boiling. Calculate the +volume of magnesium ammonium chloride solution ("magnesia mixture") +required to precipitate the phosphoric acid, assuming 40 per cent +P_{2}O_{5} in the apatite. Measure out about 5 cc. in excess of this +amount, and pour it into the acid solution. Then add slowly dilute +ammonium hydroxide (1 volume of strong ammonia (sp. gr. 0.90) and 9 +volumes of water), stirring constantly until a precipitate forms. Then +add a volume of filtered, concentrated ammonia (sp. gr. 0.90) equal to +one third of the volume of liquid in the beaker (Note 1). Allow the +whole to cool. The precipitated magnesium ammonium phosphate should +then be definitely crystalline in appearance (Note 2). (If it is +desired to hasten the precipitation, the solution may be cooled, first +in cold and then in ice-water, and stirred !constantly! for half an +hour, when precipitation will usually be complete.) + +Decant the clear liquid through a filter, and transfer the precipitate +to the filter, using as wash-water a mixture of one volume of +concentrated ammonia and three volumes of water. It is not necessary +to clean the beaker completely or to wash the precipitate thoroughly +at this point, as it is necessary to purify it by reprecipitation. + +[Note 1: Magnesium ammonium phosphate is not a wholly insoluble +substance, even under the most favorable analytical conditions. It +is least soluble in a liquid containing one fourth of its volume of +concentrated aqueous ammonia (sp. gr. 0.90) and this proportion should +be carefully maintained as prescribed in the procedure. On account of +this slight solubility the volume of solutions should be kept as small +as possible and the amount of wash-water limited to that absolutely +required. + +A large excess of the magnesium solution tends both to throw out +magnesium hydroxide (shown by a persistently flocculent precipitate) +and to cause the phosphate to carry down molybdic acid. The tendency +of the magnesium precipitate to carry down molybdic acid is also +increased if the solution is too concentrated. The volume should not +be less than 90 cc., nor more than 125 cc., at the time of the first +precipitation with the magnesia mixture.] + +[Note 2: The magnesium ammonium phosphate should be perfectly +crystalline, and will be so if the directions are followed. The slow +addition of the reagent is essential, and the stirring not less so. +Stirring promotes the separation of the precipitate and the formation +of larger crystals, and may therefore be substituted for digestion in +the cold. The stirring-rod must not be allowed to scratch the glass, +as the crystals adhere to such scratches and are removed with +difficulty.] + + +REPRECIPITATION AND IGNITION OF MAGNESIUM AMMONIUM PHOSPHATE + +A single precipitation of the magnesium compound in the presence of +molybdenum compounds rarely yields a pure product. The molybdenum can +be removed by solution of the precipitate in acid and precipitation +of the molybdenum by sulphureted hydrogen, after which the magnesium +precipitate may be again thrown down. It is usually more satisfactory +to dissolve the magnesium precipitate and reprecipitate the phosphate +as magnesium ammonium phosphate as described below. + +PROCEDURE.--Dissolve the precipitate from the filter in a little +dilute hydrochloric acid (sp. gr. 1.12), allowing the acid solution to +run into the beaker in which the original precipitation was made (Note +1). Wash the filter with water until the wash-water shows no test for +chlorides, but avoid an unnecessary amount of wash-water. Add to +the solution 2 cc. (not more) of magnesia mixture, and then dilute +ammonium hydroxide solution (sp. gr. 0.96), drop by drop, with +constant stirring, until the liquid smells distinctly of ammonia. Stir +for a few moments and then add a volume of strong ammonia (sp. gr. +0.90), equal to one third of the volume of the solution. Allow the +solution to stand for some hours, and then filter off the magnesium +ammonium phosphate, which should be distinctly crystalline in +character. Wash the precipitate with dilute ammonia water, as +prescribed above, until, finally, 3 cc. of the washings, after +acidifying with nitric acid, show no evidence of chlorides. Test both +filtrates for complete precipitation by adding a few cubic centimeters +of magnesia mixture and allowing them to stand for some time. + +Transfer the moist precipitate to a weighed porcelain or platinum +crucible and ignite, using great care to raise the temperature slowly +while drying the filter in the crucible, and to insure the ready +access of oxygen during the combustion of the filter paper, thus +guarding against a possible reduction of the phosphate, which would +result in disastrous consequences both to the crucible, if of +platinum, and the analysis. Do not raise the temperature above +moderate redness until the precipitate is white. (Keep this precaution +well in mind.) Ignite finally at the highest temperature of the +Tirrill burner, and repeat the heating until the weight is constant. +If the ignited precipitate is persistently discolored by particles of +unburned carbon, moisten the mass with a drop or two of concentrated +nitric acid and heat cautiously, finally igniting strongly. The +acid will dissolve magnesium pyrophosphate from the surface of the +particles of carbon, which will then burn away. Nitric acid also aids +as an oxidizing agent in supplying oxygen for the combustion of the +carbon. + +From the weight of magnesium pyrophosphate (Mg_{2}P_{2}O_{7}) +obtained, calculate the phosphoric anhydride (P_{2}O_{5}) in the +sample of apatite. + +[Note 1: The ionic change involved in the precipitation of the +magnesium compound is + +PO_{4}^{---} + NH_{4}^{+} + Mg^{++} --> [MgNH_{4}PO_{4}]. + +The magnesium ammonium phosphate is readily dissolved by acids, even +those which are no stronger than acetic acid. This is accounted for +by the fact that two of the ions into which phosphoric acid may +dissociate, the HPO_{4}^{--} or H_{2}PO_{4}^{-} ions, exhibit the +characteristics of very weak acids, in that they show almost no +tendency to dissociate further into H^{+} and PO_{4}^{--} ions. +Consequently the ionic changes which occur when the magnesium ammonium +phosphate is brought into contact with an acid may be typified by the +reaction: + +H^{+} + Mg^{++} + NH_{4}^{+} + PO_{4}^{---} --> Mg^{++} + NH_{4}^{+} + +HPO_{4}^{--}; + +that is, the PO_{4}^{--} ions and the H^{+} ions lose their identity +in the formation of the new ion, HPO_{4}^{--}, and this continues +until the magnesium ammonium phosphate is entirely dissolved.] + +[Note 2: During ignition the magnesium ammonium phosphate loses +ammonia and water and is converted into magnesium pyrophosphate: + +2MgNH_{4}PO_{4} --> Mg_{2}P_{2}O_{7} + 2NH_{3} + H_{2}O. + +The precautions mentioned on pages 111 and 123 must be observed with +great care during the ignition of this precipitate. The danger here +lies in a possible reduction of the phosphate by the carbon of the +filter paper, or by the ammonia evolved, which may act as a reducing +agent. The phosphorus then attacks and injures a platinum crucible, +and the determination is valueless.] + + + + +ANALYSIS OF LIMESTONE + + +Limestones vary widely in composition from a nearly pure marble +through the dolomitic limestones, containing varying amounts of +magnesium, to the impure varieties, which contain also ferrous and +manganous carbonates and siliceous compounds in variable proportions. +Many other minerals may be inclosed in limestones in small quantities, +and an exact qualitative analysis will often show the presence of +sulphides or sulphates, phosphates, and titanates, and the alkali or +even the heavy metals. No attempt is made in the following procedures +to provide a complete quantitative scheme which would take into +account all of these constituents. Such a scheme for a complete +analysis of a limestone may be found in Bulletin No. 700 of the United +States Geological Survey. It is assumed that, for these practice +determinations, a limestone is selected which contains only the more +common constituents first enumerated above. + + +DETERMINATION OF MOISTURE + +The determination of the amount of moisture in minerals or ores is +often of great importance. Ores which have been exposed to the weather +during shipment may have absorbed enough moisture to appreciably +affect the results of analysis. Since it is essential that the seller +and buyer should make their analyses upon comparable material, it is +customary for each analyst to determine the moisture in the sample +examined, and then to calculate the percentages of the various +constituents with reference to a sample dried in the air, or at a +temperature a little above 100°C., which, unless the ore has undergone +chemical change because of the wetting, should be the same before and +after shipment. + +PROCEDURE.--Spread 25 grams of the powdered sample on a weighed +watch-glass; weigh to the nearest 10 milligrams only and heat at +105°C.; weigh at intervals of an hour, after cooling in a desiccator, +until the loss of weight after an hour's heating does not exceed +10 milligrams. It should be noted that a variation in weight of 10 +milligrams in a total weight of 25 grams is no greater relatively than +a variation of 0.1 milligram when the sample taken weighs 0.25 gram + +DETERMINATION OF THE INSOLUBLE MATTER AND SILICA + +PROCEDURE.--Weigh out two portions of the original powdered sample +(not the dried sample), of about 5 grams each, into 250 cc. +casseroles, and cover each with a watch-glass (Note 1). Pour over the +powder 25 cc. of water, and then add 50 cc. of dilute hydrochloric +acid (sp. gr. 1.12) in small portions, warming gently, until nothing +further appears to dissolve (Note 2). Evaporate to dryness on the +water bath. Pour over the residue a mixture of 5 cc. of water and +5 cc. of concentrated hydrochloric acid (sp. gr. 1.2) and again +evaporate to dryness, and finally heat for at least an hour at +a temperature of 110°C. Pour over this residue 50 cc. of dilute +hydrochloric acid (one volume acid (sp. gr. 1.12) to five volumes +water), and boil for about five minutes; then filter and wash twice +with the dilute hydrochloric acid, and then with hot water until +free from chlorides. Transfer the filter and contents to a porcelain +crucible, dry carefully over a low flame, and ignite to constant +weight. The residue represents the insoluble matter and the silica +from any soluble silicates (Note 3). + +Calculate the combined percentage of these in the limestone. + +[Note 1: The relatively large weight (5 grams) taken for analysis +insures greater accuracy in the determination of the ingredients which +are present in small proportions, and is also more likely to be a +representative sample of the material analyzed.] + +[Note 2: It is plain that the amount of the insoluble residue and also +its character will often depend upon the strength of acid used for +solution of the limestone. It cannot, therefore, be regarded as +representing any well-defined constituent, and its determination is +essentially empirical.] + +[Note 3: It is probable that some of the silicates present are wholly +or partly decomposed by the acid, and the soluble silicic acid must +be converted by evaporation to dryness, and heating, into white, +insoluble silica. This change is not complete after one evaporation. +The heating at a temperature somewhat higher than that of the water +bath for a short time tends to leave the silica in the form of a +powder, which promotes subsequent filtration. The siliceous residue +is washed first with dilute acid to prevent hydrolytic changes, which +would result in the formation of appreciable quantities of insoluble +basic iron or aluminium salts on the filter when washing with hot +water. + +If it is desired to determine the percentage of silica separately, the +ignited residue should be mixed in a platinum crucible with about six +times its weight of anhydrous sodium carbonate, and the procedure +given on page 151 should be followed. The filtrate from the silica is +then added to the main filtrate from the insoluble residue.] + + + + +DETERMINATION OF FERRIC OXIDE AND ALUMINIUM OXIDE (WITH MANGANESE) + + +PROCEDURE.--To the filtrate from the insoluble residue add ammonium +hydroxide until the solution just smells distinctly of ammonia, but do +not add an excess. Then add 5 cc. of saturated bromine water (Note 1), +and boil for five minutes. If the smell of ammonia has disappeared, +again add ammonium hydroxide in slight excess, and 3 cc. of bromine +water, and heat again for a few minutes. Finally add 10 cc. of +ammonium chloride solution and keep the solution warm until it barely +smells of ammonia; then filter promptly (Note 2). Wash the filter +twice with hot water, then (after replacing the receiving beaker) pour +through it 25 cc. of hot, dilute hydrochloric acid (one volume dilute +HCl [sp. gr. 1.12] to five volumes water). A brown residue insoluble +in the acid may be allowed to remain on the filter. Wash the filter +five times with hot water, add to the filtrate ammonium hydroxide +and bromine water as described above, and repeat the precipitation. +Collect the precipitate on the filter already used, wash it free from +chlorides with hot water, and ignite and weigh as described for ferric +hydroxide on page 110. The residue after ignition consists of ferric +oxide, alumina, and mangano-manganic oxide (Mn_{3}O_{4}), if manganese +is present. These are commonly determined together (Note 3). + +Calculate the percentage of the combined oxides in the limestone. + +[Note 1: The addition of bromine water to the ammoniacal solutions +serves to oxidize any ferrous hydroxide to ferric hydroxide and to +precipitate manganese as MnO(OH)_{2}. The solution must contain not +more than a bare excess of hydroxyl ions (ammonium hydroxide) when it +is filtered, on account of the tendency of the aluminium hydroxide to +redissolve. + +The solution should not be strongly ammoniacal when the bromine is +added, as strong ammonia reacts with the bromine, with the evolution +of nitrogen.] + +[Note 2: The precipitate produced by ammonium hydroxide and bromine +should be filtered off promptly, since the alkaline solution absorbs +carbon dioxide from the air, with consequent partial precipitation +of the calcium as carbonate. This is possible even under the most +favorable conditions, and for this reason the iron precipitate is +redissolved and again precipitated to free it from calcium. When the +precipitate is small, this reprecipitation may be omitted.] + +[Note 3: In the absence of significant amounts of manganese the iron +and aluminium may be separately determined by fusion of the mixed +ignited precipitate, after weighing, with about ten times its weight +of acid potassium sulphate, solution of the cold fused mass in water, +and volumetric determination of the iron, as described on page 66. +The aluminium is then determined by difference, after subtracting the +weight of ferric oxide corresponding to the amount of iron found. + +If a separate determination of the iron, aluminium, and manganese +is desired, the mixed precipitate may be dissolved in acid before +ignition, and the separation effected by special methods (see, for +example, Fay, !Quantitative Analyses!, First Edition, pp. 15-19 and +23-27).] + + + + +DETERMINATION OF CALCIUM + + +PROCEDURE.--To the combined filtrates from the double precipitation of +the hydroxides just described, add 5 cc. of dilute ammonium hydroxide +(sp. gr. 0.96), and transfer the liquid to a 500 cc. graduated flask, +washing out the beaker carefully. Cool to laboratory temperature, and +fill the flask with distilled water until the lowest point of the +meniscus is exactly level with the mark on the neck of the flask. +Carefully remove any drops of water which are on the inside of the +neck of the flask above the graduation by means of a strip of filter +paper, make the solution uniform by pouring it out into a dry beaker +and back into the flask several times. Measure off one fifth of this +solution as follows (Note 1): Pour into a 100 cc. graduated flask +about 10 cc. of the solution, shake the liquid thoroughly over the +inner surface of the small flask, and pour it out. Repeat the same +operation. Fill the 100 cc. flask until the lowest point of the +meniscus is exactly level with the mark on its neck, remove any drops +of solution from the upper part of the neck with filter paper, and +pour the solution into a beaker (400-500 cc.). Wash out the flask with +small quantities of water until it is clean, adding these to the 100 +cc. of solution. When the duplicate portion of 100 cc. is measured out +from the solution, remember that the flask must be rinsed out twice +with that solution, as prescribed above, before the measurement is +made. (A 100 cc. pipette may be used to measure out the aliquot +portions, if preferred.) + +Dilute each of the measured portions to 250 cc. with distilled water, +heat the whole to boiling, and add ammonium oxalate solution slowly +in moderate excess, stirring well. Boil for two minutes; allow the +precipitated calcium oxalate to settle for a half-hour, and decant +through a filter. Test the filtrate for complete precipitation by +adding a few cubic centimeters of the precipitant, allowing it to +stand for fifteen minutes. If no precipitate forms, make the solution +slightly acid with hydrochloric acid (Note 2); see that it is properly +labeled and reserve it to be combined with the filtrate from the +second calcium oxalate precipitation (Notes 3 and 4). + +Redissolve the calcium oxalate in the beaker with warm hydrochloric +acid, pouring the acid through the filter. Wash the filter five times +with water, and finally pour through it aqueous ammonia. Dilute the +solution to 250 cc., bring to boiling, and add 1 cc. ammonium oxalate +solution (Note 5) and ammonia in slight excess; boil for two minutes, +and set aside for a half-hour. Filter off the calcium oxalate upon the +filter first used, and wash free from chlorides. The filtrate should +be made barely acid with hydrochloric acid and combined with the +filtrate from the first precipitation. Begin at once the evaporation +of the solutions for the determination of magnesium as described +below. + +The precipitate of calcium oxalate may be converted into calcium oxide +by ignition without previous drying. After burning the filter, it may +be ignited for three quarters of an hour in a platinum crucible at +the highest heat of the Bunsen or Tirrill burner, and finally for ten +minutes at the blast lamp (Note 6). Repeat the heating over the blast +lamp until the weight is constant. As the calcium oxide absorbs +moisture from the air, it must (after cooling) be weighed as rapidly +as possible. + +The precipitate may, if preferred, be placed in a weighted porcelain +crucible. After burning off the filter and heating for ten minutes the +calcium precipitate may be converted into calcium sulphate by placing +2 cc. of dilute sulphuric acid in the crucible (cold), heating the +covered crucible very cautiously over a low flame to drive off the +excess of acid, and finally at redness to constant weight (Note 7). + +From the weight of the oxide or sulphate, calculate the percentage of +the calcium (Ca) in the limestone, remembering that only one fifth of +the total solution is used for this determination. + +[Note 1: If the calcium were precipitated from the entire solution, +the quantity of the precipitate would be greater than could be +properly treated. The solution is, therefore, diluted to a definite +volume (500 cc.), and exactly one fifth (100 cc.) is measured off in a +graduated flask or by means of a pipette.] + +[Note 2: The filtrate from the calcium oxalate should be made slightly +acid immediately after filtration, in order to avoid the solvent +action of the alkaline liquid upon the glass.] + +[Note 3: The accurate quantitative separation of calcium and magnesium +as oxalates requires considerable care. The calcium precipitate +usually carries down with it some magnesium, and this can best +be removed by redissolving the precipitate after filtration, and +reprecipitation in the presence of only the small amount of magnesium +which was included in the first precipitate. When, however, the +proportion of magnesium is not very large, the second precipitation of +the calcium can usually be avoided by precipitating it from a rather +dilute solution (800 cc. or so) and in the presence of a considerable +excess of the precipitant, that is, rather more than enough to convert +both the magnesium and calcium into oxalates.] + +[Note 4: The ionic changes involved in the precipitation of calcium +as oxalate are exceedingly simple, and the principles discussed in +connection with the barium sulphate precipitation on page 113 also +apply here. The reaction is + +C_{2}O_{4}^{--} + Ca^{++} --> [CaC_{2}O_{4}]. + +Calcium oxalate is nearly insoluble in water, and only very slightly +soluble in acetic acid, but is readily dissolved by the strong mineral +acids. This behavior with acids is explained by the fact that oxalic +acid is a stronger acid than acetic acid; when, therefore, the oxalate +is brought into contact with the latter there is almost no tendency to +diminish the concentration of C_{2}O_{4}^{--} ions by the formation of +an acid less dissociated than the acetic acid itself, and practically +no solvent action ensues. When a strong mineral acid is present, +however, the ionization of the oxalic acid is much reduced by the high +concentration of the H^{+} ions from the strong acid, the formation +of the undissociated acid lessens the concentration of the +C_{2}O_{4}^{--} ions in solution, more of the oxalate passes into +solution to re-establish equilibrium, and this process repeats itself +until all is dissolved. + +The oxalate is immediately reprecipitated from such a solution on the +addition of OH^{-} ions, which, by uniting with the H^{+} ions of the +acids (both the mineral acid and the oxalic acid) to form water, leave +the Ca^{++} and C_{2}O_{4}^{--} ions in the solution to recombine to +form [CaC_{2}O_{4}], which is precipitated in the absence of the +H^{+} ions. It is well at this point to add a small excess of +C_{2}O_{4}^{--} ions in the form of ammonium oxalate to decrease the +solubility of the precipitate. + +The oxalate precipitate consists mainly of CaC_{2}O_{4}.H_{2}O when +thrown down.] + +[Note 5: The small quantity of ammonium oxalate solution is added +before the second precipitation of the calcium oxalate to insure +the presence of a slight excess of the reagent, which promotes the +separation of the calcium compound.] + +[Note 6: On ignition the calcium oxalate loses carbon dioxide and +carbon monoxide, leaving calcium oxide: + +CaC_{2}O_{4}.H_{2}O --> CaO + CO_{2} + CO + H_{2}O. + +For small weights of the oxalate (0.6 gram or less) this reaction may +be brought about in a platinum crucible at the highest temperature of +a Tirrill burner, but it is well to ignite larger quantities than this +over the blast lamp until the weight is constant.] + +[Note 7: The heat required to burn the filter, and that subsequently +applied as described, will convert most of the calcium oxalate to +calcium carbonate, which is changed to sulphate by the sulphuric acid. +The reactions involved are + +CaC_{2}O_{4} --> CaCO_{3} + CO, +CaCO_{3} + H_{2}SO_{4} --> CaSO_{4} + H_{2}O + CO_{2}. + +If a porcelain crucible is employed for ignition, this conversion to +sulphate is to be preferred, as a complete conversion to oxide is +difficult to accomplish.] + +[Note 8: The determination of the calcium may be completed +volumetrically by washing the calcium oxalate precipitate from +the filter into dilute sulphuric acid, warming, and titrating +the liberated oxalic acid with a standard solution of potassium +permanganate as described on page 72. When a considerable number of +analyses are to be made, this procedure will save much of the time +otherwise required for ignition and weighing.] + + + + +DETERMINATION OF MAGNESIUM + + +PROCEDURE.--Evaporate the acidified filtrates from the calcium +precipitates until the salts begin to crystallize, but do !not! +evaporate to dryness (Note 1). Dilute the solution cautiously until +the salts are brought into solution, adding a little acid if the +solution has evaporated to very small volume. The solution should be +carefully examined at this point and must be filtered if a precipitate +has appeared. Heat the clear solution to boiling; remove the burner +and add 25 cc. of a solution of disodium phosphate. Then add slowly +dilute ammonia (1 volume strong ammonia (sp. gr. 0.90) and 9 volumes +water) as long as a precipitate continues to form. Finally, add a +volume of concentrated ammonia (sp. gr. 0.90) equal to one third of +the volume of the solution, and allow the whole to stand for about +twelve hours. + +Decant the solution through a filter, wash it with dilute ammonia +water, proceeding as prescribed for the determination of phosphoric +anhydride on page 122, including; the reprecipitation (Note 2), +except that 3 cc. of disodium phosphate solution are added before the +reprecipitation of the magnesium ammonium phosphate instead of +the magnesia mixture there prescribed. From the weight of the +pyrophosphate, calculate the percentage of magnesium oxide (MgO) in +the sample of limestone. Remember that the pyrophosphate finally +obtained is from one fifth of the original sample. + +[Note 1: The precipitation of the magnesium should be made in as small +volume as possible, and the ratio of ammonia to the total volume of +solution should be carefully provided for, on account of the relative +solubility of the magnesium ammonium phosphate. This matter has +been fully discussed in connection with the phosphoric anhydride +determination.] + +[Note 2: The first magnesium ammonium phosphate precipitate is rarely +wholly crystalline, as it should be, and is not always of the proper +composition when precipitated in the presence of such large amounts of +ammonium salts. The difficulty can best be remedied by filtering the +precipitate and (without washing it) redissolving in a small quantity +of hydrochloric acid, from which it may be again thrown down by +ammonia after adding a little disodium phosphate solution. If the +flocculent character was occasioned by the presence of magnesium +hydroxide, the second precipitation, in a smaller volume containing +fewer salts, will often result more favorably. + +The removal of iron or alumina from a contaminated precipitate is +a matter involving a long procedure, and a redetermination of the +magnesium from a new sample, with additional precautions, is usually +to be preferred.] + + + + +DETERMINATION OF CARBON DIOXIDE + + +!Absorption Apparatus! + +[Illustration: Fig. 3] + +The apparatus required for the determination of the carbon dioxide +should be arranged as shown in the cut (Fig. 3). The flask (A) is +an ordinary wash bottle, which should be nearly filled with dilute +hydrochloric acid (100 cc. acid (sp. gr. 1.12) and 200 cc. of water). +The flask is connected by rubber tubing (a) with the glass tube (b) +leading nearly to the bottom of the evolution flask (B) and having its +lower end bent upward and drawn out to small bore, so that the carbon +dioxide evolved from the limestone cannot bubble back into (b). The +evolution flask should preferably be a wide-mouthed Soxhlet extraction +flask of about 150 cc. capacity because of the ease with which tubes +and stoppers may be fitted into the neck of a flask of this type. The +flask should be fitted with a two-hole rubber stopper. The condenser +(C) may consist of a tube with two or three large bulbs blown in +it, for use as an air-cooled condenser, or it may be a small +water-jacketed condenser. The latter is to be preferred if a number of +determinations are to be made in succession. + +A glass delivery tube (c) leads from the condenser to the small U-tube +(D) containing some glass beads or small pieces of glass rod and 3 cc. +of a saturated solution of silver sulphate, with 3 cc. of concentrated +sulphuric acid (sp. gr. 1.84). The short rubber tubing (d) connects +the first U-tube to a second U-tube (E) which is filled with small +dust-free lumps of dry calcium chloride, with a small, loose plug of +cotton at the top of each arm. Both tubes should be closed by cork +stoppers, the tops of which are cut off level with, or preferably +forced a little below, the top of the U-tube, and then neatly sealed +with sealing wax. + +The carbon dioxide may be absorbed in a tube containing soda lime +(F) or in a Geissler bulb (F') containing a concentrated solution +of potassium hydroxide (Note 2). The tube (F) is a glass-stoppered +side-arm U-tube in which the side toward the evolution flask and one +half of the other side are filled with small, dust-free lumps of soda +lime of good quality (Note 3). Since soda lime contains considerable +moisture, the other half of the right side of the tube is filled with +small lumps of dry, dust-free calcium chloride to retain the moisture +from the soda lime. Loose plugs of cotton are placed at the top of +each arm and between the soda lime and the calcium chloride. + +The Geissler bulb (F'), if used, should be filled with potassium +hydroxide solution (1 part of solid potassium hydroxide dissolved in +two parts of water) until each small bulb is about two thirds full +(Note 4). A small tube containing calcium chloride is connected with +the Geissler bulb proper by a ground joint and should be wired to the +bulb for safety. This is designed to retain any moisture from the +hydroxide solution. A piece of clean, fine copper wire is so attached +to the bulb that it can be hung from the hook above a balance pan, or +other support. + +The small bottle (G) with concentrated sulphuric acid (sp. gr. 1.84) +is so arranged that the tube (f) barely dips below the surface. This +will prevent the absorption of water vapor by (F) or (F') and serves +as an aid in regulating the flow of air through the apparatus. (H) is +an aspirator bottle of about four liters capacity, filled with water; +(k) is a safety tube and a means of refilling (H); (h) is a screw +clamp, and (K) a U-tube filled with soda lime. + +[Note 1: The air current, which is subsequently drawn through the +apparatus, to sweep all of the carbon dioxide into the absorption +apparatus, is likely to carry with it some hydrochloric acid from +the evolution flask. This acid is retained by the silver sulphate +solution. The addition of concentrated sulphuric acid to this solution +reduces its vapor pressure so far that very little water is carried on +by the air current, and this slight amount is absorbed by the calcium +chloride in (E). As the calcium chloride frequently contains a small +amount of a basic material which would absorb carbon dioxide, it is +necessary to pass carbon dioxide through (E) for a short time and then +drive all the gas out with a dry air current for thirty minutes before +use.] + +[Note 2: Soda-lime absorption tubes are to be preferred if a +satisfactory quality of soda lime is available and the number of +determinations to be made successively is small. The potash bulbs will +usually permit of a larger number of successive determinations without +refilling, but they require greater care in handling and in the +analytical procedure.] + +[Note 3: Soda lime is a mixture of sodium and calcium hydroxides. Both +combine with carbon dioxide to form carbonates, with the evolution +of water. Considerable heat is generated by the reaction, and the +temperature of the tube during absorption serves as a rough index of +the progress of the reaction through the mass of soda lime. + +It is essential that soda lime of good quality for analytical purposes +should be used. The tube should not contain dust, as this is likely to +be swept away.] + +[Note 4: The solution of the hydroxide for use in the Geissler bulb +must be highly concentrated to insure complete absorption of the +carbon dioxide and also to reduce the vapor pressure of the solution, +thus lessening the danger of loss of water with the air which passes +through the bulbs. The small quantity of moisture which is then +carried out of the bulbs is held by the calcium chloride in the +prolong tube. The best form of absorption bulb is that to which the +prolong tube is attached by a ground glass joint. + +After the potassium hydroxide is approximately half consumed in the +first bulb of the absorption apparatus, potassium bicarbonate is +formed, and as it is much less soluble than the carbonate, it often +precipitates. Its formation is a warning that the absorbing power of +the hydroxide is much diminished.] + + +!The Analysis! + +PROCEDURE.-- Weigh out into the flask (B) about 1 gram of limestone. +Cover it with 15 cc. of water. Weigh the absorption apparatus (F) +or (F') accurately after allowing it to stand for 30 minutes in the +balance case, and wiping it carefully with a lintless cloth, taking +care to handle it as little as possible after wiping (Note 1). Connect +the absorption apparatus with (e) and (f). If a soda-lime tube is +used, be sure that the arm containing the soda lime is next the tube +(E) and that the glass stopcocks are open. + +To be sure that the whole apparatus is airtight, disconnect the rubber +tube from the flask (A), making sure that the tubes (a) and (b) do not +contain any hydrochloric acid, close the pinchcocks (a) and (k) and +open (h). No bubbles should pass through (D) or (G) after a few +seconds. When assured that the fittings are tight, close (h) and open +(a) cautiously to admit air to restore atmospheric pressure. This +precaution is essential, as a sudden inrush of air will project liquid +from (D) or (F'). Reconnect the rubber tube with the flask (A). +Open the pinchcocks (a) and (k) and blow over about 10 cc. of the +hydrochloric acid from (A) into (B). When the action of the acid +slackens, blow over (slowly) another 10 cc. + +The rate of gas evolution should not exceed for more than a few +seconds that at which about two bubbles per second pass through (G) +(Note 2). Repeat the addition of acid in small portions until the +action upon the limestone seems to be at an end, taking care to close +(a) after each addition of acid (Note 3). Disconnect (A) and connect +the rubber tubing with the soda-lime tube (K) and open (a). Then close +(k) and open (h), regulating the flow of water from (H) in such a way +that about two bubbles per second pass through (G). Place a small +flame under (B) and !slowly! raise the contents to boiling and boil +for three minutes. Then remove the burner from under (B) and continue +to draw air through the apparatus for 20-30 minutes, or until (H) +is emptied (Note 4). Remove the absorption apparatus, closing the +stopcocks on (F) or stoppering the open ends of (F'), leave the +apparatus in the balance case for at least thirty minutes, wipe it +carefully and weigh, after opening the stopcocks (or removing plugs). +The increase in weight is due to absorption of CO_{2}, from which its +percentage in the sample may be calculated. + +After cleaning (B) and refilling (H), the apparatus is ready for the +duplicate analysis. + +[Note 1: The absorption tubes or bulbs have large surfaces on which +moisture may collect. By allowing them to remain in the balance case +for some time before weighing, the amount of moisture absorbed on the +surface is as nearly constant as practicable during two weighings, and +a uniform temperature is also assured. The stopcocks of the U-tube +should be opened, or the plugs used to close the openings of the +Geissler bulb should be removed before weighing in order that the air +contents shall always be at atmospheric pressure.] + +[Note 2: If the gas passes too rapidly into the absorption apparatus, +some carbon dioxide may be carried through, not being completely +retained by the absorbents.] + +[Note 3: The essential ionic changes involved in this procedure are +the following: It is assumed that the limestone, which is typified by +calcium carbonate, is very slightly soluble in water, and the ions +resulting are Ca^{++} and CO_{3}^{--}. In the presence of H^{+} ions +of the mineral acid, the CO_{3}^{--} ions form [H_{2}CO_{3}]. This +is not only a weak acid which, by its formation, diminishes the +concentration of the CO_{3}^{--} ions, thus causing more of the +carbonate to dissolve to re-establish equilibrium, but it is also an +unstable compound and breaks down into carbon dioxide and water.] + +[Note 4: Carbon dioxide is dissolved by cold water, but the gas is +expelled by boiling, and, together with that which is distributed +through the apparatus, is swept out into the absorption bulb by the +current of air. This air is purified by drawing it through the tube +(K) containing soda lime, which removes any carbon dioxide which may +be in it.] + + + + +DETERMINATION OF LEAD, COPPER, IRON, AND ZINC IN BRASS + +ELECTROLYTIC SEPARATIONS + + +!General Discussion! + +When a direct current of electricity passes from one electrode to +another through solutions of electrolytes, the individual ions present +in these solutions tend to move toward the electrode of opposite +electrical charge to that which each ion bears, and to be discharged +by that electrode. Whether or not such discharge actually occurs in +the case of any particular ion depends upon the potential (voltage) of +the current which is passing through the solution, since for each ion +there is, under definite conditions, a minimum potential below which +the discharge of the ion cannot be effected. By taking advantage +of differences in discharge-potentials, it is possible to effect +separations of a number of the metallic ions by electrolysis, and at +the same time to deposit the metals in forms which admit of direct +weighing. In this way the slower procedures of precipitation and +filtration may frequently be avoided. The following paragraphs present +a brief statement of the fundamental principles and conditions +underlying electro-analysis. + +The total energy of an electric current as it passes through a +solution is distributed among three factors, first, its potential, +which is measured in volts, and corresponds to what is called "head" +in a stream of water; second, current strength, which is measured +in amperes, and corresponds to the volume of water passing a +cross-section of a stream in a given time interval; and third, the +resistance of the conducting medium, which is measured in ohms. The +relation between these three factors is expressed by Ohm's law, +namely, that !I = E/R!, when I is current strength, E potential, and R +resistance. It is plain that, for a constant resistance, the +strength of the current and its potential are mutually and directly +interdependent. + +As already stated, the applied electrical potential determines whether +or not deposition of a metal upon an electrode actually occurs. The +current strength determines the rate of deposition and the physical +characteristics of the deposit. The resistance of the solution is +generally so small as to fall out of practical consideration. + +Approximate deposition-potentials have been determined for a number +of the metallic elements, and also for hydrogen and some of the +acid-forming radicals. The values given below are those required +for deposition from normal solutions at ordinary temperatures +with reference to a hydrogen electrode. They must be regarded as +approximate, since several disturbing factors and some secondary +reactions render difficult their exact application under the +conditions of analysis. They are: + + Zn Cd Fe Ni Pb H Cu Sb Hg Ag SO_{4} ++0.77 +0.42 +0.34 +0.33 +0.13 0 -0.34 -0.67 -0.76 -0.79 +1.90 + +From these data it is evident that in order to deposit copper from a +normal solution of copper sulphate a minimum potential equal to the +algebraic sum of the deposition-potentials of copper ions and sulphate +ions must be applied, that is, +1.56 volts. The deposition of zinc +from a solution of zinc sulphate would require +2.67 volts, but, since +the deposition of hydrogen from sulphuric acid solution requires only ++1.90 volts, the quantitative deposition of zinc by electrolysis from +a sulphuric acid solution of a zinc salt is not practicable. On the +other hand, silver, if present in a solution of copper sulphate, would +deposit with the copper. + +The foregoing examples suffice to illustrate the application of the +principle of deposition potentials, but it must further be noted +that the values stated apply to normal solutions of the compounds in +question, that is, to solutions of considerable concentrations. As the +concentration of the ions diminishes, and hence fewer ions approach +the electrodes, somewhat higher voltages are required to attract and +discharge them. From this it follows that the concentrations should be +kept as high as possible to effect complete deposition in the least +practicable time, or else the potentials applied must be progressively +increased as deposition proceeds. In practice, the desired result is +obtained by starting with small volumes of solution, using as large an +electrode surface as possible, and by stirring the solution to bring +the ions in contact with the electrodes. This is, in general, a more +convenient procedure than that of increasing the potential of the +current during electrolysis, although that method is also used. + +As already stated, those ions in a solution of electrolytes will first +be discharged which have the lowest deposition potentials, and so +long as these ions are present around the electrode in considerable +concentration they, almost alone, are discharged, but, as their +concentration diminishes, other ions whose deposition potentials are +higher but still within that of the current applied, will also begin +to separate. For example, from a nitric acid solution of copper +nitrate, the copper ions will first be discharged at the cathode, but +as they diminish in concentration hydrogen ions from the acid (or +water) will be also discharged. Since the hydrogen thus liberated is a +reducing agent, the nitric acid in the solution is slowly reduced to +ammonia, and it may happen that if the current is passed through for a +long time, such a solution will become alkaline. Oxygen is liberated +at the anode, but, since there is no oxidizable substance present +around that electrode, it escapes as oxygen gas. It should be noted +that, in general, the changes occurring at the cathode are reductions, +while those at the anode are oxidations. + +For analytical purposes, solutions of nitrates or sulphates of the +metals are preferable to those of the chlorides, since liberated +chlorine attacks the electrodes. In some cases, as for example, that +of silver, solution of salts forming complex ions, like that of +the double cyanide of silver and potassium, yield better metallic +deposits. + +Most metals are deposited as such upon the cathode; a few, notably +lead and manganese, separate in the form of dioxides upon the anode. +It is evidently important that the deposited material should be so +firmly adherent that it can be washed, dried, and weighed without +loss in handling. To secure these conditions it is essential that the +current density (that is, the amount of current per unit of area of +the electrodes) shall not be too high. In prescribing analytical +conditions it is customary to state the current strength in "normal +densities" expressed in amperes per 100 sq. cm. of electrode surface, +as, for example, "N.D_{100} = 2 amps." + +If deposition occurs too rapidly, the deposit is likely to be spongy +or loosely adherent and falls off on subsequent treatment. This places +a practical limit to the current density to be employed, for a given +electrode surface. The cause of the unsatisfactory character of +the deposit is apparently sometimes to be found in the coincident +liberation of considerable hydrogen and sometimes in the failure of +the rapidly deposited material to form a continuous adherent surface. +The effect of rotating electrodes upon the character of the deposit is +referred to below. + +The negative ions of an electrolyte are attracted to the anode and are +discharged on contact with it. Anions such as the chloride ion yield +chlorine atoms, from which gaseous chlorine molecules are formed +and escape. The radicals which compose such ions as NO_{3}^{-} or +SO_{4}^{--} are not capable of independent existence after discharge, +and break down into oxygen and N_{2}O_{5} and SO_{3} respectively. The +oxygen escapes and the anhydrides, reacting with water, re-form nitric +and sulphuric acids. + +The law of Faraday expresses the relation between current strength and +the quantities of the decomposition products which, under constant +conditions, appear at the electrodes, namely, that a given quantity +of electricity, acting for a given time, causes the separation of +chemically equivalent quantities of the various elements or radicals. +For example, since 107.94 grams of silver is equivalent to 1.008 grams +of hydrogen, and that in turn to 8 grams of oxygen, or 31.78 grams of +copper, the quantity of electricity which will cause the deposit of +107.94 grams of silver in a given time will also separate the weights +just indicated of the other substances. Experiments show that a +current of one ampere passing for one second, i.e., a coulomb of +electricity, causes the deposition of 0.001118 gram of silver from a +normal solution of a silver salt. The number of coulombs required to +deposit 107.94 grams is 107.94/0.001118 or 96,550 and the same number +of coulombs will also cause the separation of 1.008 grams of hydrogen, +8 grams of oxygen or 31.78 grams of copper. While it might at first +appear that Faraday's law could thus be used as a basis for the +calculation of the time required for the deposition of a given +quantity of an electrolyte from solution, it must be remembered that +the law expresses what occurs when the concentration of the ions in +the solution is kept constant, as, for example, when the anode in +a silver salt solution is a plate of metallic silver. Under the +conditions of electro-analysis the concentration of the ions is +constantly diminishing as deposition proceeds and the time actually +required for complete deposition of a given weight of material by +a current of constant strength is, therefore, greater than that +calculated on the basis of the law as stated above. + +The electrodes employed in electro-analysis are almost exclusively +of platinum, since that metal alone satisfactorily resists chemical +action of the electrolytes, and can be dried and weighed without +change in composition. The platinum electrodes may be used in the +form of dishes, foil or gauze. The last, on account of the ease of +circulation of the electrolyte, its relatively large surface in +proportion to its weight and the readiness with which it can be washed +and dried, is generally preferred. + +Many devices have been described by the use of which the electrode +upon which deposition occurs can be mechanically rotated. This has an +effect parallel to that of greatly increasing the electrode surface +and also provides a most efficient means of stirring the solution. +With such an apparatus the amperage may be increased to 5 or even 10 +amperes and depositions completed with great rapidity and accuracy. It +is desirable, whenever practicable, to provide a rotating or stirring +device, since, for example, the time consumed in the deposition of the +amount of copper usually found in analysis may be reduced from the +20 to 24 hours required with stationary electrodes, and unstirred +solutions, to about one half hour. + + + + +DETERMINATION OF COPPER AND LEAD + + +PROCEDURE.--Weigh out two portions of about 0.5 gram each (Note 1) +into tall, slender lipless beakers of about 100 cc. capacity. Dissolve +the metal in a solution of 5 cc. of dilute nitric acid (sp. gr. 1.20) +and 5 cc. of water, heating gently, and keeping the beaker covered. +When the sample has all dissolved (Note 2), wash down the sides of the +beaker and the bottom of the watch-glass with water and dilute the +solution to about 50 cc. Carefully heat to boiling and boil for a +minute or two to expel nitrous fumes. + +Meanwhile, four platinum electrodes, two anodes and two cathodes, +should be cleaned by dipping in dilute nitric acid, washing with water +and finally with 95 per cent alcohol (Note 3). The alcohol may be +ignited and burned off. The electrodes are then cooled in a desiccator +and weighed. Connect the electrodes with the binding posts (or other +device for connection with the electric circuit) in such a way that +the copper will be deposited upon the electrode with the larger +surface, which is made the cathode. The beaker containing the solution +should then be raised into place from below the electrodes until the +latter reach nearly to the bottom of the beaker. The support for the +beaker must be so arranged that it can be easily raised or lowered. + +If the electrolytic apparatus is provided with a mechanism for the +rotation of the electrode or stirring of the electrolyte, proceed as +follows: Arrange the resistance in the circuit to provide a direct +current of about one ampere. Pass this current through the solution +to be electrolyzed, and start the rotating mechanism. Keep the beaker +covered as completely as possible, using a split watch-glass (or other +device) to avoid loss by spattering. When the solution is colorless, +which is usually the case after about 35 minutes, rinse off the cover +glass, wash down the sides of the beaker, add about 0.30 gram of urea +and continue the electrolysis for another five minutes (Notes 4 and +5). + +If stationary electrodes are employed, the current strength should be +about 0.1 ampere, which may, after 12 to 15 hours, be increased to 0.2 +ampere. The time required for complete deposition is usually from 20 +to 24 hours. It is advisable to add 5 cc. of nitric acid (sp. gr. 1.2) +if the electrolysis extends over this length of time. No urea is added +in this case. + +When the deposition of the copper appears to be complete, stop the +rotating mechanism and slowly lower the beaker with the left hand, +directing at the same time a stream of water from a wash bottle on +both electrodes. Remove the beaker, shut off the current, and, if +necessary, complete the washing of the electrodes (Note 6). Rinse the +electrodes cautiously with alcohol and heat them in a hot closet until +the alcohol has just evaporated, but no longer, since the copper is +likely to oxidize at the higher temperature. (The alcohol may be +removed by ignition if care is taken to keep the electrodes in motion +in the air so that the copper deposit is not too strongly heated at +any one point.) + +Test the solution in the beaker for copper as follows, remembering +that it is to be used for subsequent determinations of iron and zinc: +Remove about 5 cc. and add a slight excess of ammonia. Compare the +mixture with some distilled water, holding both above a white surface. +The solution should not show any tinge of blue. If the presence of +copper is indicated, add the test portion to the main solution, +evaporate the whole to a volume of about 100 cc., and again +electrolyze with clean electrodes (Note 7). + +After cooling the electrodes in a desiccator, weigh them and from the +weight of copper on the cathode and of lead dioxide (PbO_{2}) on the +anode, calculate the percentage of copper (Cu) and of lead (Pb) in the +brass. + +[Note 1: It is obvious that the brass taken for analysis should be +untarnished, which can be easily assured, when wire is used, by +scouring with emery. If chips or borings are used, they should be well +mixed, and the sample for analysis taken from different parts of the +mixture.] + +[Note 2: If a white residue remains upon treatment of the alloy with +nitric acid, it indicates the presence of tin. The material is not, +therefore, a true brass. This may be treated as follows: Evaporate the +solution to dryness, moisten the residue with 5 cc. of dilute nitric +acid (sp. gr. 1.2) and add 50 cc. of hot water. Filter off the +meta-stannic acid, wash, ignite in porcelain and weigh as SnO_{2}. +This oxide is never wholly free from copper and must be purified for +an exact determination. If it does not exceed 2 per cent of the alloy, +the quantity of copper which it contains may usually be neglected.] + +[Note 3: The electrodes should be freed from all greasy matter before +using, and those portions upon which the metal will deposit should not +be touched with the fingers after cleaning.] + +[Note 4: Of the ions in solution, the H^{+}, Cu^{++}, Zn^{++}, and +Fe^{+++} ions tend to move toward the cathode. The NO_{3}^{-} ions and +the lead, probably in the form of PbO_{2}^{--} ions, move toward the +anode. At the cathode the Cu^{++} ions are discharged and plate out as +metallic copper. This alone occurs while the solution is relatively +concentrated. Later on, H^{+} ions are also discharged. In the +presence of considerable quantities of H^{+} ions, as in this acid +solution, no Zn^{++} or Fe^{+++} ions are discharged because of their +greater deposition potentials. At the anode the lead is deposited as +PbO_{2} and oxygen is evolved. + +For the reasons stated on page 141 care must be taken that the +solution does not become alkaline if the electrolysis is long +continued.] + +[Note 5: Urea reacts with nitrous acid, which may be formed in the +solution as a result of the reducing action of the liberated hydrogen. +Its removal promotes the complete precipitation of the copper. The +reaction is + +CO(NH_{2})_{2} + 2HNO_{2} --> CO_{2} + 2N_{2} + 3H_{2}O.] + +[Note 6: The electrodes must be washed nearly or quite free from +the nitric acid solution before the circuit is broken to prevent +re-solution of the copper. + +If several solutions are connected in the same circuit it is obvious +that some device must be used to close the circuit as soon as the +beaker is removed.] + +[Note 7: The electrodes upon which the copper has been deposited +may be cleaned by immersion in warm nitric acid. To remove the lead +dioxide, add a few crystals of oxalic acid to the nitric acid.] + + + + +DETERMINATION OF IRON + + +Most brasses contain small percentages of iron (usually not over 0.1 +per cent) which, unless removed, is precipitated as phosphate and +weighed with the zinc. + +PROCEDURE.--To the solution from the precipitation of copper and +lead by electrolysis, add dilute ammonia (sp. gr. 0.96) until the +precipitate of zinc hydroxide which first forms re-dissolves, leaving +only a slight red precipitate of ferric hydroxide. Filter off the +iron precipitate, using a washed filter, and wash five times with hot +water. Test a portion of the last washing with a dilute solution of +ammonium sulphide to assure complete removal of the zinc. + +The precipitate may then be ignited and weighed as ferric oxide, as +described on page 110. + +Calculate the percentage of iron (Fe) in the brass. + + + + +DETERMINATION OF ZINC + + +PROCEDURE.--Acidify the filtrate from the iron determination with +dilute nitric acid. Concentrate it to 150 cc. Add to the cold solution +dilute ammonia (sp. gr. 0.96) cautiously until it barely smells of +ammonia; then add !one drop! of a dilute solution of litmus (Note 1), +and drop in, with the aid of a dropper, dilute nitric acid until the +blue of the litmus just changes to red. It is important that this +point should not be overstepped. Heat the solution nearly to boiling +and pour into it slowly a filtered solution of di-ammonium hydrogen +phosphate[1] containing a weight of the phosphate about equal +to twelve times that of the zinc to be precipitated. (For this +calculation the approximate percentage of zinc is that found by +subtracting the sum of the percentages of the copper, lead and iron +from 100 per cent.) Keep the solution just below boiling for fifteen +minutes, stirring frequently (Note 2). If at the end of this time the +amorphous precipitate has become crystalline, allow the solution to +cool for about four hours, although a longer time does no harm (Note +3), and filter upon an asbestos filter in a porcelain Gooch crucible. +The filter is prepared as described on page 103, and should be dried +to constant weight at 105°C. + +[Footnote 1: The ammonium phosphate which is commonly obtainable +contains some mono-ammonium salt, and this is not satisfactory as a +precipitant. It is advisable, therefore, to weigh out the amount of +the salt required, dissolve it in a small volume of water, add a drop +of phenolphthalein solution, and finally add dilute ammonium hydroxide +solution cautiously until the solution just becomes pink, but do not +add an excess.] + +Wash the precipitate until free from sulphates with a warm 1 per cent +solution of the di-ammonium phosphate, and then five times with 50 per +cent alcohol (Note 4). Dry the crucible and precipitate for an hour at +105°C., and finally to constant weight (Note 5). The filtrate should +be made alkaline with ammonia and tested for zinc with a few drops of +ammonium sulphide, allowing it to stand (Notes 6, 7 and 8). + +From the weight of the zinc ammonium phosphate (ZnNH_{4}PO_{4}) +calculate the percentage of the zinc (Zn) in the brass. + +[Note 1: The zinc ammonium phosphate is soluble both in acids and in +ammonia. It is, therefore, necessary to precipitate the zinc in a +nearly neutral solution, which is more accurately obtained by adding +a drop of a litmus solution to the liquid than by the use of litmus +paper.] + +[Note 2: The precipitate which first forms is amorphous, and may have +a variable composition. On standing it becomes crystalline and then +has the composition ZnNH_{4}PO_{4}. The precipitate then settles +rapidly and is apt to occasion "bumping" if the solution is heated to +boiling. Stirring promotes the crystallization.] + +[Note 3: In a carefully neutralized solution containing a considerable +excess of the precipitant, and also ammonium salts, the separation +of the zinc is complete after standing four hours. The ionic changes +connected with the precipitation of the zinc as zinc ammonium +phosphate are similar to those described for magnesium ammonium +phosphate, except that the zinc precipitate is soluble in an excess of +ammonium hydroxide, probably as a result of the formation of complex +ions of the general character Zn(NH_{3})_{4}^{++}.] + +[Note 4: The precipitate is washed first with a dilute solution of the +phosphate to prevent a slight decomposition of the precipitate (as a +result of hydrolysis) if hot water alone is used. The alcohol is added +to the final wash-water to promote the subsequent drying.] + +[Note 5: The precipitate may be ignited and weighed as +Zn_{2}P_{2}O_{7}, by cautiously heating the porcelain Gooch crucible +within a nickel or iron crucible, used as a radiator. The heating +must be very slow at first, as the escaping ammonia may reduce the +precipitate if it is heated too quickly.] + +[Note 6: If the ammonium sulphide produced a distinct precipitate, +this should be collected on a small filter, dissolved in a few cubic +centimeters of dilute nitric acid, and the zinc reprecipitated as +phosphate, filtered off, dried, and weighed, and the weight added to +that of the main precipitate.] + +[Note 7: It has been found that some samples of asbestos are acted +upon by the phosphate solution and lose weight. An error from this +source may be avoided by determining the weight of the crucible +and filter after weighing the precipitate. For this purpose the +precipitate may be dissolved in dilute nitric acid, the asbestos +washed thoroughly, and the crucible reweighed.] + +[Note 8. The details of this method of precipitation of zinc are fully +discussed in an article by Dakin, !Ztschr. Anal. Chem.!, 39 (1900), +273.] + + + + +DETERMINATION OF SILICA IN SILICATES + + +Of the natural silicates, or artificial silicates such as slags and +some of the cements, a comparatively few can be completely decomposed +by treatment with acids, but by far the larger number require fusion +with an alkaline flux to effect decomposition and solution +for analysis. The procedure given below applies to silicates +undecomposable by acids, of which the mineral feldspar is taken as a +typical example. Modifications of the procedure, which are applicable +to silicates which are completely or partially decomposable by acids, +are given in the Notes on page 155. + + +PREPARATION OF THE SAMPLE + +Grind about 3 grams of the mineral in an agate mortar (Note 1) until +no grittiness is to be detected, or, better, until it will entirely +pass through a sieve made of fine silk bolting cloth. The sieve may be +made by placing a piece of the bolting cloth over the top of a small +beaker in which the ground mineral is placed, holding the cloth in +place by means of a rubber band below the lip of the beaker. By +inverting the beaker over clean paper and gently tapping it, the fine +particles pass through the sieve, leaving the coarser particles within +the beaker. These must be returned to the mortar and ground, and the +process of sifting and grinding repeated until the entire sample +passes through the sieve. + +[Note 1: If the sample of feldspar for analysis is in the massive or +crystalline form, it should be crushed in an iron mortar until the +pieces are about half the size of a pea, and then transferred to a +steel mortar, in which they are reduced to a coarse powder. A wooden +mallet should always be used to strike the pestle of the steel mortar, +and the blows should not be sharp. + +It is plain that final grinding in an agate mortar must be continued +until the whole of the portion of the mineral originally taken has +been ground so that it will pass the bolting cloth, otherwise the +sifted portion does not represent an average sample, the softer +ingredients, if foreign matter is present, being first reduced to +powder. For this reason it is best to start with not more than the +quantity of the feldspar needed for analysis. The mineral must be +thoroughly mixed after the grinding.] + + +FUSION AND SOLUTION + +PROCEDURE.--Weigh into platinum crucibles two portions of the ground +feldspar of about 0.8 gram each. Weigh on rough balances two portions +of anhydrous sodium carbonate, each amounting to about six times the +weight of the feldspar taken for analysis (Note 1). Pour about three +fourths of the sodium carbonate into the crucible, place the latter on +a piece of clean, glazed paper, and thoroughly mix the substance and +the flux by carefully stirring for several minutes with a dry glass +rod, the end of which has been recently heated and rounded in a flame +and slowly cooled. The rod may be wiped off with a small fragment of +filter paper, which may be placed in the crucible. Place the remaining +fourth of the carbonate on the top of the mixture. Cover the crucible, +heat it to dull redness for five minutes, and then gradually increase +the heat to the full capacity of a Bunsen or Tirrill burner for +twenty minutes, or until a quiet, liquid fusion is obtained (Note 2). +Finally, heat the sides and cover strongly until any material which +may have collected upon them is also brought to fusion. + +Allow the crucible to cool, and remove the fused mass as directed on +page 116. Disintegrate the mass by placing it in a previously prepared +mixture of 100 cc. of water and 50 cc. of dilute hydrochloric acid +(sp. gr. 1.12) in a covered casserole (Note 3). Clean the crucible and +lid by means of a little hydrochloric acid, adding this acid to the +main solution (Notes 4 and 5). + +[Note 1: Quartz, and minerals containing very high percentages of +silica, may require eight or ten parts by weight of the flux to insure +a satisfactory decomposition.] + +[Note 2: During the fusion the feldspar, which, when pure, is a +silicate of aluminium and either sodium or potassium, but usually +contains some iron, calcium, and magnesium, is decomposed by the +alkaline flux. The sodium of the latter combines with the silicic acid +of the silicate, with the evolution of carbon dioxide, while about two +thirds of the aluminium forms sodium aluminate and the remainder +is converted into basic carbonate, or the oxide. The calcium and +magnesium, if present, are changed to carbonates or oxides. + +The heat is applied gently to prevent a too violent reaction when +fusion first takes place.] + +[Note 3: The solution of a silicate by a strong acid is the result of +the combination of the H^{+} ions of the acid and the silicate ions +of the silicate to form a slightly ionized silicic acid. As a +consequence, the concentration of the silicate ions in the solution is +reduced nearly to zero, and more silicate dissolves to re-establish +the disturbed equilibrium. This process repeats itself until all of +the silicate is brought into solution. + +Whether the resulting solution of the silicate contains ortho-silicic +acid (H_{4}SiO_{4}) or whether it is a colloidal solution of some +other less hydrated acid, such as meta-silicic acid (H_{2}SiO_{3}), +is a matter that is still debatable. It is certain, however, that the +gelatinous material which readily separates from such solutions is of +the nature of a hydrogel, that is, a colloid which is insoluble in +water. This substance when heated to 100°C., or higher, is completely +dehydrated, leaving only the anhydride, SiO_{2}. The changes may be +represented by the equation: + +SiO_{3}^{--} + 2H^{+} --> [H_{2}SiO_{3}] --> H_{2}O + SiO_{2}.] + +[Note 4: A portion of the fused mass is usually projected upward by +the escaping carbon dioxide during the fusion. The crucible must +therefore be kept covered as much as possible and the lid carefully +cleaned.] + +[Note 5: A gritty residue remaining after the disintegration of +the fused mass by acid indicates that the substance has been but +imperfectly decomposed. Such a residue should be filtered, washed, +dried, ignited, and again fused with the alkaline flux; or, if the +quantity of material at hand will permit, it is better to reject the +analysis, and to use increased care in grinding the mineral and in +mixing it with the flux.] + + +DEHYDRATION AND FILTRATION + +PROCEDURE.--Evaporate the solution of the fusion to dryness, stirring +frequently until the residue is a dry powder. Moisten the residue with +5 cc. of strong hydrochloric acid (sp. gr. 1.20) and evaporate again +to dryness. Heat the residue for at least one hour at a temperature +of 110°C. (Note 1). Again moisten the residue with concentrated +hydrochloric acid, warm gently, making sure that the acid comes into +contact with the whole of the residue, dilute to about 200 cc. and +bring to boiling. Filter off the silica without much delay (Note 2), +and wash five times with warm dilute hydrochloric acid (one part +dilute acid (1.12 sp. gr.) to three parts of water). Allow the filter +to drain for a few moments, then place a clean beaker below the funnel +and wash with water until free from chlorides, discarding these +washings. Evaporate the original filtrate to dryness, dehydrate at +110°C. for one hour (Note 3), and proceed as before, using a second +filter to collect the silica after the second dehydration. Wash this +filter with warm, dilute hydrochloric acid (Note 4), and finally with +hot water until free from chlorides. + +[Note 1: The silicic acid must be freed from its combination with +a base (sodium, in this instance) before it can be dehydrated. +The excess of hydrochloric acid accomplishes this liberation. By +disintegrating the fused mass with a considerable volume of dilute +acid the silicic acid is at first held in solution to a large extent. +Immediate treatment of the fused mass with strong acid is likely +to cause a semi-gelatinous silicic acid to separate at once and to +inclose alkali salts or alumina. + +A flocculent residue will often remain after the decomposition of the +fused mass is effected. This is usually partially dehydrated silicic +acid and does not require further treatment at this point. The +progress of the dehydration is indicated by the behavior of the +solution, which as evaporation proceeds usually gelatinizes. On this +account it is necessary to allow the solution to evaporate on a steam +bath, or to stir it vigorously, to avoid loss by spattering.] + +[Note 2: To obtain an approximately pure silica, the residue after +evaporation must be thoroughly extracted by warming with hydrochloric +acid, and the solution freely diluted to prevent, as far as possible, +the inclosure of the residual salts in the particles of silica. The +filtration should take place without delay, as the dehydrated silica +slowly dissolves in hydrochloric acid on standing.] + +[Note 3: It has been shown by Hillebrand that silicic acid cannot be +completely dehydrated by a single evaporation and heating, nor by +several such treatments, unless an intermediate filtration of the +silica occurs. If, however, the silica is removed and the filtrates +are again evaporated and the residue heated, the amount of silica +remaining in solution is usually negligible, although several +evaporations and filtrations are required with some silicates to +insure absolute accuracy. + +It is probable that temperatures above 100°C. are not absolutely +necessary to dehydrate the silica; but it is recommended, as tending +to leave the silica in a better condition for filtration than when +the lower temperature of the water bath is used. This, and many other +points in the analysis of silicates, are fully discussed by Dr. +Hillebrand in the admirable monograph on "The Analysis of Silicate and +Carbonate Rocks," Bulletin No. 700 of the United States Geological +Survey. + +The double evaporation and filtration spoken of above are essential +because of the relatively large amount of alkali salts (sodium +chloride) present after evaporation. For the highest accuracy in the +determination of silica, or of iron and alumina, it is also necessary +to examine for silica the precipitate produced in the filtrate by +ammonium hydroxide by fusing it with acid potassium sulphate and +solution of the fused mass in water. The insoluble silica is filtered, +washed, and weighed, and the weight added to the weight of silica +previously obtained.] + +[Note 4: Aluminium and iron are likely to be thrown down as basic +salts from hot, very dilute solutions of their chlorides, as a result +of hydrolysis. If the silica were washed only with hot water, the +solution of these chlorides remaining in the filter after the passage +of the original filtrate would gradually become so dilute as to throw +down basic salts within the pores of the filter, which would remain +with the silica. To avoid this, an acid wash-water is used until the +aluminium and iron are practically removed. The acid is then removed +by water.] + + +IGNITION AND TESTING OF SILICA + +PROCEDURE.--Transfer the two washed filters belonging to each +determination to a platinum crucible, which need not be previously +weighed, and burn off the filter (Note 1). Ignite for thirty minutes +over the blast lamp with the cover on the crucible, and then for +periods of ten minutes, until the weight is constant. + +When a constant weight has been obtained, pour into the crucible about +3 cc. of water, and then 3 cc. of hydrofluoric acid. !This must be +done in a hood with a good draft and great care must be taken not to +come into contact with the acid or to inhale its fumes (Note 2!). + +If the precipitate has dissolved in this quantity of acid, add two +drops of concentrated sulphuric acid, and heat very slowly (always +under the hood) until all the liquid has evaporated, finally igniting +to redness. Cool in a desiccator, and weigh the crucible and residue. +Deduct this weight from the previous weight of crucible and impure +silica, and from the difference calculate the percentage of silica in +the sample (Note 3). + +[Note 1: The silica undergoes no change during the ignition beyond the +removal of all traces of water; but Hillebrand (!loc. cit.!) has shown +that the silica holds moisture so tenaciously that prolonged ignition +over the blast lamp is necessary to remove it entirely. This finely +divided, ignited silica tends to absorb moisture, and should be +weighed quickly.] + +[Note 2: Notwithstanding all precautions, the ignited precipitate of +silica is rarely wholly pure. It is tested by volatilisation of the +silica as silicon fluoride after solution in hydrofluoric acid, and, +if the analysis has been properly conducted, the residue, after +treatment with the acids and ignition, should not exceed 1 mg. + +The acid produces ulceration if brought into contact with the skin, +and its fumes are excessively harmful if inhaled.] + +[Note 3: The impurities are probably weighed with the original +precipitate in the form of oxides. The addition of the sulphuric +acid displaces the hydrofluoric acid, and it may be assumed that the +resulting sulphates (usually of iron or aluminium) are converted to +oxides by the final ignition. + +It is obvious that unless the sulphuric and hydrofluoric acids used +are known to leave no residue on evaporation, a quantity equal to that +employed in the analysis must be evaporated and a correction applied +for any residue found.] + +[Note 4: If the silicate to be analyzed is shown by a previous +qualitative examination to be completely decomposable, it may be +directly treated with hydrochloric acid, the solution evaporated to +dryness, and the silica dehydrated and further treated as described in +the case of the feldspar after fusion. + +A silicate which gelatinizes on treatment with acids should be mixed +first with a little water, and the strong acid added in small portions +with stirring, otherwise the gelatinous silicic acid incloses +particles of the original silicate and prevents decomposition. The +water, by separating the particles and slightly lessening the rapidity +of action, prevents this difficulty. This procedure is one which +applies in general to the solution of fine mineral powders in acids. + +If a small residue remains undecomposed by the treatment of the +silicate with acid, this may be filtered, washed, ignited and fused +with sodium carbonate and a solution of the fused mass added to the +original acid solution. This double procedure has an advantage, in +that it avoids adding so large a quantity of sodium salts as is +required for disintegration of the whole of the silicate by the fusion +method.] + + + + +PART IV + +STOICHIOMETRY + + +The problems with which the analytical chemist has to deal are not, as +a matter of actual fact, difficult either to solve or to understand. +That they appear difficult to many students is due to the fact that, +instead of understanding the principles which underlie each of the +small number of types into which these problems may be grouped, each +problem is approached as an individual puzzle, unrelated to others +already solved or explained. This attitude of mind should be carefully +avoided. + +It is obvious that ability to make the calculations necessary for +the interpretation of analytical data is no less important than the +manipulative skill required to obtain them, and that a moderate time +spent in the careful study of the solutions of the typical problems +which follow may save much later embarrassment. + +1. It is often necessary to calculate what is known as a "chemical +factor," or its equivalent logarithmic value called a "log factor," +for the conversion of the weight of a given chemical substance into an +equivalent weight of another substance. This is, in reality, a very +simple problem in proportion, making use of the atomic or molecular +weights of the substances in question which are chemically equivalent +to each other. One of the simplest cases of this sort is the +following: What is the factor for the conversion of a given weight of +barium sulphate (BaSO_{4}) into an equivalent weight of sulphur (S)? +The molecular weight of BaSO_{4} is 233.5. There is one atom of S in +the molecule and the atomic weight of S is 32.1. The chemical factor +is, therefore, 32.1/233.5, or 0.1375 and the weight of S corresponding +to a given weight of BaSO_{4} is found by multiplying the weight of +BaSO_{4} by this factor. If the problem takes the form, "What is +the factor for the conversion of a given weight of ferric oxide +(Fe_{2}O_{3}) into ferrous oxide (FeO), or of a given weight of +mangano-manganic oxide (Mn_{3}O_{4}) into manganese (Mn)?" the +principle involved is the same, but it must then be noted that, in the +first instance, each molecule of Fe_{2}O_{3} will be equivalent to two +molecules of FeO, and in the second instance that each molecule of +Mn_{3}O_{4} is equivalent to three atoms of Mn. The respective factors +then become + +(2FeO/Fe_{2}O_{3}) or (143.6/159.6) and (3Mn/Mn_{3}O_{4}) or +(164.7/228.7). + +It is obvious that the arithmetical processes involved in this type +of problem are extremely simple. It is only necessary to observe +carefully the chemical equivalents. It is plainly incorrect to express +the ratio of ferrous to ferric oxide as (FeO/Fe_{2}O_{3}), since each +molecule of the ferric oxide will yield two molecules of the ferrous +oxide. Mistakes of this sort are easily made and constitute one of the +most frequent sources of error. + +2. A type of problem which is slightly more complicated in appearance, +but exactly comparable in principle, is the following: "What is the +factor for the conversion of a given weight of ferrous sulphate +(FeSO_{4}), used as a reducing agent against potassium permanganate, +into the equivalent weight of sodium oxalate (Na_{2}C_{2}O_{4})?" To +determine the chemical equivalents in such an instance it is necessary +to inspect the chemical reactions involved. These are: + +10FeSO_{4} + 2KMnO_{4} + 8H_{2}SO_{4} --> 5Fe_{2}(SO_{4})_{3} + +K_{2}SO_{4} + 2MnSO_{4} + 8H_{2}O, + +5Na_{2}C_{2}O_{4} + 2KMnO_{4} + 8H_{2}SO_{4} --> 5Na_{2}SO_{4} + +10CO_{2} + K_{2}SO_{4} + 2MnSO_{4} + 8H_{2}O. + +It is evident that 10FeSO_{4} in the one case, and 5Na_{2}C_{2}O_{4} +in the other, each react with 2KMnO_{4}. These molecular +quantities are therefore equivalent, and the factor becomes +(10FeSO_{4}/5Na_{2}C_{2}O_{4}) or (2FeSO_{4}/Na_{2}C_{2}O_{4}) or +(303.8/134). + +Again, let it be assumed that it is desired to determine the +factor required for the conversion of a given weight of potassium +permanganate (KMnO_{4}) into an equivalent weight of potassium +bichromate (K_{2}Cr_{2}O_{7}), each acting as an oxidizing agent +against ferrous sulphate. The reactions involved are: + +10FeSO_{4} + 2KMnO_{4} + 8H_{2}SO_{4} --> 5Fe_{2}(SO_{4})_{3} + +K_{2}SO_{4} + 2MnSO_{4} + 8H_{2}O, + +6FeSO_{4} + K_{2}Cr_{2}O_{7} + 7H_{2}SO_{4} --> 3Fe_{2}(SO_{3})_{3} + +K_{2}SO_{4} + Cr_{2}(SO_{4})_{3} + 7H_{2}O. + +An inspection of these equations shows that 2KMO_{4} react with +10FeSO_{4}, while K_{2}Cr_{2}O_{7} reacts with 6FeSO_{4}. These are +not equivalent, but if the first equation is multiplied by 3 and the +second by 5 the number of molecules of FeSO_{4} is then the same in +both, and the number of molecules of KMnO_{4} and K_{2}Cr_{2}O_{7} +reacting with these 30 molecules become 6 and 5 respectively. These +are obviously chemically equivalent and the desired factor is +expressed by the fraction (6KMnO_{4}/5K_{2}Cr_{2}O_{7}) or +(948.0/1471.0). + +3. It is sometimes necessary to calculate the value of solutions +according to the principles just explained, when several successive +reactions are involved. Such problems may be solved by a series of +proportions, but it is usually possible to eliminate the common +factors and solve but a single one. For example, the amount of MnO_{2} +in a sample of the mineral pyrolusite may be determined by dissolving +the mineral in hydrochloric acid, absorbing the evolved chlorine in a +solution of potassium iodide, and measuring the liberated iodine +by titration with a standard solution of sodium thiosulphate. The +reactions involved are: + +MnO_{2} + 4HCl --> MnCl_{2} + 2H_{2}O + Cl_{2} +Cl_{2} + 2KI --> I_{2} + 2KCl +I_{2} + 2Na_{2}S_{2}O_{3} --> 2NaI + Na_{2}S_{4}O_{6} + +Assuming that the weight of thiosulphate corresponding to the +volume of sodium thiosulphate solution used is known, what is the +corresponding weight of manganese dioxide? From the reactions given +above, the following proportions may be stated: + +2Na_{2}S_{2}O_{3}:I_{2} = 316.4:253.9, + +I_{2}:Cl_{2} = 253.9:71, + +Cl_{2}:MnO_{2} = 71:86.9. + +After canceling the common factors, there remains +2Na_{2}S_{2}O_{3}:MnO_{2} = 316.4:86.9, and the factor for the +conversion of thiosulphate into an equivalent of manganese dioxide is +86.9/316.4. + +4. To calculate the volume of a reagent required for a specific +operation, it is necessary to know the exact reaction which is to be +brought about, and, as with the calculation of factors, to keep in +mind the molecular relations between the reagent and the substance +reacted upon. For example, to estimate the weight of barium chloride +necessary to precipitate the sulphur from 0.1 gram of pure pyrite +(FeS_{2}), the proportion should read + + 488. 120.0 + 2(BaCl_{2}.2H_{2}O):FeS_{2} = x:0.1, + +where !x! represents the weight of the chloride required. Each of the +two atoms of sulphur will form upon oxidation a molecule of sulphuric +acid or a sulphate, which, in turn, will require a molecule of the +barium chloride for precipitation. To determine the quantity of the +barium chloride required, it is necessary to include in its molecular +weight the water of crystallization, since this is inseparable from +the chloride when it is weighed. This applies equally to other similar +instances. + +If the strength of an acid is expressed in percentage by weight, due +regard must be paid to its specific gravity. For example, hydrochloric +acid (sp. gr. 1.12) contains 23.8 per cent HCl !by weight!; that is, +0.2666 gram HCl in each cubic centimeter. + +5. It is sometimes desirable to avoid the manipulation required for +the separation of the constituents of a mixture of substances by +making what is called an "indirect analysis." For example, in the +analysis of silicate rocks, the sodium and potassium present may be +obtained in the form of their chlorides and weighed together. If the +weight of such a mixture is known, and also the percentage of chlorine +present, it is possible to calculate the amount of each chloride in +the mixture. Let it be assumed that the weight of the mixed chlorides +is 0.15 gram, and that it contains 53 per cent of chlorine. + +The simplest solution of such a problem is reached through algebraic +methods. The weight of chlorine is evidently 0.15 x 0.53, or 0.0795 +gram. Let x represent the weight of sodium chloride present and y +that of potassium chloride. The molecular weight of NaCl is 58.5 and +that of KCl is 74.6. The atomic weight of chlorine is 35.5. Then + +x + y = 0.15 +(35.5/58.5)x + (35.5/74.6)y = 0.00795 + +Solving these equations for x shows the weight of NaCl to be 0.0625 +gram. The weight of KCl is found by subtracting this from 0.15. + +The above is one of the most common types of indirect analyses. Others +are more complex but they can be reduced to algebraic expressions and +solved by their aid. It should, however, be noted that the results +obtained by these indirect methods cannot be depended upon for high +accuracy, since slight errors in the determination of the common +constituent, as chlorine in the above mixture, will cause considerable +variations in the values found for the components. They should not be +employed when direct methods are applicable, if accuracy is essential. + + + + +PROBLEMS + + +(The reactions necessary for the solution of these problems are either +stated with the problem or may be found in the earlier text. In the +calculations from which the answers are derived, the atomic weights +given on page 195 have been employed, using, however, only the first +decimal but increasing this by 1 when the second decimal is 5 or +above. Thus, 39.1 has been taken as the atomic weight of potassium, +32.1 for sulphur, etc. This has been done merely to secure uniformity +of treatment, and the student should remember that it is always well +to take into account the degree of accuracy desired in a particular +instance in determining the number of decimal places to retain. +Four-place logarithms were employed in the calculations. Where four +figures are given in the answer, the last figure may vary by one or +(rarely) by two units, according to the method by which the problem is +solved.) + + +VOLUMETRIC ANALYSIS + +1. How many grams of pure potassium hydroxide are required for exactly +1 liter of normal alkali solution? + +!Answer!: 56.1 grams. + +2. Calculate the equivalent in grams (a) of sulphuric acid as an acid; +(b) of hydrochloric acid as an acid; (c) of oxalic acid as an acid; +(d) of nitric acid as an acid. + +!Answers!: (a) 49.05; (b) 36.5; (c) 63; (d) 63. + +3. Calculate the equivalent in grams of (a) potassium hydroxide; +(b) of sodium carbonate; (c) of barium hydroxide; (d) of sodium +bicarbonate when titrated with an acid. + +!Answers!: (a) 56.1; (b) 53.8; (c) 85.7; (d) 84. + +4. What is the equivalent in grams of Na_{2}HPO_{4} (a) as a +phosphate; (b) as a sodium salt? + +!Answers!: (a) 47.33; (b) 71.0. + +5. A sample of aqueous hydrochloric acid has a specific gravity +of 1.12 and contains 23.81 per cent hydrochloric acid by weight. +Calculate the grams and the milliequivalents of hydrochloric acid +(HCl) in each cubic centimeter of the aqueous acid. + +!Answers!: 0.2667 gram; 7.307 milliequivalents. + +6. How many cubic centimeters of hydrochloric acid (sp. gr. 1.20 +containing 39.80 per cent HCl by weight) are required to furnish 36.45 +grams of the gaseous compound? + +!Answer!: 76.33 cc. + +7. A given solution contains 0.1063 equivalents of hydrochloric acid +in 976 cc. What is its normal value? + +!Answer!: 0.1089 N. + +8. In standardizing a hydrochloric acid solution it is found that +47.26 cc. of hydrochloric acid are exactly equivalent to 1.216 grams +of pure sodium carbonate, using methyl orange as an indicator. What is +the normal value of the hydrochloric acid? + +!Answer!: 0.4855 N. + +9. Convert 42.75 cc. of 0.5162 normal hydrochloric acid to the +equivalent volume of normal hydrochloric acid. + +!Answer!: 22.07 cc. + +10. A solution containing 25.27 cc. of 0.1065 normal hydrochloric acid +is added to one containing 92.21 cc. of 0.5431 normal sulphuric acid +and 50 cc. of exactly normal potassium hydroxide added from a pipette. +Is the solution acid or alkaline? How many cubic centimeters of +0.1 normal acid or alkali must be added to exactly neutralize the +solution? + +!Answer!: 27.6 cc. alkali (solution is acid). + +11. By experiment the normal value of a sulphuric acid solution is +found to be 0.5172. Of this acid 39.65 cc. are exactly equivalent to +21.74 cc. of a standard alkali solution. What is the normal value of +the alkali? + +!Answer!: 0.9432 N. + +12. A solution of sulphuric acid is standardized against a sample of +calcium carbonate which has been previously accurately analyzed and +found to contain 92.44% CaCO_{3} and no other basic material. The +sample weighing 0.7423 gram was titrated by adding an excess of acid +(42.42 cc.) and titrating the excess with sodium hydroxide solution +(11.22 cc.). 1 cc. of acid is equivalent to 0.9976 cc. of sodium +hydroxide. Calculate the normal value of each. + +!Answers!: Acid 0.4398 N; alkali 0.4409 N. + +13. Given five 10 cc. portions of 0.1 normal hydrochloric acid, (a) +how many grams of silver chloride will be precipitated by a portion +when an excess of silver nitrate is added? (b) how many grams of pure +anhydrous sodium carbonate (Na_{2}CO_{3}) will be neutralized by a +portion of it? (c) how many grams of silver will there be in the +silver chloride formed when an excess of silver nitrate is added to a +portion? (d) how many grams of iron will be dissolved to FeCl_{2} by a +portion of it? (e) how many grams of magnesium chloride will be formed +and how many grams of carbon dioxide liberated when an excess of +magnesium carbonate is treated with a portion of the acid? + +!Answers!: (a) 0.1434; (b) 0.053; (c) 0.1079; (d) 0.0279; (e) 0.04765, +and 0.022. + +14. If 30.00 grams of potassium tetroxalate +(KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O) are dissolved and the solution +diluted to exactly 1 liter, and 40 cc. are neutralized with 20 cc. +of a potassium carbonate solution, what is the normal value of the +carbonate solution? + +!Answer!: 0.7084 N. + +15. How many cubic centimeters of 0.3 normal sulphuric acid will be +required to neutralize (a) 30 cc. of 0.5 normal potassium hydroxide; +(b) to neutralize 30 cc. of 0.5 normal barium hydroxide; (c) to +neutralize 20 cc. of a solution containing 10.02 grams of potassium +bicarbonate per 100 cc.; (d) to give a precipitate of barium sulphate +weighing 0.4320 gram? + +!Answers!: (a) 50 cc.; (b) 50 cc.; (c) 66.73 cc.; (d) 12.33 cc. + +16. It is desired to dilute a solution of sulphuric acid of which 1 +cc. is equivalent to 0.1027 gram of pure sodium carbonate to make it +exactly 1.250 normal. 700 cc. of the solution are available. To what +volume must it be diluted? + +!Answer!: 1084 cc. + +17. Given the following data: 1 cc. of NaOH = 1.117 cc. HCl. The HCl +is 0.4876 N. How much water must be added to 100 cc. of the alkali to +make it exactly 0.5 N.? + +!Answer!: 9.0 cc. + +18. What is the normal value of a sulphuric acid solution which has a +specific gravity of 1.839 and contains 95% H_{2}SO_{4} by weight? + +!Answer!: 35.61 N. + +19. A sample of Rochelle Salt (KNaC_{4}H_{4}O_{6}.4H_{2}O), after +ignition in platinum to convert it to the double carbonate, is +titrated with sulphuric acid, using methyl orange as an indicator. +From the following data calculate the percentage purity of the sample: + +Wt. sample = 0.9500 gram +H_{2}SO_{4} used = 43.65 cc. +NaOH used = 1.72 cc. +1 cc. H_{2}SO_{4} = 1.064 cc. NaOH +Normal value NaOH = 0.1321 N. + +!Answer!: 87.72 cc. + +20. One gram of a mixture of 50% sodium carbonate and 50% potassium +carbonate is dissolved in water, and 17.36 cc. of 1.075 N acid is +added. Is the resulting solution acid or alkaline? How many cubic +centimeters of 1.075 N acid or alkali will have to be added to make +the solution exactly neutral? + +!Answers!: Acid; 1.86 cc. alkali. + +21. In preparing an alkaline solution for use in volumetric work, an +analyst, because of shortage of chemicals, mixed exactly 46.32 grams +of pure KOH and 27.64 grams of pure NaOH, and after dissolving in +water, diluted the solution to exactly one liter. How many cubic +centimeters of 1.022 N hydrochloric acid are necessary to neutralize +50 cc. of the basic solution? + +!Answer!: 74.18 cc. + +22. One gram of crude ammonium salt is treated with strong potassium +hydroxide solution. The ammonia liberated is distilled and collected +in 50 cc. of 0.5 N acid and the excess titrated with 1.55 cc. of 0.5 N +sodium hydroxide. Calculate the percentage of NH_{3} in the sample. + +!Answer!: 41.17%. + + +23. In titrating solutions of alkali carbonates in the presence of +phenolphthalein, the color change takes place when the carbonate has +been converted to bicarbonate. In the presence of methyl orange, the +color change takes place only when the carbonate has been completely +neutralized. From the following data, calculate the percentages of +Na_{2}CO_{3} and NaOH in an impure mixture. Weight of sample, 1.500 +grams; HCl (0.5 N) required for phenolphthalein end-point, 28.85 cc.; +HCl (0.5 N) required to complete the titration after adding methyl +orange, 23.85 cc. + +!Answers!: 6.67% NaOH; 84.28% Na_{2}CO_{3}. + +24. A sample of sodium carbonate containing sodium hydroxide weighs +1.179 grams. It is titrated with 0.30 N hydrochloric acid, using +phenolphthalein in cold solution as an indicator and becomes colorless +after the addition of 48.16 cc. Methyl orange is added and 24.08 cc. +are needed for complete neutralization. What is the percentage of NaOH +and Na_{2}CO_{3}? + +!Answers!: 24.50% NaOH; 64.92% Na_{2}CO_{3}. + +25. From the following data, calculate the percentages of Na_{2}CO_{3} +and NaHCO_{3} in an impure mixture. Weight of sample 1.000 gram; +volume of 0.25 N hydrochloric acid required for phenolphthalein +end-point, 26.40 cc.; after adding an excess of acid and boiling out +the carbon dioxide, the total volume of 0.25 N hydrochloric acid +required for phenolphthalein end-point, 67.10 cc. + +!Answer!: 69.95% Na_{2}CO_{3}; 30.02% NaHCO_{3}. + +26. In the analysis of a one-gram sample of soda ash, what must be the +normality of the acid in order that the number of cubic centimeters of +acid used shall represent the percentage of carbon dioxide present? + +!Answer!: 0.4544 gram. + +27. What weight of pearl ash must be taken for analysis in order that +the number of cubic centimeters of 0.5 N acid used may be equal to one +third the percentage of K_{2}CO_{3}? + +!Answer!: 1.152 grams. + +28. What weight of cream of tartar must have been taken for analysis +in order to have obtained 97.60% KHC_{4}H_{4}O_{6} in an analysis +involving the following data: NaOH used = 30.06 cc.; H_{2}SO_{4} +solution used = 0.50 cc.; 1 cc. H_{2}SO_{4} sol. = 0.0255 gram +CaCO_{3}; 1 cc. H_{2}SO_{4} sol. = 1.02 cc. NaOH sol.? + +!Answer!: 2.846 grams. + +29. Calculate the percentage of potassium oxide in an impure sample of +potassium carbonate from the following data: Weight of sample = 1.00 +gram; HCl sol. used = 55.90 cc.; NaOH sol. used = 0.42 cc.; 1 cc. NaOH +sol. = 0.008473 gram of KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O; 2 cc. +HCl sol. = 5 cc. NaOH sol. + +!Answer!: 65.68%. + +30. Calculate the percentage purity of a sample of calcite +(CaCO_{3}) from the following data: (Standardization); Weight of +H_{2}C_{2}O_{4}.2H_{2}O = 0.2460 gram; NaOH solution used = 41.03 +cc.; HCl solution used = 0.63; 1 cc. NaOH solution = 1.190 cc. HCl +solution. (Analysis); Weight of sample 0.1200 gram; HCl used = 36.38 +cc.; NaOH used = 6.20 cc. + +!Answer!: 97.97%. + +31. It is desired to dilute a solution of hydrochloric acid to exactly +0.05 N. The following data are given: 44.97 cc. of the hydrochloric +acid are equivalent to 43.76 cc. of the NaOH solution. The NaOH +is standardized against a pure potassium tetroxalate +(KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O) weighing 0.2162 gram and +requires 49.14 cc. How many cc. of water must be added to 1000 cc. of +the aqueous hydrochloric acid? + +!Answer!: 11 cc. + +32. How many cubic centimeters of 3 N phosphoric acid must be added to +300 cc. of 0.4 N phosphoric acid in order that the resulting solution +may be 0.6 N? + +!Answer!: 25 cc. + +33. To oxidize the iron in 1 gram of +FeSO_{4}(NH_{4})_{2}SO_{4}.6H_{2}O (mol. wgt. 392) requires 3 cc. of +a given solution of HNO_{3}. What is the normality of the nitric +acid when used as an acid? 6FeSO_{4} + 2HNO_{3} + 2H_{2}SO_{4} = +3Fe_{2}(SO_{4})_{3} + 2NO + 4H_{2}O. + +!Answer!: 0.2835 N. + +34. The same volume of carbon dioxide at the same temperature and the +same pressure is liberated from a 1 gram sample of dolomite, by adding +an excess of hydrochloric acid, as can be liberated by the addition of +35 cc. of 0.5 N hydrochloric acid to an excess of any pure or impure +carbonate. Calculate the percentage of CO_{2} in the dolomite. + +!Answer!: 38.5%. + +35. How many cubic centimeters of sulphuric acid (sp. gr. 1.84, +containing 96% H_{2}SO_{4} by weight) will be required to displace the +chloride in the calcium chloride formed by the action of 100 cc. of +0.1072 N hydrochloric acid on an excess of calcium carbonate, and how +many grams of CaSO_{4} will be formed? + +!Answers!: 0.298 cc.; 0.7300 gram. + +36. Potassium hydroxide which has been exposed to the air is found on +analysis to contain 7.62% water, 2.38% K_{2}CO_{3}. and 90% KOH. What +weight of residue will be obtained if one gram of this sample is added +to 46 cc. of normal hydrochloric acid and the resulting solution, +after exact neutralization with 1.070 N potassium hydroxide solution, +is evaporated to dryness? + +!Answer!: 3.47 grams. + +37. A chemist received four different solutions, with the statement +that they contained either pure NaOH; pure Na_{2}CO_{3}; pure +NaHCO_{3}, or mixtures of these substances. From the following data +identify them: + +Sample I. On adding phenolphthalein to a solution of the substance, it +gave no color to the solution. + +Sample II. On titrating with standard acid, it required 15.26 cc. for +a change in color, using phenolphthalein, and 17.90 cc. additional, +using methyl orange as an indicator. + +Sample III. The sample was titrated with hydrochloric acid until the +pink of phenolphthalein disappeared, and on the addition of methyl +orange the solution was colored pink. + +Sample IV. On titrating with hydrochloric acid, using phenolphthalein, +15.00 cc. were required. A new sample of the same weight required +exactly 30 cc. of the same acid for neutralization, using methyl +orange. + +!Answers!: (a) NaHCO_{3}; (b) NaHCO_{3}+Na_{2}CO_{3}; (c)NaOH; (d) +Na_{2}CO_{3}. + +38. In the analysis of a sample of KHC_{4}H_{4}O_{6} the following +data are obtained: Weight sample = 0.4732 gram. NaOH solution used = +24.97 cc. 3.00 cc. NaOH = 1 cc. of H_{3}PO_{4} solution of which 1 +cc. will precipitate 0.01227 gram of magnesium as MgNH_{4}PO_{4}. +Calculate the percentage of KHC_{4}H_{4}O_{6}. + +!Answer!: 88.67%. + +39. A one-gram sample of sodium hydroxide which has been exposed to +the air for some time, is dissolved in water and diluted to exactly +500 cc. One hundred cubic centimeters of the solution, when titrated +with 0.1062 N hydrochloric acid, using methyl orange as an indicator, +requires 38.60 cc. for complete neutralization. Barium chloride in +excess is added to a second portion of 100 cc. of the solution, which +is diluted to exactly 250 cc., allowed to stand and filtered. Two +hundred cubic centimeters of this filtrate require 29.62 cc. of 0.1062 +N hydrochloric acid for neutralization, using phenolphthalein as an +indicator. Calculate percentage of NaOH, Na_{2}CO_{3}, and H_{2}O. + +!Answers!: 78.63% NaOH; 4.45% Na_{2}CO_{3}; 16.92% H_{2}O. + +40. A sodium hydroxide solution (made from solid NaOH which has been +exposed to the air) was titrated against a standard acid using methyl +orange as an indicator, and was found to be exactly 0.1 N. This +solution was used in the analysis of a material sold at 2 cents per +pound per cent of an acid constituent A, and always mixed so that +it was supposed to contain 15% of A, on the basis of the analyst's +report. Owing to the carelessness of the analyst's assistant, the +sodium hydroxide solution was used with phenolphthalein as an +indicator in cold solution in making the analyses. The concern +manufacturing this material sells 600 tons per year, and when the +mistake was discovered it was estimated that at the end of a year +the error in the use of indicators would either cost them or their +customers $6000. Who would lose and why? Assuming the impure NaOH used +originally in making the titrating solution consisted of NaOH and +Na_{2}CO_{3} only, what per cent of each was present? + +!Answers!: Customer lost; 3.94% Na_{2}CO_{3}; 96.06% NaOH. + +41. In the standardization of a K_{2}Cr_{2}O_{7} solution against iron +wire, 99.85% pure, 42.42 cc. of the solution were added. The weight of +the wire used was 0.22 gram. 3.27 cc. of a ferrous sulphate solution +having a normal value as a reducing agent of 0.1011 were added +to complete the titration. Calculate the normal value of the +K_{2}Cr_{2}O_{7}. + +!Answer!: 0.1006 N. + +42. What weight of iron ore containing 56.2% Fe should be taken to +standardize an approximately 0.1 N oxidizing solution, if not more +than 47 cc. are to be used? + +!Answer!: 0.4667 gram. + +43. One tenth gram of iron wire, 99.78% pure, is dissolved in +hydrochloric acid and the iron oxidized completely with bromine water. +How many grams of stannous chloride are there in a liter of solution +if it requires 9.47 cc. to just reduce the iron in the above? What +is the normal value of the stannous chloride solution as a reducing +agent? + +!Answer!: 17.92 grams; 0.1888 N. + +44. One gram of an oxide of iron is fused with potassium acid sulphate +and the fusion dissolved in acid. The iron is reduced with stannous +chloride, mercuric chloride is added, and the iron titrated with a +normal K_{2}Cr_{2}O_{7} solution. 12.94 cc. were used. What is the +formula of the oxide, FeO, Fe_{2}O_{3}, or Fe_{3}O_{4}? + +!Answer!: Fe_{3}O_{4}. + +45. If an element has 98 for its atomic weight, and after reduction +with stannous chloride could be oxidized by bichromate to a state +corresponding to an XO_{4}^{-} anion, compute the oxide, or valence, +corresponding to the reduced state from the following data: 0.3266 +gram of the pure element, after being dissolved, was reduced with +stannous chloride and oxidized by 40 cc. of K_{2}Cr_{2}O_{7}, of which +one cc. = 0.1960 gram of FeSO_{4}(NH_{4})_{2}SO_{4}.6H_{2}O. + +!Answer!: Monovalent. + +46. Determine the percentage of iron in a sample of limonite from the +following data: Sample = 0.5000 gram. KMnO_{4} used = 50 cc. 1 cc. +KMnO_{4} = 0.005317 gram Fe. FeSO_{4} used = 6 cc. 1 cc. FeSO_{4} = +0.009200 gram FeO. + +!Answer!: 44.60%. + +47. If 1 gram of a silicate yields 0.5000 gram of Fe_{2}O_{3} and +Al_{2}O_{3} and the iron present requires 25 cc. of 0.2 N KMnO_{4}, +calculate the percentage of FeO and Al_{2}O_{3} in the sample. + +!Answer!: 35.89% FeO; 10.03% Al_{2}O_{3}. + +48. A sample of magnesia limestone has the following composition: +Silica, 3.00%; ferric oxide and alumina, 0.20%; calcium oxide, 33.10%; +magnesium oxide, 20.70%; carbon dioxide, 43.00%. In manufacturing lime +from the above the carbon dioxide is reduced to 3.00%. How many cubic +centimeters of normal KMnO_{4} will be required to determine the +calcium oxide volumetrically in a 1 gram sample of the lime? + +!Answer!: 20.08 cc. + +49. If 100 cc. of potassium bichromate solution (10 gram +K_{2}Cr_{2}O_{7} per liter), 5 cc. of 6 N sulphuric acid, and 75 cc. +of ferrous sulphate solution (80 grams FeSO_{4}.7H_{2}O per liter) are +mixed, and the resulting solution titrated with 0.2121 N KMnO_{4}, how +many cubic centimeters of the KMnO_{4} solution will be required to +oxidize the iron? + +!Answer!: 5.70 cc. + +50. If a 0.5000 gram sample of limonite containing 59.50 per cent +Fe_{2}O_{3} requires 40 cc. of KMnO_{4} to oxidize the iron, what +is the value of 1 cc. of the permanganate in terms of (a) Fe, (b) +H_{2}C_{2}O_{4}.2H_{2}O? + +!Answers!: (a) 0.005189 gram; (b) 0.005859 gram. + +51. A sample of pyrolusite weighing 0.6000 gram is treated with 0.9000 +gram of oxalic acid. The excess oxalic acid requires 23.95 cc. of +permanganate (1 cc. = 0.03038 gram FeSO_{4}.7H_{2}O). What is the +percentage of MnO_{2}, in the sample? + +!Answer!: 84.47%. + +52. A solution contains 50 grams of +KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O per liter. What is the normal +value of the solution (a) as an acid, and (b) as a reducing agent? + +!Answers!: (a) 0.5903 N; (b) 0.7872 N. + +53. In the analysis of an iron ore containing 60% Fe_{2}O_{3}, a +sample weighing 0.5000 gram is taken and the iron is reduced with +sulphurous acid. On account of failure to boil out all the excess +SO_{2}, 38.60 cubic centimeters of 0.1 N KMnO_{4} were required to +titrate the solution. What was the error, percentage error, and what +weight of sulphur dioxide was in the solution? + +!Answers!: (a) 1.60%; (b) 2.67%; (c) 0.00322 gram. + +54. From the following data, calculate the ratio of the nitric acid as +an oxidizing agent to the tetroxalate solution as a reducing agent: +1 cc. HNO_{3} = 1.246 cc. NaOH solution; 1 cc. NaOH = 1.743 cc. +KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O solution; Normal value NaOH = +0.12. + +!Answer!: 4.885. + +55. Given the following data: 25 cc. of a hydrochloric acid, when +standardized gravimetrically as silver chloride, yields a precipitate +weighing 0.5465 gram. 24.35 cc. of the hydrochloric acid are exactly +equivalent to 30.17 cc. of KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O +solution. How much water must be added to a liter of the oxalate +solution to make it exactly 0.025 N as a reducing agent? + +!Answer!: 5564 cc. + +56. Ten grams of a mixture of pure potassium tetroxalate +(KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O) and pure oxalic acid +(H_{2}C_{2}O_{4}.2H_{2}O) are dissolved in water and diluted to +exactly 1000 cc. The normal value of the oxalate solution when used as +an acid is 0.1315. Calculate the ratio of tetroxalate to oxalate used +in making up the solution and the normal value of the solution as a +reducing agent. + +!Answers!: 2:1; 0.1577 N. + +57. A student standardized a solution of NaOH and one of KMnO_{4} +against pure KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O and found the former +to be 0.07500 N as an alkali and the latter exactly 0.1 N as an +oxidizing agent. By coincidence, exactly 47.26 cc. were used in each +standardization. Find the ratio of the oxalate used in the +NaOH standardization to the oxalate used in the permanganate +standardization. + +!Answer!: 1:1. + +58. A sample of apatite weighing 0.60 gram is analyzed for its +phosphoric anhydride content. If the phosphate is precipitated as +(NH_{4})_{3}PO_{4}.12MoO_{3}, and the precipitate (after solution and +reduction of the MoO_{3} to Mo_{24}O_{37}), requires 100 cc. of normal +KMnO_{4} to oxidize it back to MoO_{3}, what is the percentage of +P_{2}O_{5}? + +!Answer!: 33.81%. + +59. In the analysis of a sample of steel weighing 1.881 grams the +phosphorus was precipitated with ammonium molybdate and the yellow +precipitate was dissolved, reduced and titrated with KMnO_{4}. If the +sample contained 0.025 per cent P and 6.01 cc. of KMnO_{4} were used, +to what oxide was the molybdenum reduced? 1 cc. KMnO_{4} = 0.007188 +gram Na_{2}C_{2}O_{4}. + +!Answer!: Mo_{4}O_{5}. + +60. What is the value of 1 cc. of an iodine solution (1 cc. equivalent +to 0.0300 gram Na_{2}S_{2}O_{3}) in terms of As_{2}O_{3}? + +!Answer!: 0.009385 gram. + +61. 48 cc. of a solution of sodium thiosulphate are required to +titrate the iodine liberated from an excess of potassium iodide +solution by 0.3000 gram of pure KIO_{3}. (KIO_{3} + 5KI + 3H_{2}SO_{4} += 3K_{2}SO_{4} + 3I_{2} + 3H_{2}O.) What is the normal strength of the +sodium thiosulphate and the value of 1 cc. of it in terms of iodine? + +!Answers!: 0.1753 N; 0.02224 gram. + +62. One thousand cubic centimeters of 0.1079 N sodium thiosulphate +solution is allowed to stand. One per cent by weight of the +thiosulphate is decomposed by the carbonic acid present in the +solution. To what volume must the solution be diluted to make it +exactly 0.1 N as a reducing agent? (Na_{2}S_{2}O_{3} + 2H_{2}CO_{3} = +H_{2}SO_{3} + 2NaHCO_{3} + S.) + +!Answer!: 1090 cc. + +63. An analyzed sample of stibnite containing 70.05% Sb is given for +analysis. A student titrates it with a solution of iodine of which 1 +cc. is equivalent to 0.004950 gram of As_{2}O_{3}. Due to an error on +his part in standardization, the student's analysis shows the sample +to contain 70.32% Sb. Calculate the true normal value of the iodine +solution, and the percentage error in the analysis. + +!Answers!: 0.1000 N; 0.39%. + +64. A sample of pyrolusite weighing 0.5000 gram is treated with an +excess of hydrochloric acid, the liberated chlorine is passed into +potassium iodide and the liberated iodine is titrated with sodium +thiosulphate solution (49.66 grams of pure Na_{2}S_{2}O_{3}.5H_{2}O +per liter). If 38.72 cc. are required, what volume of 0.25 normal +permanganate solution will be required in an indirect determination +in which a similar sample is reduced with 0.9012 gram +H_{2}C_{2}O_{4}.2H_{2}O and the excess oxalic acid titrated? + +!Answer!: 26.22 cc. + +65. In the determination of sulphur in steel by evolving the sulphur +as hydrogen sulphide, precipitating cadmium sulphide by passing the +liberated hydrogen sulphide through ammoniacal cadmium chloride +solution, and decomposing the CdS with acid in the presence of a +measured amount of standard iodine, the following data are obtained: +Sample, 5.027 grams; cc. Na_{2}S_{2}O_{3} sol. = 12.68; cc. Iodine +sol. = 15.59; 1 cc. Iodine sol. = 1.086 cc. Na_{2}S_{2}O_{3} sol.; 1 +cc. Na_{2}S_{2}O_{3}= 0.005044 gram Cu. Calculate the percentage of +sulphur. (H_{2}S + I_{2} = 2HI + S.) + +!Answer!: 0.107%. + +66. Given the following data, calculate the percentage of iron in +a sample of crude ferric chloride weighing 1.000 gram. The iodine +liberated by the reaction 2FeCl_{3}+ 2HI = 2HCl + 2FeCl_{2} + I_{2} is +reduced by the addition of 50 cc. of sodium thiosulphate solution and +the excess thiosulphate is titrated with standard iodine and requires +7.85 cc. 45 cc. I_{2} solution = 45.95 cc. Na_{2}S_{2}O_{3} solution; +45 cc. As_{2}O_{3} solution = 45.27 cc. I_{2} solution. 1 cc. arsenite +solution = 0.005160 gram As_{2}O_{3}. + +!Answer!: 23.77%. + +67. Sulphide sulphur was determined in a sample of reduced barium +sulphate by the evolution method, in which the sulphur was evolved as +hydrogen sulphide and was passed into CdCl_{2} solution, the acidified +precipitate being titrated with iodine and thiosulphate. Sample, 5.076 +grams; cc. I_{2} = 20.83; cc. Na_{2}S_{2}O_{3} = 12.37; 43.45 cc. +Na_{2}S_{2}O_{3} = 43.42 cc. I_{2}; 8.06 cc. KMnO_{4} = 44.66 cc. +Na_{2}S_{2}O_{3}; 28.87 cc. KMnO_{4} = 0.2004 gram Na_{2}C_{2}O_{4}. +Calculate the percentage of sulphide sulphur in the sample. + +!Answer!: 0.050%. + +68. What weight of pyrolusite containing 89.21% MnO_{2} will oxidize +the same amount of oxalic acid as 37.12 cc. of a permanganate +solution, of which 1 cc. will liberate 0.0175 gram of I_{2} from KI? + +!Answer!: 0.2493 gram. + +69. A sample of pyrolusite weighs 0.2400 gram and is 92.50% pure +MnO_{2}. The iodine liberated from KI by the manganese dioxide is +sufficient to react with 46.24 cc. of Na_{2}S_{2}O_{3} sol. What is +the normal value of the thiosulphate? + +!Answer!:: 0.1105 N. + +70. In the volumetric analysis of silver coin (90% Ag), using a +0.5000 gram sample, what is the least normal value that a potassium +thiocyanate solution may have and not require more than 50 cc. of +solution in the analysis? + +!Answer!: 0.08339 N. + +71. A mixture of pure lithium chloride and barium bromide weighing +0.6 gram is treated with 45.15 cubic centimeters of 0.2017 N silver +nitrate, and the excess titrated with 25 cc. of 0.1 N KSCN solution, +using ferric alum as an indicator. Calculate the percentage of bromine +in the sample. + +!Answer!: 40.11%. + +72. A mixture of the chlorides of sodium and potassium from 0.5000 +gram of a feldspar weighs 0.1500 gram, and after solution in water +requires 22.71 cc. of 0.1012 N silver nitrate for the precipitation of +the chloride ions. What are the percentages of Na_{2}O and K_{2}O in +the feldspar? + +!Answer!: 8.24% Na_{2}O; 9.14% K_{2}O. + + +GRAVIMETRIC ANALYSIS + +73. Calculate (a) the grams of silver in one gram of silver chloride; +(b) the grams of carbon dioxide liberated by the addition of an excess +of acid to one gram of calcium carbonate; (c) the grams of MgCl_{2} +necessary to precipitate 1 gram of MgNH_{4}PO_{4}. + +!Answers!: (a) 0.7526; (b) 0.4397; (c) 0.6940. + +74. Calculate the chemical factor for (a) Sn in SnO_{2}; (b) MgO +in Mg_{2}P_{2}O_{7}; (c) P_{2}O_{5} in Mg_{2}P_{2}O_{7}; (d) Fe in +Fe_{2}O_{3}; (e) SO_{4} in BaSO_{4}. + +!Answers!: (a) 0.7879; (b) 0.3620; (c) 0.6378; (d) 0.6990; (e) 0.4115. + +75. Calculate the log factor for (a) Pb in PbCrO_{4}; (b) Cr_{2}O_{3} +in PbCrO_{4}; (c) Pb in PbO_{2} and (d) CaO in CaC_{2}O_{4}. + +!Answers!: (a) 9.8069-10, (b) 9.3713-10; (c) 9.9376-10; (d) 9.6415-10. + +76. How many grams of Mn_{3}O_{4} can be obtained from 1 gram of +MnO_{2}? + +!Answer!: 0.8774 gram. + +77. If a sample of silver coin weighing 0.2500 gram gives a +precipitate of AgCl weighing 0.2991 gram, what weight of AgI could +have been obtained from the same weight of sample, and what is the +percentage of silver in the coin? + +!Answers!: 0.4898 gr.; 90.05%. + +78. How many cubic centimeters of hydrochloric acid (sp. gr. 1.13 +containing 25.75% HCl by weight) are required to exactly neutralize +25 cc. of ammonium hydroxide (sp. gr. .90 containing 28.33% NH_{3} by +weight)? + +!Answer!: 47.03 cc. + +79. How many cubic centimeters of ammonium hydroxide solution (sp. gr. +0.96 containing 9.91% NH_{3} by weight) are required to precipitate +the aluminium as aluminium hydroxide from a two-gram sample of alum +(KAl(SO_{4})_{2}.12H_{2}O)? What will be the weight of the ignited +precipitate? + +!Answers!: 2.26 cc.; 0.2154 gram. + +80. What volume of nitric acid (sp. gr. 1.05 containing 9.0% +HNO_{3} by weight) is required to oxidize the iron in one gram of +FeSO_{4}.7H_{2}O in the presence of sulphuric acid? 6FeSO_{4} + +2HNO_{3} + 3H_{2}SO_{4} = 3Fe_{2}(SO_{4})_{3} + 2NO + 4H_{2}O. + +!Answer!: 0.80 cc. + +81. If 0.7530 gram of ferric nitrate (Fe(NO_{3})_{3}.9H_{2}O) is +dissolved in water and 1.37 cc. of HCl (sp. gr. 1.11 containing 21.92% +HCl by weight) is added, how many cubic centimeters of ammonia (sp. +gr. 0.96 containing 9.91% NH_{3} by weight) are required to neutralize +the acid and precipitate the iron as ferric hydroxide? + +!Answer!: 2.63 cc. + +82. To a suspension of 0.3100 gram of Al(OH)_{3} in water are added +13.00 cc. of aqueous ammonia (sp. gr. 0.90 containing 28.4% NH_{3} by +weight). How many cubic centimeters of sulphuric acid (sp. gr. 1.18 +containing 24.7% H_{2}SO_{4} by weight) must be added to the mixture +in order to bring the aluminium into solution? + +!Answer!: 34.8 cc. + +83. How many cubic centimeters of sulphurous acid (sp. gr. 1.04 +containing 75 grams SO_{2} per liter) are required to reduce the +iron in 1 gram of ferric alum (KFe(SO_{4})_{2}.12H_{2}O)? +Fe_{2}(SO_{4})_{3} + SO_{2} + 2H_{2}O = 2FeSO_{4} + 2H_{2}SO_{4}. + +!Answer!: 0.85 cc. + +84. How many cubic centimeters of a solution of potassium bichromate +containing 26.30 grams of K_{2}Cr_{2}O_{7} per liter must be taken +in order to yield 0.6033 gram of Cr_{2}O_{3} after reduction and +precipitation of the chromium? + +K_{2}Cr_{2}O_{7} + 3SO_{2} + H_{2}SO_{4} = K_{2}SO_{4} + +Cr_{2}(SO_{4})_{3} + H_{2}O. + +!Answer!: 44.39 cc. + +85. How many cubic centimeters of ammonium hydroxide (sp. gr. 0.946 +containing 13.88% NH_{3} by weight) are required to precipitate +the iron as Fe(OH)_{3} from a sample of pure +FeSO_{4}.(NH_{4})_{2}SO_{4}.6H_{2}O, which requires 0.34 cc. of nitric +acid (sp. gr. 1.350 containing 55.79% HNO_{3} by weight) for oxidation +of the iron? (See problem No. 80 for reaction.) + +!Answer!: 4.74 cc. + +86. In the analysis of an iron ore by solution, oxidation and +precipitation of the iron as Fe(OH)_{3}, what weight of sample must be +taken for analysis so that each one hundredth of a gram of the ignited +precipitate of Fe_{2}O_{3} shall represent one tenth of one per cent +of iron? + +!Answer!: 6.99 grams. + +87. What weight in grams of impure ferrous ammonium sulphate should +be taken for analysis so that the number of centigrams of BaSO_{4} +obtained will represent five times the percentage of sulphur in the +sample? + +!Answer!: 0.6870 gram. + +88. What weight of magnetite must be taken for analysis in order that, +after precipitating and igniting all the iron to Fe_{2}O_{3}, the +percentage of Fe_{2}O_{4} in the sample may be found by multiplying +the weight in grams of the ignited precipitate by 100? + +!Answer!: 0.9665 gram. + +89. After oxidizing the arsenic in 0.5000 gram of pure As_{2}S_{3} to +arsenic acid, it is precipitated with "magnesia mixture" (MgCl_{2} + +2NH_{4}Cl). If exactly 12.6 cc. of the mixture are required, how many +grams of MgCl_{2} per liter does the solution contain? H_{3}AsO_{4} + +MgCl_{2} + 3NH_{4}OH = MgNH_{4}AsO_{4} + 2NH_{4}Cl + 3H_{2}O. + +!Answer!: 30.71 grams. + +90. A sample is prepared for student analysis by mixing pure apatite +(Ca_{3}(PO_{4})_{2}.CaCl_{2}) with an inert material. If 1 gram of +the sample gives 0.4013 gram of Mg_{2}P_{2}O_{7}, how many cubic +centimeters of ammonium oxalate solution (containing 40 grams of +(NH_{4})_{2}C_{2}O_{4}.H_{2}O per liter) would be required to +precipitate the calcium from the same weight of sample? + +!Answer!: 25.60 cc. + +91. If 0.6742 gram of a mixture of pure magnesium carbonate and pure +calcium carbonate, when treated with an excess of hydrochloric acid, +yields 0.3117 gram of carbon dioxide, calculate the percentage of +magnesium oxide and of calcium oxide in the sample. + +!Answers!: 13.22% MgO; 40.54% CaO. 92. The calcium in a sample of +dolomite weighing 0.9380 gram is precipitated as calcium oxalate and +ignited to calcium oxide. What volume of gas, measured over water +at 20°C. and 765 mm. pressure, is given off during ignition, if the +resulting oxide weighs 0.2606 gram? (G.M.V. = 22.4 liters; V.P. water +at 20°C. = 17.4 mm.) + +!Answer!: 227 cc. + +93. A limestone is found to contain 93.05% CaCO_{3}, and 5.16 % +MgCO_{3}. Calculate the weight of CaO obtainable from 3 tons of the +limestone, assuming complete conversion to oxide. What weight of +Mg_{2}P_{2}O_{7} could be obtained from a 3-gram sample of the +limestone? + +!Answers!: 1.565 tons; 0.2044 gram. + +94. A sample of dolomite is analyzed for calcium by precipitating +as the oxalate and igniting the precipitate. The ignited product is +assumed to be CaO and the analyst reports 29.50% Ca in the sample. +Owing to insufficient ignition, the product actually contained 8% of +its weight of CaCO_{3}. What is the correct percentage of calcium in +the sample, and what is the percentage error? + +!Answers!: 28.46%; 3.65% error. + +95. What weight of impure calcite (CaCO_{3}) should be taken for +analysis so that the volume in cubic centimeters of CO_{2} obtained by +treating with acid, measured dry at 18°C. and 763 mm., shall equal the +percentage of CaO in the sample? + +!Answer!: 0.2359 gram. + +96. How many cubic centimeters of HNO_{3} (sp. gr. 1.13 containing +21.0% HNO_{3} by weight) are required to dissolve 5 grams of brass, +containing 0.61% Pb, 24.39% Zn, and 75% Cu, assuming reduction of the +nitric acid to NO by each constituent? What fraction of this volume of +acid is used for oxidation? + +!Answers!: 55.06 cc.; 25%. + +97. What weight of metallic copper will be deposited from a cupric +salt solution by a current of 1.5 amperes during a period of 45 +minutes, assuming 100% current efficiency? (1 Faraday = 96,500 +coulombs.) + +!Answer!: 1.335 grams. + +98. In the electrolysis of a 0.8000 gram sample of brass, there is +obtained 0.0030 gram of PbO_{2}, and a deposit of metallic copper +exactly equal in weight to the ignited precipitate of Zn_{2}P_{2}O_{7} +subsequently obtained from the solution. What is the percentage +composition of the brass? + +!Answers!: 69.75% Cu; 29.92% Zn; 0.33% Pb. + +99. A sample of brass (68.90% Cu; 1.10% Pb and 30.00% Zn) weighing +0.9400 gram is dissolved in nitric acid. The lead is determined by +weighing as PbSO_{4}, the copper by electrolysis and the zinc by +precipitation with (NH_{4})_{2}HPO_{4} in a neutral solution. + +(a) Calculate the cubic centimeters of nitric acid (sp. gr. 1.42 +containing 69.90% HNO_{3} by weight) required to just dissolve the +brass, assuming reduction to NO. + +!Answer!: 2.48 cc. + +(b) Calculate the cubic centimeters of sulphuric acid (sp. gr. 1.84 +containing 94% H_{2}SO_{4} by weight) to displace the nitric acid. + +!Answer!: 0.83 cc. + +(c) Calculate the weight of PbSO_{4}. + +!Answer!: 0.0152 gram. + +(d) The clean electrode weighs 10.9640 grams. Calculate the weight +after the copper has been deposited. + +!Answer!: 11.6116 grams. + +(e) Calculate the grams of (NH_{4})_{2}HPO_{4} required to precipitate +the zinc as ZnNH_{4}PO_{4}. + +!Answer!: 0.5705 gram. + +(f) Calculate the weight of ignited Zn_{2}P_{2}O_{7}. + +!Answer!: 0.6573 gram. + +100. If in the analysis of a brass containing 28.00% zinc an error is +made in weighing a 2.5 gram portion by which 0.001 gram too much is +weighed out, what percentage error in the zinc determination would +result? What volume of a solution of sodium hydrogen phosphate, +containing 90 grams of Na_{2}HPO_{4}.12H_{2}O per liter, would be +required to precipitate the zinc as ZnNH_{4}PO_{4} and what weight of +precipitate would be obtained? + +!Answers!: (a) 0.04% error; (b) 39.97 cc.; (c) 1.909 grams. + +101. A sample of magnesium carbonate, contaminated with SiO_{2} as its +only impurity, weighs 0.5000 gram and loses 0.1000 gram on ignition. +What volume of disodium phosphate solution (containing 90 grams +Na_{2}HPO_{4}.12H_{2}O per liter) will be required to precipitate the +magnesium as magnesium ammonium phosphate? + +!Answer!: 9.07 cc. + +102. 2.62 cubic centimeters of nitric acid (sp. gr. 1.42 containing +69.80% HNO_{2} by weight) are required to just dissolve a sample +of brass containing 69.27% Cu; 0.05% Pb; 0.07% Fe; and 30.61% Zn. +Assuming the acid used as oxidizing agent was reduced to NO in every +case, calculate the weight of the brass and the cubic centimeters of +acid used as acid. + +!Answer!: 0.992 gram; 1.97 cc. + +103. One gram of a mixture of silver chloride and silver bromide is +found to contain 0.6635 gram of silver. What is the percentage of +bromine? + +!Answer!: 21.30%. + +104. A precipitate of silver chloride and silver bromide weighs 0.8132 +gram. On heating in a current of chlorine, the silver bromide is +converted to silver chloride, and the mixture loses 0.1450 gram +in weight. Calculate the percentage of chlorine in the original +precipitate. + +!Answer!: 6.13%. + +105. A sample of feldspar weighing 1.000 gram is fused and the silica +determined. The weight of silica is 0.6460 gram. This is fused with 4 +grams of sodium carbonate. How many grams of the carbonate actually +combined with the silica in fusion, and what was the loss in weight +due to carbon dioxide during the fusion? + +!Answers!: 1.135 grams; 0.4715 gram. + +106. A mixture of barium oxide and calcium oxide weighing 2.2120 grams +is transformed into mixed sulphates, weighing 5.023 grams. Calculate +the grams of calcium oxide and barium oxide in the mixture. + +!Answers!: 1.824 grams CaO; 0.3877 gram BaO. + + + + +APPENDIX + + +ELECTROLYTIC DISSOCIATION THEORY + +The following brief statements concerning the ionic theory and a few +of its applications are intended for reference in connection with the +explanations which are given in the Notes accompanying the various +procedures. The reader who desires a more extended discussion of the +fundamental theory and its uses is referred to such books as Talbot +and Blanchard's !Electrolytic Dissociation Theory! (Macmillan +Company), or Alexander Smith's !Introduction to General Inorganic +Chemistry! (Century Company). + +The !electrolytic dissociation theory!, as propounded by Arrhenius in +1887, assumes that acids, bases, and salts (that is, electrolytes) +in aqueous solution are dissociated to a greater or less extent into +!ions!. These ions are assumed to be electrically charged atoms or +groups of atoms, as, for example, H^{+} and Br^{-} from hydrobromic +acid, Na^{+} and OH^{-} from sodium hydroxide, 2NH_{4}^{+} and +SO_{4}^{--} from ammonium sulphate. The unit charge is that which is +dissociated with a hydrogen ion. Those upon other ions vary in sign +and number according to the chemical character and valence of the +atoms or radicals of which the ions are composed. In any solution the +aggregate of the positive charges upon the positive ions (!cations!) +must always balance the aggregate negative charges upon the negative +ions (!anions!). + +It is assumed that the Na^{+} ion, for example, differs from the +sodium atom in behavior because of the very considerable electrical +charge which it carries and which, as just stated, must, in an +electrically neutral solution, be balanced by a corresponding negative +charge on some other ion. When an electric current is passed through a +solution of an electrolyte the ions move with and convey the current, +and when the cations come into contact with the negatively charged +cathode they lose their charges, and the resulting electrically +neutral atoms (or radicals) are liberated as such, or else enter at +once into chemical reaction with the components of the solution. + +Two ions of identically the same composition but with different +electrical charges may exhibit widely different properties. For +example, the ion MnO_{4}^{-} from permanganates yields a purple-red +solution and differs in its chemical behavior from the ion +MnO_{4}^{--} from manganates, the solutions of which are green. + +The chemical changes upon which the procedures of analytical chemistry +depend are almost exclusively those in which the reacting substances +are electrolytes, and analytical chemistry is, therefore, essentially +the chemistry of the ions. The percentage dissociation of the same +electrolyte tends to increase with increasing dilution of its +solution, although not in direct proportion. The percentage +dissociation of different electrolytes in solutions of equivalent +concentrations (such, for example, as normal solutions) varies widely, +as is indicated in the following tables, in which approximate figures +are given for tenth-normal solutions at a temperature of about 18°C. + + ACIDS +========================================================================= + | + SUBSTANCE | PERCENTAGE DISSOCIATION IN + | 0.1 EQUIVALENT SOLUTION +_____________________________________________|___________________________ + | +HCl, HBr, HI, HNO_{3} | 90 + | +HClO_{3}, HClO_{4}, HMnO_{4} | 90 + | +H_{2}SO_{4} <--> H^{+} + HSO_{4}^{-} | 90 + | +H_{2}C_{2}O_{4} <--> H^{+} + HC_{2}O_{4}^{-} | 50 + | +H_{2}SO_{3} <--> H^{+} + HSO{_}3^{-} | 20 + | +H_{3}PO_{4} <--> H^{+} + H_{2}PO_{4}^{-} | 27 + | +H_{2}PO_{4}^{-} <--> H^{+} + HPO_{4}^{--} | 0.2 + | +H_{3}AsO_{4} <--> H^{+} + H_{2}AsO_{4}^{-} | 20 + | +HF | 9 + | +HC_{2}H_{3}O_{2} | 1.4 + | +H_{2}CO_{3} <--> H^{+} + HCO_{3}^{-} | 0.12 + | +H_{2}S <--> H^{+} + HS^{-} | 0.05 + | +HCN | 0.01 + | +========================================================================= + + + BASES +========================================================================= + | + SUBSTANCE | PERCENTAGE DISSOCIATION IN + | 0.1 EQUIVALENT SOLUTION +_____________________________________________|___________________________ + | +KOH, NaOH | 86 + | +Ba(OH)_{2} | 75 + | +NH_{4}OH | 1.4 + | +========================================================================= + + + SALTS +========================================================================= + | + TYPE OF SALT | PERCENTAGE DISSOCIATION IN + | 0.1 EQUIVALENT SOLUTION +_____________________________________________|___________________________ + | +R^{+}R^{-} | 86 + | +R^{++}(R^{-})_{2} | 72 + | +(R^{+})_{2}R^{--} | 72 + | +R^{++}R^{--} | 45 + | +========================================================================= + +The percentage dissociation is determined by studying the electrical +conductivity of the solutions and by other physico-chemical methods, +and the following general statements summarize the results: + +!Salts!, as a class, are largely dissociated in aqueous solution. + +!Acids! yield H^{+} ions in water solution, and the comparative +!strength!, that is, the activity, of acids is proportional to the +concentration of the H^{+} ions and is measured by the percentage +dissociation in solutions of equivalent concentration. The common +mineral acids are largely dissociated and therefore give a relatively +high concentration of H^{+} ions, and are commonly known as "strong +acids." The organic acids, on the other hand, belong generally to the +group of "weak acids." + +!Bases! yield OH^{-} ions in water solution, and the comparative +strength of the bases is measured by their relative dissociation in +solutions of equivalent concentration. Ammonium hydroxide is a weak +base, as shown in the table above, while the hydroxides of sodium and +potassium exhibit strongly basic properties. + +Ionic reactions are all, to a greater or less degree, !reversible +reactions!. A typical example of an easily reversible reaction is that +representing the changes in ionization which an electrolyte such as +acetic acid undergoes on dilution or concentration of its solutions, +!i.e.!, HC_{2}H_{3}O_{2} <--> H^{+} + C_{2}H_{3}O_{2}^{-}. As was +stated above, the ionization increases with dilution, the reaction +then proceeding from left to right, while concentration of the +solution occasions a partial reassociation of the ions, and the +reaction proceeds from right to left. To understand the principle +underlying these changes it is necessary to consider first the +conditions which prevail when a solution of acetic acid, which has +been stirred until it is of uniform concentration throughout, has come +to a constant temperature. A careful study of such solutions has shown +that there is a definite state of equilibrium between the constituents +of the solution; that is, there is a definite relation between the +undissociated acetic acid and its ions, which is characteristic for +the prevailing conditions. It is not, however, assumed that this is a +condition of static equilibrium, but rather that there is continual +dissociation and association, as represented by the opposing +reactions, the apparent condition of rest resulting from the fact that +the amount of change in one direction during a given time is exactly +equal to that in the opposite direction. A quantitative study of +the amount of undissociated acid, and of H^{+} ions and +C_{2}H_{3}O_{2}^{-} ions actually to be found in a large number of +solutions of acetic acid of varying dilution (assuming them to be in +a condition of equilibrium at a common temperature), has shown that +there is always a definite relation between these three quantities +which may be expressed thus: + +(!Conc'n H^{+} x Conc'n C_{2}H_{3}O_{2}^{-})/Conc'n HC_{2}H_{3}O_{2} = +Constant!. + +In other words, there is always a definite and constant ratio between +the product of the concentrations of the ions and the concentration of +the undissociated acid when conditions of equilibrium prevail. + +It has been found, further, that a similar statement may be made +regarding all reversible reactions, which may be expressed in general +terms thus: The rate of chemical change is proportional to the product +of the concentrations of the substances taking part in the reaction; +or, if conditions of equilibrium are considered in which, as stated, +the rate of change in opposite directions is assumed to be equal, then +the product of the concentrations of the substances entering into +the reaction stands in a constant ratio to the product of the +concentrations of the resulting substances, as given in the expression +above for the solutions of acetic acid. This principle is called the +!Law of Mass Action!. + +It should be borne in mind that the expression above for acetic acid +applies to a wide range of dilutions, provided the temperature remains +constant. If the temperature changes the value of the constant changes +somewhat, but is again uniform for different dilutions at that +temperature. The following data are given for temperatures of about +18°C.[1] + +========================================================================== + | | | | + MOLAL | FRACTION | MOLAL CONCENTRA- | MOLAL CONCENTRA- | VALUE OF +CONCENTRATION | IONIZED | TION OF H^{+} AND| TION OF UNDIS- | CONSTANT + CONSTANT | | ACETATE^{-} IONS | SOCIATED ACID | +______________|__________|__________________|__________________|__________ + | | | | + 1.0 | .004 | .004 | .996 | .0000161 + | | | | + 0.1 | .013 | .0013 | .0987 | .0000171 + | | | | + 0.01 | .0407 | .000407 | .009593 | .0000172 + | | | | +=========================================================================== + +[Footnote 1: Alexander Smith, !General Inorganic Chemistry!, p. 579.] + +The molal concentrations given in the table refer to fractions of a +gram-molecule per liter of the undissociated acid, and to fractions of +the corresponding quantities of H^{+} and C_{2}H_{3}O_{2}^{-} ions +per liter which would result from the complete dissociation of a +gram-molecule of acetic acid. The values calculated for the constant +are subject to some variation on account of experimental errors in +determining the percentage ionized in each case, but the approximate +agreement between the values found for molal and centimolal (one +hundredfold dilution) is significant. + +The figures given also illustrate the general principle, that the +!relative! ionization of an electrolyte increases with the dilution of +its solution. If we consider what happens during the (usually) brief +period of dilution of the solution from molal to 0.1 molal, for +example, it will be seen that on the addition of water the conditions +of concentration which led to equality in the rate of change, and +hence to equilibrium in the molal solution, cease to exist; and since +the dissociating tendency increases with dilution, as just stated, +it is true at the first instant after the addition of water that the +concentration of the undissociated acid is too great to be +permanent under the new conditions of dilution, and the reaction, +HC_{2}H_{3}O_{2} <--> H^{+} + C_{2}H_{3}O_{2}^{-}, will proceed from +left to right with great rapidity until the respective concentrations +adjust themselves to the new conditions. + +That which is true of this reaction is also true of all reversible +reactions, namely, that any change of conditions which occasions +an increase or a decrease in concentration of one or more of the +components causes the reaction to proceed in one direction or the +other until a new state of equilibrium is established. This principle +is constantly applied throughout the discussion of the applications +of the ionic theory in analytical chemistry, and it should be clearly +understood that whenever an existing state of equilibrium is disturbed +as a result of changes of dilution or temperature, or as a consequence +of chemical changes which bring into action any of the constituents of +the solution, thus altering their concentrations, there is always a +tendency to re-establish this equilibrium in accordance with the law. +Thus, if a base is added to the solution of acetic acid the H^{+} ions +then unite with the OH^{-} ions from the base to form undissociated +water. The concentration of the H^{+} ions is thus diminished, and +more of the acid dissociates in an attempt to restore equilbrium, +until finally practically all the acid is dissociated and neutralized. + +Similar conditions prevail when, for example, silver ions react with +chloride ions, or barium ions react with sulphate ions. In the former +case the dissociation reaction of the silver nitrate is AgNO_{3} <--> +Ag^{+} + NO_{3}^{-}, and as soon as the Ag^{+} ions unite with the +Cl^{-} ions the concentration of the former is diminished, more of the +AgNO_{3} dissociates, and this process goes on until the Ag^{+} ions +are practically all removed from the solution, if the Cl^{-} ions are +present in sufficient quantity. + +For the sake of accuracy it should be stated that the mass law cannot +be rigidly applied to solutions of those electrolytes which are +largely dissociated. While the explanation of the deviation from +quantitative exactness in these cases is not known, the law is still +of marked service in developing analytical methods along more logical +lines than was formerly practicable. It has not seemed wise to qualify +each statement made in the Notes to indicate this lack of quantitative +exactness. The student should recognize its existence, however, and +will realize its significance better as his knowledge of physical +chemistry increases. + +If we apply the mass law to the case of a substance of small +solubility, such as the compounds usually precipitated in quantitative +analysis, we derive what is known as the !solubility product!, as +follows: Taking silver chloride as an example, and remembering that it +is not absolutely insoluble in water, the equilibrium expression for +its solution is: + +(!Conc'n Ag^{+} x Conc'n Cl^{-})/Conc'n AgCl = Constant!. + +But such a solution of silver chloride which is in contact with the +solid precipitate must be saturated for the existing temperature, and +the quantity of undissociated AgCl in the solution is definite and +constant for that temperature. Since it is a constant, it may be +eliminated, and the expression becomes !Conc'n Ag^{+} x Conc'n +Cl^{-} = Constant!, and this is known as the solubility product. No +precipitation of a specific substance will occur until the product of +the concentrations of its ions in a solution exceeds the solubility +product for that substance; whenever that product is exceeded +precipitation must follow. + +It will readily be seen that if a substance which yields an ion in +common with the precipitated compound is added to such a solution as +has just been described, the concentration of that ion is +increased, and as a result the concentration of the other ion must +proportionately decrease, which can only occur through the formation +of some of the undissociated compound which must separate from the +already saturated solution. This explains why the addition of an +excess of the precipitant is often advantageous in quantitative +procedures. Such a case is discussed at length in Note 2 on page 113. + +Similarly, the ionization of a specific substance in solution tends to +diminish on the addition of another substance with a common ion, as, +for instance, the addition of hydrochloric acid to a solution +of hydrogen sulphide. Hydrogen sulphide is a weak acid, and the +concentration of the hydrogen ions in its aqueous solutions is very +small. The equilibrium in such a solution may be represented as: + +(!(Conc'n H^{+})^{2} x Conc'n S^{--})/Conc'n H_{2}S = Constant!, and a +marked increase in the concentration of the H^{+} ions, such as would +result from the addition of even a small amount of the highly ionized +hydrochloric acid, displaces the point of equilibrium and some of the +S^{--} ions unite with H^{+} ions to form undissociated H_{2}S. This +is of much importance in studying the reactions in which hydrogen +sulphide is employed, as in qualitative analysis. By a parallel course +of reasoning it will be seen that the addition of a salt of a weak +acid or base to solutions of that acid or base make it, in effect, +still weaker because they decrease its percentage ionization. + +To understand the changes which occur when solids are dissolved where +chemical action is involved, it should be remembered that no substance +is completely insoluble in water, and that those products of a +chemical change which are least dissociated will first form. Consider, +for example, the action of hydrochloric acid upon magnesium hydroxide. +The minute quantity of dissolved hydroxide dissociates thus: +Mg(OH)_{2} <--> Mg^{++} + 2OH^{-}. When the acid is introduced, +the H^{+} ions of the acid unite with the OH^{-} ions to form +undissociated water. The concentration of the OH^{-} ions is thus +diminished, more Mg(OH)_{2} dissociates, the solution is no longer +saturated with the undissociated compound, and more of the solid +dissolves. This process repeats itself with great rapidity until, if +sufficient acid is present, the solid passes completely into solution. + +Exactly the same sort of process takes place if calcium oxalate, for +example, is dissolved in hydrochloric acid. The C_{2}O_{4}^{--} ions +unite with the H^{+} ions to form undissociated oxalic acid, the acid +being less dissociated than normally in the presence of the H^{+} ions +from the hydrochloric acid (see statements regarding hydrogen sulphide +above). As the undissociated oxalic acid forms, the concentration of +the C_{2}O_{4}^{--} ions lessens and more CaC_{2}O_{4} dissolves, +as described for the Mg(OH)_{2} above. Numerous instances of the +applications of these principles are given in the Notes. + +Water itself is slightly dissociated, and although the resulting H^{+} +and OH^{-} ions are present only in minute concentrations (1 mol. of +dissociated water in 10^{7} liters), yet under some conditions they +may give rise to important consequences. The term !hydrolysis! is +applied to the changes which result from the reaction of these ions. +Any salt which is derived from a weak base or a weak acid (or both) +is subject to hydrolytic action. Potassium cyanide, for example, when +dissolved in water gives an alkaline solution because some of the +H^{+} ions from the water unite with CN^{-} ions to form (HCN), which +is a very weak acid, and is but very slightly dissociated. Potassium +hydroxide, which might form from the OH^{-} ions, is so largely +dissociated that the OH^{-} ions remain as such in the solution. The +union of the H^{+} ions with the CN^{-} ions to form the undissociated +HCN diminishes the concentration of the H^{+} ions, and more water +dissociates (H_{2}O <--> H^{+} + OH^{-}) to restore the equilibrium. +It is clear, however, that there must be a gradual accumulation of +OH^{-} ions in the solution as a result of these changes, causing the +solution to exhibit an alkaline reaction, and also that ultimately the +further dissociation of the water will be checked by the presence of +these ions, just as the dissociation of the H_{2}S was lessened by the +addition of HCl. + +An exactly opposite result follows the solution of such a salt as +Al_{2}(SO_{4})_{3} in water. In this case the acid is strong and the +base weak, and the OH^{-} ions form the little dissociated Al(OH)_{3}, +while the H^{+} ions remain as such in the solution, sulphuric acid +being extensively dissociated. The solution exhibits an acid reaction. + +Such hydrolytic processes as the above are of great importance in +analytical chemistry, especially in the understanding of the action of +indicators in volumetric analysis. (See page 32.) + +The impelling force which causes an element to pass from the atomic +to the ionic condition is termed !electrolytic solution pressure!, or +ionization tension. This force may be measured in terms of electrical +potential, and the table below shows the relative values for a number +of elements. + +In general, an element with a greater solution pressure tends to cause +the deposition of an element of less solution pressure when placed in +a solution of its salt, as, for instance, when a strip of zinc or +iron is placed in a solution of a copper salt, with the resulting +precipitation of metallic copper. + +Hydrogen is included in the table, and its position should be noted +with reference to the other common elements. For a more extended +discussion of this topic the student should refer to other treatises. + + POTENTIAL SERIES OF THE METALS + +__________________________________________________________________ + | | | + | POTENTIAL | | POTENTIAL + | IN VOLTS | | IN VOLTS +_____________________|___________|____________________|___________ + | | | +Sodium Na^{+} | +2.44 | Lead Pb^{++} | -0.13 +Calcium Ca^{++} | | Hydrogen H^{+} | -0.28 +Magnesium Mg^{++} | | Bismuth Bi^{+++}| +Aluminum A1^{+++} | +1.00 | Antimony | -0.75 +Manganese Mn^{++} | | Arsenic | +Zinc Zn^{++} | +0.49 | Copper Cu^{++} | -0.61 +Cadmium Cd^{++} | +0.14 | Mercury Hg^{+} | -1.03 +Iron Fe^{++} | +0.063 | Silver Ag^{+} | -1.05 +Cobalt Co^{++} | -0.045 | Platinum | +Nickel Ni^{++} | -0.049 | Gold | +Tin Sn^{++} | -0.085(?) | | +_____________________|___________|____________________|__________ + + + +THE FOLDING OF A FILTER PAPER + +If a filter paper is folded along its diameter, and again folded along +the radius at right angles to the original fold, a cone is formed on +opening, the angle of which is 60°. Funnels for analytical use are +supposed to have the same angle, but are rarely accurate. It is +possible, however, with care, to fit a filter thus folded into a +funnel in such a way as to prevent air from passing down between the +paper and the funnel to break the column of liquid in the stem, +which aids greatly, by its gentle suction, in promoting the rate of +filtration. + +Such a filter has, however, the disadvantage that there are three +thicknesses of paper back of half of its filtering surface, as a +consequence of which one half of a precipitate washes or drains more +slowly. Much time may be saved in the aggregate by learning to fold a +filter in such a way as to improve its effective filtering surface. +The directions which follow, though apparently complicated on first +reading, are easily applied and easily remembered. Use a 6-inch filter +for practice. Place four dots on the filter, two each on diameters +which are at right angles to each other. Then proceed as follows: +(1) Fold the filter evenly across one of the diameters, creasing it +carefully; (2) open the paper, turn it over, rotate it 90° to the +right, bring the edges together and crease along the other diameter; +(3) open, and rotate 45° to the right, bring edges together, and +crease evenly; (4) open, and rotate 90° to the right, and crease +evenly; (5) open, turn the filter over, rotate 22-(1/2)° to the right, +and crease evenly; (6) open, rotate 45° to the right and crease +evenly; (7) open, rotate 45° to the right and crease evenly; (8) open, +rotate 45° to the right and crease evenly; (9) open the filter, and, +starting with one of the dots between thumb and forefinger of the +right hand, fold the second crease to the left over on it, and do +the same with each of the other dots. Place it, thus folded, in the +funnel, moisten it, and fit to the side of the funnel. The filter will +then have four short segments where there are three thicknesses +and four where there is one thickness, but the latter are evenly +distributed around its circumference, thus greatly aiding the passage +of liquids through the paper and hastening both filtration and washing +of the whole contents of the filter. + + +!SAMPLE PAGES FOR LABORATORY RECORDS! + +!Page A! + +Date + +CALIBRATION OF BURETTE No. + +___________________________________________________________________________ + | | | | + BURETTE | DIFFERENCE | OBSERVED | DIFFERENCE | CALCULATED + READINGS | | WEIGHTS | | CORRECTION +_______________|______________|______________|______________|______________ + 0.02 | | 16.27 | | + 10.12 | 10.10 | 26.35 | 10.08 | -.02 + 20.09 | 9.97 | 36.26 | 9.91 | -.06 + 30.16 | 10.07 | 46.34 | 10.08 | +.01 + 40.19 | 10.03 | 56.31 | 9.97 | -.06 + 50.00 | 9.81 | 66.17 | 9.86 | +.05 +_______________|______________|______________|______________|______________ + + These data to be obtained in duplicate for each burette. + + +!Page B! + +Date + + +DETERMINATION OF COMPARATIVE STRENGTH HCl vs. NaOH + +___________________________________________________________________________ + | | + DETERMINATION | I | II +_________________________|________________________|________________________ + | | + | Corrected | Corrected +Final Reading HCl | 48.17 48.08 | 43.20 43.14 +Initial Reading HCl | 0.12 .12 | .17 .17 + | ----- ----- | ----- ----- + | 47.96 | 42.97 + | | + | Corrected | Corrected +Final Reading HCl | 46.36 46.29 | 40.51 40.37 +Initial Reading HCl | 1.75 1.75 | .50 .50 + | ----- ----- | ----- ----- + | 44.54 | 39.87 + | | + log cc. NaOH | 1.6468 | 1.6008 + colog cc. HCl | 8.3192 | 8.3668 + | ------ | ------ + | 9.9680 - 10 | 9.9676 - 10 + 1 cc. HCl | .9290 cc. NaOH | .9282 cc. NaOH + Mean | .9286 | +_________________________|________________________|________________________ + + +Signed + +!Page C! +Date + + +STANDARDIZATION OF HYDROCHLORIC ACID +===================================================================== + | | +Weight sample and tube| 9.1793 | 8.1731 + | 8.1731 | 6.9187 + | ------ | ------ + Weight sample | 1.0062 | 1.2544 + | | +Final Reading HCl | 39.97 39.83 | 49.90 49.77 +Initial Reading HCl | .00 .00 | .04 .04 + | ----- ----- | ----- ----- + | 39.83 | 49.73 + | | +Final Reading NaOH | .26 .26 | .67 .67 +Initial Reading NaOH | .12 .12 | .36 .36 + | --- --- | --- --- + | .14 | .31 + | | + | .14 | .31 +Corrected cc. HCl | 39.83 - ----- = 39.68 | 49.73 - ----- = 49.40 + | .9286 | .9286 + | | +log sample | 0.0025 | 0.0983 +colog cc | 8.4014 - 10 | 8.3063 - 10 +colog milli equivalent| 1.2757 | 1.2757 + | ------ | ------ + | 9.6796 - 10 | 9.6803 - 10 + | | +Normal value HCl | .4782 | .4789 + Mean | .4786 | + | | +===================================================================== + +Signed + + +!Page D! +Date + + +DETERMINATION OF CHLORINE IN CHLORIDE, SAMPLE No. +===================================================================== + | | +Weight sample and tube| 16.1721 | 15.9976 + | 15.9976 | 15.7117 + | ------- | ------- + Weight sample | .1745 | .2859 + | | +Weight crucible | | + + precipitate | 14.4496 | 15.6915 + Constant weights | 14.4487 | 15.6915 + | 14.4485 | + | | + Weight crucible | 14.2216 | 15.3196 + Constant weight | 14.2216 | 15.3194 + | | + Weight AgCl | .2269 | .3721 + | | + log Cl | 1.5496 | 1.5496 + log weight AgCl | 9.3558 - 10 | 9.5706 - 10 + log 100 | 2.0000 | 2.0000 + colog AgCl | 7.8438 - 10 | 7.7438 - 10 + colog sample | 0.7583 | 0.5438 + | ------- | ------- + | 1.5075 | 1.5078 + | | + Cl in sample No. | 32.18% | 32.20% + | | +===================================================================== + +Signed + + +STRENGTH OF REAGENTS + +The concentrations given in this table are those suggested for use +in the procedures described in the foregoing pages. It is obvious, +however, that an exact adherence to these quantities is not essential. + + + Approx. Approx. + Grams relation relation + per to normal to molal + liter. solution solution + +Ammonium oxalate, (NH_{4})_{2}C_{2}O_{4}.H_{2}O 40 0.5N 0.25 +Barium chloride, BaCl_{2}.2H_{2}O 25 0.2N 0.1 +Magnesium ammonium chloride (of MgCl_{2}) 71 1.5N 0.75 +Mercuric chloride, HgCl_{2} 45 0.33N 0.66 +Potassium hydroxide, KOH (sp. gr. 1.27) 480 +Potassium thiocyanate, KSCN 5 0.05N 0.55 +Silver nitrate, AgNO_{3} 21 0.125N 0.125 +Sodium hydroxide, NaOH 100 2.5N 2.5 +Sodium carbonate. Na_{2}CO_{3} 159 3N 1.5 +Sodium phosphate, Na_{2}HPO_{4}.12H_{2}O 90 0.5N or 0.75N 0.25 + +Stannous chloride, SnCl_{2}, made by saturating hydrochloric acid (sp. +gr. 1.2) with tin, diluting with an equal volume of water, and adding +a slight excess of acid from time to time. A strip of metallic tin is +kept in the bottle. + +A solution of ammonium molybdate is best prepared as follows: Stir +100 grams of molybdic acid (MoO_{3}) into 400 cc. of cold, distilled +water. Add 80 cc. of concentrated ammonium hydroxide (sp. gr. 0.90). +Filter, and pour the filtrate slowly, with constant stirring, into a +mixture of 400 cc. concentrated nitric acid (sp. gr. 1.42) and 600 +cc. of water. Add to the mixture about 0.05 gram of microcosmic salt. +Filter, after allowing the whole to stand for 24 hours. + +The following data regarding the common acids and aqueous ammonia +are based upon percentages given in the Standard Tables of the +Manufacturing Chemists' Association of the United States [!J.S.C.I.!, +24 (1905), 787-790]. All gravities are taken at 15.5°C. and compared +with water at the same temperature. + +Aqueous ammonia (sp. gr. 0.96) contains 9.91 per cent NH_{3} by +weight, and corresponds to a 5.6 N and 5.6 molal solution. + +Aqueous ammonia (sp. gr. 0.90) contains 28.52 per cent NH_{3} by +weight, and corresponds to a 5.6 N and 5.6 molal solution. + +Hydrochloric acid (sp. gr. 1.12) contains 23.81 per cent HCl by +weight, and corresponds to a 7.3 N and 7.3 molal solution. + +Hydrochloric acid (sp. gr. 1.20) contains 39.80 per cent HCl by +weight, and corresponds to a 13.1 N and 13.1 molal solution. + +Nitric acid (sp. gr. 1.20) contains 32.25 per cent HNO_{3} by weight, +and corresponds to a 6.1 N and 6.1 molal solution: + +Nitric acid (sp. gr. 1.42) contains 69.96 per cent HNO_{3} by weight, +and corresponds to a 15.8 N and 15.8 molal solution. + +Sulphuric acid (sp. gr. 1.8354) contains 93.19 per cent H_{2}SO_{4} by +weight, and corresponds to a 34.8 N or 17.4 molal solution. + +Sulphuric acid (sp. gr. 1.18) contains 24.74 per cent H_{2}SO_{4} by +weight, and corresponds to a 5.9 N or 2.95 molal solution. + +The term !normal! (N), as used above, has the same significance as +in volumetric analyses. The molal solution is assumed to contain one +molecular weight in grams in a liter of solution. + +DENSITIES AND VOLUMES OF WATER AT TEMPERATURES FROM 15-30°C. + +Temperature Density. Volume. +Centigrade. + + 4° 1.000000 1.000000 + 15° 0.999126 1.000874 + 16° 0.998970 1.001031 + 17° 0.998801 1.001200 + 18° 0.998622 1.001380 + 19° 0.998432 1.001571 + 20° 0.998230 1.001773 + 21° 0.998019 1.001985 + 22° 0.997797 1.002208 + 23° 0.997565 1.002441 + 24° 0.997323 1.002685 + 25° 0.997071 1.002938 + 26° 0.996810 1.003201 + 27° 0.996539 1.003473 + 28° 0.996259 1.003755 + 29° 0.995971 1.004046 + 30° 0.995673 1.004346 + +Authority: Landolt, Börnstein, and Meyerhoffer's !Tabellen!, third +edition. + + +CORRECTIONS FOR CHANGE OF TEMPERATURE OF STANDARD SOLUTIONS + +The values below are average values computed from data relating to a +considerable number of solutions. They are sufficiently accurate for +use in chemical analyses, except in the comparatively few cases +where the highest attainable accuracy is demanded in chemical +investigations. The expansion coefficients should then be carefully +determined for the solutions employed. For a compilation of the +existing data, consult Landolt, Börnstein, and Meyerhoffer's +!Tabellen!, third edition. + + Corrections for 1 cc. + Concentration. of solution between + 15° and 35°C. + + Normal .00029 + 0.5 Normal .00025 + 0.1 Normal or more dilute solutions .00020 + +The volume of solution used should be multiplied by the values given, +and that product multiplied by the number of degrees which the +temperature of the solution varies from the standard temperature +selected for the laboratory. The total correction thus found is +subtracted from the observed burette reading if the temperature is +higher than the standard, or added, if it is lower. Corrections are +not usually necessary for variations of temperature of 2°C. or less. + + + + INTERNATIONAL ATOMIC WEIGHTS + +========================================================== + | | | + | 1920 | | 1920 +_________________|_________|___________________|__________ + | | | +Aluminium Al | 27.1 | Molybdenum Mo | 96.0 +Antimony Sb | 120.2 | Neodymium Nd | 144.3 +Argon A | 39.9 | Neon Ne | 20.2 +Arsenic As | 74.96 | Nickel Ni | 58.68 +Barium Ba | 137.37 | Nitrogen N | 14.008 +Bismuth Bi | 208.0 | Osmium Os | 190.9 +Boron B | 11.0 | Oxygen O | 16.00 +Bromine Br | 79.92 | Palladium Pd | 106.7 +Cadmium Cd | 112.40 | Phosphorus P | 31.04 +Caesium Cs | 132.81 | Platinum Pt | 195.2 +Calcium Ca | 40.07 | Potassium K | 39.10 +Carbon C | 12.005 | Praseodymium Pr | 140.9 +Cerium Ce | 140.25 | Radium Ra | 226.0 +Chlorine Cl | 35.46 | Rhodium Rh | 102.9 +Chromium Cr | 52.0 | Rubidium Rb | 85.45 +Cobalt Co | 58.97 | Ruthenium Ru | 101.7 +Columbium Cb | 93.1 | Samarium Sm | 150.4 +Copper Cu | 63.57 | Scandium Sc | 44.1 +Dysprosium Dy | 162.5 | Selenium Se | 79.2 +Erbium Er | 167.7 | Silicon Si | 28.3 +Europium Eu | 152.0 | Silver Ag | 107.88 +Fluorine Fl | 19.0 | Sodium Na | 23.00 +Gadolinium Gd | 157.3 | Strontium Sr | 87.63 +Gallium Ga | 69.9 | Sulphur S | 32.06 +Germanium Ge | 72.5 | Tantalum Ta | 181.5 +Glucinum Gl | 9.1 | Tellurium Te | 127.5 +Gold Au | 197.2 | Terbium Tb | 159.2 +Helium He | 4.00 | Thallium Tl | 204.0 +Hydrogen H | 1.008 | Thorium Th | 232.4 +Indium In | 114.8 | Thulium Tm | 168.5 +Iodine I | 126.92 | Tin Sn | 118.7 +Iridium Ir | 193.1 | Titanium Ti | 48.1 +Iron Fe | 55.84 | Tungsten W | 184.0 +Krypton Kr | 82.92 | Uranium U | 238.2 +Lanthanum La | 139.0 | Vanadium V | 51.0 +Lead Pb | 207.2 | Xenon Xe | 130.2 +Lithium Li | 6.94 | Ytterbium Yb | 173.5 +Lutecium Lu | 175.0 | Yttrium Y | 88.7 +Magnesium Mg | 24.32 | Zinc Zn | 65.37 +Manganese Mn | 54.93 | Zirconium Zr | 90.6 +Mercury Hg | 200.6 | | +========================================================== + + + + +INDEX + +Acidimetry +Acid solutions, normal + standard +Acids, definition of +Acids, weak, action of other acids on + action of salts on +Accuracy demanded +Alkalimetry +Alkali solutions, normal + standard +Alumina, determination of in stibnite +Ammonium nitrate, acid +Analytical chemistry, subdivisions of +Antimony, determination of, in stibnite +Apatite, analysis of +Asbestos filters +Atomic weights, table of + +Balances, essential features of + use and care of +Barium sulphate, determination of sulphur in +Bases, definition of +Bichromate process for iron +Bleaching powder, analysis of +Brass, analysis of +Burette, description of + calibration of + cleaning of + reading of + +Calcium, determination of, in limestone +Calibration, definition of + of burettes + of flasks +Carbon dioxide, determination of, in limestone +Chlorimetry +Chlorine, gravimetric determination of +Chrome iron ore, analysis of +Coin, determination of silver in +Colloidal solution of precipitates +Colorimetric analyses, definition of +Copper, determination of, in brass + determination of in copper ores +Crucibles, use of +Crystalline precipitates + +Densities of water +Deposition potentials +Desiccators +Direct methods +Dissociation, degree of + +Economy of time +Electrolytic dissociation, theory of +Electrolytic separations, principles of +End-point, definition of +Equilibrium, chemical +Evaporation of liquids + +Faraday's law +Feldspar, analysis of +Ferrous ammonium sulphate, analysis of +Filters, folding of + how fitted +Filtrates, testing of +Filtration +Flasks, graduation of +Funnels +Fusions, removal of from crucibles + +General directions for gravimetric analysis + volumetric analysis +Gooch filter +Gravimetric analysis, definition of + +Hydrochloric acid, standardization of +Hydrolysis + +Ignition of precipitates +Indicators, definition of + for acidimetry + preparation of +Indirect methods +Insoluble matter, determination of in limestone +Integrity +Iodimetry +Ions, definition of +Iron, gravimetric determination of + volumetric determination of + +Jones reductor + +Lead, determination of in brass +Limestone, analysis of +Limonite, determination of iron in +Liquids, evaporation of + transfer of +Litmus +Logarithms + +Magnesium, determination of +Mass action, law of +Measuring instruments +Methyl orange +Moisture, determination of in limestone + +Neutralization methods +Normal solutions, acid and alkali + oxidizing agents + reducing agents +Notebooks, sample pages of + +Oxalic acid, determination of strength of +Oxidation processes +Oxidizing power of pyrolusite + +Permanganate process for iron +Phenolphthalein +Phosphoric anhydride, determination of +Pipette, calibration of + description of +Platinum crucibles, care of +Precipitates, colloidal + crystalline + ignition of + separation from filter + washing of +Precipitation +Precipitation methods (volumetric) +Problems +Pyrolusite, oxidizing power of + +Quantitative Analyses, subdivisions of + +Reagents, strength of +Reducing solution, normal +Reductor, Jones +Reversible reactions + +Silica, determination of, in limestone + determination of, in silicates + purification of +Silicic acid, dehydration of +Silver, determination of in coin +Soda ash, alkaline strength of +Sodium chloride, determination of chlorine in +Solubility product +Solution pressure +Solutions, normal + standard +Standardization, definition of +Standard solutions, acidimetry and alkalimetry + chlorimetry + iodimetry + oxidizing and reducing agents + thiocyanate +Starch solutions +Stibnite, determination of antimony in +Stirring rods +Stoichiometry +Strength of reagents +Suction, use of +Sulphur, determination of in ferrous ammonium sulphate + in barium sulphate + +Temperature, corrections for +Testing of washings +Theory of electrolytic dissociation +Thiocyanate process for silver +Titration, definition of +Transfer of liquids + +Volumetric analysis, definition of + general directions + +Wash-bottles +Washed filters +Washing of precipitates +Washings, testing of +Water, ionization of + densities of +Weights, care of + +Zimmermann-Reinhardt method for iron +Zinc, determination of, in brass + + + + + + + +End of the Project Gutenberg EBook of An Introductory Course of Quantitative +Chemical Analysis, by Henry P. 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You may copy it, give it away or +re-use it under the terms of the Project Gutenberg License included +with this eBook or online at www.gutenberg.org + + +Title: An Introductory Course of Quantitative Chemical Analysis + With Explanatory Notes + +Author: Henry P. Talbot + +Release Date: June 30, 2004 [EBook #12787] + +Language: English + +Character set encoding: ISO-8859-1 + +*** START OF THIS PROJECT GUTENBERG EBOOK QUANTITATIVE CHEMICAL ANALYSIS *** + + + + +Produced by Kevin Handy, Dave Maddock, Josephine Paolucci and the +Online Distributed Proofreading Team. + + + + + +[Transcriber's notes: In the chemical equations, superscripts are +indicated with a ^ and subscripts are indicated with a _. The affected +item is enclosed in curly brackets {}. Examples are H^{+} for hydrogen +ion and H_{2}O for water. Since the underscore is already being used +in this project, italics are designated by an exclamation point +before and after the italicized word or phrase.] + + + +AN INTRODUCTORY COURSE + +OF + +QUANTITATIVE + +CHEMICAL ANALYSIS + +WITH + +EXPLANATORY NOTES + + +BY + +HENRY P. TALBOT + +PROFESSOR OF INORGANIC CHEMISTRY AT THE MASSACHUSETTS INSTITUTE OF +TECHNOLOGY + +SIXTH EDITION, COMPLETELY REWRITTEN + + + + +PREFACE + + +This Introductory Course of Quantitative Analysis has been prepared +to meet the needs of students who are just entering upon the subject, +after a course of qualitative analysis. It is primarily intended to +enable the student to work successfully and intelligently without the +necessity for a larger measure of personal assistance and supervision +than can reasonably be given to each member of a large class. To this +end the directions are given in such detail that there is very little +opportunity for the student to go astray; but the manual is not, the +author believes, on this account less adapted for use with small +classes, where the instructor, by greater personal influence, can +stimulate independent thought on the part of the pupil. + +The method of presentation of the subject is that suggested by +Professor A.A. Noyes' excellent manual of Qualitative Analysis. For +each analysis the procedure is given in considerable detail, and +this is accompanied by explanatory notes, which are believed to be +sufficiently expanded to enable the student to understand fully the +underlying reason for each step prescribed. The use of the book +should, nevertheless, be supplemented by classroom instruction, mainly +of the character of recitations, and the student should be taught to +consult larger works. The general directions are intended to emphasize +those matters upon which the beginner in quantitative analysis must +bestow special care, and to offer helpful suggestions. The student +can hardly be expected to appreciate the force of all the statements +contained in these directions, or, indeed, to retain them all in +the memory after a single reading; but the instructor, by frequent +reference to special paragraphs, as suitable occasion presents itself, +can soon render them familiar to the student. + +The analyses selected for practice are those comprised in the first +course of quantitative analysis at the Massachusetts Institute of +Technology, and have been chosen, after an experience of years, +as affording the best preparation for more advanced work, and as +satisfactory types of gravimetric and volumetric methods. From the +latter point of view, they also seem to furnish the best insight into +quantitative analysis for those students who can devote but a limited +time to the subject, and who may never extend their study beyond the +field covered by this manual. The author has had opportunity to test +the efficiency of the course for use with such students, and has found +the results satisfactory. + +In place of the usual custom of selecting simple salts as material for +preliminary practice, it has been found advantageous to substitute, in +most instances, approximately pure samples of appropriate minerals or +industrial products. The difficulties are not greatly enhanced, while +the student gains in practical experience. + +The analytical procedures described in the following pages have been +selected chiefly with reference to their usefulness in teaching the +subject, and with the purpose of affording as wide a variety of +processes as is practicable within an introductory course of this +character. The scope of the manual precludes any extended attempt to +indicate alternative procedures, except through general references to +larger works on analytical chemistry. The author is indebted to the +standard works for many suggestions for which it is impracticable to +make specific acknowledgment; no considerable credit is claimed by him +for originality of procedure. + +For many years, as a matter of convenience, the classes for which this +text was originally prepared were divided, one part beginning with +gravimetric processes and the other with volumetric analyses. After a +careful review of the experience thus gained the conclusion has been +reached that volumetric analysis offers the better approach to the +subject. Accordingly the arrangement of the present (the sixth) +edition of this manual has been changed to introduce volumetric +procedures first. Teachers who are familiar with earlier editions +will, however, find that the order of presentation of the material +under the various divisions is nearly the same as that previously +followed, and those who may still prefer to begin the course of +instruction with gravimetric processes will, it is believed, be able +to follow that order without difficulty. + +Procedures for the determination of sulphur in insoluble sulphates, +for the determination of copper in copper ores by iodometric methods, +for the determination of iron by permanganate in hydrochloric acid +solutions, and for the standardization of potassium permanganate +solutions using sodium oxalate as a standard, and of thiosulphate +solutions using copper as a standard, have been added. The +determination of silica in silicates decomposable by acids, as a +separate procedure, has been omitted. + +The explanatory notes have been rearranged to bring them into closer +association with the procedures to which they relate. The number of +problems has been considerably increased. + +The author wishes to renew his expressions of appreciation of the +kindly reception accorded the earlier editions of this manual. He has +received helpful suggestions from so many of his colleagues within the +Institute, and friends elsewhere, that his sense of obligation must +be expressed to them collectively. He is under special obligations +to Professor L.F. Hamilton for assistance in the preparation of the +present edition. + +HENRY P. TALBOT + +!Massachusetts Institute of Technology, September, 1921!. + + + + +CONTENTS + + +PART I. INTRODUCTION + +SUBDIVISIONS OF ANALYTICAL CHEMISTRY + +GENERAL DIRECTIONS + Accuracy and Economy of Time; Notebooks; Reagents; Wash-bottles; + Transfer of Liquids + + +PART II. VOLUMETRIC ANALYSIS + +GENERAL DISCUSSION + Subdivisions; The Analytical Balance; Weights; Burettes; + Calibration of Measuring Devices +GENERAL DIRECTIONS + Standard and Normal Solutions + +!I. Neutralization Methods! + +ALKALIMETRY AND ACIDIMETRY + Preparation and Standardization of Solutions; Indicators +STANDARDIZATION OF HYDROCHLORIC ACID +DETERMINATION OF TOTAL ALKALINE STRENGTH OF SODA ASH +DETERMINATION OF ACID STRENGTH OF OXALIC ACID + +!II. Oxidation Processes! + +GENERAL DISCUSSION +BICHROMATE PROCESS FOR THE DETERMINATION OF IRON +DETERMINATION OF IRON IN LIMONITE BY THE BICHROMATE PROCESS +DETERMINATION OF CHROMIUM IN CHROME IRON ORE +PERMANGANATE PROCESS FOR THE DETERMINATION OF IRON +DETERMINATION OF IRON IN LIMONITE BY THE PERMANGANATE PROCESS +DETERMINATION OF IRON IN LIMONITE BY THE ZIMMERMANN-REINHARDT PROCESS +DETERMINATION OF THE OXIDIZING POWER OF PYROLUSITE +IODIMETRY +DETERMINATION OF COPPER IN ORES +DETERMINATION OF ANTIMONY IN STIBNITE +CHLORIMETRY +DETERMINATION OF AVAILABLE CHLORINE IN BLEACHING POWDER + +!III. Precipitation Methods! + +DETERMINATION OF SILVER BY THE THIOCYANATE PROCESS + + +PART III. GRAVIMETRIC ANALYSIS + +GENERAL DIRECTIONS + Precipitation; Funnels and Filters; Filtration and Washing of + Precipitates; Desiccators; Crucibles and their Preparation + for Use; Ignition of Precipitates +DETERMINATION OF CHLORINE IN SODIUM CHLORIDE +DETERMINATION OF IRON AND OF SULPHUR IN FERROUS AMMONIUM SULPHATE +DETERMINATION OF SULPHUR IN BARIUM SULPHATE +DETERMINATION OF PHOSPHORIC ANHYDRIDE IN APATITE +ANALYSIS OF LIMESTONE + Determination of Moisture; Insoluble Matter and Silica; Ferric + Oxide and Alumina; Calcium; Magnesium; Carbon Dioxide +ANALYSIS OF BRASS + Electrolytic Separations; Determination of Lead, Copper, Iron + and Zinc. +DETERMINATION OF SILICA IN SILICATES + +PART IV. STOICHIOMETRY + +SOLUTIONS OF TYPICAL PROBLEMS +PROBLEMS + +APPENDIX + +ELECTROLYTIC DISSOCIATION THEORY +FOLDING OF A FILTER PAPER +SAMPLE NOTEBOOK PAGES +STRENGTH OF REAGENTS +DENSITIES AND VOLUMES OF WATER +CORRECTIONS FOR CHANGE OF TEMPERATURE OF STANDARD SOLUTIONS +ATOMIC WEIGHTS +LOGARITHM TABLES + + + + +QUANTITATIVE CHEMICAL ANALYSIS + + + + +PART I + +INTRODUCTION + +SUBDIVISIONS OF ANALYTICAL CHEMISTRY + + +A complete chemical analysis of a body of unknown composition involves +the recognition of its component parts by the methods of !qualitative +analysis!, and the determination of the proportions in which these +components are present by the processes of !quantitative analysis!. +A preliminary qualitative examination is generally indispensable, if +intelligent and proper provisions are to be made for the separation of +the various constituents under such conditions as will insure accurate +quantitative estimations. + +It is assumed that the operations of qualitative analysis are familiar +to the student, who will find that the reactions made use of in +quantitative processes are frequently the same as those employed in +qualitative analyses with respect to both precipitation and systematic +separation from interfering substances; but it should be noted that +the conditions must now be regulated with greater care, and in such +a manner as to insure the most complete separation possible. For +example, in the qualitative detection of sulphates by precipitation +as barium sulphate from acid solution it is not necessary, in most +instances, to take into account the solubility of the sulphate +in hydrochloric acid, while in the quantitative determination of +sulphates by this reaction this solubility becomes an important +consideration. The operations of qualitative analysis are, therefore, +the more accurate the nearer they are made to conform to quantitative +conditions. + +The methods of quantitative analysis are subdivided, according +to their nature, into those of !gravimetric analysis, volumetric +analysis!, and !colorimetric analysis!. In !gravimetric! processes the +constituent to be determined is sometimes isolated in elementary +form, but more commonly in the form of some compound possessing a +well-established and definite composition, which can be readily and +completely separated, and weighed either directly or after ignition. +From the weight of this substance and its known composition, the +amount of the constituent in question is determined. + +In !volumetric! analysis, instead of the final weighing of a definite +body, a well-defined reaction is caused to take place, wherein the +reagent is added from an apparatus so designed that the volume of the +solution employed to complete the reaction can be accurately measured. +The strength of this solution (and hence its value for the reaction +in question) is accurately known, and the volume employed serves, +therefore, as a measure of the substance acted upon. An example will +make clear the distinction between these two types of analysis. +The percentage of chlorine in a sample of sodium chloride may be +determined by dissolving a weighed amount of the chloride in water +and precipitating the chloride ions as silver chloride, which is +then separated by filtration, ignited, and weighed (a !gravimetric! +process); or the sodium chloride may be dissolved in water, and a +solution of silver nitrate containing an accurately known amount of +the silver salt in each cubic centimeter may be cautiously added from +a measuring device called a burette until precipitation is complete, +when the amount of chlorine may be calculated from the number of cubic +centimeters of the silver nitrate solution involved in the reaction. +This is a !volumetric! process, and is equivalent to weighing without +the use of a balance. + +Volumetric methods are generally more rapid, require less apparatus, +and are frequently capable of greater accuracy than gravimetric +methods. They are particularly useful when many determinations of the +same sort are required. + +In !colorimetric! analyses the substance to be determined is converted +into some compound which imparts to its solutions a distinct color, +the intensity of which must vary in direct proportion to the amount of +the compound in the solution. Such solutions are compared with respect +to depth of color with standard solutions containing known amounts of +the colored compound, or of other similar color-producing substance +which has been found acceptable as a color standard. Colorimetric +methods are, in general, restricted to the determinations of very +small quantities, since only in dilute solutions are accurate +comparisons of color possible. + + + + +GENERAL DIRECTIONS + + +The following paragraphs should be read carefully and thoughtfully. A +prime essential for success as an analyst is attention to details and +the avoidance of all conditions which could destroy, or even lessen, +confidence in the analyses when completed. The suggestions here given +are the outcome of much experience, and their adoption will tend to +insure permanently work of a high grade, while neglect of them will +often lead to disappointment and loss of time. + + +ACCURACY AND ECONOMY OF TIME + +The fundamental conception of quantitative analysis implies a +necessity for all possible care in guarding against loss of material +or the introduction of foreign matter. The laboratory desk, and all +apparatus, should be scrupulously neat and clean at all times. A +sponge should always be ready at hand, and desk and filter-stands +should be kept dry and in good order. Funnels should never be allowed +to drip upon the base of the stand. Glassware should always be +wiped with a clean, lintless towel just before use. All filters and +solutions should be covered to protect them from dust, just as far as +is practicable, and every drop of solution or particle of precipitate +must be regarded as invaluable for the success of the analysis. + +An economical use of laboratory hours is best secured by acquiring +a thorough knowledge of the character of the work to be done before +undertaking it, and then by so arranging the work that no time shall +be wasted during the evaporation of liquids and like time-consuming +operations. To this end the student should read thoughtfully not only +the !entire! procedure, but the explanatory notes as well, before +any step is taken in the analysis. The explanatory notes furnish, in +general, the reasons for particular steps or precautions, but they +also occasionally contain details of manipulation not incorporated, +for various reasons, in the procedure. These notes follow the +procedures at frequent intervals, and the exact points to which they +apply are indicated by references. The student should realize that a +!failure to study the notes will inevitably lead to mistakes, loss of +time, and an inadequate understanding of the subject!. + +All analyses should be made in duplicate, and in general a close +agreement of results should be expected. It should, however, be +remembered that a close concordance of results in "check analyses" is +not conclusive evidence of the accuracy of those results, although the +probability of their accuracy is, of course, considerably enhanced. +The satisfaction in obtaining "check results" in such analyses must +never be allowed to interfere with the critical examination of the +procedure employed, nor must they ever be regarded as in any measure a +substitute for absolute truth and accuracy. + +In this connection it must also be emphasized that only the operator +himself can know the whole history of an analysis, and only he can +know whether his work is worthy of full confidence. No work should be +continued for a moment after such confidence is lost, but should +be resolutely discarded as soon as a cause for distrust is fully +established. The student should, however, determine to put forth his +best efforts in each analysis; it is well not to be too ready to +condone failures and to "begin again," as much time is lost in these +fruitless attempts. Nothing less than !absolute integrity! is or can +be demanded of a quantitative analyst, and any disregard of this +principle, however slight, is as fatal to success as lack of chemical +knowledge or inaptitude in manipulation can possibly be. + + +NOTEBOOKS + +Notebooks should contain, beside the record of observations, +descriptive notes. All records of weights should be placed upon the +right-hand page, while that on the left is reserved for the notes, +calculations of factors, or the amount of reagents required. + +The neat and systematic arrangement of the records of analyses is +of the first importance, and is an evidence of careful work and an +excellent credential. Of two notebooks in which the results may be, +in fact, of equal value as legal evidence, that one which is neatly +arranged will carry with it greater weight. + +All records should be dated, and all observations should be recorded +at once in the notebook. The making of records upon loose paper is a +practice to be deprecated, as is also that of copying original entries +into a second notebook. The student should accustom himself to orderly +entries at the time of observation. Several sample pages of systematic +records are to be found in the Appendix. These are based upon +experience; but other arrangements, if clear and orderly, may prove +equally serviceable. The student is advised to follow the sample pages +until he is in a position to plan out a system of his own. + + +REAGENTS + +The habit of carefully testing reagents, including distilled water, +cannot be too early acquired or too constantly practiced; for, in +spite of all reasonable precautionary measures, inferior chemicals +will occasionally find their way into the stock room, or errors will +be made in filling reagent bottles. The student should remember that +while there may be others who share the responsibility for the purity +of materials in the laboratory of an institution, the responsibility +will later be one which he must individually assume. + +The stoppers of reagent bottles should never be laid upon the desk, +unless upon a clean watch-glass or paper. The neck and mouth of all +such bottles should be kept scrupulously clean, and care taken that no +confusion of stoppers occurs. + + +WASH-BOTTLES + +Wash-bottles for distilled water should be made from flasks of about +750 cc. capacity and be provided with gracefully bent tubes, which +should not be too long. The jet should be connected with the tube +entering the wash-bottle by a short piece of rubber tubing in such +a way as to be flexible, and should deliver a stream about one +millimeter in diameter. The neck of the flask may be wound with cord, +or covered with wash-leather, for greater comfort when hot water is +used. It is well to provide several small wash-bottles for liquids +other than distilled water, which should invariably be clearly +labeled. + + +TRANSFER OF LIQUIDS + +Liquids should never be transferred from one vessel to another, nor to +a filter, without the aid of a stirring rod held firmly against the +side or lip of the vessel. When the vessel is provided with a lip it +is not usually necessary to use other means to prevent the loss of +liquid by running down the side; whenever loss seems imminent a !very +thin! layer of vaseline, applied with the finger to the edge of the +vessel, will prevent it. The stirring rod down which the liquid runs +should never be drawn upward in such a way as to allow the solution to +collect on the under side of the rim or lip of a vessel. + +The number of transfers of liquids from one vessel to another during +an analysis should be as small as possible to avoid the risk of slight +losses. Each vessel must, of course, be completely washed to insure +the transfer of all material; but it should be remembered that this +can be accomplished better by the use of successive small portions of +wash-water (perhaps 5-10 cc.), if each wash-water is allowed to drain +away for a few seconds, than by the addition of large amounts which +unnecessarily increase the volume of the solutions, causing loss of +time in subsequent filtrations or evaporations. + +All stirring rods employed in quantitative analyses should be rounded +at the ends by holding them in the flame of a burner until they begin +to soften. If this is not done, the rods will scratch the inner +surface of beakers, causing them to crack on subsequent heating. + + +EVAPORATION OF LIQUIDS + +The greatest care must be taken to prevent loss of solutions during +processes of evaporation, either from too violent ebullition, from +evaporation to dryness and spattering, or from the evolution of gas +during the heating. In general, evaporation upon the steam bath is to +be preferred to other methods on account of the impossibility of +loss by spattering. If the steam baths are well protected from dust, +solutions should be left without covers during evaporation; but +solutions which are boiled upon the hot plate, or from which gases are +escaping, should invariably be covered. In any case a watch-glass may +be supported above the vessel by means of a glass triangle, or other +similar device, and the danger of loss of material or contamination by +dust thus be avoided. It is obvious that evaporation is promoted by +the use of vessels which admit of the exposure of a broad surface to +the air. + +Liquids which contain suspended matter (precipitates) should always +be cautiously heated, since the presence of the solid matter is +frequently the occasion of violent "bumping," with consequent risk to +apparatus and analysis. + + + + +PART II + +VOLUMETRIC ANALYSIS + + +The processes of volumetric analysis are, in general, simpler than +those of gravimetric analysis and accordingly serve best as an +introduction to the practice of quantitative analysis. For their +execution there are required, first, an accurate balance with which +to weigh the material for analysis; second, graduated instruments in +which to measure the volume of the solutions employed; third, standard +solutions, that is, solutions the value of which is accurately known; +and fourth, indicators, which will furnish accurate evidence of the +point at which the desired reaction is completed. The nature of the +indicators employed will be explained in connection with the different +analyses. + +The process whereby a !standard solution! is brought into reaction is +called !titration!, and the point at which the reaction is exactly +completed is called the !end-point!. The !indicator! should show the +!end-point! of the !titration!. The volume of the standard solution +used then furnishes the measure of the substance to be determined as +truly as if that substance had been separated and weighed. + +The processes of volumetric analysis are easily classified, according +to their character, into: + +I. NEUTRALIZATION METHODS; such, for example, as those of acidimetry +and alkalimetry. + +II. OXIDATION PROCESSES; as exemplified in the determination of +ferrous iron by its oxidation with potassium bichromate. + +III. PRECIPITATION METHODS; of which the titration for silver with +potassium thiocyanate solution is an illustration. + +From a somewhat different standpoint the methods in each case may +be subdivided into (a) DIRECT METHODS, in which the substance to be +measured is directly determined by titration to an end-point with a +standard solution; and (b) INDIRECT METHODS, in which the substance +itself is not measured, but a quantity of reagent is added which is +known to be an excess with respect to a specific reaction, and the +unused excess determined by titration. Examples of the latter class +will be pointed out as they occur in the procedures. + + +MEASURING INSTRUMENTS + + +THE ANALYTICAL BALANCE + +For a complete discussion of the physical principles underlying the +construction and use of balances, and the various methods of weighing, +the student is referred to larger manuals of Quantitative Analysis, +such as those of Fresenius, or Treadwell-Hall, and particularly to +the admirable discussion of this topic in Morse's !Exercises in +Quantitative Chemistry!. + +The statements and rules of procedure which follow are sufficient +for the intelligent use of an analytical balance in connection with +processes prescribed in this introductory manual. It is, however, +imperative that the student should make himself familiar with these +essential features of the balance, and its use. He should fully +realize that the analytical balance is a delicate instrument which +will render excellent service under careful treatment, but such +treatment is an essential condition if its accuracy is to be depended +upon. He should also understand that no set of rules, however +complete, can do away with the necessity for a sense of personal +responsibility, since by carelessness he can render inaccurate not +only his own analyses, but those of all other students using the same +balance. + +Before making any weighings the student should seat himself before a +balance and observe the following details of construction: + +1. The balance case is mounted on three brass legs, which should +preferably rest in glass cups, backed with rubber to prevent slipping. +The front legs are adjustable as to height and are used to level the +balance case; the rear leg is of permanent length. + +2. The front of the case may be raised to give access to the balance. +In some makes doors are provided also at the ends of the balance case. + +3. The balance beam is mounted upon an upright in the center of the +case on the top of which is an inlaid agate plate. To the center of +the beam there is attached a steel or agate knife-edge on which the +beam oscillates when it rests on the agate plate. + +4. The balance beam, extending to the right and left, is graduated +along its upper edge, usually on both sides, and has at its +extremities two agate or steel knife-edges from which are suspended +stirrups. Each of these stirrups has an agate plate which, when the +balance is in action, rests upon the corresponding knife-edge of the +beam. The balance pans are suspended from the stirrups. + +5. A pointer is attached to the center of the beam, and as the beam +oscillates this pointer moves in front of a scale near the base of the +post. + +6. At the base of the post, usually in the rear, is a spirit-level. + +7. Within the upright is a mechanism, controlled by a knob at the +front of the balance case, which is so arranged as to raise the entire +beam slightly above the level at which the knife-edges are in contact +with the agate plates. When the balance is not in use the beam must +be supported by this device since, otherwise, the constant jarring +to which a balance is inevitably subjected, will soon dull the +knife-edges, and lessen the sensitiveness of the balance. + +8. A small weight, or bob, is attached to the pointer (or sometimes +to the beam) by which the center of gravity of the beam and its +attachments may be regulated. The center of gravity must lie very +slightly below the level of the agate plates to secure the desired +sensitiveness of the balance. This is provided for when the balance is +set up and very rarely requires alteration. The student should never +attempt to change this adjustment. + +9. Below the balance pans are two pan-arrests operated by a button +from the front of the case. These arrests exert a very slight upward +pressure upon the pans and minimize the displacement of the beam when +objects or weights are being placed upon the pans. + +10. A movable rod, operated from one end of the balance case, extends +over the balance beam and carries a small wire weight, called a rider. +By means of this rod the rider can be placed upon any desired division +of the scale on the balance beam. Each numbered division on the beam +corresponds to one milligram, and the use of the rider obviates the +placing of very small fractional weights on the balance pan. + +If a new rider is purchased, or an old one replaced, care must be +taken that its weight corresponds to the graduations on the beam of +the balance on which it is to be used. The weight of the rider in +milligrams must be equal to the number of large divisions (5, 6, 10, +or 12) between the central knife-edge and the knife-edge at the end of +the beam. It should be noted that on some balances the last division +bears no number. Each new rider should be tested against a 5 or +10-milligram weight. + +In some of the most recent forms of the balance a chain device +replaces the smaller weights and the use of the rider as just +described. + +Before using a balance, it is always best to test its adjustment. This +is absolutely necessary if the balance is used by several workers; it +is always a wise precaution under any conditions. For this purpose, +brush off the balance pans with a soft camel's hair brush. Then note +(1) whether the balance is level; (2) that the mechanism for raising +and lowering the beams works smoothly; (3) that the pan-arrests touch +the pans when the beam is lowered; and (4) that the needle swings +equal distances on either side of the zero-point when set in motion +without any load on the pans. If the latter condition is not +fulfilled, the balance should be adjusted by means of the adjusting +screw at the end of the beam unless the variation is not more than one +division on the scale; it is often better to make a proper allowance +for this small zero error than to disturb the balance by an attempt at +correction. Unless a student thoroughly understands the construction +of a balance he should never attempt to make adjustments, but should +apply to the instructor in charge. + +The object to be weighed should be placed on the left-hand balance pan +and the weights upon the right-hand pan. Every substance which +could attack the metal of the balance pan should be weighed upon a +watch-glass, and all objects must be dry and cold. A warm body gives +rise to air currents which vitiate the accuracy of the weighing. + +The weights should be applied in the order in which they occur in the +weight-box (not at haphazard), beginning with the largest weight which +is apparently required. After a weight has been placed upon the pan +the beam should be lowered upon its knife-edges, and, if necessary, +the pan-arrests depressed. The movement of the pointer will then +indicate whether the weight applied is too great or too small. When +the weight has been ascertained, by the successive addition of small +weights, to the nearest 5 or 10 milligrams, the weighing is completed +by the use of the rider. The correct weight is that which causes the +pointer to swing an equal number of divisions to the right and left +of the zero-point, when the pointer traverses not less than five +divisions on either side. + +The balance case should always be closed during the final weighing, +while the rider is being used, to protect the pans from the effect of +air currents. + +Before the final determination of an exact weight the beam should +always be lifted from the knife-edges and again lowered into place, +as it frequently happens that the scale pans are, in spite of the +pan-arrests, slightly twisted by the impact of the weights, the beam +being thereby virtually lengthened or shortened. Lifting the beam +restores the proper alignment. + +The beam should never be set in motion by lowering it forcibly upon +the knife-edges, nor by touching the pans, but rather by lifting the +rider (unless the balance be provided with some of the newer devices +for the purpose), and the swing should be arrested only when the +needle approaches zero on the scale, otherwise the knife-edges become +dull. For the same reason the beam should never be left upon its +knife-edges, nor should weights be removed from or placed on the +pans without supporting the beam, except in the case of the small +fractional weights. + +When the process of weighing has been completed, the weight should +be recorded in the notebook by first noting the vacant spaces in the +weight-box, and then checking the weight by again noting the weights +as they are removed from the pan. This practice will often detect and +avoid errors. It is obvious that the weights should always be returned +to their proper places in the box, and be handled only with pincers. + +It should be borne in mind that if the mechanism of a balance is +deranged or if any substance is spilled upon the pans or in the +balance case, the damage should be reported at once. In many instances +serious harm can be averted by prompt action when delay might ruin the +balance. + +Samples for analysis are commonly weighed in small tubes with cork +stoppers. Since the stoppers are likely to change in weight from +the varying amounts of moisture absorbed from the atmosphere, it is +necessary to confirm the recorded weight of a tube which has been +unused for some time before weighing out a new portion of substance +from it. + + +WEIGHTS + +The sets of weights commonly used in analytical chemistry range from +20 grams to 5 milligrams. The weights from 20 grams to 1 gram are +usually of brass, lacquered or gold plated. The fractional weights +are of German silver, gold, platinum or aluminium. The rider is of +platinum or aluminium wire. + +The sets of weights purchased from reputable dealers are usually +sufficiently accurate for analytical work. It is not necessary that +such a set should be strictly exact in comparison with the absolute +standard of weight, provided they are relatively correct among +themselves, and provided the same set of weights is used in all +weighings made during a given analysis. The analyst should assure +himself that the weights in a set previously unfamiliar to him are +relatively correct by a few simple tests. For example, he should make +sure that in his set two weights of the same denomination (i.e., two +10-gram weights, or the two 100-milligram weights) are actually equal +and interchangeable, or that the 500-milligram weight is equal to +the sum of the 200, 100, 100, 50, 20, 20 and 10-milligram weights +combined, and so on. If discrepancies of more than a few tenths of a +milligram (depending upon the total weight involved) are found, the +weights should be returned for correction. The rider should also be +compared with a 5 or 10-milligram weight. + +In an instructional laboratory appreciable errors should be reported +to the instructor in charge for his consideration. + +When the highest accuracy is desired, the weights may be calibrated +and corrections applied. A calibration procedure is described in a +paper by T.W. Richards, !J. Am. Chem. Soc.!, 22, 144, and in many +large text-books. + +Weights are inevitably subject to corrosion if not properly protected +at all times, and are liable to damage unless handled with great care. +It is obvious that anything which alters the weight of a single piece +in an analytical set will introduce an error in every weighing made +in which that piece is used. This source of error is often extremely +obscure and difficult to detect. The only safeguard against such +errors is to be found in scrupulous care in handling and protection +on the part of the analyst, and an equal insistence that if several +analysts use the same set of weights, each shall realize his +responsibility for the work of others as well as his own. + + +BURETTES + +A burette is made from a glass tube which is as uniformly cylindrical +as possible, and of such a bore that the divisions which are etched +upon its surface shall correspond closely to actual contents. + +The tube is contracted at one extremity, and terminates in either a +glass stopcock and delivery-tube, or in such a manner that a piece of +rubber tubing may be firmly attached, connecting a delivery-tube of +glass. The rubber tubing is closed by means of a glass bead. Burettes +of the latter type will be referred to as "plain burettes." + +The graduations are usually numbered in cubic centimeters, and the +latter are subdivided into tenths. + +One burette of each type is desirable for the analytical procedures +which follow. + + +PREPARATION OF A BURETTE FOR USE + +The inner surface of a burette must be thoroughly cleaned in order +that the liquid as drawn out may drain away completely, without +leaving drops upon the sides. This is best accomplished by treating +the inside of the burette with a warm solution of chromic acid in +concentrated sulphuric acid, applied as follows: If the burette is of +the "plain" type, first remove the rubber tip and force the lower +end of the burette into a medium-sized cork stopper. Nearly fill the +burette with the chromic acid solution, close the upper end with a +cork stopper and tip the burette backward and forward in such a way +as to bring the solution into contact with the entire inner surface. +Remove the stopper and pour the solution into a stock bottle to be +kept for further use, and rinse out the burette with water several +times. Unless the water then runs freely from the burette without +leaving drops adhering to the sides, the process must be repeated +(Note 1). + +If the burette has a glass stopcock, this should be removed after +the cleaning and wiped, and also the inside of the ground joint. The +surface of the stopcock should then be smeared with a thin coating of +vaseline and replaced. It should be attached to the burette by means +of a wire, or elastic band, to lessen the danger of breakage. + +Fill the burettes with distilled water, and allow the water to run out +through the stopcock or rubber tip until convinced that no air +bubbles are inclosed (Note 2). Fill the burette to a point above the +zero-point and draw off the water until the meniscus is just below +that mark. It is then ready for calibration. + +[Note 1: The inner surface of the burette must be absolutely clean if +the liquid is to run off freely. Chromic acid in sulphuric acid is +usually found to be the best cleansing agent, but the mixture must be +warm and concentrated. The solution can be prepared by pouring over a +few crystals of potassium bichromate a little water and then adding +concentrated sulphuric acid.] + +[Note 2: It is always necessary to insure the absence of air bubbles +in the tips or stopcocks. The treatment described above will usually +accomplish this, but, in the case of plain burettes it is sometimes +better to allow a little of the liquid to flow out of the tip while it +is bent upwards. Any air which may be entrapped then rises with the +liquid and escapes. + +If air bubbles escape during subsequent calibration or titration, an +error is introduced which vitiates the results.] + + +READING OF A BURETTE + +All liquids when placed in a burette form what is called a meniscus at +their upper surfaces. In the case of liquids such as water or +aqueous solutions this meniscus is concave, and when the liquids are +transparent accurate readings are best obtained by observing the +position on the graduated scales of the lowest point of the meniscus. +This can best be done as follows: Wrap around the burette a piece of +colored paper, the straight, smooth edges of which are held evenly +together with the colored side next to the burette (Note 1). Hold the +paper about two small divisions below the meniscus and raise or lower +the level of the eyes until the edge of the paper at the back of the +burette is just hidden from the eye by that in front (Note 2). Note +the position of the lowest point of the curve of the meniscus, +estimating the tenths of the small divisions, thus reading its +position to hundredths of a cubic centimeter. + +[Note 1: The ends of the colored paper used as an aid to accurate +readings may be fastened together by means of a gummed label. The +paper may then remain on the burette and be ready for immediate use by +sliding it up or down, as required.] + +[Note 2: To obtain an accurate reading the eye must be very nearly on +a level with the meniscus. This is secured by the use of the paper +as described. The student should observe by trial how a reading is +affected when the meniscus is viewed from above or below. + +The eye soon becomes accustomed to estimating the tenths of the +divisions. If the paper is held as directed, two divisions below the +meniscus, one whole division is visible to correct the judgment. It is +not well to attempt to bring the meniscus exactly to a division mark +on the burette. Such readings are usually less accurate than those in +which the tenths of a division are estimated.] + + +CALIBRATION OF GLASS MEASURING DEVICES + +If accuracy of results is to be attained, the correctness of all +measuring instruments must be tested. None of the apparatus offered +for sale can be implicitly relied upon except those more expensive +instruments which are accompanied by a certificate from the !National +Bureau of Standards! at Washington, or other equally authentic source. + +The bore of burettes is subject to accidental variations, and since +the graduations are applied by machine without regard to such +variations of bore, local errors result. + +The process of testing these instruments is called !calibration!. +It is usually accomplished by comparing the actual weight of water +contained in the instrument with its apparent volume. + +There is, unfortunately, no uniform standard of volume which has been +adopted for general use in all laboratories. It has been variously +proposed to consider the volume of 1000 grams of water at 4°, 15.5°, +16°, 17.5°, and even 20°C., as a liter for practical purposes, and to +consider the cubic centimeter to be one one-thousandth of that volume. +The true liter is the volume of 1000 grams of water at 4°C.; but +this is obviously a lower temperature than that commonly found in +laboratories, and involves the constant use of corrections if taken as +a laboratory standard. Many laboratories use 15.5°C. (60° F.) as the +working standard. It is plain that any temperature which is deemed +most convenient might be chosen for a particular laboratory, but it +cannot be too strongly emphasized that all measuring instruments, +including burettes, pipettes, and flasks, should be calibrated at that +temperature in order that the contents of each burette, pipette, +etc., shall be comparable with that of every other instrument, thus +permitting general interchange and substitution. For example, it is +obvious that if it is desired to remove exactly 50 cc. from a solution +which has been diluted to 500 cc. in a graduated flask, the 50 cc. +flask or pipette used to remove the fractional portion must give +a correct reading at the same temperature as the 500 cc. flask. +Similarly, a burette used for the titration of the 50 cc. of solution +removed should be calibrated under the same conditions as the +measuring flasks or pipettes employed with it. + +The student should also keep constantly in mind the fact that all +volumetric operations, to be exact, should be carried out as nearly at +a constant temperature as is practicable. The spot selected for +such work should therefore be subject to a minimum of temperature +variations, and should have as nearly the average temperature of +the laboratory as is possible. In all work, whether of calibration, +standardization, or analysis, the temperature of the liquids employed +must be taken into account, and if the temperature of these liquids +varies more than 3° or 4° from the standard temperature chosen for the +laboratory, corrections must be applied for errors due to expansion or +contraction, since volumes of a liquid measured at different times are +comparable only under like conditions as to temperature. Data to be +used for this purpose are given in the Appendix. Neglect of this +correction is frequently an avoidable source of error and annoyance in +otherwise excellent work. The temperature of all solutions at the time +of standardization should be recorded to facilitate the application of +temperature corrections, if such are necessary at any later time. + + +CALIBRATION OF THE BURETTES + +Two burettes, one at least of which should have a glass stopper, are +required throughout the volumetric work. Both burettes should be +calibrated by the student to whom they are assigned. + +PROCEDURE.--Weigh a 50 cc., flat-bottomed flask (preferably a +light-weight flask), which must be dry on the outside, to the nearest +centigram. Record the weight in the notebook. (See Appendix for +suggestions as to records.) Place the flask under the burette and draw +out into it about 10 cc. of water, removing any drop on the tip by +touching it against the inside of the neck of the flask. Do not +attempt to stop exactly at the 10 cc. mark, but do not vary more than +0.1 cc. from it. Note the time, and at the expiration of three minutes +(or longer) read the burette accurately, and record the reading in the +notebook (Note 1). Meanwhile weigh the flask and water to centigrams +and record its weight (Note 2). Draw off the liquid from 10 cc. to +about 20 cc. into the same flask without emptying it; weigh, and at +the expiration of three minutes take the reading, and so on throughout +the length of the burette. When it is completed, refill the burette +and check the first calibration. + +The differences in readings represent the apparent volumes, the +differences in weights the true volumes. For example, if an apparent +volume of 10.05 cc. is found to weigh 10.03 grams, it may be assumed +with sufficient accuracy that the error in that 10 cc. amounts to +-0.02 cc., or -0.002 for each cubic centimeter (Note 3). + +In the calculation of corrections the temperature of the water must be +taken into account, if this varies more than 4°C. from the laboratory +standard temperature, consulting the table of densities of water in +the Appendix. + +From the final data, plot the corrections to be applied so that they +may be easily read for each cubic centimeter throughout the burette. +The total correction at each 10 cc. may also be written on the burette +with a diamond, or etching ink, for permanence of record. + +[Note 1: A small quantity of liquid at first adheres to the side of +even a clean burette. This slowly unites with the main body of liquid, +but requires an appreciable time. Three minutes is a sufficient +interval, but not too long, and should be adopted in every instance +throughout the whole volumetric practice before final readings are +recorded.] + +[Note 2: A comparatively rough balance, capable of weighing to +centigrams, is sufficiently accurate for use in calibrations, for a +moment's reflection will show that it would be useless to weigh the +water with an accuracy greater than that of the readings taken on +the burette. The latter cannot exceed 0.01 cc. in accuracy, which +corresponds to 0.01 gram. + +The student should clearly understand that !all other weighings!, +except those for calibration, should be made accurately to 0.0001 +gram, unless special directions are given to the contrary. + +Corrections for temperature variations of less than 4°C. are +negligible, as they amount to less than 0.01 gram for each 10 grams of +water withdrawn.] + +[Note 3: Should the error discovered in any interval of 10 cc. on the +burette exceed 0.10 cc., it is advisable to weigh small portions (even +1 cc.) to locate the position of the variation of bore in the +tube rather than to distribute the correction uniformly over the +corresponding 10 cc. The latter is the usual course for small +corrections, and it is convenient to calculate the correction +corresponding to each cubic centimeter and to record it in the form +of a table or calibration card, or to plot a curve representing the +values. + +Burettes may also be calibrated by drawing off the liquid in +successive portions through a 5 cc. pipette which has been accurately +calibrated, as a substitute for weighing. If many burettes are to be +tested, this is a more rapid method.] + + +PIPETTES + +A !pipette! may consist of a narrow tube, in the middle of which is +blown a bulb of a capacity a little less than that which it is desired +to measure by the pipette; or it may be a miniature burette, without +the stopcock or rubber tip at the lower extremity. In either case, the +flow of liquid is regulated by the pressure of the finger on the top, +which governs the admission of the air. + +Pipettes are usually already graduated when purchased, but they +require calibration for accurate work. + + +CALIBRATION OF PIPETTES + +PROCEDURE.--Clean the pipette. Draw distilled water into it by sucking +at the upper end until the water is well above the graduation mark. +Quickly place the forefinger over the top of the tube, thus preventing +the entrance of air and holding the water in the pipette. Cautiously +admit a little air by releasing the pressure of the finger, and allow +the level of the water to fall until the lowest point of the meniscus +is level with the graduation. Hold the water at that point by pressure +of the finger and then allow the water to run out from the pipette +into a small tared, or weighed, beaker or flask. After a definite time +interval, usually two to three minutes, touch the end of the pipette +against the side of the beaker or flask to remove any liquid adhering +to it (Note 1). The increase in weight of the flask in grams +represents the volume of the water in cubic centimeters delivered by +the pipette. Calculate the necessary correction. + +[Note 1: A definite interval must be allowed for draining, and a +definite practice adopted with respect to the removal of the liquid +which collects at the end of the tube, if the pipette is designed to +deliver a specific volume when emptied. This liquid may be removed +at the end of a definite interval either by touching the side of the +vessel or by gently blowing out the last drops. Either practice, when +adopted, must be uniformly adhered to.] + + +FLASKS + +!Graduated or measuring flasks! are similar to the ordinary +flat-bottomed flasks, but are provided with long, narrow necks in +order that slight variations in the position of the meniscus with +respect to the graduation shall represent a minimum volume of liquid. +The flasks must be of such a capacity that, when filled with the +specified volume, the liquid rises well into the neck. + + +GRADUATION OF FLASKS + +It is a general custom to purchase the flasks ungraduated and to +graduate them for use under standard conditions selected for the +laboratory in question. They may be graduated for "contents" or +"delivery." When graduated for "contents" they contain a specified +volume when filled to the graduation at a specified temperature, and +require to be washed out in order to remove all of the solution from +the flask. Flasks graduated for "delivery" will deliver the specified +volume of a liquid without rinsing. A flask may, of course, be +graduated for both contents and delivery by placing two graduation +marks upon it. + +PROCEDURE.--To calibrate a flask for !contents!, proceed as follows: +Clean the flask, using a chromic acid solution, and dry it carefully +outside and inside. Tare it accurately; pour water into the flask +until the weight of the latter counterbalances weights on the opposite +pan which equal in grams the number of cubic centimeters of water +which the flask is to contain. Remove any excess of water with the aid +of filter paper (Note 1). Take the flask from the balance, stopper +it, place it in a bath at the desired temperature, usually 15.5° +or 17.5°C., and after an hour mark on the neck with a diamond the +location of the lowest point of the meniscus (Note 2). The mark may +be etched upon the flask by hydrofluoric acid, or by the use of an +etching ink now commonly sold on the market. + +To graduate a flask which is designed to !deliver! a specified volume, +proceed as follows: Clean the flask as usual and wipe all moisture +from the outside. Fill it with distilled water. Pour out the water +and allow the water to drain from the flask for three minutes. +Counterbalance the flask with weights to the nearest centigram. +Add weights corresponding in grams to the volume desired, and add +distilled water to counterbalance these weights. An excess of water, +or water adhering to the neck of the flask, may be removed by means of +a strip of clean filter paper. Stopper the flask, place it in a bath +at 15.5°C. or 17.5°C. and, after an hour, mark the location of the +lowest point of the meniscus, as described above. + +[Note 1: The allowable error in counterbalancing the water and +weights varies with the volume of the flask. It should not exceed one +ten-thousandth of the weight of water.] + +[Note 2: Other methods are employed which involve the use of +calibrated apparatus from which the desired volume of water may be run +into the dry flask and the position of the meniscus marked directly +upon it. For a description of a procedure which is most convenient +when many flasks are to be calibrated, the student is referred to the +!Am. Chem J.!, 16, 479.] + + + + +GENERAL DIRECTIONS FOR VOLUMETRIC ANALYSES + + +It cannot be too strongly emphasized that for the success of analyses +uniformity of practice must prevail throughout all volumetric work +with respect to those factors which can influence the accuracy of the +measurement of liquids. For example, whatever conditions are imposed +during the calibration of a burette, pipette, or flask (notably the +time allowed for draining), must also prevail whenever the flask or +burette is used. + +The student should also be constantly watchful to insure parallel +conditions during both standardization and analyst with respect to the +final volume of liquid in which a titration takes place. The value +of a standard solution is only accurate under the conditions which +prevailed when it was standardized. It is plain that the standard +solutions must be scrupulously protected from concentration or +dilution, after their value has been established. Accordingly, great +care must be taken to thoroughly rinse out all burettes, flasks, etc., +with the solutions which they are to contain, in order to remove all +traces of water or other liquid which could act as a diluent. It is +best to wash out a burette at least three times with small portions of +a solution, allowing each to run out through the tip before assuming +that the burette is in a condition to be filled and used. It is, of +course, possible to dry measuring instruments in a hot closet, but +this is tedious and unnecessary. + +To the same end, all solutions should be kept stoppered and away from +direct sunlight or heat. The bottles should be shaken before use to +collect any liquid which may have distilled from the solution and +condensed on the sides. + +The student is again reminded that variations in temperature of +volumetric solutions must be carefully noted, and care should always +be taken that no source of heat is sufficiently near the solutions to +raise the temperature during use. + +Much time may be saved by estimating the approximate volume of a +standard solution which will be required for a titration (if the data +are obtainable) before beginning the operation. It is then possible to +run in rapidly approximately the required amount, after which it is +only necessary to determine the end-point slowly and with accuracy. +In such cases, however, the knowledge of the approximate amount to be +required should never be allowed to influence the judgment regarding +the actual end-point. + + +STANDARD SOLUTIONS + +The strength or value of a solution for a specific reaction is +determined by a procedure called !Standardization!, in which the +solution is brought into reaction with a definite weight of a +substance of known purity. For example, a definite weight of pure +sodium carbonate may be dissolved in water, and the volume of a +solution of hydrochloric acid necessary to exactly neutralize the +carbonate accurately determined. From these data the strength or value +of the acid is known. It is then a !standard solution!. + + +NORMAL SOLUTIONS + +Standard solutions may be made of a purely empirical strength dictated +solely by convenience of manipulation, or the concentration may +be chosen with reference to a system which is applicable to all +solutions, and based upon chemical equivalents. Such solutions are +called !Normal Solutions! and contain such an amount of the reacting +substance per liter as is equivalent in its chemical action to one +gram of hydrogen, or eight grams of oxygen. Solutions containing one +half, one tenth, or one one-hundredth of this quantity per liter are +called, respectively, half-normal, tenth-normal, or hundredth-normal +solutions. + +Since normal solutions of various reagents are all referred to a +common standard, they have an advantage not possessed by empirical +solutions, namely, that they are exactly equivalent to each other. +Thus, a liter of a normal solution of an acid will exactly neutralize +a liter of a normal alkali solution, and a liter of a normal oxidizing +solution will exactly react with a liter of a normal reducing +solution, and so on. + +Beside the advantage of uniformity, the use of normal solutions +simplifies the calculations of the results of analyses. This is +particularly true if, in connection with the normal solution, the +weight of substance for analysis is chosen with reference to the +atomic or molecular weight of the constituent to be determined. (See +problem 26.) + +The preparation of an !exactly! normal, half-normal, or tenth-normal +solution requires considerable time and care. It is usually carried +out only when a large number of analyses are to be made, or when the +analyst has some other specific purpose in view. It is, however, a +comparatively easy matter to prepare standard solutions which differ +but slightly from the normal or half-normal solution, and these have +the advantage of practical equality; that is, two approximately +half-normal solutions are more convenient to work with than two which +are widely different in strength. It is, however, true that some of +the advantage which pertains to the use of normal solutions as regards +simplicity of calculations is lost when using these approximate +solutions. + +The application of these general statements will be made clear in +connection with the use of normal solutions in the various types of +volumetric processes which follow. + + + + +I. NEUTRALIZATION METHODS + +ALKALIMETRY AND ACIDIMETRY + + + + +GENERAL DISCUSSION + + +!Standard Acid Solutions! may be prepared from either hydrochloric, +sulphuric, or oxalic acid. Hydrochloric acid has the advantage of +forming soluble compounds with the alkaline earths, but its solutions +cannot be boiled without danger of loss of strength; sulphuric acid +solutions may be boiled without loss, but the acid forms insoluble +sulphates with three of the alkaline earths; oxalic acid can be +accurately weighed for the preparation of solutions, and its solutions +may be boiled without loss, but it forms insoluble oxalates with +three of the alkaline earths and cannot be used with certain of the +indicators. + +!Standard Alkali Solutions! may be prepared from sodium or potassium +hydroxide, sodium carbonate, barium hydroxide, or ammonia. Of sodium +and potassium hydroxide, it may be said that they can be used with all +indicators, and their solutions may be boiled, but they absorb carbon +dioxide readily and attack the glass of bottles, thereby losing +strength; sodium carbonate may be weighed directly if its purity is +assured, but the presence of carbonic acid from the carbonate is a +disadvantage with many indicators; barium hydroxide solutions may +be prepared which are entirely free from carbon dioxide, and such +solutions immediately show by precipitation any contamination from +absorption, but the hydroxide is not freely soluble in water; ammonia +does not absorb carbon dioxide as readily as the caustic alkalies, +but its solutions cannot be boiled nor can they be used with all +indicators. The choice of a solution must depend upon the nature of +the work in hand. + +A !normal acid solution! should contain in one liter that quantity of +the reagent which represents 1 gram of hydrogen replaceable by a base. +For example, the normal solution of hydrochloric acid (HCl) should +contain 36.46 grams of gaseous hydrogen chloride, since that amount +furnishes the requisite 1 gram of replaceable hydrogen. On the other +hand, the normal solution of sulphuric acid (H_{2}SO_{4}) should +contain only 49.03 grams, i.e., one half of its molecular weight in +grams. + +A !normal alkali solution! should contain sufficient alkali in a liter +to replace 1 gram of hydrogen in an acid. This quantity is represented +by the molecular weight in grams (40.01) of sodium hydroxide (NaOH), +while a sodium carbonate solution (Na_{2}CO_{3}) should contain but +one half the molecular weight in grams (i.e., 53.0 grams) in a liter +of normal solution. + +Half-normal or tenth-normal solutions are employed in most analyses +(except in the case of the less soluble barium hydroxide). Solutions +of the latter strength yield more accurate results when small +percentages of acid or alkali are to be determined. + + +INDICATORS + +It has already been pointed out that the purpose of an indicator is to +mark (usually by a change of color) the point at which just enough of +the titrating solution has been added to complete the chemical change +which it is intended to bring about. In the neutralization processes +which are employed in the measurement of alkalies (!alkalimetry!) +or acids (!acidimetry!) the end-point of the reaction should, in +principle, be that of complete neutrality. Expressed in terms of ionic +reactions, it should be the point at which the H^{+} ions from an +acid[Note 1] unite with a corresponding number of OH^{-} ions from a +base to form water molecules, as in the equation + +H^{+}, Cl^{-} + Na^{+}, OH^{-} --> Na^{+}, Cl^{-} + (H_{2}O). + +It is not usually possible to realize this condition of exact +neutrality, but it is possible to approach it with sufficient +exactness for analytical purposes, since substances are known which, +in solution, undergo a sharp change of color as soon as even a minute +excess of H^{+} or OH^{-} ions are present. Some, as will be seen, +react sharply in the presence of H^{+} ions, and others with OH^{-} +ions. These substances employed as indicators are usually organic +compounds of complex structure and are closely allied to the dyestuffs +in character. + +[Note 1: A knowledge on the part of the student of the ionic theory +as applied to aqueous solutions of electrolytes is assumed. A brief +outline of the more important applications of the theory is given in +the Appendix.] + + +BEHAVIOR OF ORGANIC INDICATORS + +The indicators in most common use for acid and alkali titrations are +methyl orange, litmus, and phenolphthalein. + +In the following discussion of the principles underlying the behavior +of the indicators as a class, methyl orange and phenolphthalein will +be taken as types. It has just been pointed out that indicators are +bodies of complicated structure. In the case of the two indicators +named, the changes which they undergo have been carefully studied by +Stieglitz (!J. Am. Chem. Soc.!, 25, 1112) and others, and it appears +that the changes involved are of two sorts: First, a rearrangement +of the atoms within the molecule, such as often occurs in organic +compounds; and, second, ionic changes. The intermolecular changes +cannot appropriately be discussed here, as they involve a somewhat +detailed knowledge of the classification and general behavior of +organic compounds; they will, therefore, be merely alluded to, and +only the ionic changes followed. + +Methyl orange is a representative of the group of indicators which, +in aqueous solutions, behave as weak bases. The yellow color which it +imparts to solutions is ascribed to the presence of the undissociated +base. If an acid, such as HCl, is added to such a solution, the acid +reacts with the indicator (neutralizes it) and a salt is formed, as +indicated by the equation: + +(M.o.)^{+}, OH^{-} + H^{+}, Cl^{-} --> (M.o.)^{+} Cl^{-} + (H_{2}O). + +This salt ionizes into (M.o.)^{+} (using this abbreviation for the +positive complex) and Cl^{-}; but simultaneously with this ionization +there appears to be an internal rearrangement of the atoms which +results in the production of a cation which may be designated as +(M'.o'.)^{+}, and it is this which imparts a characteristic red color +to the solution. As these changes occur in the presence of even a +very small excess of acid (that is, of H^{+} ions), it serves as the +desired index of their presence in the solution. If, now, an alkali, +such as NaOH, is added to this reddened solution, the reverse +series of changes takes place. As soon as the free acid present is +neutralized, the slightest excess of sodium hydroxide, acting as +a strong base, sets free the weak, little-dissociated base of the +indicator, and at the moment of its formation it reverts, because of +the rearrangement of the atoms, to the yellow form: + +OH^{-} + (M'.o'.)^{+} --> [M'.o'.OH] --> [M.o.OH]. + +Phenolphthalein, on the other hand, is a very weak, little-dissociated +acid, which is colorless in neutral aqueous solution or in the +presence of free H^{+} ions. When an alkali is added to such a +solution, even in slight excess, the anion of the salt which has +formed from the acid of the indicator undergoes a rearrangement of the +atoms, and a new ion, (Ph')^{+}, is formed, which imparts a pink color +to the solution: + +H^{+}, (Ph)^{-} + Na^{+}, OH^{-} --> (H_{2}O) + Na^{+}, (Ph)^{-} +--> Na^{+}, (Ph')^{-} + +The addition of the slightest excess of an acid to this solution, on +the other hand, occasions first the reversion to the colorless ion and +then the setting free of the undissociated acid of the indicator: + +H^{+}, (Ph')^{-} --> H^{+}, (Ph)^{-} --> (HPh). + +Of the common indicators methyl orange is the most sensitive toward +alkalies and phenolphthalein toward acids; the others occupy +intermediate positions. That methyl orange should be most sensitive +toward alkalies is evident from the following considerations: Methyl +orange is a weak base and, therefore, but little dissociated. It +should, then, be formed in the undissociated condition as soon as even +a slight excess of OH^{-} ions is present in the solution, and there +should be a prompt change from red to yellow as outlined above. On the +other hand, it should be an unsatisfactory indicator for use with weak +acids (acetic acid, for example) because the salts which it forms +with such acids are, like all salts of that type, hydrolyzed to a +considerable extent. This hydrolytic change is illustrated by the +equation: + +(M.o.)^{+} C_{2}H_{3}O_{2}^{-} + H^{+}, OH^{-} --> [M.o.OH] + H^{+}, +C_{2}H_{3}O_{2}^{-}. + +Comparison of this equation with that on page 30 will make it plain +that hydrolysis is just the reverse of neutralization and must, +accordingly, interfere with it. Salts of methyl orange with weak acids +are so far hydrolyzed that the end-point is uncertain, and methyl +orange cannot be used in the titration of such acids, while with +the very weak acids, such as carbonic acid or hydrogen sulphide +(hydrosulphuric acid), the salts formed with methyl orange are, in +effect, completely hydrolyzed (i.e., no neutralization occurs), and +methyl orange is accordingly scarcely affected by these acids. This +explains its usefulness, as referred to later, for the titration of +strong acids, such as hydrochloric acid, even in the presence of +carbonates or sulphides in solution. + +Phenolphthalein, on the other hand, should be, as it is, the best of +the common indicators for use with weak acids. For, since it is +itself a weak acid, it is very little dissociated, and its nearly +undissociated, colorless molecules are promptly formed as soon as +there is any free acid (that is, free H^{+} ions) in the solution. +This indicator cannot, however, be successfully used with weak bases, +even ammonium hydroxide; for, since it is weak acid, the salts +which it forms with weak alkalies are easily hydrolyzed, and as a +consequence of this hydrolysis the change of color is not sharp. +This indicator can, however, be successfully used with strong bases, +because the salts which it forms with such bases are much less +hydrolyzed and because the excess of OH^{-} ions from these bases also +diminishes the hydrolytic action of water. + +This indicator is affected by even so weak an acid as carbonic acid, +which must be removed by boiling the solution before titration. It is +the indicator most generally employed for the titration of organic +acids. + +In general, it may be stated that when a strong acid, such as +hydrochloric, sulphuric or nitric acid, is titrated against a strong +base, such as sodium hydroxide, potassium hydroxide, or barium +hydroxide, any of these indicators may be used, since very little +hydrolysis ensues. It has been noted above that the color change does +not occur exactly at theoretical neutrality, from which it follows +that no two indicators will show exactly the same end-point when acids +and alkalis are brought together. It is plain, therefore, that the +same indicator must be employed for both standardization and analysis, +and that, if this is done, accurate results are obtainable. + +The following table (Note 1) illustrates the variations in the volume +of an alkali solution (tenth-normal sodium hydroxide) required to +produce an alkaline end-point when run into 10 cc. of tenth-normal +sulphuric acid, diluted with 50 cc. of water, using five drops of each +of the different indicator solutions. + +==================================================================== + | | | | + INDICATOR | N/10 | N/10 |COLOR IN ACID|COLOR IN ALKA- + | H_{2}SO_{4}| NaOH |SOLUTION |LINE SOLUTION +_______________|____________|__________|_____________|______________ + | cc. | cc. | cc. | +Methyl orange | 10 | 9.90 | Red | Yellow +Lacmoid | 10 | 10.00 | Red | Blue +Litmus | 10 | 10.00 | Red | Blue +Rosalic acid | 10 | 10.07 | Yellow | Pink +Phenolphthalein| 10 | 10.10 | Colorless | Pink +==================================================================== + +It should also be stated that there are occasionally secondary +changes, other than those outlined above, which depend upon the +temperature and concentration of the solutions in which the indicators +are used. These changes may influence the sensitiveness of an +indicator. It is important, therefore, to take pains to use +approximately the same volume of solution when standardizing that is +likely to be employed in analysis; and when it is necessary, as is +often the case, to titrate the solution at boiling temperature, the +standardization should take place under the same conditions. It is +also obvious that since some acid or alkali is required to react with +the indicator itself, the amount of indicator used should be uniform +and not excessive. Usually a few drops of solution will suffice. + +The foregoing statements with respect to the behavior of indicators +present the subject in its simplest terms. Many substances other than +those named may be employed, and they have been carefully studied to +determine the exact concentration of H^{+} ions at which the color +change of each occurs. It is thus possible to select an indicator +for a particular purpose with considerable accuracy. As data of this +nature do not belong in an introductory manual, reference is made to +the following papers or books in which a more extended treatment of +the subject may be found: + +Washburn, E.W., Principles of Physical Chemistry (McGraw-Hill Book +Co.), (Second Edition, 1921), pp. 380-387. + +Prideaux, E.B.R., The Theory and Use of Indicators (Constable & Co., +Ltd.), (1917). + +Salm, E., A Study of Indicators, !Z. physik. Chem.!, 57 (1906), +471-501. + +Stieglitz, J., Theories of Indicators, !J. Am. Chem. Soc.!, 25 (1903), +1112-1127. + +Noyes, A.A., Quantitative Applications of the Theory of Indicators to +Volumetric Analysis, !J. Am. Chem. Soc.!, 32 (1911), 815-861. + +Bjerrum, N., General Discussion, !Z. Anal. Chem.!, 66 (1917), 13-28 +and 81-95. + +Ostwald, W., Colloid Chemistry of Indicators, !Z. Chem. Ind. +Kolloide!, 10 (1912), 132-146. + +[Note 1: Glaser, !Indikatoren der Acidimetrie und Alkalimetrie!. +Wiesbaden, 1901.] + + +PREPARATION OF INDICATOR SOLUTIONS + +A !methyl orange solution! for use as an indicator is commonly made by +dissolving 0.05-0.1 gram of the compound (also known as Orange III) in +a few cubic centimeters of alcohol and diluting with water to 100 cc. +A good grade of material should be secured. It can be successfully +used for the titration of hydrochloric, nitric, sulphuric, phosphoric, +and sulphurous acids, and is particularly useful in the determination +of bases, such as sodium, potassium, barium, calcium, and ammonium +hydroxides, and even many of the weak organic bases. It can also be +used for the determination, by titration with a standard solution of +a strong acid, of the salts of very weak acids, such as carbonates, +sulphides, arsenites, borates, and silicates, because the weak acids +which are liberated do not affect the indicator, and the reddening of +the solution does not take place until an excess of the strong acid +is added. It should be used in cold, not too dilute, solutions. Its +sensitiveness is lessened in the presence of considerable quantities +of the salts of the alkalies. + +A !phenolphthalein solution! is prepared by dissolving 1 gram of the +pure compound in 100 cc. of 95 per cent alcohol. This indicator is +particularly valuable in the determination of weak acids, especially +organic acids. It cannot be used with weak bases, even ammonia. It +is affected by carbonic acid, which must, therefore, be removed by +boiling when other acids are to be measured. It can be used in hot +solutions. Some care is necessary to keep the volume of the solutions +to be titrated approximately uniform in standardization and in +analysis, and this volume should not in general exceed 125-150 cc. for +the best results, since the compounds formed by the indicator undergo +changes in very dilute solution which lessen its sensitiveness. + +The preparation of a !solution of litmus! which is suitable for use +as an indicator involves the separation from the commercial litmus of +azolithmine, the true coloring principle. Soluble litmus tablets are +often obtainable, but the litmus as commonly supplied to the market is +mixed with calcium carbonate or sulphate and compressed into lumps. To +prepare a solution, these are powdered and treated two or three times +with alcohol, which dissolves out certain constituents which cause a +troublesome intermediate color if not removed. The alcohol is decanted +and drained off, after which the litmus is extracted with hot water +until exhausted. The solution is allowed to settle for some time, the +clear liquid siphoned off, concentrated to one-third its volume and +acetic acid added in slight excess. It is then concentrated to a +sirup, and a large excess of 95 per cent. alcohol added to it. This +precipitates the blue coloring matter, which is filtered off, washed +with alcohol, and finally dissolved in a small volume of water and +diluted until about three drops of the solution added to 50 cc. of +water just produce a distinct color. This solution must be kept in an +unstoppered bottle. It should be protected from dust by a loose plug +of absorbent cotton. If kept in a closed bottle it soon undergoes a +reduction and loses its color, which, however, is often restored by +exposure to the air. + +Litmus can be employed successfully with the strong acids and bases, +and also with ammonium hydroxide, although the salts of the latter +influence the indicator unfavorably if present in considerable +concentration. It may be employed with some of the stronger organic +acids, but the use of phenolphthalein is to be preferred. + + +PREPARATION OF STANDARD SOLUTIONS + +!Hydrochloric Acid and Sodium Hydroxide. Approximate Strength!, 0.5 N + + +PROCEDURE.--Measure out 40 cc. of concentrated, pure hydrochloric +acid into a clean liter bottle, and dilute with distilled water to an +approximate volume of 1000 cc. Shake the solution vigorously for a +full minute to insure uniformity. Be sure that the bottle is not too +full to permit of a thorough mixing, since lack of care at this point +will be the cause of much wasted time (Note 1). + +Weigh out, upon a rough balance, 23 grams of sodium hydroxide (Note +2). Dissolve the hydroxide in water in a beaker. Pour the solution +into a liter bottle and dilute, as above, to approximately 1000 cc. +This bottle should preferably have a rubber stopper, as the hydroxide +solution attacks the glass of the ground joint of a glass stopper, and +may cement the stopper to the bottle. Shake the solution as described +above. + +[Note 1: The original solutions are prepared of a strength greater +than 0.5 N, as they are more readily diluted than strengthened if +later adjustment is desired. + +Too much care cannot be taken to insure perfect uniformity of +solutions before standardization, and thoroughness in this respect +will, as stated, often avoid much waste of time. A solution once +thoroughly mixed remains uniform.] + +[Note 2: Commercial sodium hydroxide is usually impure and always +contains more or less carbonate; an allowance is therefore made for +this impurity by placing the weight taken at 23 grams per liter. If +the hydroxide is known to be pure, a lesser amount (say 21 grams) will +suffice.] + + +COMPARISON OF ACID AND ALKALI SOLUTIONS + +PROCEDURE.--Rinse a previously calibrated burette three times with the +hydrochloric acid solution, using 10 cc. each time, and allowing the +liquid to run out through the tip to displace all water and air +from that part of the burette. Then fill the burette with the acid +solution. Carry out the same procedure with a second burette, using +the sodium hydroxide solution. + +The acid solution may be placed in a plain or in a glass-stoppered +burette as may be more convenient, but the alkaline solution should +never be allowed to remain long in a glass-stoppered burette, as it +tends to cement the stopper to the burette, rendering it useless. It +is preferable to use a plain burette for this solution. + +When the burettes are ready for use and all air bubbles displaced from +the tip (see Note 2, page 17) note the exact position of the liquid in +each, and record the readings in the notebook. (Consult page 188.) Run +out from the burette into a beaker about 40 cc. of the acid and add +two drops of a solution of methyl orange; dilute the acid to about +80 cc. and run out alkali solution from the other burette, stirring +constantly, until the pink has given place to a yellow. Wash down the +sides of the beaker with a little distilled water if the solution has +spattered upon them, return the beaker to the acid burette, and add +acid to restore the pink; continue these alternations until the point +is accurately fixed at which a single drop of either solutions served +to produce a distinct change of color. Select as the final end-point +the appearance of the faintest pink tinge which can be recognized, or +the disappearance of this tinge, leaving a pure yellow; but always +titrate to the same point (Note 1). If the titration has occupied more +than the three minutes required for draining the sides of the burette, +the final reading may be taken immediately and recorded in the +notebook. + +Refill the burettes and repeat the titration. From the records of +calibration already obtained, correct the burette readings and make +corrections for temperature, if necessary. Obtain the ratio of the +sodium hydroxide solution to that of hydrochloric acid by dividing +the number of cubic centimeters of acid used by the number of cubic +centimeters of alkali required for neutralization. The check results +of the two titrations should not vary by more than two parts in one +thousand (Note 2). If the variation in results is greater than this, +refill the burettes and repeat the titration until satisfactory values +are obtained. Use a new page in the notebook for each titration. +Inaccurate values should not be erased or discarded. They should be +retained and marked "correct" or "incorrect," as indicated by the +final outcome of the titrations. This custom should be rigidly +followed in all analytical work. + +[Note 1: The end-point should be chosen exactly at the point of +change; any darker tint is unsatisfactory, since it is impossible to +carry shades of color in the memory and to duplicate them from day to +day.] + +[Note 2: While variation of two parts in one thousand in the values +obtained by an inexperienced analyst is not excessive, the idea must +be carefully avoided that this is a standard for accurate work to be +!generally applied!. In many cases, after experience is gained, the +allowable error is less than this proportion. In a few cases a +larger variation is permissible, but these are rare and can only +be recognized by an experienced analyst. It is essential that the +beginner should acquire at least the degree of accuracy indicated if +he is to become a successful analyst.] + + + + +STANDARDIZATION OF HYDROCHLORIC ACID + +SELECTION AND PREPARATION OF STANDARD + + +The selection of the best substance to be used as a standard for acid +solutions has been the subject of much controversy. The work of Lunge +(!Ztschr. angew. Chem.! (1904), 8, 231), Ferguson (!J. Soc. Chem. +Ind.! (1905), 24, 784), and others, seems to indicate that the best +standard is sodium carbonate prepared from sodium bicarbonate by +heating the latter at temperature between 270° and 300°C. The +bicarbonate is easily prepared in a pure state, and at the +temperatures named the decomposition takes place according to the +equation + +2HNaCO_{3} --> Na_{2}CO_{3} + H_{2}O + CO_{2} + +and without loss of any carbon dioxide from the sodium carbonate, such +as may occur at higher temperatures. The process is carried out as +described below. + +PROCEDURE.--Place in a porcelain crucible about 6 grams (roughly +weighed) of the purest sodium bicarbonate obtainable. Rest the +crucible upon a triangle of iron or copper wire so placed within a +large crucible that there is an open air space of about three eighths +of an inch between them. The larger crucible may be of iron, nickel or +porcelain, as may be most convenient. Insert the bulb of a thermometer +reading to 350°C. in the bicarbonate, supporting it with a clamp so +that the bulb does not rest on the bottom of the crucible. Heat +the outside crucible, using a rather small flame, and raise the +temperature of the bicarbonate fairly rapidly to 270°C. Then regulate +the heat in such a way that the temperature rises !slowly! to 300°C. +in the course of a half-hour. The bicarbonate should be frequently +stirred with a clean, dry, glass rod, and after stirring, should be +heaped up around the bulb of the thermometer in such a way as to cover +it. This will require attention during most of the heating, as the +temperature should not be permitted to rise above 310°C. for any +length of time. At the end of the half-hour remove the thermometer and +transfer the porcelain crucible, which now contains sodium carbonate, +to a desiccator. When it is cold, transfer the carbonate to a +stoppered weighing tube or weighing-bottle. + + +STANDARDIZATION + +PROCEDURE.--Clean carefully the outside of a weighing-tube, or +weighing-bottle, containing the pure sodium carbonate, taking care +to handle it as little as possible after wiping. Weigh the tube +accurately to 0.0001 gram, and record the weight in the notebook. Hold +the tube over the top of a beaker (200-300 cc.) and cautiously remove +the stopper, making sure that no particles fall from it or from the +tube elsewhere than in the beaker. Pour out from the tube a portion +of the carbonate, replace the stopper and determine approximately how +much has been removed. Continue this procedure until 1.00 to 1.10 +grams has been taken from the tube. Then weigh the tube accurately +and record the weight under the first weight in the notebook. +The difference in the two weights is the weight of the carbonate +transferred to the beaker. Proceed in the same way to transfer a +second portion of the carbonate from the tube to another beaker of +about the same size as the first. The beakers should be labeled and +plainly marked to correspond with the entries in the notebook. + +Pour over the carbonate in each beaker about 80 cc. of water, stir +until solution is complete, and add two drops of methyl orange +solution. Fill the burettes with the standard acid and alkali +solutions, noting the initial readings of the burettes and temperature +of the solutions. Run in acid from the burette, stirring and avoiding +loss by effervescence, until the solution has become pink. Wash down +the sides of the beaker with a !little! water from a wash-bottle, and +then run in alkali from the other burette until the pink is replaced +by yellow; then finish the titration as described on page 37. Note the +readings of the burettes after the proper interval, and record them in +the notebook. Repeat the procedure, using the second portion of sodium +carbonate. Apply the necessary calibration corrections to the volumes +of the solutions used, and correct for temperature if necessary. + +From the data obtained, calculate the volume of the hydrochloric +acid solution which is equivalent to the volume of sodium hydroxide +solution used in this titration. Subtract this volume from the volume +of hydrochloric acid. The difference represents the volume of acid +used to react with the sodium carbonate. Divide the weight of sodium +carbonate by this volume in cubic centimeters, thus obtaining the +weight of sodium carbonate equivalent to each cubic centimeter of the +acid. + +From this weight it is possible to calculate the corresponding weight +of HCl in each cubic centimeter of the acid, and in turn the relation +of the acid to the normal. + +If, however, it is recalled that normal solutions are equivalent to +each other, it will be seen that the same result may be more readily +reached by dividing the weight in grams of sodium carbonate per cubic +centimeter just found by titration by the weight which would be +contained in the same volume of a normal solution of sodium carbonate. +A normal solution of sodium carbonate contains 53.0 grams per liter, +or 0.0530 gram per cc. (see page 29). The relation of the acid +solution to the normal is, therefore, calculated by dividing the +weight of the carbonate to which each cubic centimeter of the acid is +equivalent by 0.0530. The standardization must be repeated until the +values obtained agree within, at most, two parts in one thousand. + +When the standard of the acid solution has been determined, calculate, +from the known ratio of the two solutions, the relation of the sodium +hydroxide solution to a normal solution (Notes 1 and 2). + +[Note 1: In the foregoing procedure the acid solution is standardized +and the alkali solution referred to this standard by calculation. It +is equally possible, if preferred, to standardize the alkali solution. +The standards in a common use for this purpose are purified +oxalic acid (H_{2}C_{2}O_{4}.2H_{2}O), potassium acid oxalate +(KHC_{2}O_{4}.H_{2}O or KHC_{2}O_{4}), potassium tetroxalate +(KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O), or potassium acid tartrate +(KHC_{4}O_{6}), with the use of a suitable indicator. The oxalic acid +and the oxalates should be specially prepared to insure purity, +the main difficulty lying in the preservation of the water of +crystallization. + +It should be noted that the acid oxalate and the acid tartrate each +contain one hydrogen atom replaceable by a base, while the tetroxalate +contains three such atoms and the oxalic acid two. Each of the two +salts first named behave, therefore, as monobasic acids, and the +tetroxalate as a tribasic acid.] + +[Note 2: It is also possible to standardize a hydrochloric acid +solution by precipitating the chloride ions as silver chloride and +weighing the precipitate, as prescribed under the analysis of sodium +chloride to be described later. Sulphuric acid solutions may be +standardized by precipitation of the sulphate ions as barium sulphate +and weighing the ignited precipitate, but the results are not above +criticism on account of the difficulty in obtaining large precipitates +of barium sulphate which are uncontaminated by inclosures or are not +reduced on ignition.] + + + + +DETERMINATION OF THE TOTAL ALKALINE STRENGTH OF SODA ASH + + +Soda ash is crude sodium carbonate. If made by the ammonia process it +may contain also sodium chloride, sulphate, and hydroxide; when made +by the Le Blanc process it may contain sodium sulphide, silicate, and +aluminate, and other impurities. Some of these, notably the hydroxide, +combine with acids and contribute to the total alkaline strength, +but it is customary to calculate this strength in terms of sodium +carbonate; i.e., as though no other alkali were present. + +PROCEDURE.--In order to secure a sample which shall represent the +average value of the ash, it is well to take at least 5 grams. As this +is too large a quantity for convenient titration, an aliquot portion +of the solution is measured off, representing one fifth of the entire +quantity. This is accomplished as follows: Weigh out on an analytical +balance two samples of soda ash of about 5 grams each into beakers +of about 500 cc. capacity. (The weighings need be made to centigrams +only.) Dissolve the ash in 75 cc. of water, warming gently, and filter +off the insoluble residue; wash the filter by filling it at least +three times with distilled water, and allowing it to drain, adding the +washings to the main filtrate. Cool the filtrate to approximately the +standard temperature of the laboratory, and transfer it to a 250 cc. +measuring flask, washing out the beaker thoroughly. Add distilled +water of laboratory temperature until the lowest point of the meniscus +is level with the graduation on the neck of the flask and remove any +drops of water that may be on the neck above the graduation by means +of a strip of filter paper; make the solution thoroughly uniform by +pouring it out into a dry beaker and back into the flask several +times. Measure off 50 cc. of the solution in a measuring flask, or +pipette, either of which before use should, unless they are dry on the +inside, be rinsed out with at least two small portions of the soda ash +solution to displace any water. + +If a flask is used, fill it to the graduation with the soda ash +solution and remove any liquid from the neck above the graduation with +filter paper. Empty it into a beaker, and wash out the small flask, +unless it is graduated for !delivery!, using small quantities of +water, which are added to the liquid in the beaker. A second 50 cc. +portion from the main solution should be measured off into a second +beaker. Dilute the solutions in each beaker to 100 cc., add two drops +of a solution of methyl orange (Note 1) and titrate for the alkali +with the standard hydrochloric acid solution, using the alkali +solution to complete the titration as already prescribed. + +From the volumes of acid and alkali employed, corrected for burette +errors and temperature changes, and the data derived from the +standardization, calculate the percentage of alkali present, assuming +it all to be present as sodium carbonate (Note 2). + +[Note 1: The hydrochloric acid sets free carbonic acid which is +unstable and breaks down into water and carbon dioxide, most of which +escapes from the solution. Carbonic acid is a weak acid and, as such, +does not yield a sufficient concentration of H^{+} ions to cause the +indicator to change to a pink (see page 32). + +The chemical changes involved may be summarized as follows: + +2H^{+}, 2Cl^{-} + 2Na^{+}, CO_{3}^{--} --> 2Na^{+}, 2Cl^{-} + +[H_{2}CO_{3}] --> H_{2}O + CO_{2}] + +[Note 2: A determination of the alkali present as hydroxide in soda +ash may be determined by precipitating the carbonate by the addition +of barium chloride, removing the barium carbonate by filtration, and +titrating the alkali in the filtrate. + +The caustic alkali may also be determined by first using +phenolphthalein as an indicator, which will show by its change from +pink to colorless the point at which the caustic alkali has been +neutralized and the carbonate has been converted to bicarbonate, and +then adding methyl orange and completing the titration. The amount of +acid necessary to change the methyl orange to pink is a measure of one +half of the carbonate present. The results of the double titration +furnish the data necessary for the determination of the caustic alkali +and of the carbonate in the sample.] + + + + +DETERMINATION OF THE ACID STRENGTH OF OXALIC ACID + + +PROCEDURE.--Weigh out two portions of the acid of about 1 gram +each. Dissolve these in 50 cc. of warm water. Add two drops of +phenolphthalein solution, and run in alkali from the burette until the +solution is pink; add acid from the other burette until the pink is +just destroyed, and then add 0.3 cc. (not more) in excess. Heat the +solution to boiling for three minutes. If the pink returns during the +boiling, discharge it with acid and again add 0.3 cc. in excess and +repeat the boiling (Note 1). If the color does not then reappear, add +alkali until it does, and a !drop or two! of acid in excess and boil +again for one minute (Note 2). If no color reappears during this time, +complete the titration in the hot solution. The end-point should be +the faintest visible shade of color (or its disappearance), as the +same difficulty would exist here as with methyl orange if an attempt +were made to match shades of pink. + +From the corrected volume of alkali required to react with the +oxalic acid, calculate the percentage of the crystallized acid +(H_{2}C_{2}O_{4}.2H_{2}O) in the sample (Note 3). + +[Note 1: All commercial caustic soda such as that from which the +standard solution was made contains some sodium carbonate. This reacts +with the oxalic acid, setting free carbonic acid, which, in turn, +forms sodium bicarbonate with the remaining carbonate: + +H_{2}CO_{3} + Na_{2}CO_{3} --> 2HNaCO_{3}. + +This compound does not hydrolyze sufficiently to furnish enough OH^{-} +ions to cause phenolphthalein to remain pink; hence, the color of +the indicator is discharged in cold solutions at the point at which +bicarbonate is formed. If, however, the solution is heated to boiling, +the bicarbonate loses carbon dioxide and water, and reverts to sodium +carbonate, which causes the indicator to become again pink: + +2HNaCO_{3} --> H_{2}O + CO_{2} + Na_{2}CO_{3}. + +By adding successive portions of hydrochloric acid and boiling, the +carbonate is ultimately all brought into reaction. + +The student should make sure that the difference in behavior of the +two indicators, methyl orange and phenolphthalein, is understood.] + +[Note 2: Hydrochloric acid is volatilized from aqueous solutions, +except such as are very dilute. If the directions in the procedure +are strictly followed, no loss of acid need be feared, but the amount +added in excess should not be greater than 0.3-0.4 cc.] + +[Note 3: Attention has already been called to the fact that the color +changes in the different indicators occur at varying concentrations +of H^{+} or OH^{-} ions. They do not indicate exact theoretical +neutrality, but a particular indicator always shows its color change +at a particular concentration of H^{+} or OH^{-} ions. The results +of titration with a given indicator are, therefore, comparable. As a +matter of fact, a small error is involved in the procedure as outlined +above. The comparison of the acid and alkali solutions was made, using +methyl orange as an indicator, while the titration of the oxalic acid +is made with the use of phenolphthalein. For our present purposes the +small error may be neglected but, if time permits, the student is +recommended to standardize the alkali solution against one of the +substances named in Note 1, page 41, and also to ascertain +the comparative value of the acid and alkali solutions, using +phenolphthalein as indicator throughout, and conducting the titrations +as described above. This will insure complete accuracy.] + + + + +II. OXIDATION PROCESSES + +GENERAL DISCUSSION + + +In the oxidation processes of volumetric analysis standard solutions +of oxidizing agents and of reducing agents take the place of the acid +and alkali solutions of the neutralization processes already studied. +Just as an acid solution was the principal reagent in alkalimetry, and +the alkali solution used only to make certain of the end-point, the +solution of the oxidizing agent is the principal reagent for the +titration of substances exerting a reducing action. It is, in general, +true that oxidizable substances are determined by !direct! titration, +while oxidizing substances are determined by !indirect! titration. + +The important oxidizing agents employed in volumetric solutions are +potassium bichromate, potassium permangenate, potassium ferricyanide, +iodine, ferric chloride, and sodium hypochlorite. + +The important reducing agents which are used in the form of standard +solutions are ferrous sulphate (or ferrous ammonium sulphate), oxalic +acid, sodium thiosulphate, stannous chloride, arsenious acid, and +potassium cyanide. Other reducing agents, as sulphurous acid, +sulphureted hydrogen, and zinc (nascent hydrogen), may take part in +the processes, but not as standard solutions. + +The most important combinations among the foregoing are: Potassium +bichromate and ferrous salts; potassium permanganate and ferrous +salts; potassium permanganate and oxalic acid, or its derivatives; +iodine and sodium thiosulphate; hypochlorites and arsenious acid. + + + + +BICHROMATE PROCESS FOR THE DETERMINATION OF IRON + + +Ferrous salts may be promptly and completely oxidized to ferric salts, +even in cold solution, by the addition of potassium bichromate, +provided sufficient acid is present to hold in solution the ferric and +chromic compounds which are formed. + +The acid may be either hydrochloric or sulphuric, but the former is +usually preferred, since it is by far the best solvent for iron and +its compounds. The reaction in the presence of hydrochloric acid is as +follows: + +6FeCl_{2} + K_{2}Cr_{2}O_{7} + 14HCl --> 6FeCl_{3} + 2CrCl_{3} + 2KCl ++ 7H_{2}O. + + +NORMAL SOLUTIONS OF OXIDIZING OR REDUCING AGENTS + +It will be recalled that the system of normal solutions is based upon +the equivalence of the reagents which they contain to 8 grams of +oxygen or 1 gram of hydrogen. A normal solution of an oxidizing agent +should, therefore, contain that amount per liter which is equivalent +in oxidizing power to 8 grams of oxygen; a normal reducing solution +must be equivalent in reducing power to 1 gram of hydrogen. In order +to determine what the amount per liter will be it is necessary to know +how the reagents enter into reaction. The two solutions to be employed +in the process under consideration are those of potassium bichromate +and ferrous sulphate. The reaction between them, in the presence of an +excess of sulphuric acid, may be expressed as follows: + +6FeSO_{4} + K_{2}Cr_{2}O_{7} + 7H_{2}SO_{4} --> 3Fe_{2}(SO_{4})_{3} + +K_{2}SO_{4} + Cr_{2}(SO_{4})_{3} + 7H_{2}O. + +If the compounds of iron and chromium, with which alone we are now +concerned, be written in such a way as to show the oxides of these +elements in each, they would appear as follows: On the left-hand side +of the equation 6(FeO.SO_{3}) and K_{2}O.2CrO_{3}; on the right-hand +side, 3(Fe_{2}O_{3}.3SO_{3}) and Cr_{2}O_{3}.3SO_{3}. A careful +inspection shows that there are three less oxygen atoms associated +with chromium atoms on the right-hand side of the equation than on the +left-hand, but there are three more oxygen atoms associated with iron +atoms on the right than on the left. In other words, a molecule of +potassium bichromate has given up three atoms of oxygen for oxidation +purposes; i.e., a molecular weight in grams of the bichromate (294.2) +will furnish 3 X 16 or 48 grams of oxygen for oxidation purposes. +As this 48 grams is six times 8 grams, the basis of the system, the +normal solution of potassium bichromate should contain per liter one +sixth of 294.2 grams or 49.03 grams. + +A further inspection of the dissected compounds above shows that six +molecules of FeO.SO_{3} were required to react with the three atoms of +oxygen from the bichromate. From the two equations + +3H_{2} + 3O --> 3H_{2}O +6(FeO.SO_{3}) + 3O --> 3(Fe_{2}O_{3}.3SO_{3}) + +it is plain that one molecule of ferrous sulphate is equivalent to one +atom of hydrogen in reducing power; therefore one molecular weight in +grams of ferrous sulphate (151.9) is equivalent to 1 gram of +hydrogen. Since the ferrous sulphate crystalline form has the formula +FeSO_{4}.7H_{2}O, a normal reducing solution of this crystalline salt +should contain 277.9 grams per liter. + + +PREPARATION OF SOLUTIONS + +!Approximate Strength 0.1 N! + +It is possible to purify commercial potassium bichromate by +recrystallization from hot water. It must then be dried and cautiously +heated to fusion to expel the last traces of moisture, but not +sufficiently high to expel any oxygen. The pure salt thus prepared, +may be weighed out directly, dissolved, and the solution diluted in a +graduated flask to a definite volume. In this case no standardization +is made, as the normal value can be calculated directly. It is, +however, more generally customary to standardize a solution of +the commercial salt by comparison with some substance of definite +composition, as described below. + +PROCEDURE.--Pulverize about 5 grams of potassium bichromate of good +quality. Dissolve the bichromate in distilled water, transfer the +solution to a liter bottle, and dilute to approximately 1000 cc. Shake +thoroughly until the solution is uniform. + +To prepare the solution of the reducing agent, pulverize about 28 +grams of ferrous sulphate (FeSO_{4}.7H_{2}O) or about 40 grams of +ferrous ammonium sulphate (FeSO_{4}.(NH_{4})_{2}SO_{4}.6H_{2}O) and +dissolve in distilled water containing 5 cc. of concentrated sulphuric +acid. Transfer the solution to a liter bottle, add 5 cc. concentrated +sulphuric acid, make up to about 1000 cc. and shake vigorously to +insure uniformity. + + +INDICATOR SOLUTION + +No indicator is known which, like methyl orange, can be used within +the solution, to show when the oxidation process is complete. Instead, +an outside indicator solution is employed to which drops of the +titrated solution are transferred for testing. The reagent used is +potassium ferricyanide, which produces a blue precipitate (or color) +with ferrous compounds as long as there are unoxidized ferrous ions in +the titrated solution. Drops of the indicator solution are placed upon +a glazed porcelain tile, or upon white cardboard which has been coated +with paraffin to render it waterproof, and drops of the titrated +solution are transferred to the indicator on the end of a stirring +rod. When the oxidation is nearly completed only very small amounts +of the ferrous compounds remain unoxidized and the reaction with the +indicator is no longer instantaneous. It is necessary to allow a brief +time to elapse before determining that no blue color is formed. Thirty +seconds is a sufficient interval, and should be adopted throughout the +analytical procedure. If left too long, the combined effect of light +and dust from the air will cause a reduction of the ferric compounds +already formed and a resultant blue will appear which misleads the +observer with respect to the true end-point. + +The indicator solution must be highly diluted, otherwise its own color +interferes with accurate observation. Prepare a fresh solution, as +needed each day, by dissolving a crystal of potassium ferricyanide +about the size of a pin's head in 25 cc. of distilled water. The salt +should be carefully tested with ferric chloride for the presence of +ferrocyanides, which give a blue color with ferric salts. + +In case of need, the ferricyanide can be purified by adding to its +solution a little bromine water and recrystallizing the compound. + + +COMPARISON OF OXIDIZING AND REDUCING SOLUTIONS + +PROCEDURE.--Fill one burette with each of the solutions, observing +the general procedure with respect to cleaning and rinsing already +prescribed. The bichromate solution is preferably to be placed in a +glass-stoppered burette. + +Run out from a burette into a beaker of about 300 cc. capacity nearly +40 cc. of the ferrous solution, add 15 cc. of dilute hydrochloric acid +(sp. gr. 1.12) and 150 cc. of water and run in the bichromate +solution from another burette. Since both solutions are approximately +tenth-normal, 35 cc. of the bichromate solution may be added without +testing. Test at that point by removing a very small drop of the +iron solution on the end of a stirring rod, mixing it with a drop of +indicator on the tile (Note 1). If a blue precipitate appears at once, +0.5 cc. of the bichromate solution may be added before testing again. +The stirring rod which has touched the indicator should be dipped in +distilled water before returning it to the iron solution. As soon as +the blue appears to be less intense, add the bichromate solution in +small portions, finally a single drop at a time, until the point is +reached at which no blue color appears after the lapse of thirty +seconds from the time of mixing solution and indicator. At the close +of the titration a large drop of the iron solution should be taken for +the test. To determine the end-point beyond any question, as soon as +the thirty seconds have elapsed remove another drop of the solution +of the same size as that last taken and mix it with the indicator, +placing it beside the last previous test. If this last previous test +shows a blue tint in comparison with the fresh mixture, the end-point +has not been reached; if no difference can be noted the reaction is +complete. Should the end-point be overstepped, a little more of the +ferrous solution may be added and the end-point definitely fixed. + +From the volumes of the solutions used, after applying corrections for +burette readings, and, if need be, for the temperature of solutions, +calculate the value of the ferrous solution in terms of the oxidizing +solution. + +[Note 1: The accuracy of the work may be much impaired by the removal +of unnecessarily large quantities of solution for the tests. At the +beginning of the titration, while much ferrous iron is still present, +the end of the stirring rod need only be moist with the solution; but +at the close of the titration drops of considerable size may properly +be taken for the final tests. The stirring rod should be washed to +prevent transfer of indicator to the main solution. This cautious +removal of solution does not seriously affect the accuracy of the +determination, as it will be noted that the volume of the titrated +solution is about 200 cc. and the portions removed are very +small. Moreover, if the procedure is followed as prescribed, the +concentration of unoxidized iron decreases very rapidly as the +titration is carried out so that when the final tests are made, though +large drops may be taken, the amount of ferrous iron is not sufficient +to produce any appreciable error in results. + +If the end-point is determined as prescribed, it can be as accurately +fixed as that of other methods; and if a ferrous solution is at +hand, the titration need consume hardly more time than that of the +permanganate process to be described later on.] + + +STANDARDIZATION OF POTASSIUM BICHROMATE SOLUTIONS + +!Selection of a Standard! + +A substance which will serve satisfactorily as a standard for +oxidizing solutions must possess certain specific properties: It must +be of accurately known composition and definite in its behavior as a +reducing agent, and it must be permanent against oxidation in the air, +at least for considerable periods. Such standards may take the form of +pure crystalline salts, such as ferrous ammonium sulphate, or may be +in the form of iron wire or an iron ore of known iron content. It is +not necessary that the standard should be of 100 per cent purity, +provided the content of the active reducing agent is known and no +interfering substances are present. + +The two substances most commonly used as standards for a bichromate +solution are ferrous ammonium sulphate and iron wire. A standard wire +is to be purchased in the market which answers the purpose well, and +its iron content may be determined for each lot purchased by a number +of gravimetric determinations. It may best be preserved in jars +containing calcium chloride, but this must not be allowed to come +into contact with the wire. It should, however, even then be examined +carefully for rust before use. + +If pure ferrous ammonium sulphate is used as the standard, clear +crystals only should be selected. It is perhaps even better to +determine by gravimetric methods once for all the iron content of a +large commercial sample which has been ground and well mixed. This +salt is permanent over long periods if kept in stoppered containers. + + +STANDARDIZATION + +PROCEDURE.--Weigh out two portions of iron wire of about 0.24-0.26 +gram each, examining the wire carefully for rust. It should be handled +and wiped with filter paper (not touched by the fingers), should +be weighed on a watch-glass, and be bent in such a way as not to +interfere with the movement of the balance. + +Place 30 cc. of hydrochloric acid (sp. gr. 1.12) in each of two 300 +cc. Erlenmeyer flasks, cover them with watch-glasses, and bring the +acid just to boiling. Remove them from the flame and drop in the +portions of wire, taking great care to avoid loss of liquid during +solution. Boil for two or three minutes, keeping the flasks covered +(Note 1), then wash the sides of the flasks and the watch-glass with +a little water and add stannous chloride solution to the hot liquid +!from a dropper! until the solution is colorless, but avoid more than +a drop or two in excess (Note 2). Dilute with 150 cc. of water and +cool !completely!. When cold, add rapidly about 30 cc. of mercuric +chloride solution. Allow the solutions to stand about three minutes +and then titrate without further delay (Note 3), add about 35 cc. of +the standard solution at once and finish the titration as prescribed +above, making use of the ferrous solution if the end-point should be +passed. + +From the corrected volumes of the bichromate solution required to +oxidize the iron actually know to be present in the wire, calculate +the relation of the standard solution to the normal. + +Repeat the standardization until the results are concordant within at +least two parts in one thousand. + + +[Note 1: The hydrochloric acid is added to the ferrous solution +to insure the presence of at least sufficient free acid for the +titration, as required by the equation on page 48. + +The solution of the wire in hot acid and the short boiling insure the +removal of compounds of hydrogen and carbon which are formed from the +small amount of carbon in the iron. These might be acted upon by the +bichromate if not expelled.] + +[Note 2: It is plain that all the iron must be reduced to the ferrous +condition before the titration begins, as some oxidation may have +occurred from the oxygen of the air during solution. It is also +evident that any excess of the agent used to reduce the iron must be +removed; otherwise it will react with the bichromate added later. + +The reagents available for the reduction of iron are stannous +chloride, sulphurous acid, sulphureted hydrogen, and zinc; of these +stannous chloride acts most readily, the completion of the reaction +is most easily noted, and the excess of the reagent is most readily +removed. The latter object is accomplished by oxidation to stannic +chloride by means of mercuric chloride added in excess, as the +mercuric salts have no effect upon ferrous iron or the bichromate. The +reactions involved are: + +2FeCl_{3} + SnCl_{2} --> 2FeCl_{2} + SnCl_{4} +SnCl_{2} + 2HgCl_{2} --> SnCl_{4} + 2HgCl + +The mercurous chloride is precipitated. + +It is essential that the solution should be cold and that the stannous +chloride should not be present in great excess, otherwise a secondary +reaction takes place, resulting in the reduction of the mercurous +chloride to metallic mercury: + +SnCl_{2} + 2HgCl --> SnCl_{4} + 2Hg. + +The occurrence of this secondary reaction is indicated by the +darkening of the precipitate; and, since potassium bichromate oxidizes +this mercury slowly, solutions in which it has been precipitated are +worthless as iron determinations.] + +[Note 3: The solution should be allowed to stand about three minutes +after the addition of mercuric chloride to permit the complete +deposition of mercurous chloride. It should then be titrated without +delay to avoid possible reoxidation of the iron by the oxygen of the +air.] + + + + +DETERMINATION OF IRON IN LIMONITE + + +PROCEDURE.--Grind the mineral (Note 1) to a fine powder. Weigh out +accurately two portions of about 0.5 gram (Note 2) into porcelain +crucibles; heat these crucibles to dull redness for ten minutes, +allow them to cool, and place them, with their contents, in beakers +containing 30 cc. of dilute hydrochloric acid (sp. gr. 1.12). Heat +at a temperature just below boiling until the undissolved residue is +white or until solvent action has ceased. If the residue is white, +or known to be free from iron, it may be neglected and need not be +removed by filtration. If a dark residue remains, collect it on a +filter, wash free from hydrochloric acid, and ignite the filter in a +platinum crucible (Note 3). Mix the ash with five times its weight of +sodium carbonate and heat to fusion; cool, and disintegrate the fused +mass with boiling water in the crucible. Unite this solution and +precipitate (if any) with the acid solution, taking care to avoid loss +by effervescence. Wash out the crucible, heat the acid solution +to boiling, add stannous chloride solution until it is colorless, +avoiding a large excess (Note 4); cool, and when !cold!, add 40 cc. of +mercuric chloride solution, dilute to 200 cc., and proceed with the +titration as already described. + +From the standardization data already obtained, and the known weight +of the sample, calculate the percentage of iron (Fe) in the limonite. + +[Note 1: Limonite is selected as a representative of iron ores in +general. It is a native, hydrated oxide of iron. It frequently occurs +in or near peat beds and contains more or less organic matter which, +if brought into solution, would be acted upon by the potassium +bichromate. This organic matter is destroyed by roasting. Since a high +temperature tends to lessen the solubility of ferric oxide, the heat +should not be raised above low redness.] + +[Note 2: It is sometimes advantageous to dissolve a large portion--say +5 grams--and to take one tenth of it for titration. The sample will +then represent more closely the average value of the ore.] + +[Note 3: A platinum crucible may be used for the roasting of the +limonite and must be used for the fusion of the residue. When used, it +must not be allowed to remain in the acid solution of ferric chloride +for any length of time, since the platinum is attacked and dissolved, +and the platinic chloride is later reduced by the stannous chloride, +and in the reduced condition reacts with the bichromate, thus +introducing an error. It should also be noted that copper and antimony +interfere with the determination of iron by the bichromate process.] + +[Note 4: The quantity of stannous chloride required for the reduction +of the iron in the limonite will be much larger than that added to the +solution of iron wire, in which the iron was mainly already in the +ferrous condition. It should, however, be added from a dropper to +avoid an unnecessary excess.] + + + + +DETERMINATION OF CHROMIUM IN CHROME IRON ORE + + +PROCEDURE.--Grind the chrome iron ore (Note 1) in an agate mortar +until no grit is perceptible under the pestle. Weigh out two portions +of 0.5 gram each into iron crucibles which have been scoured inside +until bright (Note 2). Weigh out on a watch-glass (Note 3), using the +rough balances, 5 grams of dry sodium peroxide for each portion, and +pour about three quarters of the peroxide upon the ore. Mix ore and +flux by thorough stirring with a dry glass rod. Then cover the mixture +with the remainder of the peroxide. Place the crucible on a triangle +and raise the temperature !slowly! to the melting point of the flux, +using a low flame, and holding the lamp in the hand (Note 4). Maintain +the fusion for five minutes, and stir constantly with a stout iron +wire, but do not raise the temperature above moderate redness (Notes 5 +and 6). + +Allow the crucible to cool until it can be comfortably handled (Note +7) and then place it in a 300 cc. beaker, and cover it with distilled +water (Note 8). The beaker must be carefully covered to avoid loss +during the disintegration of the fused mass. When the evolution of +gas ceases, rinse off and remove the crucible; then heat the solution +!while still alkaline! to boiling for fifteen minutes. Allow the +liquid to cool for a few minutes; then acidify with dilute sulphuric +acid (1:5), adding 10 cc. in excess of the amount necessary to +dissolve the ferric hydroxide (Note 9). Dilute to 200 cc., cool, add +from a burette an excess of a standard ferrous solution, and titrate +for the excess with a standard solution of potassium bichromate, using +the outside indicator (Note 10). + +From the corrected volumes of the two standard solutions, and their +relations to normal solutions, calculate the percentage of chromium in +the ore. + +[Note 1: Chrome iron ore is essentially a ferrous chromite, or +combination of FeO and Cr_{2}O_{3}. It must be reduced to a state of +fine subdivision to ensure a prompt reaction with the flux.] + +[Note 2: The scouring of the iron crucible is rendered much easier if +it is first heated to bright redness and plunged into cold water. In +this process oily matter is burned off and adhering scale is caused to +chip off when the hot crucible contracts rapidly in the cold water.] + +[Note 3: Sodium peroxide must be kept off of balance pans and should +not be weighed out on paper, as is the usual practice in the rough +weighing of chemicals. If paper to which the peroxide is adhering is +exposed to moist air it is likely to take fire as a result of +the absorption of moisture, and consequent evolution of heat and +liberation of oxygen.] + +[Note 4: The lamp should never be allowed to remain under the +crucible, as this will raise the temperature to a point at which the +crucible itself is rapidly attacked by the flux and burned through.] + +[Note 5: The sodium peroxide acts as both a flux and an oxidizing +agent. The chromic oxide is dissolved by the flux and oxidized to +chromic anhydride (CrO_{3}) which combines with the alkali to form +sodium chromate. The iron is oxidized to ferric oxide.] + +[Note 6: The sodium peroxide cannot be used in porcelain, platinum, or +silver crucibles. It attacks iron and nickel as well; but crucibles +made from these metals may be used if care is exercised to keep the +temperature as low as possible. Preference is here given to iron +crucibles, because the resulting ferric hydroxide is more readily +brought into solution than the nickelic oxide from a nickel crucible. +The peroxide must be dry, and must be protected from any admixture of +dust, paper, or of organic matter of any kind, otherwise explosions +may ensue.] + +[Note 7: When an iron crucible is employed it is desirable to allow +the fusion to become nearly cold before it is placed in water, +otherwise scales of magnetic iron oxide may separate from the +crucible, which by slowly dissolving in acid form ferrous sulphate, +which reduces the chromate.] + +[Note 8: Upon treatment with water the chromate passes into solution, +the ferric hydroxide remains undissolved, and the excess of peroxide +is decomposed with the evolution of oxygen. The subsequent boiling +insures the complete decomposition of the peroxide. Unless this is +complete, hydrogen peroxide is formed when the solution is acidified, +and this reacts with the bichromate, reducing it and introducing a +serious error.] + +[Note 9: The addition of the sulphuric acid converts the sodium +chromate to bichromate, which behaves exactly like potassium +bichromate in acid solution.] + +[Note 10: If a standard solution of a ferrous salt is not at hand, a +weight of iron wire somewhat in excess of the amount which would be +required if the chromite were pure FeO.Cr_{2}O_{3} may be weighed out +and dissolved in sulphuric acid; after reduction of all the iron by +stannous chloride and the addition of mercuric chloride, this solution +may be poured into the chromate solution and the excess of iron +determined by titration with standard bichromate solution.] + + + + +PERMANGANATE PROCESS FOR THE DETERMINATION OF IRON + + +Potassium permanganate oxidizes ferrous salts in cold, acid solution +promptly and completely to the ferric condition, while in hot acid +solution it also enters into a definite reaction with oxalic acid, by +which the latter is oxidized to carbon dioxide and water. + +The reactions involved are these: + +10FeSO_{4} + 2KMnO_{4} + 8H_{2}S_{4} --> 5Fe_{2}(SO_{4})_{3} + +K_{2}SO_{4} + 2MnSO_{4} + 8H_{2}O + +5C_{2}H_{2}O_{4}(2H_{2}O) + 2KMnO_{4} +3H_{2}SO_{4} --> K_{2}SO_{4} + +2MnSO_{4} + 10CO_{2} + 1 H_{2}O. + +These are the fundamental reactions upon which the extensive use of +potassium permanganate depends; but besides iron and oxalic acid the +permanganate enters into reaction with antimony, tin, copper, mercury, +and manganese (the latter only in neutral solution), by which these +metals are changed from a lower to a higher state of oxidation; and it +also reacts with sulphurous acid, sulphureted hydrogen, nitrous acid, +ferrocyanides, and most soluble organic bodies. It should be noted, +however, that very few of these organic compounds react quantitatively +with the permanganate, as is the case with oxalic acid and the +oxalates. + +Potassium permanganate is acted upon by hydrochloric acid; the action +is rapid in hot or concentrated solution (particularly in the presence +of iron salts, which appear to act as catalyzers, increasing the +velocity of the reaction), but slow in cold, dilute solutions. +However, the greater solubility of iron compounds in hydrochloric acid +makes it desirable to use this acid as a solvent, and experiments made +with this end in view have shown that in cold, dilute hydrochloric +acid solution, to which considerable quantities of manganous sulphate +and an excess of phosphoric acid have been added, it is possible to +obtain satisfactory results. + +It is also possible to replace the hydrochloric acid by evaporating +the solutions with an excess of sulphuric acid until the latter fumes. +This procedure is somewhat more time-consuming, but the end-point of +the permanganate titration is more permanent. Both procedures are +described below. + +Potassium permanganate has an intense coloring power, and since the +solution resulting from the oxidation of the iron and the reduction of +the permanganate is colorless, the latter becomes its own indicator. +The slightest excess is indicated with great accuracy by the pink +color of the solution. + + +PREPARATION OF A STANDARD SOLUTION + +!Approximate Strength 0.1 N! + +A study of the reactions given above which represent the oxidation of +ferrous compounds by potassium permanganate, shows that there are 2 +molecules of KMnO_{4} and 10 molecules of FeSO_{4} on the +left-hand side, and 2 molecules of MnSO_{4} and 5 molecules of +Fe_{2}(SO_{4})_{5} on the right-hand side. Considering only these +compounds, and writing the formulas in such a way as to show the +oxides of the elements in each, the equation becomes: + +K_{2}O.Mn_{2}O_{7} + 10(FeO.SO_{3}) --> K_{2}O.SO_{3} + 2(MnO.SO_{3}) ++ 5(Fe_{2}O_{3}.3SO_{3}). + +From this it appears that two molecules of KMnO_{4} (or 316.0 grams) +have given up five atoms (or 80 grams) of oxygen to oxidize the +ferrous compound. Since 8 grams of oxygen is the basis of normal +oxidizing solutions and 80 grams of oxygen are supplied by 316.0 grams +of KMnO_{4}, the normal solution of the permanganate should contain, +per liter, 316.0/10 grams, or 31.60 grams (Note 1). + +The preparation of an approximately tenth-normal solution of the +reagent may be carried out as follows: + +PROCEDURE.--Dissolve about 3.25 grams of potassium permanganate +crystals in approximately 1000 cc. of distilled water in a large +beaker, or casserole. Heat slowly and when the crystals have +dissolved, boil the solution for 10-15 minutes. Cover the solution +with a watch-glass; allow it to stand until cool, or preferably over +night. Filter the solution through a layer of asbestos. Transfer the +filtrate to a liter bottle and mix thoroughly (Note 2). + +[Note 1: The reactions given on page 61 are those which take place in +the presence of an excess of acid. In neutral solutions the reduction +of the permanganate is less complete, and, under these conditions, +two gram-molecular weights of KMnO_{4} will furnish only 48 grams +of oxygen. A normal solution for use under these conditions should, +therefore, contain 316.0/6 grams, or 52.66 grams.] + +[Note 2: Potassium permanganate solutions are not usually stable for +long periods, and change more rapidly when first prepared than after +standing some days. This change is probably caused by interaction +with the organic matter contained in all distilled water, except that +redistilled from an alkaline permanganate solution. The solutions +should be protected from light and heat as far as possible, since both +induce decomposition with a deposition of manganese dioxide, and it +has been shown that decomposition proceeds with considerable rapidity, +with the evolution of oxygen, after the dioxide has begun to form. As +commercial samples of the permanganate are likely to be contaminated +by the dioxide, it is advisable to boil and filter solutions through +asbestos before standardization, as prescribed above. Such solutions +are relatively stable.] + + +COMPARISON OF PERMANGANATE AND FERROUS SOLUTIONS + +PROCEDURE.--Fill a glass-stoppered burette with the permanganate +solution, observing the usual precautions, and fill a second burette +with the ferrous sulphate solution prepared for use with the potassium +bichromate. The permanganate solution cannot be used in burettes with +rubber tips, as a reduction takes place upon contact with the rubber. +The solution has so deep a color that the lower line of the meniscus +cannot be detected; readings must therefore be made from the upper +edge. Run out into a beaker about 40 cc. of the ferrous solution, +dilute to about 100 cc., add 10 cc. of dilute sulphuric acid, and run +in the permanganate solution to a slight permanent pink. Repeat, until +the ratio of the two solutions is satisfactorily established. + + +STANDARDIZATION OF A POTASSIUM PERMANGANATE SOLUTION + +!Selection of a Standard! + +Commercial potassium permanganate is rarely sufficiently pure to admit +of its direct weighing as a standard. On this account, and because +of the uncertainties as to the permanence of its solutions, it is +advisable to standardize them against substances of known value. Those +in most common use are iron wire, ferrous ammonium sulphate, sodium +oxalate, oxalic acid, and some other derivatives of oxalic acid. +With the exception of sodium oxalate, these all contain water of +crystallization which may be lost on standing. They should, therefore, +be freshly prepared, and with great care. At present, sodium oxalate +is considered to be one of the most satisfactory standards. + + +!Method A! + + +!Iron Standards! + +The standardization processes employed when iron or its compounds are +selected as standards differ from those applicable in connection with +oxalate standards. The procedure which immediately follows is that in +use with iron standards. + +As in the case of the bichromate process, it is necessary to reduce +the iron completely to the ferrous condition before titration. The +reducing agents available are zinc, sulphurous acid, or sulphureted +hydrogen. Stannous chloride may also be used when the titration is +made in the presence of hydrochloric acid. Since the excess of both +the gaseous reducing agents can only be expelled by boiling, with +consequent uncertainty regarding both the removal of the excess and +the reoxidation of the iron, zinc or stannous chlorides are the most +satisfactory agents. For prompt and complete reduction it is essential +that the iron solution should be brought into ultimate contact with +the zinc. This is brought about by the use of a modified Jones +reductor, as shown in Figure 1. This reductor is a standard apparatus +and is used in other quantitative processes. + +[Illustration: Fig. 1] + +The tube A has an inside diameter of 18 mm. and is 300 mm. long; the +small tube has an inside diameter of 6 mm. and extends 100 mm. below +the stopcock. At the base of the tube A are placed some pieces of +broken glass or porcelain, covered by a plug of glass wool about 8 mm. +thick, and upon this is placed a thin layer of asbestos, such as is +used for Gooch filters, 1 mm. thick. The tube is then filled with the +amalgamated zinc (Note 1) to within 50 mm. of the top, and on the zinc +is placed a plug of glass wool. If the top of the tube is not already +shaped like the mouth of a thistle-tube (B), a 60 mm. funnel is fitted +into the tube with a rubber stopper and the reductor is connected +with a suction bottle, F. The bottle D is a safety bottle to +prevent contamination of the solution by water from the pump. After +preparation for use, or when left standing, the tube A should be +filled with water, to prevent clogging of the zinc. + +[Note 1: The use of fine zinc in the reductor is not necessary and +tends to clog the tube. Particles which will pass a 10-mesh sieve, but +are retained by one of 20 meshes to the inch, are most satisfactory. +The zinc can be amalgamated by stirring or shaking it in a mixture of +25 cc. of normal mercuric chloride solution, 25 cc. of hydrochloric +acid (sp. gr. 1.12) and 250 cc. of water for two minutes. The solution +should then be poured off and the zinc thoroughly washed. It is then +ready for bottling and preservation under water. A small quantity of +glass wool is placed in the neck of the funnel to hold back foreign +material when the reductor is in use.] + + +STANDARDIZATION + +PROCEDURE.--Weigh out into Erlenmeyer flasks two portions of iron wire +of about 0.25 gram each. Dissolve these in hot dilute sulphuric acid +(5 cc. of concentrated acid and 100 cc. of water), using a covered +flask to avoid loss by spattering. Boil the solution for two or +three minutes after the iron has dissolved to remove any volatile +hydrocarbons. Meanwhile prepare the reductor for use as follows: +Connect the vacuum bottle with the suction pump and pour into the +funnel at the top warm, dilute sulphuric acid, prepared by adding 5 +cc. of concentrated sulphuric acid to 100 cc. of distilled water. See +that the stopcock (C) is open far enough to allow the acid to run +through slowly. Continue to pour in acid until 200 cc. have passed +through, then close the stopcock !while a small quantity of liquid +is still left in the funnel!. Discard the filtrate, and again +pass through 100 cc. of the warm, dilute acid. Test this with the +permanganate solution. A single drop should color it permanently; if +it does not, repeat the washing, until assured that the zinc is not +contaminated with appreciable quantities of reducing substances. Be +sure that no air enters the reductor (Note 1). + +Pour the iron solution while hot (but not boiling) through the +reductor at a rate not exceeding 50 cc. per minute (Notes 2 and 3). +Wash out the beaker with dilute sulphuric acid, and follow the iron +solution without interruption with 175 cc. of the warm acid and +finally with 75 cc. of distilled water, leaving the funnel partially +filled. Remove the filter bottle and cool the solution quickly under +the water tap (Note 4), avoiding unnecessary exposure to the oxygen of +the air. Add 10 cc. of dilute sulphuric acid and titrate to a faint +pink with the permanganate solution, adding it directly to the +contents of the vacuum flask. Should the end-point be overstepped, the +ferrous sulphate solution may be added. + +From the volume of the solution required to oxidize the iron in +the wire, calculate the relation to the normal of the permanganate +solution. The duplicate results should be concordant within two parts +in one thousand. + +[Note 1: The funnel of the reductor must never be allowed to empty. +If it is left partially filled with water the reductor is ready for +subsequent use after a very little washing; but a preliminary test is +always necessary to safeguard against error. + +If more than a small drop of permanganate solution is required to +color 100 cc. of the dilute acid after the reductor is well washed, an +allowance must be made for the iron in the zinc. !Great care! must be +used to prevent the access of air to the reductor after it has been +washed out ready for use. If air enters, hydrogen peroxide forms, +which reacts with the permanganate, and the results are worthless.] + +[Note 2: The iron is reduced to the ferrous condition by contact with +the zinc. The active agent may be considered to be !nascent! hydrogen, +and it must be borne in mind that the visible bubbles are produced by +molecular hydrogen, which is without appreciable effect upon ferric +iron. + +The rate at which the iron solution passes through the zinc should not +exceed that prescribed, but the rate may be increased somewhat when +the wash-water is added. It is well to allow the iron solution to run +nearly, but not entirely, out of the funnel before the wash-water +is added. If it is necessary to interrupt the process, the complete +emptying of the funnel can always be avoided by closing the stopcock. + +It is also possible to reduce the iron by treatment with zinc in a +flask from which air is excluded. The zinc must be present in excess +of the quantity necessary to reduce the iron and is finally completely +dissolved. This method is, however, less convenient and more tedious +than the use of the reductor.] + +[Note 3: The dilute sulphuric acid for washing must be warmed ready +for use before the reduction of the iron begins, and it is of the +first importance that the volume of acid and of wash-water should +be measured, and the volume used should always be the same in the +standardizations and all subsequent analyses.] + +[Note 4: The end-point is more permanent in cold than hot solutions, +possibly because of a slight action of the permanganate upon the +manganous sulphate formed during titration. If the solution turns +brown, it is an evidence of insufficient acid, and more should be +immediately added. The results are likely to be less accurate in this +case, however, as a consequence of secondary reactions between the +ferrous iron and the manganese dioxide thrown down. It is wiser to +discard such results and repeat the process.] + +[Note 5: The potassium permanganate may, of course, be diluted and +brought to an exactly 0.1 N solution from the data here obtained. The +percentage of iron in the iron wire must be taken into account in all +calculations.] + + +!Method B! + +!Oxalate Standards! + +PROCEDURE.--Weigh out two portions of pure sodium oxalate of 0.25-0.3 +gram each into beakers of about 600 cc. capacity. Add about 400 cc. of +boiling water and 20 cc. of manganous sulphate solution (Note 1). +When the solution of the oxalate is complete, heat the liquid, if +necessary, until near its boiling point (70-90°C.) and run in the +standard permanganate solution drop by drop from a burette, stirring +constantly until an end-point is reached (Note 2). Make a blank test +with 20 cc. of manganous sulphate solution and a volume of distilled +water equal to that of the titrated solution to determine the volume +of the permanganate solution required to produce a very slight pink. +Deduct this volume from the amount of permanganate solution used in +the titration. + +From the data obtained, calculate the relation of the permanganate +solution to the normal. The reaction involved is: + +5Na_{2}C_{2}O_{4} + 2KMnO_{4} + 8H_{2}SO_{4} --> 5Na_{2}SO_{4} + +K_{2}SO_{4} + 2MnSO_{4} + 10CO_{2} + 8H_{2}O + +[Note 1: The manganous sulphate titrating solution is made by +dissolving 20 grams of MnSO_{4} in 200 cubic centimeters of water and +adding 40 cc. of concentrated sulphuric acid (sp. gr. 1.84) and 40 cc. +or phosphoric acid (85%).] + +[Note 2: The reaction between oxalates and permanganates takes place +quantitatively only in hot acid solutions. The temperatures must not +fall below 70°C.] + + + + +DETERMINATION OF IRON IN LIMONITE + + +!Method A! + +The procedures, as here prescribed, are applicable to iron ores in +general, provided these ores contain no constituents which are reduced +by zinc or stannous chloride and reoxidized by permanganates. Many +iron ores contain titanium, and this element among others does +interfere with the determination of iron by the process described. +If, however, the solutions of such ores are treated with sulphureted +hydrogen or sulphurous acid, instead of zinc or stannous chloride to +reduce the iron, and the excess reducing agent removed by boiling, an +accurate determination of the iron can be made. + +PROCEDURE.--Grind the mineral to a fine powder. Weigh out two portions +of about 0.5 gram each into small porcelain crucibles. Roast the ore +at dull redness for ten minutes (Note 1), allow the crucibles to cool, +and place them and their contents in casseroles containing 30 cc. of +dilute hydrochloric acid (sp. gr. 1.12). + +Proceed with the solution of the ore, and the treatment of the +residue, if necessary, exactly as described for the bichromate process +on page 56. When solution is complete, add 6 cc. of concentrated +sulphuric acid to each casserole, and evaporate on the steam bath +until the solution is nearly colorless (Note 2). Cover the casseroles +and heat over the flame of the burner, holding the casserole in +the hand and rotating it slowly to hasten evaporation and prevent +spattering, until the heavy white fumes of sulphuric anhydride are +freely evolved (Note 3). Cool the casseroles, add 100 cc. of water +(measured), and boil gently until the ferric sulphate is dissolved; +pour the warm solution through the reductor which has been previously +washed; proceed as described under standardization, taking pains +to use the same volume and strength of acid and the same volume of +wash-water as there prescribed, and titrate with the permanganate +solution in the reductor flask, using the ferrous sulphate solution if +the end-point should be overstepped. + +From the corrected volume of permanganate solution used, calculate the +percentage of iron (Fe) in the limonite. + +[Note 1: The preliminary roasting is usually necessary because, even +though the sulphuric acid would subsequently char the carbonaceous +matter, certain nitrogenous bodies are not thereby rendered insoluble +in the acid, and would be oxidized by the permanganate.] + +[Note 2: The temperature of the steam bath is not sufficient to +volatilize sulphuric acid. Solutions may, therefore, be left to +evaporate overnight without danger of evaporation to dryness.] + +[Note 3: The hydrochloric acid, both free and combined, is displaced +by the less volatile sulphuric acid at its boiling point. Ferric +sulphate separates at this point, since there is no water to hold +it in solution and care is required to prevent bumping. The ferric +sulphate usually has a silky appearance and is easily distinguished +from the flocculent silica which often remains undissolved.] + + +!Zimmermann-Reinhardt Procedure! + + +!Method (B)! + +PROCEDURE.--Grind the mineral to a fine powder. Weigh out two portions +of about 0.5 gram each into small porcelain crucibles. Proceed with +the solution of the ore, treat the residue, if necessary, and reduce +the iron by the addition of stannous chloride, followed by mercuric +chloride, as described for the bichromate process on page 56. Dilute +the solution to about 400 cc. with cold water, add 10 cc. of the +manganous sulphate titrating solution (Note 1, page 68) and titrate +with the standard potassium permanganate solution to a faint pink +(Note 1). + +From the standardization data already obtained calculate the +percentage of iron (Fe) in the limonite. + +[Note 1: It has already been noted that hydrochloric acid reacts +slowly in cold solutions with potassium permanganate. It is, however, +possible to obtain a satisfactory, although somewhat fugitive +end-point in the presence of manganous sulphate and phosphoric acid. +The explanation of the part played by these reagents is somewhat +obscure as yet. It is possible that an intermediate manganic compound +is formed which reacts rapidly with the ferrous compounds--thus in +effect catalyzing the oxidizing process. + +While an excess of hydrochloric acid is necessary for the successful +reduction of the iron by stannous chloride, too large an amount +should be avoided in order to lessen the chance of reduction of the +permanganate by the acid during titration.] + + + + +DETERMINATION OF THE OXIDIZING POWER OF PYROLUSITE + +INDIRECT OXIDATION + + +Pyrolusite, when pure, consists of manganese dioxide. Its value as an +oxidizing agent, and for the production of chlorine, depends upon the +percentage of MnO_{2} in the sample. This percentage is determined +by an indirect method, in which the manganese dioxide is reduced and +dissolved by an excess of ferrous sulphate or oxalic acid in the +presence of sulphuric acid, and the unused excess determined by +titration with standard permanganate solution. + +PROCEDURE.--Grind the mineral in an agate mortar until no grit +whatever can be detected under the pestle (Note 1). Transfer it to a +stoppered weighing-tube, and weigh out two portions of about 0.5 gram +into beakers (400-500 cc.) Read Note 2, and then calculate in each +case the weight of oxalic acid (H_{2}C_{2}O_{4}.2H_{2}O) required to +react with the weights of pyrolusite taken. The reaction involved is + +MnO_{2} + H_{2}C_{2}O_{4}(2H_{2}O) + H_{2}SO_{4} --> MnSO_{4} + +2CO_{2} + 4H_{2}O. + +Weigh out about 0.2 gram in excess of this quantity of !pure! oxalic +acid into the corresponding beakers, weighing the acid accurately and +recording the weight in the notebook. Pour into each beaker 25 cc. of +water and 50 cc. of dilute sulphuric acid (1:5), cover and warm the +beaker and its contents gently until the evolution of carbon dioxide +ceases (Note 3). If a residue remains which is sufficiently colored to +obscure the end-reaction of the permanganate, it must be removed by +filtration. + +Finally, dilute the solution to 200-300 cc., heat the solution to a +temperature just below boiling, add 15 cc. of a manganese sulphate +solution and while hot, titrate for the excess of the oxalic acid with +standard permanganate solution (Notes 4 and 5). + +From the corrected volume of the solution required, calculate the +amount of oxalic acid undecomposed by the pyrolusite; subtract this +from the total quantity of acid used, and calculate the weight of +manganese dioxide which would react with the balance of the acid, and +from this the percentage in the sample. + +[Note 1: The success of the analysis is largely dependent upon the +fineness of the powdered mineral. If properly ground, solution should +be complete in fifteen minutes or less.] + +[Note 2: A moderate excess of oxalic acid above that required to react +with the pyrolusite is necessary to promote solution; otherwise the +residual quantity of oxalic acid would be so small that the last +particles of the mineral would scarcely dissolve. It is also desirable +that a sufficient excess of the acid should be present to react with a +considerable volume of the permanganate solution during the titration, +thus increasing the accuracy of the process. On the other hand, the +excess of oxalic acid should not be so large as to react with more of +the permanganate solution than is contained in a 50 cc. burette. If +the pyrolusite under examination is known to be of high grade, say 80 +per cent pure, or above the calculation of the oxalic acid needed may +be based upon an assumption that the mineral is all MnO_{2}. If the +quality of the mineral is unknown, it is better to weigh out three +portions instead of two and to add to one of these the amount of +oxalic prescribed, assuming complete purity of the mineral. Then run +in the permanganate solution from a pipette or burette to determine +roughly the amount required. If the volume exceeds the contents of a +burette, the amount of oxalic acid added to the other two portions is +reduced accordingly.] + +[Note 3: Care should be taken that the sides of the beaker are not +overheated, as oxalic acid would be decomposed by heat alone if +crystallization should occur on the sides of the vessel. Strong +sulphuric acid also decomposes the oxalic acid. The dilute acid +should, therefore, be prepared before it is poured into the beaker.] + +[Note 4: Ferrous ammonium sulphate, ferrous sulphate, or iron wire +may be substituted for the oxalic acid. The reaction is then the +following: + +2 FeSO_{4} + MnO_{2} + 2H_{2}SO_{4} --> Fe_{2}(SO_{4})_{3} + 2H_{2}O + +The excess of ferrous iron may also be determined by titration with +potassium bichromate, if desired. Care is required to prevent the +oxidation of the iron by the air, if ferrous salts are employed.] + +[Note 5: The oxidizing power of pyrolusite may be determined by other +volumetric processes, one of which is outlined in the following +reactions: + +MnO_{2} + 4HCl --> MnCl_{2} + Cl_{2} + 2H_{2}O +Cl_{2} + 2KI --> I_{2} + 2KCl +I_{2} + 2Na_{2}S_{2}O_{3} --> Na_{2}S_{4}O_{6} + 2NaI. + +The chlorine generated by the pyrolusite is passed into a solution of +potassium iodide. The liberated iodine is then determined by titration +with sodium thiosulphate, as described on page 78. This is a direct +process, although it involves three steps.] + + + + +IODIMETRY + + +The titration of iodine against sodium thiosulphate, with starch as an +indicator, may perhaps be regarded as the most accurate of volumetric +processes. The thiosulphate solution may be used in both acid and +neutral solutions to measure free iodine and the latter may, in turn, +serve as a measure of any substance capable of liberating iodine from +potassium iodide under suitable conditions for titration, as, for +example, in the process outlined in Note 5 on page 74. + +The fundamental reaction upon which iodometric processes are based is +the following: + +I_{2} + 2 Na_{2}S_{2}O_{3} --> 2 NaI + Na_{2}S_{4}O_{6}. + +This reaction between iodine and sodium thiosulphate, resulting in +the formation of the compound Na_{2}S_{4}O_{6}, called sodium +tetrathionate, is quantitatively exact, and differs in that +respect from the action of chlorine or bromine, which oxidize the +thiosulphate, but not quantitatively. + +NORMAL SOLUTIONS OF IODINE AND SODIUM THIOSULPHATE + +If the formulas of sodium thiosulphate and sodium tetrathionate are +written in a manner to show the atoms of oxygen associated +with sulphur atoms in each, thus, 2(Na_{2}).S_{2}O_{2} and +Na_{2}O.S_{4}O_{5}, it is plain that in the tetrathionate there are +five atoms of oxygen associated with sulphur, instead of the four +in the two molecules of the thiosulphate taken together. Although, +therefore, the iodine contains no oxygen, the two atoms of iodine +have, in effect, brought about the addition of one oxygen atoms to the +sulphur atoms. That is the same thing as saying that 253.84 grams of +iodine (I_{2}) are equivalent to 16 grams of oxygen; hence, since 8 +grams of oxygen is the basis of normal solutions, 253.84/2 or 126.97 +grams of iodine should be contained in one liter of normal iodine +solution. By a similar course of reasoning the conclusion is reached +that the normal solution of sodium thiosulphate should contain, +per liter, its molecular weight in grams. As the thiosulphate in +crystalline form has the formula Na_{2}S_{2}O_{3}.5H_{2}O, this weight +is 248.12 grams. Tenth-normal or hundredth-normal solutions are +generally used. + + +PREPARATION OF STANDARD SOLUTIONS + +!Approximate Strength, 0.1 N! + +PROCEDURE.--Weigh out on the rough balances 13 grams of commercial +iodine. Place it in a mortar with 18 grams of potassium iodide and +triturate with small portions of water until all is dissolved. Dilute +the solution to 1000 cc. and transfer to a liter bottle and mix +thoroughly (Note 1).[1] + +[Footnote 1: It will be found more economical to have a considerable +quantity of the solution prepared by a laboratory attendant, and to +have all unused solutions returned to the common stock.] + +Weigh out 25 grams of sodium thiosulphate, dissolve it in water which +has been previously boiled and cooled, and dilute to 1000 cc., also +with boiled water. Transfer the solution to a liter bottle and mix +thoroughly (Note 2). + +[Note 1: Iodine solutions react with water to form hydriodic acid +under the influence of the sunlight, and even at low room temperatures +the iodine tends to volatilize from solution. They should, therefore, +be protected from light and heat. Iodine solutions are not stable for +long periods under the best of conditions. They cannot be used in +burettes with rubber tips, since they attack the rubber.] + +[Note 2: Sodium thiosulphate (Na_{2}S_{2}O_{3}.5H_{2}O) is +rarely wholly pure as sold commercially, but may be purified by +recrystallization. The carbon dioxide absorbed from the air by +distilled water decomposes the salt, with the separation of sulphur. +Boiled water which has been cooled out of contact with the air should +be used in preparing solutions.] + + +INDICATOR SOLUTION + +The starch solution for use as an indicator must be freshly prepared. +A soluble starch is obtainable which serves well, and a solution of +0.5 gram of this starch in 25 cc. of boiling water is sufficient. The +solution should be filtered while hot and is ready for use when cold. + +If soluble starch is not at hand, potato starch may be used. Mix about +1 gram with 5 cc. of cold water to a smooth paste, pour 150 cc. of +!boiling! water over it, warm for a moment on the hot plate, and put +it aside to settle. Decant the supernatant liquid through a filter +and use the clear filtrate; 5 cc. of this solution are needed for a +titration. + +The solution of potato starch is less stable than the soluble starch. +The solid particles of the starch, if not removed by filtration, +become so colored by the iodine that they are not readily decolorized +by the thiosulphate (Note 1). + +[Note 1: The blue color which results when free iodine and starch +are brought together is probably not due to the formation of a true +chemical compound. It is regarded as a "solid solution" of iodine in +starch. Although it is unstable, and easily destroyed by heat, it +serves as an indicator for the presence of free iodine of remarkable +sensitiveness, and makes the iodometric processes the most +satisfactory of any in the field of volumetric analysis.] + + +COMPARISON OF IODINE AND THIOSULPHATE SOLUTIONS + +PROCEDURE.--Place the solutions in burettes (the iodine in a +glass-stoppered burette), observing the usual precautions. Run out 40 +cc. of the thiosulphate solution into a beaker, dilute with 150 cc. of +water, add 1 cc. to 2 cc. of the soluble starch solution, and titrate +with the iodine to the appearance of the blue of the iodo-starch. +Repeat until the ratio of the two solutions is established, +remembering all necessary corrections for burettes and for temperature +changes. + + +STANDARDIZATION OF SOLUTIONS + +Commercial iodine is usually not sufficiently pure to permit of its +use as a standard for thiosulphate solutions or the direct preparation +of a standard solution of iodine. It is likely to contain, beside +moisture, some iodine chloride, if chlorine was used to liberate the +iodine when it was prepared. It may be purified by sublimation after +mixing it with a little potassium iodide, which reacts with the iodine +chloride, forming potassium chloride and setting free the iodine. The +sublimed iodine is then dried by placing it in a closed container over +concentrated sulphuric acid. It may then be weighed in a stoppered +weighing-tube and dissolved in a solution of potassium iodide in a +stoppered flask to prevent loss of iodine by volatilization. About 18 +grams of the iodide and twelve grams of iodine per liter are required +for an approximately tenth-normal solution. + +An iodine solution made from commercial iodine may also be +standardized against arsenious oxide (As_{4}O_{6}). This substance +also usually requires purification by sublimation before use. + +The substances usually employed for the standardization of a +thiosulphate solution are potassium bromate and metallic copper. The +former is obtainable in pure condition or may be easily purified by +re-crystallization. Copper wire of high grade is sufficiently pure +to serve as a standard. Both potassium bromate and cupric salts in +solution will liberate iodine from an iodide, which is then titrated +with the thiosulphate solution. + +The reactions involved are the following: + +(a) KBrO_{3} + 6KI + 3H_{2}SO_{4} --> KBr + 3I_{2} + 3K_{2}SO_{4} + 3H_{2}O, + +(b) 3Cu + 8HNO_{3} --> 3Cu(NO_{3})_{2} + 2NO + 4H_{2}O, + 2Cu(NO_{3})_{2} + 4KI --> 2CuI + 4KNO_{3} + I_{2}. + +Two methods for the direct standardization of the sodium thiosulphate +solution are here described, and one for the direct standardization of +the iodine solution. + + +!Method A! + +PROCEDURE.--Weigh out into 500 cc. beakers two portions of about +0.150-0.175 gram of potassium bromate. Dissolve each of these in 50 +cc. of water, and add 10 cc. of a potassium iodide solution containing +3 grams of the salt in that volume (Note 1). Add to the mixture 10 cc. +of dilute sulphuric acid (1 volume of sulphuric acid with 5 volumes of +water), allow the solution to stand for three minutes, and dilute to +150 cc. (Note 2). Run in thiosulphate solution from a burette until +the color of the liberated iodine is nearly destroyed, and then add 1 +cc. or 2 cc. of starch solution, titrate to the disappearance of the +iodo-starch blue, and finally add iodine solution until the color +is just restored. Make a blank test for the amount of thiosulphate +solution required to react with the iodine liberated by the iodate +which is generally present in the potassium iodide solution, and +deduct this from the total volume used in the titration. + +From the data obtained, calculate the relation of the thiosulphate +solution to a normal solution, and subsequently calculate the similar +value for the iodine solution. + +[Note 1:--Potassium iodide usually contains small amounts of potassium +iodate as impurity which, when the iodide is brought into an acid +solution, liberates iodine, just as does the potassium bromate used as +a standard. It is necessary to determine the amount of thiosulphate +which reacts with the iodine thus liberated by making a "blank test" +with the iodide and acid alone. As the iodate is not always uniformly +distributed throughout the iodide, it is better to make up a +sufficient volume of a solution of the iodide for the purposes of the +work in hand, and to make the blank test by using the same volume of +the iodide solution as is added in the standardizing process. The +iodide solution should contain about 3 grams of the salt in 10 cc.] + +[Note 2: The color of the iodo-starch is somewhat less satisfactory in +concentrated solutions of the alkali salts, notably the iodides. The +dilution prescribed obviates this difficulty.] + + +!Method B! + +PROCEDURE.--Weigh out two portions of 0.25-0.27 gram of clean copper +wire into 250 cc. Erlenmeyer flasks (Note 1). Add to each 5 cc. of +concentrated nitric acid (sp. gr. 1.42) and 25 cc. of water, cover, +and warm until solution is complete. Add 5 cc. of bromine water and +boil until the excess of bromine is expelled. Cool, and add strong +ammonia (sp. gr. 0.90) drop by drop until a deep blue color indicates +the presence of an excess. Boil the solution until the deep blue is +replaced by a light bluish green, or a brown stain appears on the +sides of the flask (Note 2). Add 10 cc. of strong acetic acid (sp. +gr. 1.04), cool under the water tap, and add a solution of potassium +iodide (Note 3) containing about 3 grams of the salt, and titrate +with thiosulphate solution until the color of the liberated iodine +is nearly destroyed. Then add 1-2 cc. of freshly prepared starch +solution, and add thiosulphate solution, drop by drop, until the blue +color is discharged. + +From the data obtained, including the "blank test" of the iodide, +calculate the relation of the thiosulphate solution to the normal. + +[Note 1: While copper wire of commerce is not absolutely pure, the +requirements for its use as a conductor of electricity are such that +the impurities constitute only a few hundredths of one per cent and +are negligible for analytical purposes.] + +[Note 2: Ammonia neutralizes the free nitric acid. It should be added +in slight excess only, since the excess must be removed by boiling, +which is tedious. If too much ammonia is present when acetic acid is +added, the resulting ammonium acetate is hydrolyzed, and the ammonium +hydroxide reacts with the iodine set free.] + +[Note 3: A considerable excess of potassium iodide is necessary for +the prompt liberation of iodine. While a large excess will do no harm, +the cost of this reagent is so great that waste should be avoided.] + + +!Method C! + +PROCEDURE.--Weigh out into 500 cc. beakers two portions of 0.175-0.200 +gram each of pure arsenious oxide. Dissolve each of these in 10 cc. of +sodium hydroxide solution, with stirring. Dilute the solutions to 150 +cc. and add dilute hydrochloric acid until the solutions contain a few +drops in excess, and finally add to each a concentrated solution of +5 grams of pure sodium bicarbonate (NaHCO_{3}) in water. Cover the +beakers before adding the bicarbonate, to avoid loss. Add the starch +solution and titrate with the iodine to the appearance of the blue of +the iodo-starch, taking care not to pass the end-point by more than a +few drops (Note 1). + +From the corrected volume of the iodine solution used to oxidize the +arsenious oxide, calculate its relation to the normal. From the +ratio between the solutions, calculate the similar value for the +thiosulphate solution. + +[Note 1: Arsenious oxide dissolves more readily in caustic alkali than +in a bicarbonate solution, but the presence of caustic alkali during +the titration is not admissible. It is therefore destroyed by the +addition of acid, and the solution is then made neutral with the +solution of bicarbonate, part of which reacts with the acid, the +excess remaining in solution. + +The reaction during titration is the following: + +Na_{3}AsO_{3} + I_{2} + 2NaHCO_{3} --> Na_{3}AsO_{4} + 2NaI + 2CO_{2} ++ H_{2}O + +As the reaction between sodium thiosulphate and iodine is not always +free from secondary reactions in the presence of even the weakly +alkaline bicarbonate, it is best to avoid the addition of any +considerable excess of iodine. Should the end-point be passed by a few +drops, the thiosulphate may be used to correct it.] + + + + +DETERMINATION OF COPPER IN ORES + + +Copper ores vary widely in composition from the nearly pure copper +minerals, such as malachite and copper sulphide, to very low grade +materials which contain such impurities as silica, lead, iron, silver, +sulphur, arsenic, and antimony. In nearly all varieties there will be +found a siliceous residue insoluble in acids. The method here given, +which is a modification of that described by A.H. Low (!J. Am. Chem. +Soc.! (1902), 24, 1082), provides for the extraction of the copper +from commonly occurring ores, and for the presence of their common +impurities. For practice analyses it is advisable to select an ore of +a fair degree of purity. + +PROCEDURE.-- Weigh out two portions of about 0.5 gram each of the +ore (which should be ground until no grit is detected) into 250 cc. +Erlenmeyer flasks or small beakers. Add 10 cc. of concentrated nitric +acid (sp. gr. 1.42) and heat very gently until the ore is decomposed +and the acid evaporated nearly to dryness (Note 1). Add 5 cc. of +concentrated hydrochloric acid (sp. gr. 1.2) and warm gently. Then +add about 7 cc. of concentrated sulphuric acid (sp. gr. 1.84) and +evaporate over a free flame until the sulphuric acid fumes freely +(Note 2). It has then displaced nitric and hydrochloric acid from +their compounds. + +Cool the flask or beaker, add 25 cc. of water, heat the solution +to boiling, and boil for two minutes. Filter to remove insoluble +sulphates, silica and any silver that may have been precipitated as +silver chloride, and receive the filtrate in a small beaker, washing +the precipitate and filter paper with warm water until the filtrate +and washings amount to 75 cc. Bend a strip of aluminium foil (5 cm. x +12 cm.) into triangular form and place it on edge in the beaker. Cover +the beaker and boil the solution (being careful to avoid loss of +liquid by spattering) for ten minutes, but do not evaporate to small +volume. + +Wash the cover glass and sides of the beaker. The copper should now be +in the form of a precipitate at the bottom of the beaker or adhering +loosely to the aluminium sheet. Remove the sheet, wash it carefully +with hydrogen sulphide water and place it in a small beaker. Decant +the solution through a filter, wash the precipitated copper twice by +decantation with hydrogen sulphide water, and finally transfer the +copper to the filter paper, where it is again washed thoroughly, being +careful at all times to keep the precipitated copper covered with the +wash water. Remove and discard the filtrate and place an Erlenmeyer +flask under the funnel. Pour 15 cc. of dilute nitric acid (sp. gr. +1.20) over the aluminium foil in the beaker, thus dissolving any +adhering copper. Wash the foil with hot water and remove it. Warm this +nitric acid solution and pour it slowly through the filter paper, +thereby dissolving the copper on the paper, receiving the acid +solution in the Erlenmeyer flask. Before washing the paper, pour 5 cc. +of saturated bromine water (Note 3) through it and finally wash the +paper carefully with hot water and transfer any particles of copper +which may be left on it to the Erlenmeyer flask. Boil to expel the +bromine. Add concentrated ammonia drop by drop until the appearance of +a deep blue coloration indicates an excess. Boil until the deep blue +is displaced by a light bluish green coloration, or until brown stains +form on the sides of the flask. Add 10 cc. of strong acetic acid (Note +4) and cool under the water tap. Add a solution containing about 3 +grams of potassium iodide, as in the standardization, and titrate with +thiosulphate solution until the yellow of the liberated iodine is +nearly discharged. Add 1-2 cc. of freshly prepared starch solution and +titrate to the disappearance of the blue color. + +From the data obtained, calculate the percentage of copper (Cu) in the +ore. + +[Note 1: Nitric acid, because of its oxidizing power, is used as a +solvent for the sulphide ores. As a strong acid it will also dissolve +the copper from carbonate ores. The hydrochloric acid is added to +dissolve oxides of iron and to precipitate silver and lead. The +sulphuric acid displaces the other acids, leaving a solution +containing sulphates only. It also, by its dehydrating action, renders +silica from silicates insoluble.] + +[Note 2: Unless proper precautions are taken to insure the correct +concentrations of acid the copper will not precipitate quantitatively +on the aluminium foil; hence care must be taken to follow directions +carefully at this point. Lead and silver have been almost completely +removed as sulphate and chloride respectively, or they too would +be precipitated on the aluminium. Bismuth, though precipitated on +aluminium, has no effect on the analysis. Arsenic and antimony +precipitate on aluminium and would interfere with the titration if +allowed to remain in the lower state of oxidation.] + +[Note 3: Bromine is added to oxidize arsenious and antimonious +compounds from the original sample, and to oxidize nitrous acid formed +by the action of nitric acid on copper and copper sulphide.] + +[Note 4: This reaction can be carried out in the presence of sulphuric +and hydrochloric acids as well as acetic acid, but in the presence +of these strong acids arsenic and antimonic acids may react with the +hydriodic acid produced with the liberation of free iodine, thereby +reversing the process and introducing an error.] + + + + +DETERMINATION OF ANTIMONY IN STIBNITE + + +Stibnite is native antimony sulphide. Nearly pure samples of this +mineral are easily obtainable and should be used for practice, since +many impurities, notably iron, seriously interfere with the accurate +determination of the antimony by iodometric methods. It is, moreover, +essential that the directions with respect to amounts of reagents +employed and concentration of solutions should be followed closely. + +PROCEDURE.--Grind the mineral with great care, and weigh out two +portions of 0.35-0.40 gram into small, dry beakers (100 cc.). +Cover the beakers and pour over the stibnite 5 cc. of concentrated +hydrochloric acid (sp. gr. 1.20) and warm gently on the water bath +(Note 1). When the residue is white, add to each beaker 2 grams of +powdered tartaric acid (Note 2). Warm the solution on the water bath +for ten minutes longer, dilute the solution very cautiously by adding +water in portions of 5 cc., stopping if the solution turns red. It +is possible that no coloration will appear, in which case cautiously +continue the dilution to 125 cc. If a red precipitate or coloration +does appear, warm the solution until it is colorless, and again dilute +cautiously to a total volume of 125 cc. and boil for a minute (Note +3). + +If a white precipitate of the oxychloride separates during dilution +(which should not occur if the directions are followed), it is best to +discard the determination and to start anew. + +Carefully neutralize most of the acid with ammonium hydroxide solution +(sp. gr. 0.96), but leave it distinctly acid (Note 4). Dissolve 3 +grams of sodium bicarbonate in 200 cc. of water in a 500 cc. beaker, +and pour the cold solution of the antimony chloride into this, +avoiding loss by effervescence. Make sure that the solution contains +an excess of the bicarbonate, and then add 1 cc. or 2 cc. of starch +solution and titrate with iodine solution to the appearance of the +blue, avoiding excess (Notes 5 and 6). + +From the corrected volume of the iodine solution required to oxidize +the antimony, calculate the percentage of antimony (Sb) in the +stibnite. + +[Note 1: Antimony chloride is volatile with steam from its +concentrated solutions; hence these solutions must not be boiled until +they have been diluted.] + +[Note 2: Antimony salts, such as the chloride, are readily hydrolyzed, +and compounds such as SbOCl are formed which are often relatively +insoluble; but in the presence of tartaric acid compounds with complex +ions are formed, and these are soluble. An excess of hydrochloric acid +also prevents precipitation of the oxychloride because the H^{+} ions +from the acid lessen the dissociation of the water and thus prevent +any considerable hydrolysis.] + +[Note 3: The action of hydrochloric acid upon the sulphide sets free +sulphureted hydrogen, a part of which is held in solution by the acid. +This is usually expelled by the heating upon the water bath; but if it +is not wholly driven out, a point is reached during dilution at which +the antimony sulphide, being no longer held in solution by the acid, +separates. If the dilution is immediately stopped and the solution +warmed, this sulphide is again brought into solution and at the same +time more of the sulphureted hydrogen is expelled. This procedure must +be continued until the sulphureted hydrogen is all removed, since it +reacts with iodine. If no precipitation of the sulphide occurs, it +is an indication that the sulphureted hydrogen was all expelled on +solution of the stibnite.] + +[Note 4: Ammonium hydroxide is added to neutralize most of the acid, +thus lessening the amount of sodium bicarbonate to be added. The +ammonia should not neutralize all of the acid.] + +[Note 5: The reaction which takes place during titration may be +expressed thus: + +Na_{3}SbO_{3} + 2NaHCO_{3} + I_{2} --> Na_{3}SbO_{4} + 2NaI + H_{2}O + +2CO_{2}.] + +[Note 6: If the end-point is not permanent, that is, if the blue of +the iodo-starch is discharged after standing a few moments, the cause +may be an insufficient quantity of sodium bicarbonate, leaving the +solution slightly acid, or a very slight precipitation of an antimony +compound which is slowly acted upon by the iodine when the latter is +momentarily present in excess. In either case it is better to discard +the analysis and to repeat the process, using greater care in the +amounts of reagents employed.] + + + + +CHLORIMETRY + + +The processes included under the term !chlorimetry! comprise +those employed to determine chlorine, hypochlorites, bromine, and +hypobromites. The reagent employed is sodium arsenite in the presence +of sodium bicarbonate. The reaction in the case of the hypochlorites +is + +NaClO + Na_{3}AsO_{3} --> Na_{3}AsO_{4} + NaCl. + +The sodium arsenite may be prepared from pure arsenious oxide, +as described below, and is stable for considerable periods; but +commercial oxide requires resublimation to remove arsenic sulphide, +which may be present in small quantity. To prepare the solution, +dissolve about 5 grams of the powdered oxide, accurately weighed, +in 10 cc. of a concentrated sodium hydroxide solution, dilute the +solution to 300 cc., and make it faintly acid with dilute hydrochloric +acid. Add 30 grams of sodium bicarbonate dissolved in a little water, +and dilute the solution to exactly 1000 cc. in a measuring flask. +Transfer the solution to a dry liter bottle and mix thoroughly. + +It is possible to dissolve the arsenious oxide directly in a solution +of sodium bicarbonate, with gentle warming, but solution in sodium +hydroxide takes place much more rapidly, and the excess of the +hydroxide is readily neutralized by hydrochloric acid, with subsequent +addition of the bicarbonate to maintain neutrality during the +titration. + +The indicator required for this process is made by dipping strips of +filter paper in a starch solution prepared as described on page 76, +to which 1 gram of potassium iodide has been added. These strips are +allowed to drain and spread upon a watch-glass until dry. When touched +by a drop of the solution the paper turns blue until the hypochlorite +has all been reduced and an excess of the arsenite has been added. + + + + +DETERMINATION OF THE AVAILABLE CHLORINE IN BLEACHING POWDER + + +Bleaching powder consists mainly of a calcium compound which is a +derivative of both hydrochloric and hypochlorous acids. Its formula is +CaClOCl. Its use as a bleaching or disinfecting agent, or as a source +of chlorine, depends upon the amount of hypochlorous acid which it +yields when treated with a stronger acid. It is customary to express +the value of bleaching powder in terms of "available chlorine," by +which is meant the chlorine present as hypochlorite, but not the +chlorine present as chloride. + +PROCEDURE.--Weigh out from a stoppered test tube into a porcelain +mortar about 3.5 grams of bleaching powder (Note 1). Triturate the +powder in the mortar with successive portions of water until it is +well ground and wash the contents into a 500 cc. measuring flask +(Note 2). Fill the flask to the mark with water and shake thoroughly. +Measure off 25 cc. of this semi-solution in a measuring flask, or +pipette, observing the precaution that the liquid removed shall +contain approximately its proportion of suspended matter. + +Empty the flask or pipette into a beaker and wash it out. Run in the +arsenite solution from a burette until no further reaction takes place +on the starch-iodide paper when touched by a drop of the solution of +bleaching powder. Repeat the titration, using a second 25 cc. portion. + +From the volume of solution required to react with the bleaching +powder, calculate the percentage of available chlorine in the latter, +assuming the titration reaction to be that between chlorine and +arsenious oxide: + +As_{4}O_{6} + 4Cl_{2} + 4H_{2}O --> 2As_{2}O_{5} + 8HCl + +Note that only one twentieth of the original weight of bleaching +powder enters into the reaction. + +[Note 1: The powder must be triturated until it is fine, otherwise the +lumps will inclose calcium hypochlorite, which will fail to react with +the arsenious acid. The clear supernatant liquid gives percentages +which are below, and the sediment percentages which are above, the +average. The liquid measured off should, therefore, carry with it its +proper proportion of the sediment, so far as that can be brought about +by shaking the solution just before removal of the aliquot part for +titration.] + +[Note 2: Bleaching powder is easily acted upon by the carbonic acid in +the air, which liberates the weak hypochlorous acid. This, of course, +results in a loss of available chlorine. The original material for +analysis should be kept in a closed container and protected form the +air as far as possible. It is difficult to obtain analytical samples +which are accurately representative of a large quantity of the +bleaching powder. The procedure, as outlined, will yield results which +are sufficiently exact for technical purposes.] + + + + +III. PRECIPITATION METHODS + + + + +DETERMINATION OF SILVER BY THE THIOCYANATE PROCESS + + +The addition of a solution of potassium or ammonium thiocyanate to one +of silver in nitric acid causes a deposition of silver thiocyanate as +a white, curdy precipitate. If ferric nitrate is also present, the +slightest excess of the thiocyanate over that required to combine with +the silver is indicated by the deep red which is characteristic of the +thiocyanate test for iron. + +The reactions involved are: + +AgNO_{3} + KSCN --> AgSCN + KNO_{3}, +3KSCN + Fe(NO_{3})_{3} --> Fe(SCN)_{3} + 3KNO_{3}. + +The ferric thiocyanate differs from the great majority of salts in +that it is but very little dissociated in aqueous solutions, and the +characteristic color appears to be occasioned by the formation of the +un-ionized ferric salt. + +The normal solution of potassium thiocyanate should contain an amount +of the salt per liter of solution which would yield sufficient +(CNS)^{-} to combine with one gram of hydrogen to form HCNS, i.e., +a gram-molecular weight of the salt or 97.17 grams. If the ammonium +thiocyanate is used, the amount is 76.08 grams. To prepare the +solution for this determination, which should be approximately 0.05 +N, dissolve about 5 grams of potassium thiocyanate, or 4 grams of +ammonium thiocyanate, in a small amount of water; dilute this solution +to 1000 cc. in a liter bottle and mix as usual. + +Prepare 20 cc. of a saturated solution of ferric alum and add 5 cc. of +dilute nitric acid (sp. gr. 1.20). About 5 cc. of this solution should +be used as an indicator. + + +STANDARDIZATION + +PROCEDURE.--Crush a small quantity of silver nitrate crystals in a +mortar (Note 1). Transfer them to a watch-glass and dry them for an +hour at 110°C., protecting them from dust or other organic matter +(Note 2). Weigh out two portions of about 0.5 gram each and dissolve +them in 50 cc. of water. Add 10 cc. of dilute nitric acid which has +been recently boiled to expel the lower oxides of nitrogen, if any, +and then add 5 cc. of the indicator solution. Run in the thiocyanate +solution from a burette, with constant stirring, allowing the +precipitate to settle occasionally to obtain an exact recognition +of the end-point, until a faint red tinge can be detected in the +solution. + +From the data obtained, calculate the relation of the thiocyanate +solution to the normal. + +[Note 1: The thiocyanate cannot be accurately weighed; its solutions +must, therefore, be standardized against silver nitrate (or pure +silver), either in the form of a standard solution or in small, +weighed portions.] + +[Note 2: The crystals of silver nitrate sometimes inclose water which +is expelled on drying. If the nitrate has come into contact with +organic bodies it suffers a reduction and blackens during the heating. + +It is plain that a standard solution of silver nitrate (made by +weighing out the crystals) is convenient or necessary if many +titrations of this nature are to be made. In the absence of such a +solution the liability of passing the end-point is lessened by setting +aside a small fraction of the silver solution, to be added near the +close of the titration.] + + +DETERMINATION OF SILVER IN COIN + +PROCEDURE.-- Weigh out two portions of the coin of about 0.5 gram +each. Dissolve them in 15 cc. of dilute nitric acid (sp. gr. 1.2) and +boil until all the nitrous compounds are expelled (Note 1). Cool the +solution, dilute to 50 cc., and add 5 cc. of the indicator solution, +and titrate with the thiocyanate to the appearance of the faint red +coloration (Note 2). + +From the corrected volume of the thiocyanate solution required, +calculate the percentage of silver in the coin. + +[Note 1: The reaction with silver may be carried out in nitric acid +solutions and in the presence of copper, if the latter does not exceed +70 per cent. Above that percentage it is necessary to add silver in +known quantity to the solution. The liquid must be cold at the time of +titration and entirely free from nitrous compounds, as these sometimes +cause a reddening of the indicator solution. All utensils, distilled +water, the nitric acid and the beakers must be free from chlorides, +as the presence of these will cause precipitation of silver chloride, +thereby introducing an error.] + +[Note 2: The solution containing the silver precipitate, as well as +those from the standardization, should be placed in the receptacle for +"silver residues" as a matter of economy.] + + + + +PART III + +GRAVIMETRIC ANALYSIS + + + + +GENERAL DIRECTIONS + + +Gravimetric analyses involve the following principal steps: first, the +weighing of the sample; second, the solution of the sample; third, the +separation of some substance from solution containing, or bearing a +definite relation to, the constituent to be measured, under conditions +which render this separation as complete as possible; and finally, +the segregation of that substance, commonly by filtration, and the +determination of its weight, or that of some stable product formed +from it on ignition. For example, the gravimetric determination of +aluminium is accomplished by solution of the sample, by precipitation +in the form of hydroxide, collection of the hydroxide upon a filter, +complete removal by washing of all foreign soluble matter, and the +burning of the filter and ignition of the precipitate to aluminium +oxide, in which condition it is weighed. + +Among the operations which are common to nearly all gravimetric +analyses are precipitation, washing of precipitates, ignition of +precipitates, and the use of desiccators. In order to avoid burdensome +repetitions in the descriptions of the various gravimetric procedures +which follow, certain general instructions are introduced at this +point. These instructions must, therefore, be considered to be as much +a part of all subsequent procedures as the description of apparatus, +reagents, or manipulations. + +The analytical balance, the fundamentally important instrument in +gravimetric analysis, has already been described on pages 11 to 15. + + +PRECIPITATION + +For successful quantitative precipitations those substances are +selected which are least soluble under conditions which can be easily +established, and which separate from solution in such a state that +they can be filtered readily and washed free from admixed material. +In general, the substances selected are the same as those already +familiar to the student of Qualitative Analysis. + +When possible, substances are selected which separate in crystalline +form, since such substances are less likely to clog the pores of +filter paper and can be most quickly washed. In order to increase the +size of the crystals, which further promotes filtration and washing, +it is often desirable to allow a precipitate to remain for some time +in contact with the solution from which it has separated. The solution +is often kept warm during this period of "digestion." The small +crystals gradually disappear and the larger crystals increase in size, +probably as the result of the force known as surface tension, which +tends to reduce the surface of a given mass of material to a minimum, +combined with a very slightly greater solubility of small crystals as +compared with the larger ones. + +Amorphous substances, such as ferric hydroxide, aluminium hydroxide, +or silicic acid, separate in a gelatinous form and are relatively +difficult to filter and wash. Substances of this class also exhibit +a tendency to form, with pure water, what are known as colloidal +solutions. To prevent this as far as possible, they are washed with +solutions of volatile salts, as will be described in some of the +following procedures. + +In all precipitations the reagent should be added slowly, with +constant stirring, and should be hot when circumstances permit. +The slow addition is less likely to occasion contamination of the +precipitate by the inclosure of other substances which may be in the +solution, or of the reagent itself. + + +FUNNELS AND FILTERS + +Filtration in analytical processes is most commonly effected through +paper filters. In special cases these may be advantageously replaced +by an asbestos filter in a perforated porcelain or platinum crucible, +commonly known, from its originator, as a "Gooch filter." The +operation and use of a filter of this type is described on page 103. +Porous crucibles of a material known as alundum may also be employed +to advantage in special cases. + +The glass funnels selected for use with paper filters should have an +angle as near 60° as possible, and a narrow stem about six inches in +length. The filters employed should be washed filters, i.e., those +which have been treated with hydrochloric and hydrofluoric acids, and +which on incineration leave a very small and definitely known weight +of ash, generally about .00003 gram. Such filters are readily +obtainable on the market. + +The filter should be carefully folded to fit the funnel according to +either of the two well-established methods described in the Appendix. +It should always be placed so that the upper edge of the paper +is about one fourth inch below the top of the funnel. Under no +circumstances should the filter extend above the edge of the funnel, +as it is then utterly impossible to effect complete washing. + +To test the efficiency of the filter, fill it with distilled water. +This water should soon fill the stem completely, forming a continuous +column of liquid which, by its hydrostatic pressure, produces a gentle +suction, thus materially promoting the rapidity of filtration. Unless +the filter allows free passage of water under these conditions, it is +likely to give much trouble when a precipitate is placed upon it. + +The use of a suction pump to promote filtration is rarely altogether +advantageous in quantitative analysis, if paper filters are employed. +The tendency of the filter to break, unless the point of the filter +paper is supported by a perforated porcelain cone or a small "hardened +filter" of parchment, and the tendency of the precipitates to pass +through the pores of the filter, more than compensate for the possible +gain in time. On the other hand, filtration by suction may be useful +in the case of precipitates which do not require ignition before +weighing, or in the case of precipitates which are to be discarded +without weighing. This is best accomplished with the aid of the +special apparatus called a Gooch filter referred to above. + + +FILTRATION AND WASHING OF PRECIPITATES + +Solutions should be filtered while hot, as far as possible, since +the passage of a liquid through the pores of a filter is retarded by +friction, and this, for water at 100°C., is less than one sixth of the +resistance at 0°C. + +When the filtrate is received in a beaker, the stem of the funnel +should touch the side of the receiving vessel to avoid loss by +spattering. Neglect of this precaution is a frequent source of error. + +The vessels which contain the initial filtrate should !always! be +replaced by clean ones, properly labeled, before the washing of a +precipitate begins. In many instances a finely divided precipitate +which shows no tendency to pass through the filter at first, while the +solution is relatively dense, appears at once in the washings. Under +such conditions the advantages accruing from the removal of the first +filtrate are obvious, both as regards the diminished volume requiring +refiltration, and also the smaller number of washings subsequently +required. + +Much time may often be saved by washing precipitates by decantation, +i.e., by pouring over them, while still in the original vessel, +considerable volumes of wash-water and allowing them to settle. The +supernatant, clear wash-water is then decanted through the filter, +so far as practicable without disturbing the precipitate, and a new +portion of wash-water is added. This procedure can be employed to +special advantage with gelatinous precipitates, which fill up the +pores of the filter paper. As the medium from which the precipitate +is to settle becomes less dense it subsides less readily, and it +ultimately becomes necessary to transfer it to the filter and complete +the washing there. + +A precipitate should never completely fill a filter. The wash-water +should be applied at the top of the filter, above the precipitate. +It may be shown mathematically that the washing is most !rapidly! +accomplished by filling the filter well to the top with wash-water +each time, and allowing it to drain completely after each addition; +but that when a precipitate is to be washed with the !least possible +volume! of liquid the latter should be applied in repeated !small! +quantities. + +Gelatinous precipitates should not be allowed to dry before complete +removal of foreign matter is effected. They are likely to shrink and +crack, and subsequent additions of wash-water pass through these +channels only. + +All filtrates and wash-waters without exception must be properly +tested. !This lies at the foundation of accurate work!, and the +student should clearly understand that it is only by the invariable +application of this rule that assurance of ultimate reliability can +be secured. Every original filtrate must be tested to prove complete +precipitation of the compound to be separated, and the wash-waters +must also be tested to assure complete removal of foreign material. In +testing the latter, the amount first taken should be but a few +drops if the filtrate contains material which is to be subsequently +determined. When, however, the washing of the filter and precipitate +is nearly completed the amount should be increased, and for the final +test not less than 3 cc. should be used. + +It is impossible to trust to one's judgment with regard to the washing +of precipitates; the washings from !each precipitate! of a series +simultaneously treated must be tested, since the rate of washing will +often differ materially under apparently similar conditions, !No +exception can ever be made to this rule!. + +The habit of placing a clean common filter paper under the receiving +beaker during filtration is one to be commended. On this paper a +record of the number of washings can very well be made as the portions +of wash-water are added. + +It is an excellent practice, when possible, to retain filtrates and +precipitates until the completion of an analysis, in order that, in +case of question, they may be examined to discover sources of error. + +For the complete removal of precipitates from containing vessels, it +is often necessary to rub the sides of these vessels to loosen the +adhering particles. This can best be done by slipping over the end of +a stirring rod a soft rubber device sometimes called a "policeman." + + +DESICCATORS + +Desiccators should be filled with fused, anhydrous calcium chloride, +over which is placed a clay triangle, or an iron triangle covered with +silica tubes, to support the crucible or other utensils. The cover of +the desiccator should be made air-tight by the use of a thin coating +of vaseline. + +Pumice moistened with concentrated sulphuric acid may be used in place +of the calcium chloride, and is essential in special cases; but for +most purposes the calcium chloride, if renewed occasionally and not +allowed to cake together, is practically efficient and does not slop +about when the desiccator is moved. + +Desiccators should never remain uncovered for any length of time. The +dehydrating agents rapidly lose their efficiency on exposure to the +air. + + +CRUCIBLES + +It is often necessary in quantitative analysis to employ fluxes to +bring into solution substances which are not dissolved by acids. The +fluxes in most common use are sodium carbonate and sodium or potassium +acid sulphate. In gravimetric analysis it is usually necessary to +ignite the separated substance after filtration and washing, in order +to remove moisture, or to convert it through physical or chemical +changes into some definite and stable form for weighing. Crucibles +to be used in fusion processes must be made of materials which will +withstand the action of the fluxes employed, and crucibles to be used +for ignitions must be made of material which will not undergo any +permanent change during the ignition, since the initial weight of the +crucible must be deducted from the final weight of the crucible and +product to obtain the weight of the ignited substance. The three +materials which satisfy these conditions, in general, are platinum, +porcelain, and silica. + +Platinum crucibles have the advantage that they can be employed at +high temperatures, but, on the other hand, these crucibles can never +be used when there is a possibility of the reduction to the metallic +state of metals like lead, copper, silver, or gold, which would alloy +with and ruin the crucible. When platinum crucibles are used with +compounds of arsenic or phosphorus, special precautions are necessary +to prevent damage. This statement applies to both fusions and +ignitions. + +Fusions with sodium carbonate can be made only in platinum, since +porcelain or silica crucibles are attacked by this reagent. Acid +sulphate fusions, which require comparatively low temperatures, can +sometimes be made in platinum, although platinum is slightly attacked +by the flux. Porcelain or silica crucibles may be used with acid +fluxes. + +Silica crucibles are less likely to crack on heating than porcelain +crucibles on account of their smaller coefficient of expansion. +Ignition of substances not requiring too high a temperature may be +made in porcelain or silica crucibles. + +Iron, nickel or silver crucibles are used in special cases. + +In general, platinum crucibles should be used whenever such use is +practicable, and this is the custom in private, research or commercial +laboratories. Platinum has, however, become so valuable that it is +liable to theft unless constantly under the protection of the user. As +constant protection is often difficult in instructional laboratories, +it is advisable, in order to avoid serious monetary losses, to use +porcelain or silica crucibles whenever these will give satisfactory +service. When platinum utensils are used the danger of theft should +always be kept in mind. + + +PREPARATION OF CRUCIBLES FOR USE + +All crucibles, of whatever material, must always be cleaned, ignited +and allowed to cool in a desiccator before weighing, since all bodies +exposed to the air condense on their surfaces a layer of moisture +which increases their weight. The amount and weight of this moisture +varies with the humidity of the atmosphere, and the latter may change +from hour to hour. The air in the desiccator (see above) is kept at +a constant and low humidity by the drying agent which it contains. +Bodies which remain in a desiccator for a sufficient time (usually +20-30 minutes) retain, therefore, on their surfaces a constant weight +of moisture which is the same day after day, thus insuring constant +conditions. + +Hot objects, such as ignited crucibles, should be allowed to cool in +the air until, when held near the skin, but little heat is noticeable. +If this precaution is not taken, the air within the desiccator is +strongly heated and expands before the desiccator is covered. As the +temperature falls, the air contracts, causing a reduction of air +pressure within the covered vessel. When the cover is removed (which +is often rendered difficult) the inrush of air from the outside may +sweep light particles out of a crucible, thus ruining an entire +analysis. + +Constant heating of platinum causes a slight crystallization of the +surface which, if not removed, penetrates into the crucible. Gentle +polishing of the surface destroys the crystalline structure and +prevents further damage. If sea sand is used for this purpose, great +care is necessary to keep it from the desk, since beakers are easily +scratched by it, and subsequently crack on heating. + +Platinum crucibles stained in use may often be cleaned by the fusion +in them of potassium or sodium acid sulphate, or by heating with +ammonium chloride. If the former is used, care should be taken not +to heat so strongly as to expel all of the sulphuric acid, since the +normal sulphates sometimes expand so rapidly on cooling as to split +the crucible. The fused material should be poured out, while hot, on +to a !dry! tile or iron surface. + + +IGNITION OF PRECIPITATES + +Most precipitates may, if proper precautions are taken, be ignited +without previous drying. If, however, such precipitates can be dried +without loss of time to the analyst (as, for example, over night), it +is well to submit them to this process. It should, nevertheless, be +remembered that a partially dried precipitate often requires more care +during ignition than a thoroughly moist one. + +The details of the ignition of precipitates vary so much with the +character of the precipitate, its moisture content, and temperature to +which it is to be heated, that these details will be given under the +various procedures which follow. + + + + +DETERMINATION OF CHLORINE IN SODIUM CHLORIDE + + +!Method A. With the Use of a Gooch Filter! + +PROCEDURE.--Carefully clean a weighing-tube containing the sodium +chloride, handling it as little as possible with the moist fingers, +and weigh it accurately to 0.0001 gram, recording the weight at once +in the notebook (see Appendix). Hold the tube over the top of a beaker +(200-300 cc.), and cautiously remove the stopper, noting carefully +that no particles fall from it, or from the tube, elsewhere than into +the beaker. Pour out a small portion of the chloride, replace the +stopper, and determine by approximate weighing how much has been +removed. Continue this procedure until 0.25-0.30 gram has been taken +from the tube, then weigh accurately and record the weight beneath the +first in the notebook. The difference of the two weights represents +the weight of the chloride taken for analysis. Again weigh a second +portion of 0.25-0.30 gram into a second beaker of the same size as the +first. The beakers should be plainly marked to correspond with the +entries in the notebook. Dissolve each portion of the chloride in 150 +cc. of distilled water and add about ten drops of dilute nitric acid +(sp. gr. 1.20) (Note 2). Calculate the volume of silver nitrate +solution required to effect complete precipitation in each case, +and add slowly about 5 cc. in excess of that amount, with constant +stirring. Heat the solutions cautiously to boiling, stirring +occasionally, and continue the heating and stirring until the +precipitates settle promptly, leaving a nearly clear supernatant +liquid (Note 3). This heating should not take place in direct sunlight +(Note 4). The beaker should be covered with a watch-glass, and both +boiling and stirring so regulated as to preclude any possibility of +loss of material. Add to the clear liquid one or two drops of silver +nitrate solution, to make sure that an excess of the reagent is +present. If a precipitate, or cloudiness, appears as the drops fall +into the solution, heat again, and stir until the whole precipitate +has coagulated. The solution is then ready for filtration. + +Prepare a Gooch filter as follows: Fold over the top of a Gooch funnel +(Fig. 2) a piece of rubber-band tubing, such as is known as "bill-tie" +tubing, and fit into the mouth of the funnel a perforated porcelain +crucible (Gooch crucible), making sure that when the crucible is +gently forced into the mouth of the funnel an airtight joint results. +(A small 1 or 1-1/4-inch glass funnel may be used, in which case the +rubber tubing is stretched over the top of the funnel and then drawn +up over the side of the crucible until an air-tight joint is secured.) + +[ILLUSTRATION: FIG. 2] + +Fit the funnel into the stopper of a filter bottle, and connect the +filter bottle with the suction pump. Suspend some finely divided +asbestos, which has been washed with acid, in 20 to 30 cc. of water +(Note 1); allow this to settle, pour off the very fine particles, and +then pour some of the mixture cautiously into the crucible until an +even felt of asbestos, not over 1/32 inch in thickness, is formed. A +gentle suction must be applied while preparing this felt. Wash the +felt thoroughly by passing through it distilled water until all fine +or loose particles are removed, increasing the suction at the last +until no more water can be drawn out of it; place on top of the felt +the small, perforated porcelain disc and hold it in place by pouring a +very thin layer of asbestos over it, washing the whole carefully; +then place the crucible in a small beaker, and place both in a drying +closet at 100-110°C. for thirty to forty minutes. Cool the crucible +in a desiccator, and weigh. Heat again for twenty to thirty minutes, +cool, and again weigh, repeating this until the weight is constant +within 0.0003 gram. The filter is then ready for use. + +Place the crucible in the funnel, and apply a gentle suction, !after +which! the solution to be filtered may be poured in without disturbing +the asbestos felt. When pouring liquid onto a Gooch filter hold the +stirring-rod at first well down in the crucible, so that the liquid +does not fall with any force upon the asbestos, and afterward keep the +crucible will filled with the solution. + +Pour the liquid above the silver chloride slowly onto the filter, +leaving the precipitate in the beaker as far as possible. Wash the +precipitate twice by decantation with warm water; then transfer it +to the filter with the aid of a stirring-rod with a rubber tip and a +stream from the wash-bottle. + +Examine the first portions of the filtrate which pass through the +filter with great care for asbestos fibers, which are most likely to +be lost at this point. Refilter the liquid if any fibers are visible. +Finally, wash the precipitate thoroughly with warm water until free +from soluble silver salts. To test the washings, disconnect the +suction at the flask and remove the funnel or filter tube from the +suction flask. Hold the end of the tube over the mouth of a small test +tube and add from a wash-bottle 2-3 cc. of water. Allow the water to +drip through into the test tube and add a drop of dilute hydrochloric +acid. No precipitate or cloud should form in the wash-water (Note 16). +Dry the filter and contents at 100-110°C. until the weight is constant +within 0.0003 gram, as described for the preparation of the filter. +Deduct the weight of the dry crucible from the final weight, and from +the weight of silver chloride thus obtained calculate the percentage +of chlorine in the sample of sodium chloride. + +[Note 1: The washed asbestos for this type of filter is prepared by +digesting in concentrated hydrochloric acid, long-fibered asbestos +which has been cut in pieces of about 0.5 cm. in length. After +digestion, the asbestos is filtered off on a filter plate and washed +with hot, distilled water until free from chlorides. A small portion +of the asbestos is shaken with water, forming a thin suspension, which +is bottled and kept for use.] + +[Note 2: The nitric acid is added before precipitation to lessen the +tendency of the silver chloride to carry down with it other substances +which might be precipitated from a neutral solution. A large excess of +the acid would exert a slight solvent action upon the chloride.] + +[Note 3: The solution should not be boiled after the addition of the +nitric acid before the presence of an excess of silver nitrate is +assured, since a slight interaction between the nitric acid and the +sodium chloride is possible, by which a loss of chlorine, either as +such or as hydrochloric acid, might ensue. The presence of an excess +of the precipitant can usually be recognized at the time of its +addition, by the increased readiness with which the precipitate +coagulates and settles.] + +[Note 4: The precipitate should not be exposed to strong sunlight, +since under those conditions a reduction of the silver chloride ensues +which is accompanied by a loss of chlorine. The superficial alteration +which the chloride undergoes in diffused daylight is not sufficient +to materially affect the accuracy of the determination. It should be +noted, however, that a slight error does result from the effect of +light upon the silver chloride precipitate and in cases in which the +greatest obtainable accuracy is required, the procedure described +under "Method B" should be followed, in which this slight reduction of +the silver chloride is corrected by subsequent treatment with nitric +and hydrochloric acids.] + +[Note 5: The asbestos used in the Gooch filter should be of the finest +quality and capable of division into minute fibrous particles. A +coarse felt is not satisfactory.] + +[Note 6: The precipitate must be washed with warm water until it is +absolutely free from silver and sodium nitrates. It may be assumed +that the sodium salt is completely removed when the wash-water shows +no evidence of silver. It must be borne in mind that silver chloride +is somewhat soluble in hydrochloric acid, and only a single drop +should be added. The washing should be continued until no cloudiness +whatever can be detected in 3 cc. of the washings. + +Silver chloride is but slightly soluble in water. The solubility +varies with its physical condition within small limits, and is +about 0.0018 gram per liter at 18°C. for the curdy variety usually +precipitated. The chloride is also somewhat soluble in solutions of +many chlorides, in solutions of silver nitrate, and in concentrated +nitric acid. + +As a matter of economy, the filtrate, which contains whatever silver +nitrate was added in excess, may be set aside. The silver can be +precipitated as chloride and later converted into silver nitrate.] + +[Note 7: The use of the Gooch filter commends itself strongly when a +considerable number of halogen determinations are to be made, since +successive portions of the silver halides may be filtered on the same +filter, without the removal of the preceding portions, until the +crucible is about two thirds filled. If the felt is properly prepared, +filtration and washing are rapidly accomplished on this filter, and +this, combined with the possibility of collecting several precipitates +on the same filter, is a strong argument in favor of its use with any +but gelatinous precipitates.] + + +!Method B. With the Use of a Paper Filter! + +PROCEDURE.--Weigh out two portions of sodium chloride of about +0.25-0.3 gram each and proceed with the precipitation of the silver +chloride as described under Method A above. When the chloride is ready +for filtration prepare two 9 cm. washed paper filters (see Appendix). +Pour the liquid above the precipitates through the filters, wash twice +by decantation and transfer the precipitates to the filters, finally +washing them until free from silver solution as described. The funnel +should then be covered with a moistened filter paper by stretching it +over the top and edges, to which it will adhere on drying. It should +be properly labeled with the student's name and desk number, and then +placed in a drying closet, at a temperature of about 100-110°C., until +completely dry. + +The perfectly dry filter is then opened over a circular piece of +clean, smooth, glazed paper about six inches in diameter, placed upon +a larger piece about twelve inches in diameter. The precipitate is +removed from the filter as completely as possible by rubbing the sides +gently together, or by scraping them cautiously with a feather which +has been cut close to the quill and is slightly stiff (Note 1). In +either case, care must be taken not to rub off any considerable +quantity of the paper, nor to lose silver chloride in the form of +dust. Cover the precipitate on the glazed paper with a watch-glass to +prevent loss of fine particles and to protect it from dust from the +air. Fold the filter paper carefully, roll it into a small cone, and +wind loosely around !the top! a piece of small platinum wire (Note 2). +Hold the filter by the wire over a small porcelain crucible (which has +been cleaned, ignited, cooled in a desiccator, and weighed), ignite +it, and allow the ash to fall into the crucible. Place the crucible +upon a clean clay triangle, on its side, and ignite, with a low +flame well at its base, until all the carbon of the filter has been +consumed. Allow the crucible to cool, add two drops of concentrated +nitric acid and one drop of concentrated hydrochloric acid, and heat +!very cautiously!, to avoid spattering, until the acids have been +expelled; then transfer the main portion of the precipitate from the +glazed paper to the cooled crucible, placing the latter on the larger +piece of glazed paper and brushing the precipitate from the +smaller piece into it, sweeping off all particles belonging to the +determination. + +Moisten the precipitate with two drops of concentrated nitric acid and +one drop of concentrated hydrochloric acid, and again heat with great +caution until the acids are expelled and the precipitate is white, +when the temperature is slowly raised until the silver chloride just +begins to fuse at the edges (Note 3). The crucible is then cooled in a +desiccator and weighed, after which the heating (without the addition +of acids) is repeated, and it is again weighed. This must be continued +until the weight is constant within 0.0003 gram in two consecutive +weighings. Deduct the weight of the crucible, and calculate the +percentage of chlorine in the sample of sodium chloride taken for +analysis. + +[Note 1: The separation of the silver chloride from the filter is +essential, since the burning carbon of the paper would reduce a +considerable quantity of the precipitate to metallic silver, and its +complete reconversion to the chloride within the crucible, by means of +acids, would be accompanied by some difficulty. The small amount of +silver reduced from the chloride adhering to the filter paper after +separating the bulk of the precipitate, and igniting the paper +as prescribed, can be dissolved in nitric acid, and completely +reconverted to chloride by hydrochloric acid. The subsequent addition +of the two acids to the main portion of the precipitate restores the +chlorine to any chloride which may have been partially reduced by the +sunlight. The excess of the acids is volatilized by heating.] + +[Note 2: The platinum wire is wrapped around the top of the filter +during its incineration to avoid contact with any reduced silver from +the reduction of the precipitate. If the wire were placed nearer the +apex, such contact could hardly be avoided.] + +[Note 3: Silver chloride should not be heated to complete fusion, +since a slight loss by volatilization is possible at high +temperatures. The temperature of fusion is not always sufficient +to destroy filter shreds; hence these should not be allowed to +contaminate the precipitate.] + + + + +DETERMINATION OF IRON AND OF SULPHUR IN FERROUS AMMONIUM SULPHATE, + +FESO_{4}.(NH_{4})_{2}SO_{4}.6H_{2}O + + +DETERMINATION OF IRON + +PROCEDURE.--Weigh out into beakers (200-250 cc.) two portions of the +sample (Note 1) of about 1 gram each and dissolve these in 50 cc. of +water, to which 1 cc. of dilute hydrochloric acid (sp. gr. 1.12) has +been added (Note 2). Heat the solution to boiling, and while at the +boiling point add concentrated nitric acid (sp. gr. 1.42), !drop by +drop! (noting the volume used), until the brown coloration, which +appears after the addition of a part of the nitric acid, gives place +to a yellow or red (Note 3). Avoid a large excess of nitric acid, but +be sure that the action is complete. Pour this solution cautiously +into about 200 cc. of water, containing a slight excess of ammonia. +Calculate for this purpose the amount of aqueous ammonia required to +neutralize the hydrochloric and nitric acids added (see Appendix for +data), and also to precipitate the iron as ferric hydroxide from the +weight of the ferrous ammonium sulphate taken for analysis, assuming +it to be pure (Note 4). The volume thus calculated will be in excess +of that actually required for precipitation, since the acids are in +part consumed in the oxidation process, or are volatilized. Heat the +solution to boiling, and allow the precipitated ferric hydroxide to +settle. Decant the clear liquid through a washed filter (9 cm.), +keeping as much of the precipitate in the beaker as possible. Wash +twice by decantation with 100 cc. of hot water. Reserve the filtrate. +Dissolve the iron from the filter with hot, dilute hydrochloric acid +(sp. gr. 1.12), adding it in small portions, using as little as +possible and noting the volume used. Collect the solution in the +beaker in which precipitation took place. Add 1 cc. of nitric acid +(sp. gr. 1.42), boil for a few moments, and again pour into a +calculated excess of ammonia. + +Wash the precipitate twice by decantation, and finally transfer it to +the original filter. Wash continuously with hot water until finally +3 cc. of the washings, acidified with nitric acid (Note 5), show +no evidences of the presence of chlorides when tested with silver +nitrate. The filtrate and washings are combined with those from the +first precipitation and treated for the determination of sulphur, as +prescribed on page 112. + +[Note 1: If a selection of pure material for analysis is to be made, +crystals which are cloudy are to be avoided on account of loss of +water of crystallization; and also those which are red, indicating +the presence of ferric iron. If, on the other hand, the value of an +average sample of material is desired, it is preferable to grind the +whole together, mix thoroughly, and take a sample from the mixture for +analysis.] + +[Note 2: When aqueous solutions of ferrous compounds are heated in the +air, oxidation of the Fe^{++} ions to Fe^{+++} ions readily occurs in +the absence of free acid. The H^{+} and OH^{-} ions from water are +involved in the oxidation process and the result is, in effect, the +formation of some ferric hydroxide which tends to separate. Moreover, +at the boiling temperature, the ferric sulphate produced by the +oxidation hydrolyzes in part with the formation of a basic ferric +sulphate, which also tends to separate from solution. The addition of +the hydrochloric acid prevents the formation of ferric hydroxide, and +so far reduces the ionization of the water that the hydrolysis of the +ferric sulphate is also prevented, and no precipitation occurs on +heating.] + +[Note 3: The nitric acid, after attaining a moderate strength, +oxidizes the Fe^{++} ions to Fe^{+++} ions with the formation of an +intermediate nitroso-compound similar in character to that formed in +the "ring-test" for nitrates. The nitric oxide is driven out by heat, +and the solution then shows by its color the presence of ferric +compounds. A drop of the oxidized solution should be tested on +a watch-glass with potassium ferricyanide, to insure a complete +oxidation. This oxidation of the iron is necessary, since Fe^{++} ions +are not completely precipitated by ammonia. + +The ionic changes which are involved in this oxidation are perhaps +most simply expressed by the equation + +3Fe^{++} + NO_{3}^{-}+ 4H^{+} --> 3Fe^{+++} + 2H_{2}O + NO, + +the H^{+} ions coming from the acid in the solution, in this case +either the nitric or the hydrochloric acid. The full equation on which +this is based may be written thus: + +6FeSO_{4} + 2HNO_{3} + 6HCl --> 2Fe_{2}(SO_{4})_{3} + 2FeCl_{3} + 2NO ++ 4H_{2}O, + +assuming that only enough nitric acid is added to complete the +oxidation.] + +[Note 4: The ferric hydroxide precipitate tends to carry down some +sulphuric acid in the form of basic ferric sulphate. This tendency is +lessened if the solution of the iron is added to an excess of OH^{-} +ions from the ammonium hydroxide, since under these conditions +immediate and complete precipitation of the ferric hydroxide ensues. +A gradual neutralization with ammonia would result in the local +formation of a neutral solution within the liquid, and subsequent +deposition of a basic sulphate as a consequence of a local deficiency +of OH^{-} ions from the NH_{4}OH and a partial hydrolysis of the +ferric salt. Even with this precaution the entire absence of sulphates +from the first iron precipitate is not assured. It is, therefore, +redissolved and again thrown down by ammonia. The organic matter of +the filter paper may occasion a partial reduction of the iron during +solution, with consequent possibility of incomplete subsequent +precipitation with ammonia. The nitric acid is added to reoxidize this +iron. + +To avoid errors arising from the solvent action of ammoniacal +liquids upon glass, the iron precipitate should be filtered without +unnecessary delay.] + +[Note 5: The washings from the ferric hydroxide are acidified with +nitric acid, before testing with silver nitrate, to destroy the +ammonia which is a solvent of silver chloride. + +The use of suction to promote filtration and washing is permissible, +though not prescribed. The precipitate should not be allowed to dry +during the washing.] + + +!Ignition of the Iron Precipitate! + +Heat a platinum or porcelain crucible, cool it in a desiccator and +weigh, repeating until a constant weight is obtained. + +Fold the top of the filter paper over the moist precipitate of ferric +hydroxide and transfer it cautiously to the crucible. Wipe the inside +of the funnel with a small fragment of washed filter paper, if +necessary, and place the paper in the crucible. + +Incline the crucible on its side, on a triangle supported on a +ring-stand, and stand the cover on edge at the mouth of the crucible. +Place a burner below the front edge of the crucible, using a low flame +and protecting it from drafts of air by means of a chimney. The heat +from the burner is thus reflected into the crucible and dries +the precipitate without danger of loss as the result of a sudden +generation of steam within the mass of ferric hydroxide. As the drying +progresses the burner may be gradually moved toward the base of the +crucible and the flame increased until the paper of the filter begins +to char and finally to smoke, as the volatile matter is expelled. This +is known as "smoking off" a filter, and the temperature should not be +raised sufficiently high during this process to cause the paper to +ignite, as the air currents produced by the flame of the blazing paper +may carry away particles of the precipitate. + +When the paper is fully charred, move the burner to the base of the +crucible and raise the temperature to the full heat of the burner for +fifteen minutes, with the crucible still inclined on its side, but +without the cover (Note 1). Finally set the crucible upright in the +triangle, cover it, and heat at the full temperature of a blast lamp +or other high temperature burner. Cool and weigh in the usual manner +(Note 2). Repeat the strong heating until the weight is constant +within 0.0003 gram. + +From the weight of ferric oxide (Fe_{2}O_{3}) calculate the percentage +of iron (Fe) in the sample (Note 3). + +[Note 1: These directions for the ignition of the precipitate must be +closely followed. A ready access of atmospheric oxygen is of special +importance to insure the reoxidation to ferric oxide of any iron which +may be reduced to magnetic oxide (Fe_{3}O_{4}) during the combustion +of the filter. The final heating over the blast lamp is essential +for the complete expulsion of the last traces of water from the +hydroxide.] + +[Note 2: Ignited ferric oxide is somewhat hygroscopic. On this account +the weighings must be promptly completed after removal from the +desiccator. In all weighings after the first it is well to place the +weights upon the balance-pan before removing the crucible from the +desiccator. It is then only necessary to move the rider to obtain the +weight.] + +[Note 3: The gravimetric determination of aluminium or chromium is +comparable with that of iron just described, with the additional +precaution that the solution must be boiled until it contains but a +very slight excess of ammonia, since the hydroxides of aluminium and +chromium are more soluble than ferric hydroxide. + +The most important properties of these hydroxides, from a quantitative +standpoint, other than those mentioned, are the following: All are +precipitable by the hydroxides of sodium and potassium, but always +inclose some of the precipitant, and should be reprecipitated with +ammonium hydroxide before ignition to oxides. Chromium and aluminium +hydroxides dissolve in an excess of the caustic alkalies and form +anions, probably of the formula AlO_2^{-} and CrO_{2}^{-}. Chromium +hydroxide is reprecipitated from this solution on boiling. When first +precipitated the hydroxides are all readily soluble in acids, but +aluminium hydroxide dissolves with considerable difficulty after +standing or boiling for some time. The precipitation of the hydroxides +is promoted by the presence of ammonium chloride, but is partially +or entirely prevented by the presence of tartaric or citric acids, +glycerine, sugars, and some other forms of soluble organic matter. +The hydroxides yield on ignition an oxide suitable for weighing +(Al_{2}O_{3}, Cr_{2}O_{3}, Fe_{2}O_{3}).] + + + + +DETERMINATION OF SULPHUR + + +PROCEDURE.--Add to the combined filtrates from the ferric hydroxide +about 0.6 gram of anhydrous sodium carbonate; cover the beaker, and +then add dilute hydrochloric acid (sp. gr. 1.12) in moderate excess +and evaporate to dryness on the water bath. Add 10 cc. of concentrated +hydrochloric acid (sp. gr. 1.20) to the residue, and again evaporate +to dryness on the bath. Dissolve the residue in water, filter if not +clear, transfer to a 700 cc. beaker, dilute to about 400 cc., and +cautiously add hydrochloric acid until the solution shows a distinctly +acid reaction (Note 1). Heat the solution to boiling, and add !very +slowly! and with constant stirring, 20 cc. in excess of the calculated +amount of a hot barium chloride solution, containing about 20 grams +BaCl_{2}.2H_{2}O per liter (Notes 2 and 3). Continue the boiling for +about two minutes, allow the precipitate to settle, and decant the +liquid at the end of half an hour (Note 4). Replace the beaker +containing the original filtrate by a clean beaker, wash the +precipitated sulphate by decantation with hot water, and subsequently +upon the filter until it is freed from chlorides, testing the washings +as described in the determination of iron. The filter is then +transferred to a platinum or porcelain crucible and ignited, as +described above, until the weight is constant (Note 5). From the +weight of barium sulphate (BaSO_{4}) obtained, calculate the +percentage of sulphur (S) in the sample. + +[Note 1: Barium sulphate is slightly soluble in hydrochloric acid, +even dilute, probably as a result of the reduction in the degree of +dissociation of sulphuric acid in the presence of the H^{+} ions of +the hydrochloric acid, and possibly because of the formation of a +complex anion made up of barium and chlorine; hence only the smallest +excess should be added over the amount required to acidify the +solution.] + +[Note 2: The ionic changes involved in the precipitation of barium +sulphate are very simple: + +Ba^{++} + SO_{4}^{--} --> [BaSO_{4}] + +This case affords one of the best illustrations of the effect of an +excess of a precipitant in decreasing the solubility of a precipitate. +If the conditions are considered which exist at the moment when just +enough of the Ba^{++} ions have been added to correspond to the +SO_{4}^{--} ions in the solution, it will be seen that nearly all of +the barium sulphate has been precipitated, and that the small amount +which then remains in the solution which is in contact with the +precipitate must represent a saturated solution for the existing +temperature, and that this solution is comparable with a solution of +sugar to which more sugar has been added than will dissolve. It +should be borne in mind that the quantity of barium sulphate in +this !saturated solution is a constant quantity! for the existing +conditions. The dissolved barium sulphate, like any electrolyte, is +dissociated, and the equilibrium conditions may be expressed thus: + +(!Conc'n Ba^{++} x Conc'n SO_{4}^{--})/(Conc'n BaSO_{4}) = Const.!, + +and since !Conc'n BaSO_{4}! for the saturated solution has a constant +value (which is very small), it may be eliminated, when the expression +becomes !Conc'n Ba^{++} x Conc'n SO_{4}^{--} = Const.!, which is +the "solubility product" of BaSO_{4}. If, now, an excess of the +precipitant, a soluble barium salt, is added in the form of a +relatively concentrated solution (the slight change of volume of a few +cubic centimeters may be disregarded for the present discussion) +the concentration of the Ba^{++} ions is much increased, and as a +consequence the !Conc'n SO_{4}! must decrease in proportion if the +value of the expression is to remain constant, which is a requisite +condition if the law of mass action upon which our argument depends +holds true. In other words, SO_{4}^{--} ions must combine with some +of the added Ba^{++} ions to form [BaSO_{4}]; but it will be recalled +that the solution is already saturated with BaSO_{4}, and this freshly +formed quantity must, therefore, separate and add itself to the +precipitate. This is exactly what is desired in order to insure +more complete precipitation and greater accuracy, and leads to the +conclusion that the larger the excess of the precipitant added the +more successful the analysis; but a practical limit is placed upon +the quantity of the precipitant which may be properly added by other +conditions, as stated in the following note.] + +[Note 3: Barium sulphate, in a larger measure than most compounds, +tends to carry down other substances which are present in the solution +from which it separates, even when these other substances are +relatively soluble, and including the barium chloride used as the +precipitant. This is also notably true in the case of nitrates and +chlorates of the alkalies, and of ferric compounds; and, since in this +analysis ammonium nitrate has resulted from the neutralization of the +excess of the nitric acid added to oxidize the iron, it is essential +that this should be destroyed by repeated evaporation with a +relatively large quantity of hydrochloric acid. During evaporation a +mutual decomposition of the two acids takes place, and the nitric acid +is finally decomposed and expelled by the excess of hydrochloric acid. + +Iron is usually found in the precipitate of barium sulphate when +thrown down from hot solutions in the presence of ferric salts. This, +according to Kuster and Thiel (!Zeit. anorg. Chem.!, 22, 424), is due +to the formation of a complex ion (Fe(SO_{4})_{2}) which precipitates +with the Ba^{++} ion, while Richards (!Zeit. anorg. Chem.!, 23, 383) +ascribes it to hydrolytic action, which causes the formation of a +basic ferric complex which is occluded in the barium precipitate. +Whatever the character of the compound may be, it has been shown that +it loses sulphuric anhydride upon ignition, causing low results, even +though the precipitate contains iron. + +The contamination of the barium sulphate by iron is much less in the +presence of ferrous than ferric salts. If, therefore, the sulphur +alone were to be determined in the ferrous ammonium sulphate, the +precipitation by barium might be made directly from an aqueous +solution of the salt, which had been made slightly acid with +hydrochloric acid.] + +[Note 4: The precipitation of the barium sulphate is probably complete +at the end of a half-hour, and the solution may safely be filtered at +the expiration of that time if it is desired to hasten the analysis. + +As already noted, many precipitates of the general character of this +sulphate tend to grow more coarsely granular if digested for some time +with the liquid from which they have separated. It is therefore well +to allow the precipitate to stand in a warm place for several hours, +if practicable, to promote ease of filtration. The filtrate and +washings should always be carefully examined for minute quantities of +the sulphate which may pass through the pores of the filter. This is +best accomplished by imparting to the filtrate a gentle rotary motion, +when the sulphate, if present, will collect at the center of the +bottom of the beaker.] + +[Note 5: A reduction of barium sulphate to the sulphide may very +readily be caused by the reducing action of the burning carbon of the +filter, and much care should be taken to prevent any considerable +reduction from this cause. Subsequent ignition, with ready access +of air, reconverts the sulphide to sulphate unless a considerable +reduction has occurred. In the latter case it is expedient to add one +or two drops of sulphuric acid and to heat cautiously until the excess +of acid is expelled.] + +[Note 6: Barium sulphate requires about 400,000 parts of water for +its solution. It is not decomposed at a red heat but suffers loss, +probably of sulphur trioxide, at a temperature above 900°C.] + + + + +DETERMINATION OF SULPHUR IN BARIUM SULPHATE + + +PROCEDURE.--Weigh out, into platinum crucibles, two portions of about +0.5 gram of the sulphate. Mix each in the crucible with five to six +times its weight of anhydrous sodium carbonate. This can best be done +by placing the crucible on a piece of glazed paper and stirring the +mixture with a clean, dry stirring-rod, which may finally be wiped off +with a small fragment of filter paper, the latter being placed in the +crucible. Cover the crucible and heat until a quiet, liquid fusion +ensues. Remove the burner, and tip the crucible until the fused mass +flows nearly to its mouth. Hold it in that position until the mass has +solidified. When cold, the material may usually be detached in a lump +by tapping the crucible or gently pressing it near its upper edge. If +it still adheres, a cubic centimeter or so of water may be placed in +the cold crucible and cautiously brought to boiling, when the cake +will become loosened and may be removed and placed in about 250 cc. +of hot, distilled water to dissolve. Clean the crucible completely, +rubbing the sides with a rubber-covered stirring-rod, if need be. + +When the fused mass has completely disintegrated and nothing further +will dissolve, decant the solution from the residue of barium +carbonate (Note 1). Pour over the residue 20 cc. of a solution of +sodium carbonate and 10 cc. of water and heat to gentle boiling for +about three minutes (Note 2). Filter off the carbonate and wash it +with hot water, testing the slightly acidified washings for sulphate +and preserving any precipitates which appear in these tests. Acidify +the filtrate with hydrochloric acid until just acid, bring to boiling, +and slowly add hot barium chloride solution, as in the preceding +determination. Add also any tests from the washings in which +precipitates have appeared. Filter, wash, ignite, and weigh. + +From the weight of barium sulphate, calculate the percentage of +sulphur (S) in the sample. + +[Note 1: This alkaline fusion is much employed to disintegrate +substances ordinarily insoluble in acids into two components, one +of which is water soluble and the other acid soluble. The reaction +involved is: + +BaSO_{4} + Na_{2}CO_{3}, --> BaCO_{3}, + Na_{2}SO_{4}. + +As the sodium sulphate is soluble in water, and the barium carbonate +insoluble, a separation between them is possible and the sulphur can +be determined in the water-soluble portion. + +It should be noted that this method can be applied to the purification +of a precipitate of barium sulphate if contaminated by most of the +substances mentioned in Note 3 on page 114. The impurities pass into +the water solution together with the sodium sulphate, but, being +present in such minute amounts, do not again precipitate with the +barium sulphate.] + +[Note 2: The barium carbonate is boiled with sodium carbonate solution +before filtration because the reaction above is reversible; and it is +only by keeping the sodium carbonate present in excess until nearly +all of the sodium sulphate solution has been removed by filtration +that the reversion of some of the barium carbonate to barium sulphate +is prevented. This is an application of the principle of mass action, +in which the concentration of the reagent (the carbonate ion) is +kept as high as practicable and that of the sulphate ion as low as +possible, in order to force the reaction in the desired direction (see +Appendix).] + + + + +DETERMINATION OF PHOSPHORIC ANHYDRIDE IN APATITE + + +The mineral apatite is composed of calcium phosphate, associated with +calcium chloride, or fluoride. Specimens are easily obtainable which +are nearly pure and leave on treatment with acid only a slight +siliceous residue. + +For the purpose of gravimetric determination, phosphoric acid is +usually precipitated from ammoniacal solutions in the form of +magnesium ammonium phosphate which, on ignition, is converted into +magnesium pyrophosphate. Since the calcium phosphate of the apatite +is also insoluble in ammoniacal solutions, this procedure cannot be +applied directly. The separation of the phosphoric acid from the +calcium must first be accomplished by precipitation in the form of +ammonium phosphomolybdate in nitric acid solution, using ammonium +molybdate as the precipitant. The "yellow precipitate," as it is often +called, is not always of a definite composition, and therefore not +suitable for direct weighing, but may be dissolved in ammonia, and the +phosphoric acid thrown out as magnesium ammonium phosphate from the +solution. + +Of the substances likely to occur in apatite, silicic acid alone +interferes with the precipitation of the phosphoric acid in nitric +acid solution. + + +PRECIPITATION OF AMMONIUM PHOSPHOMOLYBDATE + +PROCEDURE.--Grind the mineral in an agate mortar until no grit is +perceptible. Transfer the substance to a weighing-tube, and weigh out +two portions, not exceeding 0.20 gram each (Note 1) into two beakers +of about 200 cc. capacity. Pour over them 20 cc. of dilute nitric acid +(sp. gr. 1.2) and warm gently until solvent action has apparently +ceased. Evaporate the solution cautiously to dryness, heat the residue +for about an hour at 100-110°C., and treat it again with nitric acid +as described above; separate the residue of silica by filtration on +a small filter (7 cm.) and wash with warm water, using as little as +possible (Note 2). Receive the filtrate in a beaker (200-500 cc.). +Test the washings with ammonia for calcium phosphate, but add all such +tests in which a precipitate appears to the original nitrate (Note 3). +The filtrate and washings must be kept as small as possible and should +not exceed 100 cc. in volume. Add aqueous ammonia (sp. gr. 0.96) until +the precipitate of calcium phosphate first produced just fails to +redissolve, and then add a few drops of nitric acid until this is +again brought into solution (Note 4). Warm the solution until it +cannot be comfortably held in the hand (about 60°C.) and, after +removal of the burner, add 75 cc. of ammonium molybdate solution which +has been !gently! warmed, but which must be perfectly clear. Allow +the mixture to stand at a temperature of about 50 or 60°C. for twelve +hours (Notes 5 and 6). Filter off the yellow precipitate on a 9 cm. +filter, and wash by decantation with a solution of ammonium nitrate +made acid with nitric acid.[1] Allow the precipitate to remain in the +beaker as far as possible. Test the washings for calcium with ammonia +and ammonium oxalate (Note 3). + +[Footnote 1: This solution is prepared as follows: Mix 100 cc. of +ammonia solution (sp. gr. 0.96) with 325 cc. of nitric acid (sp. gr. +1.2) and dilute with 100 cc. of water.] + +Add 10 cc. of molybdate solution to the nitrate, and leave it for +a few hours. It should then be carefully examined for a !yellow! +precipitate; a white precipitate may be neglected. + +[Note 1: Magnesium ammonium phosphate, as noted below, is slightly +soluble under the conditions of operation. Consequently the +unavoidable errors of analysis are greater in this determination than +in those which have preceded it, and some divergence may be expected +in duplicate analyses. It is obvious that the larger the amount of +substance taken for analysis the less will be the relative loss or +gain due to unavoidable experimental errors; but, in this instance, a +check is placed upon the amount of material which may be taken both by +the bulk of the resulting precipitate of ammonium phosphomolybdate +and by the excessive amount of ammonium molybdate required to effect +complete separation of the phosphoric acid, since a liberal excess +above the theoretical quantity is demanded. Molybdic acid is one of +the more expensive reagents.] + +[Note 2: Soluble silicic acid would, if present, partially separate +with the phosphomolybdate, although not in combination with +molybdenum. Its previous removal by dehydration is therefore +necessary.] + +[Note 3: When washing the siliceous residue the filtrate may be tested +for calcium by adding ammonia, since that reagent neutralizes the +acid which holds the calcium phosphate in solution and causes +precipitation; but after the removal of the phosphoric acid in +combination with the molybdenum, the addition of an oxalate is +required to show the presence of calcium.] + +[Note 4: An excess of nitric acid exerts a slight solvent +action, while ammonium nitrate lessens the solubility; hence the +neutralization of the former by ammonia.] + +[Note 5: The precipitation of the phosphomolybdate takes place more +promptly in warm than in cold solutions, but the temperature should +not exceed 60°C. during precipitation; a higher temperature tends to +separate molybdic acid from the solution. This acid is nearly white, +and its deposition in the filtrate on long standing should not be +mistaken for a second precipitation of the yellow precipitate. The +addition of 75 cc. of ammonium molybdate solution insures the presence +of a liberal excess of the reagent, but the filtrate should be tested +as in all quantitative procedures. + +The precipitation is probably complete in many cases in less than +twelve hours; but it is better, when practicable, to allow the +solution to stand for this length of time. Vigorous shaking or +stirring promotes the separation of the precipitate.] + +[Note 6: The composition of the "yellow precipitate" undoubtedly +varies slightly with varying conditions at the time of its formation. +Its composition may probably fairly be represented by the formula, +(NH_{4})_{3}PO_{4}.12MoO_{3}.H_{2}O, when precipitated under the +conditions prescribed in the procedure. Whatever other variations may +occur in its composition, the ratio of 12 MoO_{3}:1 P seems to +hold, and this fact is utilized in volumetric processes for the +determination of phosphorus, in which the molybdenum is reduced to +a lower oxide and reoxidized by a standard solution of potassium +permanganate. In principle, the procedure is comparable with that +described for the determination of iron by permanganate.] + + +PRECIPITATION OF MAGNESIUM AMMONIUM PHOSPHATE + +PROCEDURE.--Dissolve the precipitate of phosphomolybdate upon the +filter by pouring through it dilute aqueous ammonia (one volume of +dilute ammonia (sp. gr. 0.96) and three volumes of water, which +should be carefully measured), and receive the solution in the beaker +containing the bulk of the precipitate. The total volume of nitrate +and washings should not much exceed 100 cc. Acidify the solution with +dilute hydrochloric acid, and heat it nearly to boiling. Calculate the +volume of magnesium ammonium chloride solution ("magnesia mixture") +required to precipitate the phosphoric acid, assuming 40 per cent +P_{2}O_{5} in the apatite. Measure out about 5 cc. in excess of this +amount, and pour it into the acid solution. Then add slowly dilute +ammonium hydroxide (1 volume of strong ammonia (sp. gr. 0.90) and 9 +volumes of water), stirring constantly until a precipitate forms. Then +add a volume of filtered, concentrated ammonia (sp. gr. 0.90) equal to +one third of the volume of liquid in the beaker (Note 1). Allow the +whole to cool. The precipitated magnesium ammonium phosphate should +then be definitely crystalline in appearance (Note 2). (If it is +desired to hasten the precipitation, the solution may be cooled, first +in cold and then in ice-water, and stirred !constantly! for half an +hour, when precipitation will usually be complete.) + +Decant the clear liquid through a filter, and transfer the precipitate +to the filter, using as wash-water a mixture of one volume of +concentrated ammonia and three volumes of water. It is not necessary +to clean the beaker completely or to wash the precipitate thoroughly +at this point, as it is necessary to purify it by reprecipitation. + +[Note 1: Magnesium ammonium phosphate is not a wholly insoluble +substance, even under the most favorable analytical conditions. It +is least soluble in a liquid containing one fourth of its volume of +concentrated aqueous ammonia (sp. gr. 0.90) and this proportion should +be carefully maintained as prescribed in the procedure. On account of +this slight solubility the volume of solutions should be kept as small +as possible and the amount of wash-water limited to that absolutely +required. + +A large excess of the magnesium solution tends both to throw out +magnesium hydroxide (shown by a persistently flocculent precipitate) +and to cause the phosphate to carry down molybdic acid. The tendency +of the magnesium precipitate to carry down molybdic acid is also +increased if the solution is too concentrated. The volume should not +be less than 90 cc., nor more than 125 cc., at the time of the first +precipitation with the magnesia mixture.] + +[Note 2: The magnesium ammonium phosphate should be perfectly +crystalline, and will be so if the directions are followed. The slow +addition of the reagent is essential, and the stirring not less so. +Stirring promotes the separation of the precipitate and the formation +of larger crystals, and may therefore be substituted for digestion in +the cold. The stirring-rod must not be allowed to scratch the glass, +as the crystals adhere to such scratches and are removed with +difficulty.] + + +REPRECIPITATION AND IGNITION OF MAGNESIUM AMMONIUM PHOSPHATE + +A single precipitation of the magnesium compound in the presence of +molybdenum compounds rarely yields a pure product. The molybdenum can +be removed by solution of the precipitate in acid and precipitation +of the molybdenum by sulphureted hydrogen, after which the magnesium +precipitate may be again thrown down. It is usually more satisfactory +to dissolve the magnesium precipitate and reprecipitate the phosphate +as magnesium ammonium phosphate as described below. + +PROCEDURE.--Dissolve the precipitate from the filter in a little +dilute hydrochloric acid (sp. gr. 1.12), allowing the acid solution to +run into the beaker in which the original precipitation was made (Note +1). Wash the filter with water until the wash-water shows no test for +chlorides, but avoid an unnecessary amount of wash-water. Add to +the solution 2 cc. (not more) of magnesia mixture, and then dilute +ammonium hydroxide solution (sp. gr. 0.96), drop by drop, with +constant stirring, until the liquid smells distinctly of ammonia. Stir +for a few moments and then add a volume of strong ammonia (sp. gr. +0.90), equal to one third of the volume of the solution. Allow the +solution to stand for some hours, and then filter off the magnesium +ammonium phosphate, which should be distinctly crystalline in +character. Wash the precipitate with dilute ammonia water, as +prescribed above, until, finally, 3 cc. of the washings, after +acidifying with nitric acid, show no evidence of chlorides. Test both +filtrates for complete precipitation by adding a few cubic centimeters +of magnesia mixture and allowing them to stand for some time. + +Transfer the moist precipitate to a weighed porcelain or platinum +crucible and ignite, using great care to raise the temperature slowly +while drying the filter in the crucible, and to insure the ready +access of oxygen during the combustion of the filter paper, thus +guarding against a possible reduction of the phosphate, which would +result in disastrous consequences both to the crucible, if of +platinum, and the analysis. Do not raise the temperature above +moderate redness until the precipitate is white. (Keep this precaution +well in mind.) Ignite finally at the highest temperature of the +Tirrill burner, and repeat the heating until the weight is constant. +If the ignited precipitate is persistently discolored by particles of +unburned carbon, moisten the mass with a drop or two of concentrated +nitric acid and heat cautiously, finally igniting strongly. The +acid will dissolve magnesium pyrophosphate from the surface of the +particles of carbon, which will then burn away. Nitric acid also aids +as an oxidizing agent in supplying oxygen for the combustion of the +carbon. + +From the weight of magnesium pyrophosphate (Mg_{2}P_{2}O_{7}) +obtained, calculate the phosphoric anhydride (P_{2}O_{5}) in the +sample of apatite. + +[Note 1: The ionic change involved in the precipitation of the +magnesium compound is + +PO_{4}^{---} + NH_{4}^{+} + Mg^{++} --> [MgNH_{4}PO_{4}]. + +The magnesium ammonium phosphate is readily dissolved by acids, even +those which are no stronger than acetic acid. This is accounted for +by the fact that two of the ions into which phosphoric acid may +dissociate, the HPO_{4}^{--} or H_{2}PO_{4}^{-} ions, exhibit the +characteristics of very weak acids, in that they show almost no +tendency to dissociate further into H^{+} and PO_{4}^{--} ions. +Consequently the ionic changes which occur when the magnesium ammonium +phosphate is brought into contact with an acid may be typified by the +reaction: + +H^{+} + Mg^{++} + NH_{4}^{+} + PO_{4}^{---} --> Mg^{++} + NH_{4}^{+} + +HPO_{4}^{--}; + +that is, the PO_{4}^{--} ions and the H^{+} ions lose their identity +in the formation of the new ion, HPO_{4}^{--}, and this continues +until the magnesium ammonium phosphate is entirely dissolved.] + +[Note 2: During ignition the magnesium ammonium phosphate loses +ammonia and water and is converted into magnesium pyrophosphate: + +2MgNH_{4}PO_{4} --> Mg_{2}P_{2}O_{7} + 2NH_{3} + H_{2}O. + +The precautions mentioned on pages 111 and 123 must be observed with +great care during the ignition of this precipitate. The danger here +lies in a possible reduction of the phosphate by the carbon of the +filter paper, or by the ammonia evolved, which may act as a reducing +agent. The phosphorus then attacks and injures a platinum crucible, +and the determination is valueless.] + + + + +ANALYSIS OF LIMESTONE + + +Limestones vary widely in composition from a nearly pure marble +through the dolomitic limestones, containing varying amounts of +magnesium, to the impure varieties, which contain also ferrous and +manganous carbonates and siliceous compounds in variable proportions. +Many other minerals may be inclosed in limestones in small quantities, +and an exact qualitative analysis will often show the presence of +sulphides or sulphates, phosphates, and titanates, and the alkali or +even the heavy metals. No attempt is made in the following procedures +to provide a complete quantitative scheme which would take into +account all of these constituents. Such a scheme for a complete +analysis of a limestone may be found in Bulletin No. 700 of the United +States Geological Survey. It is assumed that, for these practice +determinations, a limestone is selected which contains only the more +common constituents first enumerated above. + + +DETERMINATION OF MOISTURE + +The determination of the amount of moisture in minerals or ores is +often of great importance. Ores which have been exposed to the weather +during shipment may have absorbed enough moisture to appreciably +affect the results of analysis. Since it is essential that the seller +and buyer should make their analyses upon comparable material, it is +customary for each analyst to determine the moisture in the sample +examined, and then to calculate the percentages of the various +constituents with reference to a sample dried in the air, or at a +temperature a little above 100°C., which, unless the ore has undergone +chemical change because of the wetting, should be the same before and +after shipment. + +PROCEDURE.--Spread 25 grams of the powdered sample on a weighed +watch-glass; weigh to the nearest 10 milligrams only and heat at +105°C.; weigh at intervals of an hour, after cooling in a desiccator, +until the loss of weight after an hour's heating does not exceed +10 milligrams. It should be noted that a variation in weight of 10 +milligrams in a total weight of 25 grams is no greater relatively than +a variation of 0.1 milligram when the sample taken weighs 0.25 gram + +DETERMINATION OF THE INSOLUBLE MATTER AND SILICA + +PROCEDURE.--Weigh out two portions of the original powdered sample +(not the dried sample), of about 5 grams each, into 250 cc. +casseroles, and cover each with a watch-glass (Note 1). Pour over the +powder 25 cc. of water, and then add 50 cc. of dilute hydrochloric +acid (sp. gr. 1.12) in small portions, warming gently, until nothing +further appears to dissolve (Note 2). Evaporate to dryness on the +water bath. Pour over the residue a mixture of 5 cc. of water and +5 cc. of concentrated hydrochloric acid (sp. gr. 1.2) and again +evaporate to dryness, and finally heat for at least an hour at +a temperature of 110°C. Pour over this residue 50 cc. of dilute +hydrochloric acid (one volume acid (sp. gr. 1.12) to five volumes +water), and boil for about five minutes; then filter and wash twice +with the dilute hydrochloric acid, and then with hot water until +free from chlorides. Transfer the filter and contents to a porcelain +crucible, dry carefully over a low flame, and ignite to constant +weight. The residue represents the insoluble matter and the silica +from any soluble silicates (Note 3). + +Calculate the combined percentage of these in the limestone. + +[Note 1: The relatively large weight (5 grams) taken for analysis +insures greater accuracy in the determination of the ingredients which +are present in small proportions, and is also more likely to be a +representative sample of the material analyzed.] + +[Note 2: It is plain that the amount of the insoluble residue and also +its character will often depend upon the strength of acid used for +solution of the limestone. It cannot, therefore, be regarded as +representing any well-defined constituent, and its determination is +essentially empirical.] + +[Note 3: It is probable that some of the silicates present are wholly +or partly decomposed by the acid, and the soluble silicic acid must +be converted by evaporation to dryness, and heating, into white, +insoluble silica. This change is not complete after one evaporation. +The heating at a temperature somewhat higher than that of the water +bath for a short time tends to leave the silica in the form of a +powder, which promotes subsequent filtration. The siliceous residue +is washed first with dilute acid to prevent hydrolytic changes, which +would result in the formation of appreciable quantities of insoluble +basic iron or aluminium salts on the filter when washing with hot +water. + +If it is desired to determine the percentage of silica separately, the +ignited residue should be mixed in a platinum crucible with about six +times its weight of anhydrous sodium carbonate, and the procedure +given on page 151 should be followed. The filtrate from the silica is +then added to the main filtrate from the insoluble residue.] + + + + +DETERMINATION OF FERRIC OXIDE AND ALUMINIUM OXIDE (WITH MANGANESE) + + +PROCEDURE.--To the filtrate from the insoluble residue add ammonium +hydroxide until the solution just smells distinctly of ammonia, but do +not add an excess. Then add 5 cc. of saturated bromine water (Note 1), +and boil for five minutes. If the smell of ammonia has disappeared, +again add ammonium hydroxide in slight excess, and 3 cc. of bromine +water, and heat again for a few minutes. Finally add 10 cc. of +ammonium chloride solution and keep the solution warm until it barely +smells of ammonia; then filter promptly (Note 2). Wash the filter +twice with hot water, then (after replacing the receiving beaker) pour +through it 25 cc. of hot, dilute hydrochloric acid (one volume dilute +HCl [sp. gr. 1.12] to five volumes water). A brown residue insoluble +in the acid may be allowed to remain on the filter. Wash the filter +five times with hot water, add to the filtrate ammonium hydroxide +and bromine water as described above, and repeat the precipitation. +Collect the precipitate on the filter already used, wash it free from +chlorides with hot water, and ignite and weigh as described for ferric +hydroxide on page 110. The residue after ignition consists of ferric +oxide, alumina, and mangano-manganic oxide (Mn_{3}O_{4}), if manganese +is present. These are commonly determined together (Note 3). + +Calculate the percentage of the combined oxides in the limestone. + +[Note 1: The addition of bromine water to the ammoniacal solutions +serves to oxidize any ferrous hydroxide to ferric hydroxide and to +precipitate manganese as MnO(OH)_{2}. The solution must contain not +more than a bare excess of hydroxyl ions (ammonium hydroxide) when it +is filtered, on account of the tendency of the aluminium hydroxide to +redissolve. + +The solution should not be strongly ammoniacal when the bromine is +added, as strong ammonia reacts with the bromine, with the evolution +of nitrogen.] + +[Note 2: The precipitate produced by ammonium hydroxide and bromine +should be filtered off promptly, since the alkaline solution absorbs +carbon dioxide from the air, with consequent partial precipitation +of the calcium as carbonate. This is possible even under the most +favorable conditions, and for this reason the iron precipitate is +redissolved and again precipitated to free it from calcium. When the +precipitate is small, this reprecipitation may be omitted.] + +[Note 3: In the absence of significant amounts of manganese the iron +and aluminium may be separately determined by fusion of the mixed +ignited precipitate, after weighing, with about ten times its weight +of acid potassium sulphate, solution of the cold fused mass in water, +and volumetric determination of the iron, as described on page 66. +The aluminium is then determined by difference, after subtracting the +weight of ferric oxide corresponding to the amount of iron found. + +If a separate determination of the iron, aluminium, and manganese +is desired, the mixed precipitate may be dissolved in acid before +ignition, and the separation effected by special methods (see, for +example, Fay, !Quantitative Analyses!, First Edition, pp. 15-19 and +23-27).] + + + + +DETERMINATION OF CALCIUM + + +PROCEDURE.--To the combined filtrates from the double precipitation of +the hydroxides just described, add 5 cc. of dilute ammonium hydroxide +(sp. gr. 0.96), and transfer the liquid to a 500 cc. graduated flask, +washing out the beaker carefully. Cool to laboratory temperature, and +fill the flask with distilled water until the lowest point of the +meniscus is exactly level with the mark on the neck of the flask. +Carefully remove any drops of water which are on the inside of the +neck of the flask above the graduation by means of a strip of filter +paper, make the solution uniform by pouring it out into a dry beaker +and back into the flask several times. Measure off one fifth of this +solution as follows (Note 1): Pour into a 100 cc. graduated flask +about 10 cc. of the solution, shake the liquid thoroughly over the +inner surface of the small flask, and pour it out. Repeat the same +operation. Fill the 100 cc. flask until the lowest point of the +meniscus is exactly level with the mark on its neck, remove any drops +of solution from the upper part of the neck with filter paper, and +pour the solution into a beaker (400-500 cc.). Wash out the flask with +small quantities of water until it is clean, adding these to the 100 +cc. of solution. When the duplicate portion of 100 cc. is measured out +from the solution, remember that the flask must be rinsed out twice +with that solution, as prescribed above, before the measurement is +made. (A 100 cc. pipette may be used to measure out the aliquot +portions, if preferred.) + +Dilute each of the measured portions to 250 cc. with distilled water, +heat the whole to boiling, and add ammonium oxalate solution slowly +in moderate excess, stirring well. Boil for two minutes; allow the +precipitated calcium oxalate to settle for a half-hour, and decant +through a filter. Test the filtrate for complete precipitation by +adding a few cubic centimeters of the precipitant, allowing it to +stand for fifteen minutes. If no precipitate forms, make the solution +slightly acid with hydrochloric acid (Note 2); see that it is properly +labeled and reserve it to be combined with the filtrate from the +second calcium oxalate precipitation (Notes 3 and 4). + +Redissolve the calcium oxalate in the beaker with warm hydrochloric +acid, pouring the acid through the filter. Wash the filter five times +with water, and finally pour through it aqueous ammonia. Dilute the +solution to 250 cc., bring to boiling, and add 1 cc. ammonium oxalate +solution (Note 5) and ammonia in slight excess; boil for two minutes, +and set aside for a half-hour. Filter off the calcium oxalate upon the +filter first used, and wash free from chlorides. The filtrate should +be made barely acid with hydrochloric acid and combined with the +filtrate from the first precipitation. Begin at once the evaporation +of the solutions for the determination of magnesium as described +below. + +The precipitate of calcium oxalate may be converted into calcium oxide +by ignition without previous drying. After burning the filter, it may +be ignited for three quarters of an hour in a platinum crucible at +the highest heat of the Bunsen or Tirrill burner, and finally for ten +minutes at the blast lamp (Note 6). Repeat the heating over the blast +lamp until the weight is constant. As the calcium oxide absorbs +moisture from the air, it must (after cooling) be weighed as rapidly +as possible. + +The precipitate may, if preferred, be placed in a weighted porcelain +crucible. After burning off the filter and heating for ten minutes the +calcium precipitate may be converted into calcium sulphate by placing +2 cc. of dilute sulphuric acid in the crucible (cold), heating the +covered crucible very cautiously over a low flame to drive off the +excess of acid, and finally at redness to constant weight (Note 7). + +From the weight of the oxide or sulphate, calculate the percentage of +the calcium (Ca) in the limestone, remembering that only one fifth of +the total solution is used for this determination. + +[Note 1: If the calcium were precipitated from the entire solution, +the quantity of the precipitate would be greater than could be +properly treated. The solution is, therefore, diluted to a definite +volume (500 cc.), and exactly one fifth (100 cc.) is measured off in a +graduated flask or by means of a pipette.] + +[Note 2: The filtrate from the calcium oxalate should be made slightly +acid immediately after filtration, in order to avoid the solvent +action of the alkaline liquid upon the glass.] + +[Note 3: The accurate quantitative separation of calcium and magnesium +as oxalates requires considerable care. The calcium precipitate +usually carries down with it some magnesium, and this can best +be removed by redissolving the precipitate after filtration, and +reprecipitation in the presence of only the small amount of magnesium +which was included in the first precipitate. When, however, the +proportion of magnesium is not very large, the second precipitation of +the calcium can usually be avoided by precipitating it from a rather +dilute solution (800 cc. or so) and in the presence of a considerable +excess of the precipitant, that is, rather more than enough to convert +both the magnesium and calcium into oxalates.] + +[Note 4: The ionic changes involved in the precipitation of calcium +as oxalate are exceedingly simple, and the principles discussed in +connection with the barium sulphate precipitation on page 113 also +apply here. The reaction is + +C_{2}O_{4}^{--} + Ca^{++} --> [CaC_{2}O_{4}]. + +Calcium oxalate is nearly insoluble in water, and only very slightly +soluble in acetic acid, but is readily dissolved by the strong mineral +acids. This behavior with acids is explained by the fact that oxalic +acid is a stronger acid than acetic acid; when, therefore, the oxalate +is brought into contact with the latter there is almost no tendency to +diminish the concentration of C_{2}O_{4}^{--} ions by the formation of +an acid less dissociated than the acetic acid itself, and practically +no solvent action ensues. When a strong mineral acid is present, +however, the ionization of the oxalic acid is much reduced by the high +concentration of the H^{+} ions from the strong acid, the formation +of the undissociated acid lessens the concentration of the +C_{2}O_{4}^{--} ions in solution, more of the oxalate passes into +solution to re-establish equilibrium, and this process repeats itself +until all is dissolved. + +The oxalate is immediately reprecipitated from such a solution on the +addition of OH^{-} ions, which, by uniting with the H^{+} ions of the +acids (both the mineral acid and the oxalic acid) to form water, leave +the Ca^{++} and C_{2}O_{4}^{--} ions in the solution to recombine to +form [CaC_{2}O_{4}], which is precipitated in the absence of the +H^{+} ions. It is well at this point to add a small excess of +C_{2}O_{4}^{--} ions in the form of ammonium oxalate to decrease the +solubility of the precipitate. + +The oxalate precipitate consists mainly of CaC_{2}O_{4}.H_{2}O when +thrown down.] + +[Note 5: The small quantity of ammonium oxalate solution is added +before the second precipitation of the calcium oxalate to insure +the presence of a slight excess of the reagent, which promotes the +separation of the calcium compound.] + +[Note 6: On ignition the calcium oxalate loses carbon dioxide and +carbon monoxide, leaving calcium oxide: + +CaC_{2}O_{4}.H_{2}O --> CaO + CO_{2} + CO + H_{2}O. + +For small weights of the oxalate (0.6 gram or less) this reaction may +be brought about in a platinum crucible at the highest temperature of +a Tirrill burner, but it is well to ignite larger quantities than this +over the blast lamp until the weight is constant.] + +[Note 7: The heat required to burn the filter, and that subsequently +applied as described, will convert most of the calcium oxalate to +calcium carbonate, which is changed to sulphate by the sulphuric acid. +The reactions involved are + +CaC_{2}O_{4} --> CaCO_{3} + CO, +CaCO_{3} + H_{2}SO_{4} --> CaSO_{4} + H_{2}O + CO_{2}. + +If a porcelain crucible is employed for ignition, this conversion to +sulphate is to be preferred, as a complete conversion to oxide is +difficult to accomplish.] + +[Note 8: The determination of the calcium may be completed +volumetrically by washing the calcium oxalate precipitate from +the filter into dilute sulphuric acid, warming, and titrating +the liberated oxalic acid with a standard solution of potassium +permanganate as described on page 72. When a considerable number of +analyses are to be made, this procedure will save much of the time +otherwise required for ignition and weighing.] + + + + +DETERMINATION OF MAGNESIUM + + +PROCEDURE.--Evaporate the acidified filtrates from the calcium +precipitates until the salts begin to crystallize, but do !not! +evaporate to dryness (Note 1). Dilute the solution cautiously until +the salts are brought into solution, adding a little acid if the +solution has evaporated to very small volume. The solution should be +carefully examined at this point and must be filtered if a precipitate +has appeared. Heat the clear solution to boiling; remove the burner +and add 25 cc. of a solution of disodium phosphate. Then add slowly +dilute ammonia (1 volume strong ammonia (sp. gr. 0.90) and 9 volumes +water) as long as a precipitate continues to form. Finally, add a +volume of concentrated ammonia (sp. gr. 0.90) equal to one third of +the volume of the solution, and allow the whole to stand for about +twelve hours. + +Decant the solution through a filter, wash it with dilute ammonia +water, proceeding as prescribed for the determination of phosphoric +anhydride on page 122, including; the reprecipitation (Note 2), +except that 3 cc. of disodium phosphate solution are added before the +reprecipitation of the magnesium ammonium phosphate instead of +the magnesia mixture there prescribed. From the weight of the +pyrophosphate, calculate the percentage of magnesium oxide (MgO) in +the sample of limestone. Remember that the pyrophosphate finally +obtained is from one fifth of the original sample. + +[Note 1: The precipitation of the magnesium should be made in as small +volume as possible, and the ratio of ammonia to the total volume of +solution should be carefully provided for, on account of the relative +solubility of the magnesium ammonium phosphate. This matter has +been fully discussed in connection with the phosphoric anhydride +determination.] + +[Note 2: The first magnesium ammonium phosphate precipitate is rarely +wholly crystalline, as it should be, and is not always of the proper +composition when precipitated in the presence of such large amounts of +ammonium salts. The difficulty can best be remedied by filtering the +precipitate and (without washing it) redissolving in a small quantity +of hydrochloric acid, from which it may be again thrown down by +ammonia after adding a little disodium phosphate solution. If the +flocculent character was occasioned by the presence of magnesium +hydroxide, the second precipitation, in a smaller volume containing +fewer salts, will often result more favorably. + +The removal of iron or alumina from a contaminated precipitate is +a matter involving a long procedure, and a redetermination of the +magnesium from a new sample, with additional precautions, is usually +to be preferred.] + + + + +DETERMINATION OF CARBON DIOXIDE + + +!Absorption Apparatus! + +[Illustration: Fig. 3] + +The apparatus required for the determination of the carbon dioxide +should be arranged as shown in the cut (Fig. 3). The flask (A) is +an ordinary wash bottle, which should be nearly filled with dilute +hydrochloric acid (100 cc. acid (sp. gr. 1.12) and 200 cc. of water). +The flask is connected by rubber tubing (a) with the glass tube (b) +leading nearly to the bottom of the evolution flask (B) and having its +lower end bent upward and drawn out to small bore, so that the carbon +dioxide evolved from the limestone cannot bubble back into (b). The +evolution flask should preferably be a wide-mouthed Soxhlet extraction +flask of about 150 cc. capacity because of the ease with which tubes +and stoppers may be fitted into the neck of a flask of this type. The +flask should be fitted with a two-hole rubber stopper. The condenser +(C) may consist of a tube with two or three large bulbs blown in +it, for use as an air-cooled condenser, or it may be a small +water-jacketed condenser. The latter is to be preferred if a number of +determinations are to be made in succession. + +A glass delivery tube (c) leads from the condenser to the small U-tube +(D) containing some glass beads or small pieces of glass rod and 3 cc. +of a saturated solution of silver sulphate, with 3 cc. of concentrated +sulphuric acid (sp. gr. 1.84). The short rubber tubing (d) connects +the first U-tube to a second U-tube (E) which is filled with small +dust-free lumps of dry calcium chloride, with a small, loose plug of +cotton at the top of each arm. Both tubes should be closed by cork +stoppers, the tops of which are cut off level with, or preferably +forced a little below, the top of the U-tube, and then neatly sealed +with sealing wax. + +The carbon dioxide may be absorbed in a tube containing soda lime +(F) or in a Geissler bulb (F') containing a concentrated solution +of potassium hydroxide (Note 2). The tube (F) is a glass-stoppered +side-arm U-tube in which the side toward the evolution flask and one +half of the other side are filled with small, dust-free lumps of soda +lime of good quality (Note 3). Since soda lime contains considerable +moisture, the other half of the right side of the tube is filled with +small lumps of dry, dust-free calcium chloride to retain the moisture +from the soda lime. Loose plugs of cotton are placed at the top of +each arm and between the soda lime and the calcium chloride. + +The Geissler bulb (F'), if used, should be filled with potassium +hydroxide solution (1 part of solid potassium hydroxide dissolved in +two parts of water) until each small bulb is about two thirds full +(Note 4). A small tube containing calcium chloride is connected with +the Geissler bulb proper by a ground joint and should be wired to the +bulb for safety. This is designed to retain any moisture from the +hydroxide solution. A piece of clean, fine copper wire is so attached +to the bulb that it can be hung from the hook above a balance pan, or +other support. + +The small bottle (G) with concentrated sulphuric acid (sp. gr. 1.84) +is so arranged that the tube (f) barely dips below the surface. This +will prevent the absorption of water vapor by (F) or (F') and serves +as an aid in regulating the flow of air through the apparatus. (H) is +an aspirator bottle of about four liters capacity, filled with water; +(k) is a safety tube and a means of refilling (H); (h) is a screw +clamp, and (K) a U-tube filled with soda lime. + +[Note 1: The air current, which is subsequently drawn through the +apparatus, to sweep all of the carbon dioxide into the absorption +apparatus, is likely to carry with it some hydrochloric acid from +the evolution flask. This acid is retained by the silver sulphate +solution. The addition of concentrated sulphuric acid to this solution +reduces its vapor pressure so far that very little water is carried on +by the air current, and this slight amount is absorbed by the calcium +chloride in (E). As the calcium chloride frequently contains a small +amount of a basic material which would absorb carbon dioxide, it is +necessary to pass carbon dioxide through (E) for a short time and then +drive all the gas out with a dry air current for thirty minutes before +use.] + +[Note 2: Soda-lime absorption tubes are to be preferred if a +satisfactory quality of soda lime is available and the number of +determinations to be made successively is small. The potash bulbs will +usually permit of a larger number of successive determinations without +refilling, but they require greater care in handling and in the +analytical procedure.] + +[Note 3: Soda lime is a mixture of sodium and calcium hydroxides. Both +combine with carbon dioxide to form carbonates, with the evolution +of water. Considerable heat is generated by the reaction, and the +temperature of the tube during absorption serves as a rough index of +the progress of the reaction through the mass of soda lime. + +It is essential that soda lime of good quality for analytical purposes +should be used. The tube should not contain dust, as this is likely to +be swept away.] + +[Note 4: The solution of the hydroxide for use in the Geissler bulb +must be highly concentrated to insure complete absorption of the +carbon dioxide and also to reduce the vapor pressure of the solution, +thus lessening the danger of loss of water with the air which passes +through the bulbs. The small quantity of moisture which is then +carried out of the bulbs is held by the calcium chloride in the +prolong tube. The best form of absorption bulb is that to which the +prolong tube is attached by a ground glass joint. + +After the potassium hydroxide is approximately half consumed in the +first bulb of the absorption apparatus, potassium bicarbonate is +formed, and as it is much less soluble than the carbonate, it often +precipitates. Its formation is a warning that the absorbing power of +the hydroxide is much diminished.] + + +!The Analysis! + +PROCEDURE.-- Weigh out into the flask (B) about 1 gram of limestone. +Cover it with 15 cc. of water. Weigh the absorption apparatus (F) +or (F') accurately after allowing it to stand for 30 minutes in the +balance case, and wiping it carefully with a lintless cloth, taking +care to handle it as little as possible after wiping (Note 1). Connect +the absorption apparatus with (e) and (f). If a soda-lime tube is +used, be sure that the arm containing the soda lime is next the tube +(E) and that the glass stopcocks are open. + +To be sure that the whole apparatus is airtight, disconnect the rubber +tube from the flask (A), making sure that the tubes (a) and (b) do not +contain any hydrochloric acid, close the pinchcocks (a) and (k) and +open (h). No bubbles should pass through (D) or (G) after a few +seconds. When assured that the fittings are tight, close (h) and open +(a) cautiously to admit air to restore atmospheric pressure. This +precaution is essential, as a sudden inrush of air will project liquid +from (D) or (F'). Reconnect the rubber tube with the flask (A). +Open the pinchcocks (a) and (k) and blow over about 10 cc. of the +hydrochloric acid from (A) into (B). When the action of the acid +slackens, blow over (slowly) another 10 cc. + +The rate of gas evolution should not exceed for more than a few +seconds that at which about two bubbles per second pass through (G) +(Note 2). Repeat the addition of acid in small portions until the +action upon the limestone seems to be at an end, taking care to close +(a) after each addition of acid (Note 3). Disconnect (A) and connect +the rubber tubing with the soda-lime tube (K) and open (a). Then close +(k) and open (h), regulating the flow of water from (H) in such a way +that about two bubbles per second pass through (G). Place a small +flame under (B) and !slowly! raise the contents to boiling and boil +for three minutes. Then remove the burner from under (B) and continue +to draw air through the apparatus for 20-30 minutes, or until (H) +is emptied (Note 4). Remove the absorption apparatus, closing the +stopcocks on (F) or stoppering the open ends of (F'), leave the +apparatus in the balance case for at least thirty minutes, wipe it +carefully and weigh, after opening the stopcocks (or removing plugs). +The increase in weight is due to absorption of CO_{2}, from which its +percentage in the sample may be calculated. + +After cleaning (B) and refilling (H), the apparatus is ready for the +duplicate analysis. + +[Note 1: The absorption tubes or bulbs have large surfaces on which +moisture may collect. By allowing them to remain in the balance case +for some time before weighing, the amount of moisture absorbed on the +surface is as nearly constant as practicable during two weighings, and +a uniform temperature is also assured. The stopcocks of the U-tube +should be opened, or the plugs used to close the openings of the +Geissler bulb should be removed before weighing in order that the air +contents shall always be at atmospheric pressure.] + +[Note 2: If the gas passes too rapidly into the absorption apparatus, +some carbon dioxide may be carried through, not being completely +retained by the absorbents.] + +[Note 3: The essential ionic changes involved in this procedure are +the following: It is assumed that the limestone, which is typified by +calcium carbonate, is very slightly soluble in water, and the ions +resulting are Ca^{++} and CO_{3}^{--}. In the presence of H^{+} ions +of the mineral acid, the CO_{3}^{--} ions form [H_{2}CO_{3}]. This +is not only a weak acid which, by its formation, diminishes the +concentration of the CO_{3}^{--} ions, thus causing more of the +carbonate to dissolve to re-establish equilibrium, but it is also an +unstable compound and breaks down into carbon dioxide and water.] + +[Note 4: Carbon dioxide is dissolved by cold water, but the gas is +expelled by boiling, and, together with that which is distributed +through the apparatus, is swept out into the absorption bulb by the +current of air. This air is purified by drawing it through the tube +(K) containing soda lime, which removes any carbon dioxide which may +be in it.] + + + + +DETERMINATION OF LEAD, COPPER, IRON, AND ZINC IN BRASS + +ELECTROLYTIC SEPARATIONS + + +!General Discussion! + +When a direct current of electricity passes from one electrode to +another through solutions of electrolytes, the individual ions present +in these solutions tend to move toward the electrode of opposite +electrical charge to that which each ion bears, and to be discharged +by that electrode. Whether or not such discharge actually occurs in +the case of any particular ion depends upon the potential (voltage) of +the current which is passing through the solution, since for each ion +there is, under definite conditions, a minimum potential below which +the discharge of the ion cannot be effected. By taking advantage +of differences in discharge-potentials, it is possible to effect +separations of a number of the metallic ions by electrolysis, and at +the same time to deposit the metals in forms which admit of direct +weighing. In this way the slower procedures of precipitation and +filtration may frequently be avoided. The following paragraphs present +a brief statement of the fundamental principles and conditions +underlying electro-analysis. + +The total energy of an electric current as it passes through a +solution is distributed among three factors, first, its potential, +which is measured in volts, and corresponds to what is called "head" +in a stream of water; second, current strength, which is measured +in amperes, and corresponds to the volume of water passing a +cross-section of a stream in a given time interval; and third, the +resistance of the conducting medium, which is measured in ohms. The +relation between these three factors is expressed by Ohm's law, +namely, that !I = E/R!, when I is current strength, E potential, and R +resistance. It is plain that, for a constant resistance, the +strength of the current and its potential are mutually and directly +interdependent. + +As already stated, the applied electrical potential determines whether +or not deposition of a metal upon an electrode actually occurs. The +current strength determines the rate of deposition and the physical +characteristics of the deposit. The resistance of the solution is +generally so small as to fall out of practical consideration. + +Approximate deposition-potentials have been determined for a number +of the metallic elements, and also for hydrogen and some of the +acid-forming radicals. The values given below are those required +for deposition from normal solutions at ordinary temperatures +with reference to a hydrogen electrode. They must be regarded as +approximate, since several disturbing factors and some secondary +reactions render difficult their exact application under the +conditions of analysis. They are: + + Zn Cd Fe Ni Pb H Cu Sb Hg Ag SO_{4} ++0.77 +0.42 +0.34 +0.33 +0.13 0 -0.34 -0.67 -0.76 -0.79 +1.90 + +From these data it is evident that in order to deposit copper from a +normal solution of copper sulphate a minimum potential equal to the +algebraic sum of the deposition-potentials of copper ions and sulphate +ions must be applied, that is, +1.56 volts. The deposition of zinc +from a solution of zinc sulphate would require +2.67 volts, but, since +the deposition of hydrogen from sulphuric acid solution requires only ++1.90 volts, the quantitative deposition of zinc by electrolysis from +a sulphuric acid solution of a zinc salt is not practicable. On the +other hand, silver, if present in a solution of copper sulphate, would +deposit with the copper. + +The foregoing examples suffice to illustrate the application of the +principle of deposition potentials, but it must further be noted +that the values stated apply to normal solutions of the compounds in +question, that is, to solutions of considerable concentrations. As the +concentration of the ions diminishes, and hence fewer ions approach +the electrodes, somewhat higher voltages are required to attract and +discharge them. From this it follows that the concentrations should be +kept as high as possible to effect complete deposition in the least +practicable time, or else the potentials applied must be progressively +increased as deposition proceeds. In practice, the desired result is +obtained by starting with small volumes of solution, using as large an +electrode surface as possible, and by stirring the solution to bring +the ions in contact with the electrodes. This is, in general, a more +convenient procedure than that of increasing the potential of the +current during electrolysis, although that method is also used. + +As already stated, those ions in a solution of electrolytes will first +be discharged which have the lowest deposition potentials, and so +long as these ions are present around the electrode in considerable +concentration they, almost alone, are discharged, but, as their +concentration diminishes, other ions whose deposition potentials are +higher but still within that of the current applied, will also begin +to separate. For example, from a nitric acid solution of copper +nitrate, the copper ions will first be discharged at the cathode, but +as they diminish in concentration hydrogen ions from the acid (or +water) will be also discharged. Since the hydrogen thus liberated is a +reducing agent, the nitric acid in the solution is slowly reduced to +ammonia, and it may happen that if the current is passed through for a +long time, such a solution will become alkaline. Oxygen is liberated +at the anode, but, since there is no oxidizable substance present +around that electrode, it escapes as oxygen gas. It should be noted +that, in general, the changes occurring at the cathode are reductions, +while those at the anode are oxidations. + +For analytical purposes, solutions of nitrates or sulphates of the +metals are preferable to those of the chlorides, since liberated +chlorine attacks the electrodes. In some cases, as for example, that +of silver, solution of salts forming complex ions, like that of +the double cyanide of silver and potassium, yield better metallic +deposits. + +Most metals are deposited as such upon the cathode; a few, notably +lead and manganese, separate in the form of dioxides upon the anode. +It is evidently important that the deposited material should be so +firmly adherent that it can be washed, dried, and weighed without +loss in handling. To secure these conditions it is essential that the +current density (that is, the amount of current per unit of area of +the electrodes) shall not be too high. In prescribing analytical +conditions it is customary to state the current strength in "normal +densities" expressed in amperes per 100 sq. cm. of electrode surface, +as, for example, "N.D_{100} = 2 amps." + +If deposition occurs too rapidly, the deposit is likely to be spongy +or loosely adherent and falls off on subsequent treatment. This places +a practical limit to the current density to be employed, for a given +electrode surface. The cause of the unsatisfactory character of +the deposit is apparently sometimes to be found in the coincident +liberation of considerable hydrogen and sometimes in the failure of +the rapidly deposited material to form a continuous adherent surface. +The effect of rotating electrodes upon the character of the deposit is +referred to below. + +The negative ions of an electrolyte are attracted to the anode and are +discharged on contact with it. Anions such as the chloride ion yield +chlorine atoms, from which gaseous chlorine molecules are formed +and escape. The radicals which compose such ions as NO_{3}^{-} or +SO_{4}^{--} are not capable of independent existence after discharge, +and break down into oxygen and N_{2}O_{5} and SO_{3} respectively. The +oxygen escapes and the anhydrides, reacting with water, re-form nitric +and sulphuric acids. + +The law of Faraday expresses the relation between current strength and +the quantities of the decomposition products which, under constant +conditions, appear at the electrodes, namely, that a given quantity +of electricity, acting for a given time, causes the separation of +chemically equivalent quantities of the various elements or radicals. +For example, since 107.94 grams of silver is equivalent to 1.008 grams +of hydrogen, and that in turn to 8 grams of oxygen, or 31.78 grams of +copper, the quantity of electricity which will cause the deposit of +107.94 grams of silver in a given time will also separate the weights +just indicated of the other substances. Experiments show that a +current of one ampere passing for one second, i.e., a coulomb of +electricity, causes the deposition of 0.001118 gram of silver from a +normal solution of a silver salt. The number of coulombs required to +deposit 107.94 grams is 107.94/0.001118 or 96,550 and the same number +of coulombs will also cause the separation of 1.008 grams of hydrogen, +8 grams of oxygen or 31.78 grams of copper. While it might at first +appear that Faraday's law could thus be used as a basis for the +calculation of the time required for the deposition of a given +quantity of an electrolyte from solution, it must be remembered that +the law expresses what occurs when the concentration of the ions in +the solution is kept constant, as, for example, when the anode in +a silver salt solution is a plate of metallic silver. Under the +conditions of electro-analysis the concentration of the ions is +constantly diminishing as deposition proceeds and the time actually +required for complete deposition of a given weight of material by +a current of constant strength is, therefore, greater than that +calculated on the basis of the law as stated above. + +The electrodes employed in electro-analysis are almost exclusively +of platinum, since that metal alone satisfactorily resists chemical +action of the electrolytes, and can be dried and weighed without +change in composition. The platinum electrodes may be used in the +form of dishes, foil or gauze. The last, on account of the ease of +circulation of the electrolyte, its relatively large surface in +proportion to its weight and the readiness with which it can be washed +and dried, is generally preferred. + +Many devices have been described by the use of which the electrode +upon which deposition occurs can be mechanically rotated. This has an +effect parallel to that of greatly increasing the electrode surface +and also provides a most efficient means of stirring the solution. +With such an apparatus the amperage may be increased to 5 or even 10 +amperes and depositions completed with great rapidity and accuracy. It +is desirable, whenever practicable, to provide a rotating or stirring +device, since, for example, the time consumed in the deposition of the +amount of copper usually found in analysis may be reduced from the +20 to 24 hours required with stationary electrodes, and unstirred +solutions, to about one half hour. + + + + +DETERMINATION OF COPPER AND LEAD + + +PROCEDURE.--Weigh out two portions of about 0.5 gram each (Note 1) +into tall, slender lipless beakers of about 100 cc. capacity. Dissolve +the metal in a solution of 5 cc. of dilute nitric acid (sp. gr. 1.20) +and 5 cc. of water, heating gently, and keeping the beaker covered. +When the sample has all dissolved (Note 2), wash down the sides of the +beaker and the bottom of the watch-glass with water and dilute the +solution to about 50 cc. Carefully heat to boiling and boil for a +minute or two to expel nitrous fumes. + +Meanwhile, four platinum electrodes, two anodes and two cathodes, +should be cleaned by dipping in dilute nitric acid, washing with water +and finally with 95 per cent alcohol (Note 3). The alcohol may be +ignited and burned off. The electrodes are then cooled in a desiccator +and weighed. Connect the electrodes with the binding posts (or other +device for connection with the electric circuit) in such a way that +the copper will be deposited upon the electrode with the larger +surface, which is made the cathode. The beaker containing the solution +should then be raised into place from below the electrodes until the +latter reach nearly to the bottom of the beaker. The support for the +beaker must be so arranged that it can be easily raised or lowered. + +If the electrolytic apparatus is provided with a mechanism for the +rotation of the electrode or stirring of the electrolyte, proceed as +follows: Arrange the resistance in the circuit to provide a direct +current of about one ampere. Pass this current through the solution +to be electrolyzed, and start the rotating mechanism. Keep the beaker +covered as completely as possible, using a split watch-glass (or other +device) to avoid loss by spattering. When the solution is colorless, +which is usually the case after about 35 minutes, rinse off the cover +glass, wash down the sides of the beaker, add about 0.30 gram of urea +and continue the electrolysis for another five minutes (Notes 4 and +5). + +If stationary electrodes are employed, the current strength should be +about 0.1 ampere, which may, after 12 to 15 hours, be increased to 0.2 +ampere. The time required for complete deposition is usually from 20 +to 24 hours. It is advisable to add 5 cc. of nitric acid (sp. gr. 1.2) +if the electrolysis extends over this length of time. No urea is added +in this case. + +When the deposition of the copper appears to be complete, stop the +rotating mechanism and slowly lower the beaker with the left hand, +directing at the same time a stream of water from a wash bottle on +both electrodes. Remove the beaker, shut off the current, and, if +necessary, complete the washing of the electrodes (Note 6). Rinse the +electrodes cautiously with alcohol and heat them in a hot closet until +the alcohol has just evaporated, but no longer, since the copper is +likely to oxidize at the higher temperature. (The alcohol may be +removed by ignition if care is taken to keep the electrodes in motion +in the air so that the copper deposit is not too strongly heated at +any one point.) + +Test the solution in the beaker for copper as follows, remembering +that it is to be used for subsequent determinations of iron and zinc: +Remove about 5 cc. and add a slight excess of ammonia. Compare the +mixture with some distilled water, holding both above a white surface. +The solution should not show any tinge of blue. If the presence of +copper is indicated, add the test portion to the main solution, +evaporate the whole to a volume of about 100 cc., and again +electrolyze with clean electrodes (Note 7). + +After cooling the electrodes in a desiccator, weigh them and from the +weight of copper on the cathode and of lead dioxide (PbO_{2}) on the +anode, calculate the percentage of copper (Cu) and of lead (Pb) in the +brass. + +[Note 1: It is obvious that the brass taken for analysis should be +untarnished, which can be easily assured, when wire is used, by +scouring with emery. If chips or borings are used, they should be well +mixed, and the sample for analysis taken from different parts of the +mixture.] + +[Note 2: If a white residue remains upon treatment of the alloy with +nitric acid, it indicates the presence of tin. The material is not, +therefore, a true brass. This may be treated as follows: Evaporate the +solution to dryness, moisten the residue with 5 cc. of dilute nitric +acid (sp. gr. 1.2) and add 50 cc. of hot water. Filter off the +meta-stannic acid, wash, ignite in porcelain and weigh as SnO_{2}. +This oxide is never wholly free from copper and must be purified for +an exact determination. If it does not exceed 2 per cent of the alloy, +the quantity of copper which it contains may usually be neglected.] + +[Note 3: The electrodes should be freed from all greasy matter before +using, and those portions upon which the metal will deposit should not +be touched with the fingers after cleaning.] + +[Note 4: Of the ions in solution, the H^{+}, Cu^{++}, Zn^{++}, and +Fe^{+++} ions tend to move toward the cathode. The NO_{3}^{-} ions and +the lead, probably in the form of PbO_{2}^{--} ions, move toward the +anode. At the cathode the Cu^{++} ions are discharged and plate out as +metallic copper. This alone occurs while the solution is relatively +concentrated. Later on, H^{+} ions are also discharged. In the +presence of considerable quantities of H^{+} ions, as in this acid +solution, no Zn^{++} or Fe^{+++} ions are discharged because of their +greater deposition potentials. At the anode the lead is deposited as +PbO_{2} and oxygen is evolved. + +For the reasons stated on page 141 care must be taken that the +solution does not become alkaline if the electrolysis is long +continued.] + +[Note 5: Urea reacts with nitrous acid, which may be formed in the +solution as a result of the reducing action of the liberated hydrogen. +Its removal promotes the complete precipitation of the copper. The +reaction is + +CO(NH_{2})_{2} + 2HNO_{2} --> CO_{2} + 2N_{2} + 3H_{2}O.] + +[Note 6: The electrodes must be washed nearly or quite free from +the nitric acid solution before the circuit is broken to prevent +re-solution of the copper. + +If several solutions are connected in the same circuit it is obvious +that some device must be used to close the circuit as soon as the +beaker is removed.] + +[Note 7: The electrodes upon which the copper has been deposited +may be cleaned by immersion in warm nitric acid. To remove the lead +dioxide, add a few crystals of oxalic acid to the nitric acid.] + + + + +DETERMINATION OF IRON + + +Most brasses contain small percentages of iron (usually not over 0.1 +per cent) which, unless removed, is precipitated as phosphate and +weighed with the zinc. + +PROCEDURE.--To the solution from the precipitation of copper and +lead by electrolysis, add dilute ammonia (sp. gr. 0.96) until the +precipitate of zinc hydroxide which first forms re-dissolves, leaving +only a slight red precipitate of ferric hydroxide. Filter off the +iron precipitate, using a washed filter, and wash five times with hot +water. Test a portion of the last washing with a dilute solution of +ammonium sulphide to assure complete removal of the zinc. + +The precipitate may then be ignited and weighed as ferric oxide, as +described on page 110. + +Calculate the percentage of iron (Fe) in the brass. + + + + +DETERMINATION OF ZINC + + +PROCEDURE.--Acidify the filtrate from the iron determination with +dilute nitric acid. Concentrate it to 150 cc. Add to the cold solution +dilute ammonia (sp. gr. 0.96) cautiously until it barely smells of +ammonia; then add !one drop! of a dilute solution of litmus (Note 1), +and drop in, with the aid of a dropper, dilute nitric acid until the +blue of the litmus just changes to red. It is important that this +point should not be overstepped. Heat the solution nearly to boiling +and pour into it slowly a filtered solution of di-ammonium hydrogen +phosphate[1] containing a weight of the phosphate about equal +to twelve times that of the zinc to be precipitated. (For this +calculation the approximate percentage of zinc is that found by +subtracting the sum of the percentages of the copper, lead and iron +from 100 per cent.) Keep the solution just below boiling for fifteen +minutes, stirring frequently (Note 2). If at the end of this time the +amorphous precipitate has become crystalline, allow the solution to +cool for about four hours, although a longer time does no harm (Note +3), and filter upon an asbestos filter in a porcelain Gooch crucible. +The filter is prepared as described on page 103, and should be dried +to constant weight at 105°C. + +[Footnote 1: The ammonium phosphate which is commonly obtainable +contains some mono-ammonium salt, and this is not satisfactory as a +precipitant. It is advisable, therefore, to weigh out the amount of +the salt required, dissolve it in a small volume of water, add a drop +of phenolphthalein solution, and finally add dilute ammonium hydroxide +solution cautiously until the solution just becomes pink, but do not +add an excess.] + +Wash the precipitate until free from sulphates with a warm 1 per cent +solution of the di-ammonium phosphate, and then five times with 50 per +cent alcohol (Note 4). Dry the crucible and precipitate for an hour at +105°C., and finally to constant weight (Note 5). The filtrate should +be made alkaline with ammonia and tested for zinc with a few drops of +ammonium sulphide, allowing it to stand (Notes 6, 7 and 8). + +From the weight of the zinc ammonium phosphate (ZnNH_{4}PO_{4}) +calculate the percentage of the zinc (Zn) in the brass. + +[Note 1: The zinc ammonium phosphate is soluble both in acids and in +ammonia. It is, therefore, necessary to precipitate the zinc in a +nearly neutral solution, which is more accurately obtained by adding +a drop of a litmus solution to the liquid than by the use of litmus +paper.] + +[Note 2: The precipitate which first forms is amorphous, and may have +a variable composition. On standing it becomes crystalline and then +has the composition ZnNH_{4}PO_{4}. The precipitate then settles +rapidly and is apt to occasion "bumping" if the solution is heated to +boiling. Stirring promotes the crystallization.] + +[Note 3: In a carefully neutralized solution containing a considerable +excess of the precipitant, and also ammonium salts, the separation +of the zinc is complete after standing four hours. The ionic changes +connected with the precipitation of the zinc as zinc ammonium +phosphate are similar to those described for magnesium ammonium +phosphate, except that the zinc precipitate is soluble in an excess of +ammonium hydroxide, probably as a result of the formation of complex +ions of the general character Zn(NH_{3})_{4}^{++}.] + +[Note 4: The precipitate is washed first with a dilute solution of the +phosphate to prevent a slight decomposition of the precipitate (as a +result of hydrolysis) if hot water alone is used. The alcohol is added +to the final wash-water to promote the subsequent drying.] + +[Note 5: The precipitate may be ignited and weighed as +Zn_{2}P_{2}O_{7}, by cautiously heating the porcelain Gooch crucible +within a nickel or iron crucible, used as a radiator. The heating +must be very slow at first, as the escaping ammonia may reduce the +precipitate if it is heated too quickly.] + +[Note 6: If the ammonium sulphide produced a distinct precipitate, +this should be collected on a small filter, dissolved in a few cubic +centimeters of dilute nitric acid, and the zinc reprecipitated as +phosphate, filtered off, dried, and weighed, and the weight added to +that of the main precipitate.] + +[Note 7: It has been found that some samples of asbestos are acted +upon by the phosphate solution and lose weight. An error from this +source may be avoided by determining the weight of the crucible +and filter after weighing the precipitate. For this purpose the +precipitate may be dissolved in dilute nitric acid, the asbestos +washed thoroughly, and the crucible reweighed.] + +[Note 8. The details of this method of precipitation of zinc are fully +discussed in an article by Dakin, !Ztschr. Anal. Chem.!, 39 (1900), +273.] + + + + +DETERMINATION OF SILICA IN SILICATES + + +Of the natural silicates, or artificial silicates such as slags and +some of the cements, a comparatively few can be completely decomposed +by treatment with acids, but by far the larger number require fusion +with an alkaline flux to effect decomposition and solution +for analysis. The procedure given below applies to silicates +undecomposable by acids, of which the mineral feldspar is taken as a +typical example. Modifications of the procedure, which are applicable +to silicates which are completely or partially decomposable by acids, +are given in the Notes on page 155. + + +PREPARATION OF THE SAMPLE + +Grind about 3 grams of the mineral in an agate mortar (Note 1) until +no grittiness is to be detected, or, better, until it will entirely +pass through a sieve made of fine silk bolting cloth. The sieve may be +made by placing a piece of the bolting cloth over the top of a small +beaker in which the ground mineral is placed, holding the cloth in +place by means of a rubber band below the lip of the beaker. By +inverting the beaker over clean paper and gently tapping it, the fine +particles pass through the sieve, leaving the coarser particles within +the beaker. These must be returned to the mortar and ground, and the +process of sifting and grinding repeated until the entire sample +passes through the sieve. + +[Note 1: If the sample of feldspar for analysis is in the massive or +crystalline form, it should be crushed in an iron mortar until the +pieces are about half the size of a pea, and then transferred to a +steel mortar, in which they are reduced to a coarse powder. A wooden +mallet should always be used to strike the pestle of the steel mortar, +and the blows should not be sharp. + +It is plain that final grinding in an agate mortar must be continued +until the whole of the portion of the mineral originally taken has +been ground so that it will pass the bolting cloth, otherwise the +sifted portion does not represent an average sample, the softer +ingredients, if foreign matter is present, being first reduced to +powder. For this reason it is best to start with not more than the +quantity of the feldspar needed for analysis. The mineral must be +thoroughly mixed after the grinding.] + + +FUSION AND SOLUTION + +PROCEDURE.--Weigh into platinum crucibles two portions of the ground +feldspar of about 0.8 gram each. Weigh on rough balances two portions +of anhydrous sodium carbonate, each amounting to about six times the +weight of the feldspar taken for analysis (Note 1). Pour about three +fourths of the sodium carbonate into the crucible, place the latter on +a piece of clean, glazed paper, and thoroughly mix the substance and +the flux by carefully stirring for several minutes with a dry glass +rod, the end of which has been recently heated and rounded in a flame +and slowly cooled. The rod may be wiped off with a small fragment of +filter paper, which may be placed in the crucible. Place the remaining +fourth of the carbonate on the top of the mixture. Cover the crucible, +heat it to dull redness for five minutes, and then gradually increase +the heat to the full capacity of a Bunsen or Tirrill burner for +twenty minutes, or until a quiet, liquid fusion is obtained (Note 2). +Finally, heat the sides and cover strongly until any material which +may have collected upon them is also brought to fusion. + +Allow the crucible to cool, and remove the fused mass as directed on +page 116. Disintegrate the mass by placing it in a previously prepared +mixture of 100 cc. of water and 50 cc. of dilute hydrochloric acid +(sp. gr. 1.12) in a covered casserole (Note 3). Clean the crucible and +lid by means of a little hydrochloric acid, adding this acid to the +main solution (Notes 4 and 5). + +[Note 1: Quartz, and minerals containing very high percentages of +silica, may require eight or ten parts by weight of the flux to insure +a satisfactory decomposition.] + +[Note 2: During the fusion the feldspar, which, when pure, is a +silicate of aluminium and either sodium or potassium, but usually +contains some iron, calcium, and magnesium, is decomposed by the +alkaline flux. The sodium of the latter combines with the silicic acid +of the silicate, with the evolution of carbon dioxide, while about two +thirds of the aluminium forms sodium aluminate and the remainder +is converted into basic carbonate, or the oxide. The calcium and +magnesium, if present, are changed to carbonates or oxides. + +The heat is applied gently to prevent a too violent reaction when +fusion first takes place.] + +[Note 3: The solution of a silicate by a strong acid is the result of +the combination of the H^{+} ions of the acid and the silicate ions +of the silicate to form a slightly ionized silicic acid. As a +consequence, the concentration of the silicate ions in the solution is +reduced nearly to zero, and more silicate dissolves to re-establish +the disturbed equilibrium. This process repeats itself until all of +the silicate is brought into solution. + +Whether the resulting solution of the silicate contains ortho-silicic +acid (H_{4}SiO_{4}) or whether it is a colloidal solution of some +other less hydrated acid, such as meta-silicic acid (H_{2}SiO_{3}), +is a matter that is still debatable. It is certain, however, that the +gelatinous material which readily separates from such solutions is of +the nature of a hydrogel, that is, a colloid which is insoluble in +water. This substance when heated to 100°C., or higher, is completely +dehydrated, leaving only the anhydride, SiO_{2}. The changes may be +represented by the equation: + +SiO_{3}^{--} + 2H^{+} --> [H_{2}SiO_{3}] --> H_{2}O + SiO_{2}.] + +[Note 4: A portion of the fused mass is usually projected upward by +the escaping carbon dioxide during the fusion. The crucible must +therefore be kept covered as much as possible and the lid carefully +cleaned.] + +[Note 5: A gritty residue remaining after the disintegration of +the fused mass by acid indicates that the substance has been but +imperfectly decomposed. Such a residue should be filtered, washed, +dried, ignited, and again fused with the alkaline flux; or, if the +quantity of material at hand will permit, it is better to reject the +analysis, and to use increased care in grinding the mineral and in +mixing it with the flux.] + + +DEHYDRATION AND FILTRATION + +PROCEDURE.--Evaporate the solution of the fusion to dryness, stirring +frequently until the residue is a dry powder. Moisten the residue with +5 cc. of strong hydrochloric acid (sp. gr. 1.20) and evaporate again +to dryness. Heat the residue for at least one hour at a temperature +of 110°C. (Note 1). Again moisten the residue with concentrated +hydrochloric acid, warm gently, making sure that the acid comes into +contact with the whole of the residue, dilute to about 200 cc. and +bring to boiling. Filter off the silica without much delay (Note 2), +and wash five times with warm dilute hydrochloric acid (one part +dilute acid (1.12 sp. gr.) to three parts of water). Allow the filter +to drain for a few moments, then place a clean beaker below the funnel +and wash with water until free from chlorides, discarding these +washings. Evaporate the original filtrate to dryness, dehydrate at +110°C. for one hour (Note 3), and proceed as before, using a second +filter to collect the silica after the second dehydration. Wash this +filter with warm, dilute hydrochloric acid (Note 4), and finally with +hot water until free from chlorides. + +[Note 1: The silicic acid must be freed from its combination with +a base (sodium, in this instance) before it can be dehydrated. +The excess of hydrochloric acid accomplishes this liberation. By +disintegrating the fused mass with a considerable volume of dilute +acid the silicic acid is at first held in solution to a large extent. +Immediate treatment of the fused mass with strong acid is likely +to cause a semi-gelatinous silicic acid to separate at once and to +inclose alkali salts or alumina. + +A flocculent residue will often remain after the decomposition of the +fused mass is effected. This is usually partially dehydrated silicic +acid and does not require further treatment at this point. The +progress of the dehydration is indicated by the behavior of the +solution, which as evaporation proceeds usually gelatinizes. On this +account it is necessary to allow the solution to evaporate on a steam +bath, or to stir it vigorously, to avoid loss by spattering.] + +[Note 2: To obtain an approximately pure silica, the residue after +evaporation must be thoroughly extracted by warming with hydrochloric +acid, and the solution freely diluted to prevent, as far as possible, +the inclosure of the residual salts in the particles of silica. The +filtration should take place without delay, as the dehydrated silica +slowly dissolves in hydrochloric acid on standing.] + +[Note 3: It has been shown by Hillebrand that silicic acid cannot be +completely dehydrated by a single evaporation and heating, nor by +several such treatments, unless an intermediate filtration of the +silica occurs. If, however, the silica is removed and the filtrates +are again evaporated and the residue heated, the amount of silica +remaining in solution is usually negligible, although several +evaporations and filtrations are required with some silicates to +insure absolute accuracy. + +It is probable that temperatures above 100°C. are not absolutely +necessary to dehydrate the silica; but it is recommended, as tending +to leave the silica in a better condition for filtration than when +the lower temperature of the water bath is used. This, and many other +points in the analysis of silicates, are fully discussed by Dr. +Hillebrand in the admirable monograph on "The Analysis of Silicate and +Carbonate Rocks," Bulletin No. 700 of the United States Geological +Survey. + +The double evaporation and filtration spoken of above are essential +because of the relatively large amount of alkali salts (sodium +chloride) present after evaporation. For the highest accuracy in the +determination of silica, or of iron and alumina, it is also necessary +to examine for silica the precipitate produced in the filtrate by +ammonium hydroxide by fusing it with acid potassium sulphate and +solution of the fused mass in water. The insoluble silica is filtered, +washed, and weighed, and the weight added to the weight of silica +previously obtained.] + +[Note 4: Aluminium and iron are likely to be thrown down as basic +salts from hot, very dilute solutions of their chlorides, as a result +of hydrolysis. If the silica were washed only with hot water, the +solution of these chlorides remaining in the filter after the passage +of the original filtrate would gradually become so dilute as to throw +down basic salts within the pores of the filter, which would remain +with the silica. To avoid this, an acid wash-water is used until the +aluminium and iron are practically removed. The acid is then removed +by water.] + + +IGNITION AND TESTING OF SILICA + +PROCEDURE.--Transfer the two washed filters belonging to each +determination to a platinum crucible, which need not be previously +weighed, and burn off the filter (Note 1). Ignite for thirty minutes +over the blast lamp with the cover on the crucible, and then for +periods of ten minutes, until the weight is constant. + +When a constant weight has been obtained, pour into the crucible about +3 cc. of water, and then 3 cc. of hydrofluoric acid. !This must be +done in a hood with a good draft and great care must be taken not to +come into contact with the acid or to inhale its fumes (Note 2!). + +If the precipitate has dissolved in this quantity of acid, add two +drops of concentrated sulphuric acid, and heat very slowly (always +under the hood) until all the liquid has evaporated, finally igniting +to redness. Cool in a desiccator, and weigh the crucible and residue. +Deduct this weight from the previous weight of crucible and impure +silica, and from the difference calculate the percentage of silica in +the sample (Note 3). + +[Note 1: The silica undergoes no change during the ignition beyond the +removal of all traces of water; but Hillebrand (!loc. cit.!) has shown +that the silica holds moisture so tenaciously that prolonged ignition +over the blast lamp is necessary to remove it entirely. This finely +divided, ignited silica tends to absorb moisture, and should be +weighed quickly.] + +[Note 2: Notwithstanding all precautions, the ignited precipitate of +silica is rarely wholly pure. It is tested by volatilisation of the +silica as silicon fluoride after solution in hydrofluoric acid, and, +if the analysis has been properly conducted, the residue, after +treatment with the acids and ignition, should not exceed 1 mg. + +The acid produces ulceration if brought into contact with the skin, +and its fumes are excessively harmful if inhaled.] + +[Note 3: The impurities are probably weighed with the original +precipitate in the form of oxides. The addition of the sulphuric +acid displaces the hydrofluoric acid, and it may be assumed that the +resulting sulphates (usually of iron or aluminium) are converted to +oxides by the final ignition. + +It is obvious that unless the sulphuric and hydrofluoric acids used +are known to leave no residue on evaporation, a quantity equal to that +employed in the analysis must be evaporated and a correction applied +for any residue found.] + +[Note 4: If the silicate to be analyzed is shown by a previous +qualitative examination to be completely decomposable, it may be +directly treated with hydrochloric acid, the solution evaporated to +dryness, and the silica dehydrated and further treated as described in +the case of the feldspar after fusion. + +A silicate which gelatinizes on treatment with acids should be mixed +first with a little water, and the strong acid added in small portions +with stirring, otherwise the gelatinous silicic acid incloses +particles of the original silicate and prevents decomposition. The +water, by separating the particles and slightly lessening the rapidity +of action, prevents this difficulty. This procedure is one which +applies in general to the solution of fine mineral powders in acids. + +If a small residue remains undecomposed by the treatment of the +silicate with acid, this may be filtered, washed, ignited and fused +with sodium carbonate and a solution of the fused mass added to the +original acid solution. This double procedure has an advantage, in +that it avoids adding so large a quantity of sodium salts as is +required for disintegration of the whole of the silicate by the fusion +method.] + + + + +PART IV + +STOICHIOMETRY + + +The problems with which the analytical chemist has to deal are not, as +a matter of actual fact, difficult either to solve or to understand. +That they appear difficult to many students is due to the fact that, +instead of understanding the principles which underlie each of the +small number of types into which these problems may be grouped, each +problem is approached as an individual puzzle, unrelated to others +already solved or explained. This attitude of mind should be carefully +avoided. + +It is obvious that ability to make the calculations necessary for +the interpretation of analytical data is no less important than the +manipulative skill required to obtain them, and that a moderate time +spent in the careful study of the solutions of the typical problems +which follow may save much later embarrassment. + +1. It is often necessary to calculate what is known as a "chemical +factor," or its equivalent logarithmic value called a "log factor," +for the conversion of the weight of a given chemical substance into an +equivalent weight of another substance. This is, in reality, a very +simple problem in proportion, making use of the atomic or molecular +weights of the substances in question which are chemically equivalent +to each other. One of the simplest cases of this sort is the +following: What is the factor for the conversion of a given weight of +barium sulphate (BaSO_{4}) into an equivalent weight of sulphur (S)? +The molecular weight of BaSO_{4} is 233.5. There is one atom of S in +the molecule and the atomic weight of S is 32.1. The chemical factor +is, therefore, 32.1/233.5, or 0.1375 and the weight of S corresponding +to a given weight of BaSO_{4} is found by multiplying the weight of +BaSO_{4} by this factor. If the problem takes the form, "What is +the factor for the conversion of a given weight of ferric oxide +(Fe_{2}O_{3}) into ferrous oxide (FeO), or of a given weight of +mangano-manganic oxide (Mn_{3}O_{4}) into manganese (Mn)?" the +principle involved is the same, but it must then be noted that, in the +first instance, each molecule of Fe_{2}O_{3} will be equivalent to two +molecules of FeO, and in the second instance that each molecule of +Mn_{3}O_{4} is equivalent to three atoms of Mn. The respective factors +then become + +(2FeO/Fe_{2}O_{3}) or (143.6/159.6) and (3Mn/Mn_{3}O_{4}) or +(164.7/228.7). + +It is obvious that the arithmetical processes involved in this type +of problem are extremely simple. It is only necessary to observe +carefully the chemical equivalents. It is plainly incorrect to express +the ratio of ferrous to ferric oxide as (FeO/Fe_{2}O_{3}), since each +molecule of the ferric oxide will yield two molecules of the ferrous +oxide. Mistakes of this sort are easily made and constitute one of the +most frequent sources of error. + +2. A type of problem which is slightly more complicated in appearance, +but exactly comparable in principle, is the following: "What is the +factor for the conversion of a given weight of ferrous sulphate +(FeSO_{4}), used as a reducing agent against potassium permanganate, +into the equivalent weight of sodium oxalate (Na_{2}C_{2}O_{4})?" To +determine the chemical equivalents in such an instance it is necessary +to inspect the chemical reactions involved. These are: + +10FeSO_{4} + 2KMnO_{4} + 8H_{2}SO_{4} --> 5Fe_{2}(SO_{4})_{3} + +K_{2}SO_{4} + 2MnSO_{4} + 8H_{2}O, + +5Na_{2}C_{2}O_{4} + 2KMnO_{4} + 8H_{2}SO_{4} --> 5Na_{2}SO_{4} + +10CO_{2} + K_{2}SO_{4} + 2MnSO_{4} + 8H_{2}O. + +It is evident that 10FeSO_{4} in the one case, and 5Na_{2}C_{2}O_{4} +in the other, each react with 2KMnO_{4}. These molecular +quantities are therefore equivalent, and the factor becomes +(10FeSO_{4}/5Na_{2}C_{2}O_{4}) or (2FeSO_{4}/Na_{2}C_{2}O_{4}) or +(303.8/134). + +Again, let it be assumed that it is desired to determine the +factor required for the conversion of a given weight of potassium +permanganate (KMnO_{4}) into an equivalent weight of potassium +bichromate (K_{2}Cr_{2}O_{7}), each acting as an oxidizing agent +against ferrous sulphate. The reactions involved are: + +10FeSO_{4} + 2KMnO_{4} + 8H_{2}SO_{4} --> 5Fe_{2}(SO_{4})_{3} + +K_{2}SO_{4} + 2MnSO_{4} + 8H_{2}O, + +6FeSO_{4} + K_{2}Cr_{2}O_{7} + 7H_{2}SO_{4} --> 3Fe_{2}(SO_{3})_{3} + +K_{2}SO_{4} + Cr_{2}(SO_{4})_{3} + 7H_{2}O. + +An inspection of these equations shows that 2KMO_{4} react with +10FeSO_{4}, while K_{2}Cr_{2}O_{7} reacts with 6FeSO_{4}. These are +not equivalent, but if the first equation is multiplied by 3 and the +second by 5 the number of molecules of FeSO_{4} is then the same in +both, and the number of molecules of KMnO_{4} and K_{2}Cr_{2}O_{7} +reacting with these 30 molecules become 6 and 5 respectively. These +are obviously chemically equivalent and the desired factor is +expressed by the fraction (6KMnO_{4}/5K_{2}Cr_{2}O_{7}) or +(948.0/1471.0). + +3. It is sometimes necessary to calculate the value of solutions +according to the principles just explained, when several successive +reactions are involved. Such problems may be solved by a series of +proportions, but it is usually possible to eliminate the common +factors and solve but a single one. For example, the amount of MnO_{2} +in a sample of the mineral pyrolusite may be determined by dissolving +the mineral in hydrochloric acid, absorbing the evolved chlorine in a +solution of potassium iodide, and measuring the liberated iodine +by titration with a standard solution of sodium thiosulphate. The +reactions involved are: + +MnO_{2} + 4HCl --> MnCl_{2} + 2H_{2}O + Cl_{2} +Cl_{2} + 2KI --> I_{2} + 2KCl +I_{2} + 2Na_{2}S_{2}O_{3} --> 2NaI + Na_{2}S_{4}O_{6} + +Assuming that the weight of thiosulphate corresponding to the +volume of sodium thiosulphate solution used is known, what is the +corresponding weight of manganese dioxide? From the reactions given +above, the following proportions may be stated: + +2Na_{2}S_{2}O_{3}:I_{2} = 316.4:253.9, + +I_{2}:Cl_{2} = 253.9:71, + +Cl_{2}:MnO_{2} = 71:86.9. + +After canceling the common factors, there remains +2Na_{2}S_{2}O_{3}:MnO_{2} = 316.4:86.9, and the factor for the +conversion of thiosulphate into an equivalent of manganese dioxide is +86.9/316.4. + +4. To calculate the volume of a reagent required for a specific +operation, it is necessary to know the exact reaction which is to be +brought about, and, as with the calculation of factors, to keep in +mind the molecular relations between the reagent and the substance +reacted upon. For example, to estimate the weight of barium chloride +necessary to precipitate the sulphur from 0.1 gram of pure pyrite +(FeS_{2}), the proportion should read + + 488. 120.0 + 2(BaCl_{2}.2H_{2}O):FeS_{2} = x:0.1, + +where !x! represents the weight of the chloride required. Each of the +two atoms of sulphur will form upon oxidation a molecule of sulphuric +acid or a sulphate, which, in turn, will require a molecule of the +barium chloride for precipitation. To determine the quantity of the +barium chloride required, it is necessary to include in its molecular +weight the water of crystallization, since this is inseparable from +the chloride when it is weighed. This applies equally to other similar +instances. + +If the strength of an acid is expressed in percentage by weight, due +regard must be paid to its specific gravity. For example, hydrochloric +acid (sp. gr. 1.12) contains 23.8 per cent HCl !by weight!; that is, +0.2666 gram HCl in each cubic centimeter. + +5. It is sometimes desirable to avoid the manipulation required for +the separation of the constituents of a mixture of substances by +making what is called an "indirect analysis." For example, in the +analysis of silicate rocks, the sodium and potassium present may be +obtained in the form of their chlorides and weighed together. If the +weight of such a mixture is known, and also the percentage of chlorine +present, it is possible to calculate the amount of each chloride in +the mixture. Let it be assumed that the weight of the mixed chlorides +is 0.15 gram, and that it contains 53 per cent of chlorine. + +The simplest solution of such a problem is reached through algebraic +methods. The weight of chlorine is evidently 0.15 x 0.53, or 0.0795 +gram. Let x represent the weight of sodium chloride present and y +that of potassium chloride. The molecular weight of NaCl is 58.5 and +that of KCl is 74.6. The atomic weight of chlorine is 35.5. Then + +x + y = 0.15 +(35.5/58.5)x + (35.5/74.6)y = 0.00795 + +Solving these equations for x shows the weight of NaCl to be 0.0625 +gram. The weight of KCl is found by subtracting this from 0.15. + +The above is one of the most common types of indirect analyses. Others +are more complex but they can be reduced to algebraic expressions and +solved by their aid. It should, however, be noted that the results +obtained by these indirect methods cannot be depended upon for high +accuracy, since slight errors in the determination of the common +constituent, as chlorine in the above mixture, will cause considerable +variations in the values found for the components. They should not be +employed when direct methods are applicable, if accuracy is essential. + + + + +PROBLEMS + + +(The reactions necessary for the solution of these problems are either +stated with the problem or may be found in the earlier text. In the +calculations from which the answers are derived, the atomic weights +given on page 195 have been employed, using, however, only the first +decimal but increasing this by 1 when the second decimal is 5 or +above. Thus, 39.1 has been taken as the atomic weight of potassium, +32.1 for sulphur, etc. This has been done merely to secure uniformity +of treatment, and the student should remember that it is always well +to take into account the degree of accuracy desired in a particular +instance in determining the number of decimal places to retain. +Four-place logarithms were employed in the calculations. Where four +figures are given in the answer, the last figure may vary by one or +(rarely) by two units, according to the method by which the problem is +solved.) + + +VOLUMETRIC ANALYSIS + +1. How many grams of pure potassium hydroxide are required for exactly +1 liter of normal alkali solution? + +!Answer!: 56.1 grams. + +2. Calculate the equivalent in grams (a) of sulphuric acid as an acid; +(b) of hydrochloric acid as an acid; (c) of oxalic acid as an acid; +(d) of nitric acid as an acid. + +!Answers!: (a) 49.05; (b) 36.5; (c) 63; (d) 63. + +3. Calculate the equivalent in grams of (a) potassium hydroxide; +(b) of sodium carbonate; (c) of barium hydroxide; (d) of sodium +bicarbonate when titrated with an acid. + +!Answers!: (a) 56.1; (b) 53.8; (c) 85.7; (d) 84. + +4. What is the equivalent in grams of Na_{2}HPO_{4} (a) as a +phosphate; (b) as a sodium salt? + +!Answers!: (a) 47.33; (b) 71.0. + +5. A sample of aqueous hydrochloric acid has a specific gravity +of 1.12 and contains 23.81 per cent hydrochloric acid by weight. +Calculate the grams and the milliequivalents of hydrochloric acid +(HCl) in each cubic centimeter of the aqueous acid. + +!Answers!: 0.2667 gram; 7.307 milliequivalents. + +6. How many cubic centimeters of hydrochloric acid (sp. gr. 1.20 +containing 39.80 per cent HCl by weight) are required to furnish 36.45 +grams of the gaseous compound? + +!Answer!: 76.33 cc. + +7. A given solution contains 0.1063 equivalents of hydrochloric acid +in 976 cc. What is its normal value? + +!Answer!: 0.1089 N. + +8. In standardizing a hydrochloric acid solution it is found that +47.26 cc. of hydrochloric acid are exactly equivalent to 1.216 grams +of pure sodium carbonate, using methyl orange as an indicator. What is +the normal value of the hydrochloric acid? + +!Answer!: 0.4855 N. + +9. Convert 42.75 cc. of 0.5162 normal hydrochloric acid to the +equivalent volume of normal hydrochloric acid. + +!Answer!: 22.07 cc. + +10. A solution containing 25.27 cc. of 0.1065 normal hydrochloric acid +is added to one containing 92.21 cc. of 0.5431 normal sulphuric acid +and 50 cc. of exactly normal potassium hydroxide added from a pipette. +Is the solution acid or alkaline? How many cubic centimeters of +0.1 normal acid or alkali must be added to exactly neutralize the +solution? + +!Answer!: 27.6 cc. alkali (solution is acid). + +11. By experiment the normal value of a sulphuric acid solution is +found to be 0.5172. Of this acid 39.65 cc. are exactly equivalent to +21.74 cc. of a standard alkali solution. What is the normal value of +the alkali? + +!Answer!: 0.9432 N. + +12. A solution of sulphuric acid is standardized against a sample of +calcium carbonate which has been previously accurately analyzed and +found to contain 92.44% CaCO_{3} and no other basic material. The +sample weighing 0.7423 gram was titrated by adding an excess of acid +(42.42 cc.) and titrating the excess with sodium hydroxide solution +(11.22 cc.). 1 cc. of acid is equivalent to 0.9976 cc. of sodium +hydroxide. Calculate the normal value of each. + +!Answers!: Acid 0.4398 N; alkali 0.4409 N. + +13. Given five 10 cc. portions of 0.1 normal hydrochloric acid, (a) +how many grams of silver chloride will be precipitated by a portion +when an excess of silver nitrate is added? (b) how many grams of pure +anhydrous sodium carbonate (Na_{2}CO_{3}) will be neutralized by a +portion of it? (c) how many grams of silver will there be in the +silver chloride formed when an excess of silver nitrate is added to a +portion? (d) how many grams of iron will be dissolved to FeCl_{2} by a +portion of it? (e) how many grams of magnesium chloride will be formed +and how many grams of carbon dioxide liberated when an excess of +magnesium carbonate is treated with a portion of the acid? + +!Answers!: (a) 0.1434; (b) 0.053; (c) 0.1079; (d) 0.0279; (e) 0.04765, +and 0.022. + +14. If 30.00 grams of potassium tetroxalate +(KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O) are dissolved and the solution +diluted to exactly 1 liter, and 40 cc. are neutralized with 20 cc. +of a potassium carbonate solution, what is the normal value of the +carbonate solution? + +!Answer!: 0.7084 N. + +15. How many cubic centimeters of 0.3 normal sulphuric acid will be +required to neutralize (a) 30 cc. of 0.5 normal potassium hydroxide; +(b) to neutralize 30 cc. of 0.5 normal barium hydroxide; (c) to +neutralize 20 cc. of a solution containing 10.02 grams of potassium +bicarbonate per 100 cc.; (d) to give a precipitate of barium sulphate +weighing 0.4320 gram? + +!Answers!: (a) 50 cc.; (b) 50 cc.; (c) 66.73 cc.; (d) 12.33 cc. + +16. It is desired to dilute a solution of sulphuric acid of which 1 +cc. is equivalent to 0.1027 gram of pure sodium carbonate to make it +exactly 1.250 normal. 700 cc. of the solution are available. To what +volume must it be diluted? + +!Answer!: 1084 cc. + +17. Given the following data: 1 cc. of NaOH = 1.117 cc. HCl. The HCl +is 0.4876 N. How much water must be added to 100 cc. of the alkali to +make it exactly 0.5 N.? + +!Answer!: 9.0 cc. + +18. What is the normal value of a sulphuric acid solution which has a +specific gravity of 1.839 and contains 95% H_{2}SO_{4} by weight? + +!Answer!: 35.61 N. + +19. A sample of Rochelle Salt (KNaC_{4}H_{4}O_{6}.4H_{2}O), after +ignition in platinum to convert it to the double carbonate, is +titrated with sulphuric acid, using methyl orange as an indicator. +From the following data calculate the percentage purity of the sample: + +Wt. sample = 0.9500 gram +H_{2}SO_{4} used = 43.65 cc. +NaOH used = 1.72 cc. +1 cc. H_{2}SO_{4} = 1.064 cc. NaOH +Normal value NaOH = 0.1321 N. + +!Answer!: 87.72 cc. + +20. One gram of a mixture of 50% sodium carbonate and 50% potassium +carbonate is dissolved in water, and 17.36 cc. of 1.075 N acid is +added. Is the resulting solution acid or alkaline? How many cubic +centimeters of 1.075 N acid or alkali will have to be added to make +the solution exactly neutral? + +!Answers!: Acid; 1.86 cc. alkali. + +21. In preparing an alkaline solution for use in volumetric work, an +analyst, because of shortage of chemicals, mixed exactly 46.32 grams +of pure KOH and 27.64 grams of pure NaOH, and after dissolving in +water, diluted the solution to exactly one liter. How many cubic +centimeters of 1.022 N hydrochloric acid are necessary to neutralize +50 cc. of the basic solution? + +!Answer!: 74.18 cc. + +22. One gram of crude ammonium salt is treated with strong potassium +hydroxide solution. The ammonia liberated is distilled and collected +in 50 cc. of 0.5 N acid and the excess titrated with 1.55 cc. of 0.5 N +sodium hydroxide. Calculate the percentage of NH_{3} in the sample. + +!Answer!: 41.17%. + + +23. In titrating solutions of alkali carbonates in the presence of +phenolphthalein, the color change takes place when the carbonate has +been converted to bicarbonate. In the presence of methyl orange, the +color change takes place only when the carbonate has been completely +neutralized. From the following data, calculate the percentages of +Na_{2}CO_{3} and NaOH in an impure mixture. Weight of sample, 1.500 +grams; HCl (0.5 N) required for phenolphthalein end-point, 28.85 cc.; +HCl (0.5 N) required to complete the titration after adding methyl +orange, 23.85 cc. + +!Answers!: 6.67% NaOH; 84.28% Na_{2}CO_{3}. + +24. A sample of sodium carbonate containing sodium hydroxide weighs +1.179 grams. It is titrated with 0.30 N hydrochloric acid, using +phenolphthalein in cold solution as an indicator and becomes colorless +after the addition of 48.16 cc. Methyl orange is added and 24.08 cc. +are needed for complete neutralization. What is the percentage of NaOH +and Na_{2}CO_{3}? + +!Answers!: 24.50% NaOH; 64.92% Na_{2}CO_{3}. + +25. From the following data, calculate the percentages of Na_{2}CO_{3} +and NaHCO_{3} in an impure mixture. Weight of sample 1.000 gram; +volume of 0.25 N hydrochloric acid required for phenolphthalein +end-point, 26.40 cc.; after adding an excess of acid and boiling out +the carbon dioxide, the total volume of 0.25 N hydrochloric acid +required for phenolphthalein end-point, 67.10 cc. + +!Answer!: 69.95% Na_{2}CO_{3}; 30.02% NaHCO_{3}. + +26. In the analysis of a one-gram sample of soda ash, what must be the +normality of the acid in order that the number of cubic centimeters of +acid used shall represent the percentage of carbon dioxide present? + +!Answer!: 0.4544 gram. + +27. What weight of pearl ash must be taken for analysis in order that +the number of cubic centimeters of 0.5 N acid used may be equal to one +third the percentage of K_{2}CO_{3}? + +!Answer!: 1.152 grams. + +28. What weight of cream of tartar must have been taken for analysis +in order to have obtained 97.60% KHC_{4}H_{4}O_{6} in an analysis +involving the following data: NaOH used = 30.06 cc.; H_{2}SO_{4} +solution used = 0.50 cc.; 1 cc. H_{2}SO_{4} sol. = 0.0255 gram +CaCO_{3}; 1 cc. H_{2}SO_{4} sol. = 1.02 cc. NaOH sol.? + +!Answer!: 2.846 grams. + +29. Calculate the percentage of potassium oxide in an impure sample of +potassium carbonate from the following data: Weight of sample = 1.00 +gram; HCl sol. used = 55.90 cc.; NaOH sol. used = 0.42 cc.; 1 cc. NaOH +sol. = 0.008473 gram of KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O; 2 cc. +HCl sol. = 5 cc. NaOH sol. + +!Answer!: 65.68%. + +30. Calculate the percentage purity of a sample of calcite +(CaCO_{3}) from the following data: (Standardization); Weight of +H_{2}C_{2}O_{4}.2H_{2}O = 0.2460 gram; NaOH solution used = 41.03 +cc.; HCl solution used = 0.63; 1 cc. NaOH solution = 1.190 cc. HCl +solution. (Analysis); Weight of sample 0.1200 gram; HCl used = 36.38 +cc.; NaOH used = 6.20 cc. + +!Answer!: 97.97%. + +31. It is desired to dilute a solution of hydrochloric acid to exactly +0.05 N. The following data are given: 44.97 cc. of the hydrochloric +acid are equivalent to 43.76 cc. of the NaOH solution. The NaOH +is standardized against a pure potassium tetroxalate +(KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O) weighing 0.2162 gram and +requires 49.14 cc. How many cc. of water must be added to 1000 cc. of +the aqueous hydrochloric acid? + +!Answer!: 11 cc. + +32. How many cubic centimeters of 3 N phosphoric acid must be added to +300 cc. of 0.4 N phosphoric acid in order that the resulting solution +may be 0.6 N? + +!Answer!: 25 cc. + +33. To oxidize the iron in 1 gram of +FeSO_{4}(NH_{4})_{2}SO_{4}.6H_{2}O (mol. wgt. 392) requires 3 cc. of +a given solution of HNO_{3}. What is the normality of the nitric +acid when used as an acid? 6FeSO_{4} + 2HNO_{3} + 2H_{2}SO_{4} = +3Fe_{2}(SO_{4})_{3} + 2NO + 4H_{2}O. + +!Answer!: 0.2835 N. + +34. The same volume of carbon dioxide at the same temperature and the +same pressure is liberated from a 1 gram sample of dolomite, by adding +an excess of hydrochloric acid, as can be liberated by the addition of +35 cc. of 0.5 N hydrochloric acid to an excess of any pure or impure +carbonate. Calculate the percentage of CO_{2} in the dolomite. + +!Answer!: 38.5%. + +35. How many cubic centimeters of sulphuric acid (sp. gr. 1.84, +containing 96% H_{2}SO_{4} by weight) will be required to displace the +chloride in the calcium chloride formed by the action of 100 cc. of +0.1072 N hydrochloric acid on an excess of calcium carbonate, and how +many grams of CaSO_{4} will be formed? + +!Answers!: 0.298 cc.; 0.7300 gram. + +36. Potassium hydroxide which has been exposed to the air is found on +analysis to contain 7.62% water, 2.38% K_{2}CO_{3}. and 90% KOH. What +weight of residue will be obtained if one gram of this sample is added +to 46 cc. of normal hydrochloric acid and the resulting solution, +after exact neutralization with 1.070 N potassium hydroxide solution, +is evaporated to dryness? + +!Answer!: 3.47 grams. + +37. A chemist received four different solutions, with the statement +that they contained either pure NaOH; pure Na_{2}CO_{3}; pure +NaHCO_{3}, or mixtures of these substances. From the following data +identify them: + +Sample I. On adding phenolphthalein to a solution of the substance, it +gave no color to the solution. + +Sample II. On titrating with standard acid, it required 15.26 cc. for +a change in color, using phenolphthalein, and 17.90 cc. additional, +using methyl orange as an indicator. + +Sample III. The sample was titrated with hydrochloric acid until the +pink of phenolphthalein disappeared, and on the addition of methyl +orange the solution was colored pink. + +Sample IV. On titrating with hydrochloric acid, using phenolphthalein, +15.00 cc. were required. A new sample of the same weight required +exactly 30 cc. of the same acid for neutralization, using methyl +orange. + +!Answers!: (a) NaHCO_{3}; (b) NaHCO_{3}+Na_{2}CO_{3}; (c)NaOH; (d) +Na_{2}CO_{3}. + +38. In the analysis of a sample of KHC_{4}H_{4}O_{6} the following +data are obtained: Weight sample = 0.4732 gram. NaOH solution used = +24.97 cc. 3.00 cc. NaOH = 1 cc. of H_{3}PO_{4} solution of which 1 +cc. will precipitate 0.01227 gram of magnesium as MgNH_{4}PO_{4}. +Calculate the percentage of KHC_{4}H_{4}O_{6}. + +!Answer!: 88.67%. + +39. A one-gram sample of sodium hydroxide which has been exposed to +the air for some time, is dissolved in water and diluted to exactly +500 cc. One hundred cubic centimeters of the solution, when titrated +with 0.1062 N hydrochloric acid, using methyl orange as an indicator, +requires 38.60 cc. for complete neutralization. Barium chloride in +excess is added to a second portion of 100 cc. of the solution, which +is diluted to exactly 250 cc., allowed to stand and filtered. Two +hundred cubic centimeters of this filtrate require 29.62 cc. of 0.1062 +N hydrochloric acid for neutralization, using phenolphthalein as an +indicator. Calculate percentage of NaOH, Na_{2}CO_{3}, and H_{2}O. + +!Answers!: 78.63% NaOH; 4.45% Na_{2}CO_{3}; 16.92% H_{2}O. + +40. A sodium hydroxide solution (made from solid NaOH which has been +exposed to the air) was titrated against a standard acid using methyl +orange as an indicator, and was found to be exactly 0.1 N. This +solution was used in the analysis of a material sold at 2 cents per +pound per cent of an acid constituent A, and always mixed so that +it was supposed to contain 15% of A, on the basis of the analyst's +report. Owing to the carelessness of the analyst's assistant, the +sodium hydroxide solution was used with phenolphthalein as an +indicator in cold solution in making the analyses. The concern +manufacturing this material sells 600 tons per year, and when the +mistake was discovered it was estimated that at the end of a year +the error in the use of indicators would either cost them or their +customers $6000. Who would lose and why? Assuming the impure NaOH used +originally in making the titrating solution consisted of NaOH and +Na_{2}CO_{3} only, what per cent of each was present? + +!Answers!: Customer lost; 3.94% Na_{2}CO_{3}; 96.06% NaOH. + +41. In the standardization of a K_{2}Cr_{2}O_{7} solution against iron +wire, 99.85% pure, 42.42 cc. of the solution were added. The weight of +the wire used was 0.22 gram. 3.27 cc. of a ferrous sulphate solution +having a normal value as a reducing agent of 0.1011 were added +to complete the titration. Calculate the normal value of the +K_{2}Cr_{2}O_{7}. + +!Answer!: 0.1006 N. + +42. What weight of iron ore containing 56.2% Fe should be taken to +standardize an approximately 0.1 N oxidizing solution, if not more +than 47 cc. are to be used? + +!Answer!: 0.4667 gram. + +43. One tenth gram of iron wire, 99.78% pure, is dissolved in +hydrochloric acid and the iron oxidized completely with bromine water. +How many grams of stannous chloride are there in a liter of solution +if it requires 9.47 cc. to just reduce the iron in the above? What +is the normal value of the stannous chloride solution as a reducing +agent? + +!Answer!: 17.92 grams; 0.1888 N. + +44. One gram of an oxide of iron is fused with potassium acid sulphate +and the fusion dissolved in acid. The iron is reduced with stannous +chloride, mercuric chloride is added, and the iron titrated with a +normal K_{2}Cr_{2}O_{7} solution. 12.94 cc. were used. What is the +formula of the oxide, FeO, Fe_{2}O_{3}, or Fe_{3}O_{4}? + +!Answer!: Fe_{3}O_{4}. + +45. If an element has 98 for its atomic weight, and after reduction +with stannous chloride could be oxidized by bichromate to a state +corresponding to an XO_{4}^{-} anion, compute the oxide, or valence, +corresponding to the reduced state from the following data: 0.3266 +gram of the pure element, after being dissolved, was reduced with +stannous chloride and oxidized by 40 cc. of K_{2}Cr_{2}O_{7}, of which +one cc. = 0.1960 gram of FeSO_{4}(NH_{4})_{2}SO_{4}.6H_{2}O. + +!Answer!: Monovalent. + +46. Determine the percentage of iron in a sample of limonite from the +following data: Sample = 0.5000 gram. KMnO_{4} used = 50 cc. 1 cc. +KMnO_{4} = 0.005317 gram Fe. FeSO_{4} used = 6 cc. 1 cc. FeSO_{4} = +0.009200 gram FeO. + +!Answer!: 44.60%. + +47. If 1 gram of a silicate yields 0.5000 gram of Fe_{2}O_{3} and +Al_{2}O_{3} and the iron present requires 25 cc. of 0.2 N KMnO_{4}, +calculate the percentage of FeO and Al_{2}O_{3} in the sample. + +!Answer!: 35.89% FeO; 10.03% Al_{2}O_{3}. + +48. A sample of magnesia limestone has the following composition: +Silica, 3.00%; ferric oxide and alumina, 0.20%; calcium oxide, 33.10%; +magnesium oxide, 20.70%; carbon dioxide, 43.00%. In manufacturing lime +from the above the carbon dioxide is reduced to 3.00%. How many cubic +centimeters of normal KMnO_{4} will be required to determine the +calcium oxide volumetrically in a 1 gram sample of the lime? + +!Answer!: 20.08 cc. + +49. If 100 cc. of potassium bichromate solution (10 gram +K_{2}Cr_{2}O_{7} per liter), 5 cc. of 6 N sulphuric acid, and 75 cc. +of ferrous sulphate solution (80 grams FeSO_{4}.7H_{2}O per liter) are +mixed, and the resulting solution titrated with 0.2121 N KMnO_{4}, how +many cubic centimeters of the KMnO_{4} solution will be required to +oxidize the iron? + +!Answer!: 5.70 cc. + +50. If a 0.5000 gram sample of limonite containing 59.50 per cent +Fe_{2}O_{3} requires 40 cc. of KMnO_{4} to oxidize the iron, what +is the value of 1 cc. of the permanganate in terms of (a) Fe, (b) +H_{2}C_{2}O_{4}.2H_{2}O? + +!Answers!: (a) 0.005189 gram; (b) 0.005859 gram. + +51. A sample of pyrolusite weighing 0.6000 gram is treated with 0.9000 +gram of oxalic acid. The excess oxalic acid requires 23.95 cc. of +permanganate (1 cc. = 0.03038 gram FeSO_{4}.7H_{2}O). What is the +percentage of MnO_{2}, in the sample? + +!Answer!: 84.47%. + +52. A solution contains 50 grams of +KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O per liter. What is the normal +value of the solution (a) as an acid, and (b) as a reducing agent? + +!Answers!: (a) 0.5903 N; (b) 0.7872 N. + +53. In the analysis of an iron ore containing 60% Fe_{2}O_{3}, a +sample weighing 0.5000 gram is taken and the iron is reduced with +sulphurous acid. On account of failure to boil out all the excess +SO_{2}, 38.60 cubic centimeters of 0.1 N KMnO_{4} were required to +titrate the solution. What was the error, percentage error, and what +weight of sulphur dioxide was in the solution? + +!Answers!: (a) 1.60%; (b) 2.67%; (c) 0.00322 gram. + +54. From the following data, calculate the ratio of the nitric acid as +an oxidizing agent to the tetroxalate solution as a reducing agent: +1 cc. HNO_{3} = 1.246 cc. NaOH solution; 1 cc. NaOH = 1.743 cc. +KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O solution; Normal value NaOH = +0.12. + +!Answer!: 4.885. + +55. Given the following data: 25 cc. of a hydrochloric acid, when +standardized gravimetrically as silver chloride, yields a precipitate +weighing 0.5465 gram. 24.35 cc. of the hydrochloric acid are exactly +equivalent to 30.17 cc. of KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O +solution. How much water must be added to a liter of the oxalate +solution to make it exactly 0.025 N as a reducing agent? + +!Answer!: 5564 cc. + +56. Ten grams of a mixture of pure potassium tetroxalate +(KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O) and pure oxalic acid +(H_{2}C_{2}O_{4}.2H_{2}O) are dissolved in water and diluted to +exactly 1000 cc. The normal value of the oxalate solution when used as +an acid is 0.1315. Calculate the ratio of tetroxalate to oxalate used +in making up the solution and the normal value of the solution as a +reducing agent. + +!Answers!: 2:1; 0.1577 N. + +57. A student standardized a solution of NaOH and one of KMnO_{4} +against pure KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O and found the former +to be 0.07500 N as an alkali and the latter exactly 0.1 N as an +oxidizing agent. By coincidence, exactly 47.26 cc. were used in each +standardization. Find the ratio of the oxalate used in the +NaOH standardization to the oxalate used in the permanganate +standardization. + +!Answer!: 1:1. + +58. A sample of apatite weighing 0.60 gram is analyzed for its +phosphoric anhydride content. If the phosphate is precipitated as +(NH_{4})_{3}PO_{4}.12MoO_{3}, and the precipitate (after solution and +reduction of the MoO_{3} to Mo_{24}O_{37}), requires 100 cc. of normal +KMnO_{4} to oxidize it back to MoO_{3}, what is the percentage of +P_{2}O_{5}? + +!Answer!: 33.81%. + +59. In the analysis of a sample of steel weighing 1.881 grams the +phosphorus was precipitated with ammonium molybdate and the yellow +precipitate was dissolved, reduced and titrated with KMnO_{4}. If the +sample contained 0.025 per cent P and 6.01 cc. of KMnO_{4} were used, +to what oxide was the molybdenum reduced? 1 cc. KMnO_{4} = 0.007188 +gram Na_{2}C_{2}O_{4}. + +!Answer!: Mo_{4}O_{5}. + +60. What is the value of 1 cc. of an iodine solution (1 cc. equivalent +to 0.0300 gram Na_{2}S_{2}O_{3}) in terms of As_{2}O_{3}? + +!Answer!: 0.009385 gram. + +61. 48 cc. of a solution of sodium thiosulphate are required to +titrate the iodine liberated from an excess of potassium iodide +solution by 0.3000 gram of pure KIO_{3}. (KIO_{3} + 5KI + 3H_{2}SO_{4} += 3K_{2}SO_{4} + 3I_{2} + 3H_{2}O.) What is the normal strength of the +sodium thiosulphate and the value of 1 cc. of it in terms of iodine? + +!Answers!: 0.1753 N; 0.02224 gram. + +62. One thousand cubic centimeters of 0.1079 N sodium thiosulphate +solution is allowed to stand. One per cent by weight of the +thiosulphate is decomposed by the carbonic acid present in the +solution. To what volume must the solution be diluted to make it +exactly 0.1 N as a reducing agent? (Na_{2}S_{2}O_{3} + 2H_{2}CO_{3} = +H_{2}SO_{3} + 2NaHCO_{3} + S.) + +!Answer!: 1090 cc. + +63. An analyzed sample of stibnite containing 70.05% Sb is given for +analysis. A student titrates it with a solution of iodine of which 1 +cc. is equivalent to 0.004950 gram of As_{2}O_{3}. Due to an error on +his part in standardization, the student's analysis shows the sample +to contain 70.32% Sb. Calculate the true normal value of the iodine +solution, and the percentage error in the analysis. + +!Answers!: 0.1000 N; 0.39%. + +64. A sample of pyrolusite weighing 0.5000 gram is treated with an +excess of hydrochloric acid, the liberated chlorine is passed into +potassium iodide and the liberated iodine is titrated with sodium +thiosulphate solution (49.66 grams of pure Na_{2}S_{2}O_{3}.5H_{2}O +per liter). If 38.72 cc. are required, what volume of 0.25 normal +permanganate solution will be required in an indirect determination +in which a similar sample is reduced with 0.9012 gram +H_{2}C_{2}O_{4}.2H_{2}O and the excess oxalic acid titrated? + +!Answer!: 26.22 cc. + +65. In the determination of sulphur in steel by evolving the sulphur +as hydrogen sulphide, precipitating cadmium sulphide by passing the +liberated hydrogen sulphide through ammoniacal cadmium chloride +solution, and decomposing the CdS with acid in the presence of a +measured amount of standard iodine, the following data are obtained: +Sample, 5.027 grams; cc. Na_{2}S_{2}O_{3} sol. = 12.68; cc. Iodine +sol. = 15.59; 1 cc. Iodine sol. = 1.086 cc. Na_{2}S_{2}O_{3} sol.; 1 +cc. Na_{2}S_{2}O_{3}= 0.005044 gram Cu. Calculate the percentage of +sulphur. (H_{2}S + I_{2} = 2HI + S.) + +!Answer!: 0.107%. + +66. Given the following data, calculate the percentage of iron in +a sample of crude ferric chloride weighing 1.000 gram. The iodine +liberated by the reaction 2FeCl_{3}+ 2HI = 2HCl + 2FeCl_{2} + I_{2} is +reduced by the addition of 50 cc. of sodium thiosulphate solution and +the excess thiosulphate is titrated with standard iodine and requires +7.85 cc. 45 cc. I_{2} solution = 45.95 cc. Na_{2}S_{2}O_{3} solution; +45 cc. As_{2}O_{3} solution = 45.27 cc. I_{2} solution. 1 cc. arsenite +solution = 0.005160 gram As_{2}O_{3}. + +!Answer!: 23.77%. + +67. Sulphide sulphur was determined in a sample of reduced barium +sulphate by the evolution method, in which the sulphur was evolved as +hydrogen sulphide and was passed into CdCl_{2} solution, the acidified +precipitate being titrated with iodine and thiosulphate. Sample, 5.076 +grams; cc. I_{2} = 20.83; cc. Na_{2}S_{2}O_{3} = 12.37; 43.45 cc. +Na_{2}S_{2}O_{3} = 43.42 cc. I_{2}; 8.06 cc. KMnO_{4} = 44.66 cc. +Na_{2}S_{2}O_{3}; 28.87 cc. KMnO_{4} = 0.2004 gram Na_{2}C_{2}O_{4}. +Calculate the percentage of sulphide sulphur in the sample. + +!Answer!: 0.050%. + +68. What weight of pyrolusite containing 89.21% MnO_{2} will oxidize +the same amount of oxalic acid as 37.12 cc. of a permanganate +solution, of which 1 cc. will liberate 0.0175 gram of I_{2} from KI? + +!Answer!: 0.2493 gram. + +69. A sample of pyrolusite weighs 0.2400 gram and is 92.50% pure +MnO_{2}. The iodine liberated from KI by the manganese dioxide is +sufficient to react with 46.24 cc. of Na_{2}S_{2}O_{3} sol. What is +the normal value of the thiosulphate? + +!Answer!:: 0.1105 N. + +70. In the volumetric analysis of silver coin (90% Ag), using a +0.5000 gram sample, what is the least normal value that a potassium +thiocyanate solution may have and not require more than 50 cc. of +solution in the analysis? + +!Answer!: 0.08339 N. + +71. A mixture of pure lithium chloride and barium bromide weighing +0.6 gram is treated with 45.15 cubic centimeters of 0.2017 N silver +nitrate, and the excess titrated with 25 cc. of 0.1 N KSCN solution, +using ferric alum as an indicator. Calculate the percentage of bromine +in the sample. + +!Answer!: 40.11%. + +72. A mixture of the chlorides of sodium and potassium from 0.5000 +gram of a feldspar weighs 0.1500 gram, and after solution in water +requires 22.71 cc. of 0.1012 N silver nitrate for the precipitation of +the chloride ions. What are the percentages of Na_{2}O and K_{2}O in +the feldspar? + +!Answer!: 8.24% Na_{2}O; 9.14% K_{2}O. + + +GRAVIMETRIC ANALYSIS + +73. Calculate (a) the grams of silver in one gram of silver chloride; +(b) the grams of carbon dioxide liberated by the addition of an excess +of acid to one gram of calcium carbonate; (c) the grams of MgCl_{2} +necessary to precipitate 1 gram of MgNH_{4}PO_{4}. + +!Answers!: (a) 0.7526; (b) 0.4397; (c) 0.6940. + +74. Calculate the chemical factor for (a) Sn in SnO_{2}; (b) MgO +in Mg_{2}P_{2}O_{7}; (c) P_{2}O_{5} in Mg_{2}P_{2}O_{7}; (d) Fe in +Fe_{2}O_{3}; (e) SO_{4} in BaSO_{4}. + +!Answers!: (a) 0.7879; (b) 0.3620; (c) 0.6378; (d) 0.6990; (e) 0.4115. + +75. Calculate the log factor for (a) Pb in PbCrO_{4}; (b) Cr_{2}O_{3} +in PbCrO_{4}; (c) Pb in PbO_{2} and (d) CaO in CaC_{2}O_{4}. + +!Answers!: (a) 9.8069-10, (b) 9.3713-10; (c) 9.9376-10; (d) 9.6415-10. + +76. How many grams of Mn_{3}O_{4} can be obtained from 1 gram of +MnO_{2}? + +!Answer!: 0.8774 gram. + +77. If a sample of silver coin weighing 0.2500 gram gives a +precipitate of AgCl weighing 0.2991 gram, what weight of AgI could +have been obtained from the same weight of sample, and what is the +percentage of silver in the coin? + +!Answers!: 0.4898 gr.; 90.05%. + +78. How many cubic centimeters of hydrochloric acid (sp. gr. 1.13 +containing 25.75% HCl by weight) are required to exactly neutralize +25 cc. of ammonium hydroxide (sp. gr. .90 containing 28.33% NH_{3} by +weight)? + +!Answer!: 47.03 cc. + +79. How many cubic centimeters of ammonium hydroxide solution (sp. gr. +0.96 containing 9.91% NH_{3} by weight) are required to precipitate +the aluminium as aluminium hydroxide from a two-gram sample of alum +(KAl(SO_{4})_{2}.12H_{2}O)? What will be the weight of the ignited +precipitate? + +!Answers!: 2.26 cc.; 0.2154 gram. + +80. What volume of nitric acid (sp. gr. 1.05 containing 9.0% +HNO_{3} by weight) is required to oxidize the iron in one gram of +FeSO_{4}.7H_{2}O in the presence of sulphuric acid? 6FeSO_{4} + +2HNO_{3} + 3H_{2}SO_{4} = 3Fe_{2}(SO_{4})_{3} + 2NO + 4H_{2}O. + +!Answer!: 0.80 cc. + +81. If 0.7530 gram of ferric nitrate (Fe(NO_{3})_{3}.9H_{2}O) is +dissolved in water and 1.37 cc. of HCl (sp. gr. 1.11 containing 21.92% +HCl by weight) is added, how many cubic centimeters of ammonia (sp. +gr. 0.96 containing 9.91% NH_{3} by weight) are required to neutralize +the acid and precipitate the iron as ferric hydroxide? + +!Answer!: 2.63 cc. + +82. To a suspension of 0.3100 gram of Al(OH)_{3} in water are added +13.00 cc. of aqueous ammonia (sp. gr. 0.90 containing 28.4% NH_{3} by +weight). How many cubic centimeters of sulphuric acid (sp. gr. 1.18 +containing 24.7% H_{2}SO_{4} by weight) must be added to the mixture +in order to bring the aluminium into solution? + +!Answer!: 34.8 cc. + +83. How many cubic centimeters of sulphurous acid (sp. gr. 1.04 +containing 75 grams SO_{2} per liter) are required to reduce the +iron in 1 gram of ferric alum (KFe(SO_{4})_{2}.12H_{2}O)? +Fe_{2}(SO_{4})_{3} + SO_{2} + 2H_{2}O = 2FeSO_{4} + 2H_{2}SO_{4}. + +!Answer!: 0.85 cc. + +84. How many cubic centimeters of a solution of potassium bichromate +containing 26.30 grams of K_{2}Cr_{2}O_{7} per liter must be taken +in order to yield 0.6033 gram of Cr_{2}O_{3} after reduction and +precipitation of the chromium? + +K_{2}Cr_{2}O_{7} + 3SO_{2} + H_{2}SO_{4} = K_{2}SO_{4} + +Cr_{2}(SO_{4})_{3} + H_{2}O. + +!Answer!: 44.39 cc. + +85. How many cubic centimeters of ammonium hydroxide (sp. gr. 0.946 +containing 13.88% NH_{3} by weight) are required to precipitate +the iron as Fe(OH)_{3} from a sample of pure +FeSO_{4}.(NH_{4})_{2}SO_{4}.6H_{2}O, which requires 0.34 cc. of nitric +acid (sp. gr. 1.350 containing 55.79% HNO_{3} by weight) for oxidation +of the iron? (See problem No. 80 for reaction.) + +!Answer!: 4.74 cc. + +86. In the analysis of an iron ore by solution, oxidation and +precipitation of the iron as Fe(OH)_{3}, what weight of sample must be +taken for analysis so that each one hundredth of a gram of the ignited +precipitate of Fe_{2}O_{3} shall represent one tenth of one per cent +of iron? + +!Answer!: 6.99 grams. + +87. What weight in grams of impure ferrous ammonium sulphate should +be taken for analysis so that the number of centigrams of BaSO_{4} +obtained will represent five times the percentage of sulphur in the +sample? + +!Answer!: 0.6870 gram. + +88. What weight of magnetite must be taken for analysis in order that, +after precipitating and igniting all the iron to Fe_{2}O_{3}, the +percentage of Fe_{2}O_{4} in the sample may be found by multiplying +the weight in grams of the ignited precipitate by 100? + +!Answer!: 0.9665 gram. + +89. After oxidizing the arsenic in 0.5000 gram of pure As_{2}S_{3} to +arsenic acid, it is precipitated with "magnesia mixture" (MgCl_{2} + +2NH_{4}Cl). If exactly 12.6 cc. of the mixture are required, how many +grams of MgCl_{2} per liter does the solution contain? H_{3}AsO_{4} + +MgCl_{2} + 3NH_{4}OH = MgNH_{4}AsO_{4} + 2NH_{4}Cl + 3H_{2}O. + +!Answer!: 30.71 grams. + +90. A sample is prepared for student analysis by mixing pure apatite +(Ca_{3}(PO_{4})_{2}.CaCl_{2}) with an inert material. If 1 gram of +the sample gives 0.4013 gram of Mg_{2}P_{2}O_{7}, how many cubic +centimeters of ammonium oxalate solution (containing 40 grams of +(NH_{4})_{2}C_{2}O_{4}.H_{2}O per liter) would be required to +precipitate the calcium from the same weight of sample? + +!Answer!: 25.60 cc. + +91. If 0.6742 gram of a mixture of pure magnesium carbonate and pure +calcium carbonate, when treated with an excess of hydrochloric acid, +yields 0.3117 gram of carbon dioxide, calculate the percentage of +magnesium oxide and of calcium oxide in the sample. + +!Answers!: 13.22% MgO; 40.54% CaO. 92. The calcium in a sample of +dolomite weighing 0.9380 gram is precipitated as calcium oxalate and +ignited to calcium oxide. What volume of gas, measured over water +at 20°C. and 765 mm. pressure, is given off during ignition, if the +resulting oxide weighs 0.2606 gram? (G.M.V. = 22.4 liters; V.P. water +at 20°C. = 17.4 mm.) + +!Answer!: 227 cc. + +93. A limestone is found to contain 93.05% CaCO_{3}, and 5.16 % +MgCO_{3}. Calculate the weight of CaO obtainable from 3 tons of the +limestone, assuming complete conversion to oxide. What weight of +Mg_{2}P_{2}O_{7} could be obtained from a 3-gram sample of the +limestone? + +!Answers!: 1.565 tons; 0.2044 gram. + +94. A sample of dolomite is analyzed for calcium by precipitating +as the oxalate and igniting the precipitate. The ignited product is +assumed to be CaO and the analyst reports 29.50% Ca in the sample. +Owing to insufficient ignition, the product actually contained 8% of +its weight of CaCO_{3}. What is the correct percentage of calcium in +the sample, and what is the percentage error? + +!Answers!: 28.46%; 3.65% error. + +95. What weight of impure calcite (CaCO_{3}) should be taken for +analysis so that the volume in cubic centimeters of CO_{2} obtained by +treating with acid, measured dry at 18°C. and 763 mm., shall equal the +percentage of CaO in the sample? + +!Answer!: 0.2359 gram. + +96. How many cubic centimeters of HNO_{3} (sp. gr. 1.13 containing +21.0% HNO_{3} by weight) are required to dissolve 5 grams of brass, +containing 0.61% Pb, 24.39% Zn, and 75% Cu, assuming reduction of the +nitric acid to NO by each constituent? What fraction of this volume of +acid is used for oxidation? + +!Answers!: 55.06 cc.; 25%. + +97. What weight of metallic copper will be deposited from a cupric +salt solution by a current of 1.5 amperes during a period of 45 +minutes, assuming 100% current efficiency? (1 Faraday = 96,500 +coulombs.) + +!Answer!: 1.335 grams. + +98. In the electrolysis of a 0.8000 gram sample of brass, there is +obtained 0.0030 gram of PbO_{2}, and a deposit of metallic copper +exactly equal in weight to the ignited precipitate of Zn_{2}P_{2}O_{7} +subsequently obtained from the solution. What is the percentage +composition of the brass? + +!Answers!: 69.75% Cu; 29.92% Zn; 0.33% Pb. + +99. A sample of brass (68.90% Cu; 1.10% Pb and 30.00% Zn) weighing +0.9400 gram is dissolved in nitric acid. The lead is determined by +weighing as PbSO_{4}, the copper by electrolysis and the zinc by +precipitation with (NH_{4})_{2}HPO_{4} in a neutral solution. + +(a) Calculate the cubic centimeters of nitric acid (sp. gr. 1.42 +containing 69.90% HNO_{3} by weight) required to just dissolve the +brass, assuming reduction to NO. + +!Answer!: 2.48 cc. + +(b) Calculate the cubic centimeters of sulphuric acid (sp. gr. 1.84 +containing 94% H_{2}SO_{4} by weight) to displace the nitric acid. + +!Answer!: 0.83 cc. + +(c) Calculate the weight of PbSO_{4}. + +!Answer!: 0.0152 gram. + +(d) The clean electrode weighs 10.9640 grams. Calculate the weight +after the copper has been deposited. + +!Answer!: 11.6116 grams. + +(e) Calculate the grams of (NH_{4})_{2}HPO_{4} required to precipitate +the zinc as ZnNH_{4}PO_{4}. + +!Answer!: 0.5705 gram. + +(f) Calculate the weight of ignited Zn_{2}P_{2}O_{7}. + +!Answer!: 0.6573 gram. + +100. If in the analysis of a brass containing 28.00% zinc an error is +made in weighing a 2.5 gram portion by which 0.001 gram too much is +weighed out, what percentage error in the zinc determination would +result? What volume of a solution of sodium hydrogen phosphate, +containing 90 grams of Na_{2}HPO_{4}.12H_{2}O per liter, would be +required to precipitate the zinc as ZnNH_{4}PO_{4} and what weight of +precipitate would be obtained? + +!Answers!: (a) 0.04% error; (b) 39.97 cc.; (c) 1.909 grams. + +101. A sample of magnesium carbonate, contaminated with SiO_{2} as its +only impurity, weighs 0.5000 gram and loses 0.1000 gram on ignition. +What volume of disodium phosphate solution (containing 90 grams +Na_{2}HPO_{4}.12H_{2}O per liter) will be required to precipitate the +magnesium as magnesium ammonium phosphate? + +!Answer!: 9.07 cc. + +102. 2.62 cubic centimeters of nitric acid (sp. gr. 1.42 containing +69.80% HNO_{2} by weight) are required to just dissolve a sample +of brass containing 69.27% Cu; 0.05% Pb; 0.07% Fe; and 30.61% Zn. +Assuming the acid used as oxidizing agent was reduced to NO in every +case, calculate the weight of the brass and the cubic centimeters of +acid used as acid. + +!Answer!: 0.992 gram; 1.97 cc. + +103. One gram of a mixture of silver chloride and silver bromide is +found to contain 0.6635 gram of silver. What is the percentage of +bromine? + +!Answer!: 21.30%. + +104. A precipitate of silver chloride and silver bromide weighs 0.8132 +gram. On heating in a current of chlorine, the silver bromide is +converted to silver chloride, and the mixture loses 0.1450 gram +in weight. Calculate the percentage of chlorine in the original +precipitate. + +!Answer!: 6.13%. + +105. A sample of feldspar weighing 1.000 gram is fused and the silica +determined. The weight of silica is 0.6460 gram. This is fused with 4 +grams of sodium carbonate. How many grams of the carbonate actually +combined with the silica in fusion, and what was the loss in weight +due to carbon dioxide during the fusion? + +!Answers!: 1.135 grams; 0.4715 gram. + +106. A mixture of barium oxide and calcium oxide weighing 2.2120 grams +is transformed into mixed sulphates, weighing 5.023 grams. Calculate +the grams of calcium oxide and barium oxide in the mixture. + +!Answers!: 1.824 grams CaO; 0.3877 gram BaO. + + + + +APPENDIX + + +ELECTROLYTIC DISSOCIATION THEORY + +The following brief statements concerning the ionic theory and a few +of its applications are intended for reference in connection with the +explanations which are given in the Notes accompanying the various +procedures. The reader who desires a more extended discussion of the +fundamental theory and its uses is referred to such books as Talbot +and Blanchard's !Electrolytic Dissociation Theory! (Macmillan +Company), or Alexander Smith's !Introduction to General Inorganic +Chemistry! (Century Company). + +The !electrolytic dissociation theory!, as propounded by Arrhenius in +1887, assumes that acids, bases, and salts (that is, electrolytes) +in aqueous solution are dissociated to a greater or less extent into +!ions!. These ions are assumed to be electrically charged atoms or +groups of atoms, as, for example, H^{+} and Br^{-} from hydrobromic +acid, Na^{+} and OH^{-} from sodium hydroxide, 2NH_{4}^{+} and +SO_{4}^{--} from ammonium sulphate. The unit charge is that which is +dissociated with a hydrogen ion. Those upon other ions vary in sign +and number according to the chemical character and valence of the +atoms or radicals of which the ions are composed. In any solution the +aggregate of the positive charges upon the positive ions (!cations!) +must always balance the aggregate negative charges upon the negative +ions (!anions!). + +It is assumed that the Na^{+} ion, for example, differs from the +sodium atom in behavior because of the very considerable electrical +charge which it carries and which, as just stated, must, in an +electrically neutral solution, be balanced by a corresponding negative +charge on some other ion. When an electric current is passed through a +solution of an electrolyte the ions move with and convey the current, +and when the cations come into contact with the negatively charged +cathode they lose their charges, and the resulting electrically +neutral atoms (or radicals) are liberated as such, or else enter at +once into chemical reaction with the components of the solution. + +Two ions of identically the same composition but with different +electrical charges may exhibit widely different properties. For +example, the ion MnO_{4}^{-} from permanganates yields a purple-red +solution and differs in its chemical behavior from the ion +MnO_{4}^{--} from manganates, the solutions of which are green. + +The chemical changes upon which the procedures of analytical chemistry +depend are almost exclusively those in which the reacting substances +are electrolytes, and analytical chemistry is, therefore, essentially +the chemistry of the ions. The percentage dissociation of the same +electrolyte tends to increase with increasing dilution of its +solution, although not in direct proportion. The percentage +dissociation of different electrolytes in solutions of equivalent +concentrations (such, for example, as normal solutions) varies widely, +as is indicated in the following tables, in which approximate figures +are given for tenth-normal solutions at a temperature of about 18°C. + + ACIDS +========================================================================= + | + SUBSTANCE | PERCENTAGE DISSOCIATION IN + | 0.1 EQUIVALENT SOLUTION +_____________________________________________|___________________________ + | +HCl, HBr, HI, HNO_{3} | 90 + | +HClO_{3}, HClO_{4}, HMnO_{4} | 90 + | +H_{2}SO_{4} <--> H^{+} + HSO_{4}^{-} | 90 + | +H_{2}C_{2}O_{4} <--> H^{+} + HC_{2}O_{4}^{-} | 50 + | +H_{2}SO_{3} <--> H^{+} + HSO{_}3^{-} | 20 + | +H_{3}PO_{4} <--> H^{+} + H_{2}PO_{4}^{-} | 27 + | +H_{2}PO_{4}^{-} <--> H^{+} + HPO_{4}^{--} | 0.2 + | +H_{3}AsO_{4} <--> H^{+} + H_{2}AsO_{4}^{-} | 20 + | +HF | 9 + | +HC_{2}H_{3}O_{2} | 1.4 + | +H_{2}CO_{3} <--> H^{+} + HCO_{3}^{-} | 0.12 + | +H_{2}S <--> H^{+} + HS^{-} | 0.05 + | +HCN | 0.01 + | +========================================================================= + + + BASES +========================================================================= + | + SUBSTANCE | PERCENTAGE DISSOCIATION IN + | 0.1 EQUIVALENT SOLUTION +_____________________________________________|___________________________ + | +KOH, NaOH | 86 + | +Ba(OH)_{2} | 75 + | +NH_{4}OH | 1.4 + | +========================================================================= + + + SALTS +========================================================================= + | + TYPE OF SALT | PERCENTAGE DISSOCIATION IN + | 0.1 EQUIVALENT SOLUTION +_____________________________________________|___________________________ + | +R^{+}R^{-} | 86 + | +R^{++}(R^{-})_{2} | 72 + | +(R^{+})_{2}R^{--} | 72 + | +R^{++}R^{--} | 45 + | +========================================================================= + +The percentage dissociation is determined by studying the electrical +conductivity of the solutions and by other physico-chemical methods, +and the following general statements summarize the results: + +!Salts!, as a class, are largely dissociated in aqueous solution. + +!Acids! yield H^{+} ions in water solution, and the comparative +!strength!, that is, the activity, of acids is proportional to the +concentration of the H^{+} ions and is measured by the percentage +dissociation in solutions of equivalent concentration. The common +mineral acids are largely dissociated and therefore give a relatively +high concentration of H^{+} ions, and are commonly known as "strong +acids." The organic acids, on the other hand, belong generally to the +group of "weak acids." + +!Bases! yield OH^{-} ions in water solution, and the comparative +strength of the bases is measured by their relative dissociation in +solutions of equivalent concentration. Ammonium hydroxide is a weak +base, as shown in the table above, while the hydroxides of sodium and +potassium exhibit strongly basic properties. + +Ionic reactions are all, to a greater or less degree, !reversible +reactions!. A typical example of an easily reversible reaction is that +representing the changes in ionization which an electrolyte such as +acetic acid undergoes on dilution or concentration of its solutions, +!i.e.!, HC_{2}H_{3}O_{2} <--> H^{+} + C_{2}H_{3}O_{2}^{-}. As was +stated above, the ionization increases with dilution, the reaction +then proceeding from left to right, while concentration of the +solution occasions a partial reassociation of the ions, and the +reaction proceeds from right to left. To understand the principle +underlying these changes it is necessary to consider first the +conditions which prevail when a solution of acetic acid, which has +been stirred until it is of uniform concentration throughout, has come +to a constant temperature. A careful study of such solutions has shown +that there is a definite state of equilibrium between the constituents +of the solution; that is, there is a definite relation between the +undissociated acetic acid and its ions, which is characteristic for +the prevailing conditions. It is not, however, assumed that this is a +condition of static equilibrium, but rather that there is continual +dissociation and association, as represented by the opposing +reactions, the apparent condition of rest resulting from the fact that +the amount of change in one direction during a given time is exactly +equal to that in the opposite direction. A quantitative study of +the amount of undissociated acid, and of H^{+} ions and +C_{2}H_{3}O_{2}^{-} ions actually to be found in a large number of +solutions of acetic acid of varying dilution (assuming them to be in +a condition of equilibrium at a common temperature), has shown that +there is always a definite relation between these three quantities +which may be expressed thus: + +(!Conc'n H^{+} x Conc'n C_{2}H_{3}O_{2}^{-})/Conc'n HC_{2}H_{3}O_{2} = +Constant!. + +In other words, there is always a definite and constant ratio between +the product of the concentrations of the ions and the concentration of +the undissociated acid when conditions of equilibrium prevail. + +It has been found, further, that a similar statement may be made +regarding all reversible reactions, which may be expressed in general +terms thus: The rate of chemical change is proportional to the product +of the concentrations of the substances taking part in the reaction; +or, if conditions of equilibrium are considered in which, as stated, +the rate of change in opposite directions is assumed to be equal, then +the product of the concentrations of the substances entering into +the reaction stands in a constant ratio to the product of the +concentrations of the resulting substances, as given in the expression +above for the solutions of acetic acid. This principle is called the +!Law of Mass Action!. + +It should be borne in mind that the expression above for acetic acid +applies to a wide range of dilutions, provided the temperature remains +constant. If the temperature changes the value of the constant changes +somewhat, but is again uniform for different dilutions at that +temperature. The following data are given for temperatures of about +18°C.[1] + +========================================================================== + | | | | + MOLAL | FRACTION | MOLAL CONCENTRA- | MOLAL CONCENTRA- | VALUE OF +CONCENTRATION | IONIZED | TION OF H^{+} AND| TION OF UNDIS- | CONSTANT + CONSTANT | | ACETATE^{-} IONS | SOCIATED ACID | +______________|__________|__________________|__________________|__________ + | | | | + 1.0 | .004 | .004 | .996 | .0000161 + | | | | + 0.1 | .013 | .0013 | .0987 | .0000171 + | | | | + 0.01 | .0407 | .000407 | .009593 | .0000172 + | | | | +=========================================================================== + +[Footnote 1: Alexander Smith, !General Inorganic Chemistry!, p. 579.] + +The molal concentrations given in the table refer to fractions of a +gram-molecule per liter of the undissociated acid, and to fractions of +the corresponding quantities of H^{+} and C_{2}H_{3}O_{2}^{-} ions +per liter which would result from the complete dissociation of a +gram-molecule of acetic acid. The values calculated for the constant +are subject to some variation on account of experimental errors in +determining the percentage ionized in each case, but the approximate +agreement between the values found for molal and centimolal (one +hundredfold dilution) is significant. + +The figures given also illustrate the general principle, that the +!relative! ionization of an electrolyte increases with the dilution of +its solution. If we consider what happens during the (usually) brief +period of dilution of the solution from molal to 0.1 molal, for +example, it will be seen that on the addition of water the conditions +of concentration which led to equality in the rate of change, and +hence to equilibrium in the molal solution, cease to exist; and since +the dissociating tendency increases with dilution, as just stated, +it is true at the first instant after the addition of water that the +concentration of the undissociated acid is too great to be +permanent under the new conditions of dilution, and the reaction, +HC_{2}H_{3}O_{2} <--> H^{+} + C_{2}H_{3}O_{2}^{-}, will proceed from +left to right with great rapidity until the respective concentrations +adjust themselves to the new conditions. + +That which is true of this reaction is also true of all reversible +reactions, namely, that any change of conditions which occasions +an increase or a decrease in concentration of one or more of the +components causes the reaction to proceed in one direction or the +other until a new state of equilibrium is established. This principle +is constantly applied throughout the discussion of the applications +of the ionic theory in analytical chemistry, and it should be clearly +understood that whenever an existing state of equilibrium is disturbed +as a result of changes of dilution or temperature, or as a consequence +of chemical changes which bring into action any of the constituents of +the solution, thus altering their concentrations, there is always a +tendency to re-establish this equilibrium in accordance with the law. +Thus, if a base is added to the solution of acetic acid the H^{+} ions +then unite with the OH^{-} ions from the base to form undissociated +water. The concentration of the H^{+} ions is thus diminished, and +more of the acid dissociates in an attempt to restore equilbrium, +until finally practically all the acid is dissociated and neutralized. + +Similar conditions prevail when, for example, silver ions react with +chloride ions, or barium ions react with sulphate ions. In the former +case the dissociation reaction of the silver nitrate is AgNO_{3} <--> +Ag^{+} + NO_{3}^{-}, and as soon as the Ag^{+} ions unite with the +Cl^{-} ions the concentration of the former is diminished, more of the +AgNO_{3} dissociates, and this process goes on until the Ag^{+} ions +are practically all removed from the solution, if the Cl^{-} ions are +present in sufficient quantity. + +For the sake of accuracy it should be stated that the mass law cannot +be rigidly applied to solutions of those electrolytes which are +largely dissociated. While the explanation of the deviation from +quantitative exactness in these cases is not known, the law is still +of marked service in developing analytical methods along more logical +lines than was formerly practicable. It has not seemed wise to qualify +each statement made in the Notes to indicate this lack of quantitative +exactness. The student should recognize its existence, however, and +will realize its significance better as his knowledge of physical +chemistry increases. + +If we apply the mass law to the case of a substance of small +solubility, such as the compounds usually precipitated in quantitative +analysis, we derive what is known as the !solubility product!, as +follows: Taking silver chloride as an example, and remembering that it +is not absolutely insoluble in water, the equilibrium expression for +its solution is: + +(!Conc'n Ag^{+} x Conc'n Cl^{-})/Conc'n AgCl = Constant!. + +But such a solution of silver chloride which is in contact with the +solid precipitate must be saturated for the existing temperature, and +the quantity of undissociated AgCl in the solution is definite and +constant for that temperature. Since it is a constant, it may be +eliminated, and the expression becomes !Conc'n Ag^{+} x Conc'n +Cl^{-} = Constant!, and this is known as the solubility product. No +precipitation of a specific substance will occur until the product of +the concentrations of its ions in a solution exceeds the solubility +product for that substance; whenever that product is exceeded +precipitation must follow. + +It will readily be seen that if a substance which yields an ion in +common with the precipitated compound is added to such a solution as +has just been described, the concentration of that ion is +increased, and as a result the concentration of the other ion must +proportionately decrease, which can only occur through the formation +of some of the undissociated compound which must separate from the +already saturated solution. This explains why the addition of an +excess of the precipitant is often advantageous in quantitative +procedures. Such a case is discussed at length in Note 2 on page 113. + +Similarly, the ionization of a specific substance in solution tends to +diminish on the addition of another substance with a common ion, as, +for instance, the addition of hydrochloric acid to a solution +of hydrogen sulphide. Hydrogen sulphide is a weak acid, and the +concentration of the hydrogen ions in its aqueous solutions is very +small. The equilibrium in such a solution may be represented as: + +(!(Conc'n H^{+})^{2} x Conc'n S^{--})/Conc'n H_{2}S = Constant!, and a +marked increase in the concentration of the H^{+} ions, such as would +result from the addition of even a small amount of the highly ionized +hydrochloric acid, displaces the point of equilibrium and some of the +S^{--} ions unite with H^{+} ions to form undissociated H_{2}S. This +is of much importance in studying the reactions in which hydrogen +sulphide is employed, as in qualitative analysis. By a parallel course +of reasoning it will be seen that the addition of a salt of a weak +acid or base to solutions of that acid or base make it, in effect, +still weaker because they decrease its percentage ionization. + +To understand the changes which occur when solids are dissolved where +chemical action is involved, it should be remembered that no substance +is completely insoluble in water, and that those products of a +chemical change which are least dissociated will first form. Consider, +for example, the action of hydrochloric acid upon magnesium hydroxide. +The minute quantity of dissolved hydroxide dissociates thus: +Mg(OH)_{2} <--> Mg^{++} + 2OH^{-}. When the acid is introduced, +the H^{+} ions of the acid unite with the OH^{-} ions to form +undissociated water. The concentration of the OH^{-} ions is thus +diminished, more Mg(OH)_{2} dissociates, the solution is no longer +saturated with the undissociated compound, and more of the solid +dissolves. This process repeats itself with great rapidity until, if +sufficient acid is present, the solid passes completely into solution. + +Exactly the same sort of process takes place if calcium oxalate, for +example, is dissolved in hydrochloric acid. The C_{2}O_{4}^{--} ions +unite with the H^{+} ions to form undissociated oxalic acid, the acid +being less dissociated than normally in the presence of the H^{+} ions +from the hydrochloric acid (see statements regarding hydrogen sulphide +above). As the undissociated oxalic acid forms, the concentration of +the C_{2}O_{4}^{--} ions lessens and more CaC_{2}O_{4} dissolves, +as described for the Mg(OH)_{2} above. Numerous instances of the +applications of these principles are given in the Notes. + +Water itself is slightly dissociated, and although the resulting H^{+} +and OH^{-} ions are present only in minute concentrations (1 mol. of +dissociated water in 10^{7} liters), yet under some conditions they +may give rise to important consequences. The term !hydrolysis! is +applied to the changes which result from the reaction of these ions. +Any salt which is derived from a weak base or a weak acid (or both) +is subject to hydrolytic action. Potassium cyanide, for example, when +dissolved in water gives an alkaline solution because some of the +H^{+} ions from the water unite with CN^{-} ions to form (HCN), which +is a very weak acid, and is but very slightly dissociated. Potassium +hydroxide, which might form from the OH^{-} ions, is so largely +dissociated that the OH^{-} ions remain as such in the solution. The +union of the H^{+} ions with the CN^{-} ions to form the undissociated +HCN diminishes the concentration of the H^{+} ions, and more water +dissociates (H_{2}O <--> H^{+} + OH^{-}) to restore the equilibrium. +It is clear, however, that there must be a gradual accumulation of +OH^{-} ions in the solution as a result of these changes, causing the +solution to exhibit an alkaline reaction, and also that ultimately the +further dissociation of the water will be checked by the presence of +these ions, just as the dissociation of the H_{2}S was lessened by the +addition of HCl. + +An exactly opposite result follows the solution of such a salt as +Al_{2}(SO_{4})_{3} in water. In this case the acid is strong and the +base weak, and the OH^{-} ions form the little dissociated Al(OH)_{3}, +while the H^{+} ions remain as such in the solution, sulphuric acid +being extensively dissociated. The solution exhibits an acid reaction. + +Such hydrolytic processes as the above are of great importance in +analytical chemistry, especially in the understanding of the action of +indicators in volumetric analysis. (See page 32.) + +The impelling force which causes an element to pass from the atomic +to the ionic condition is termed !electrolytic solution pressure!, or +ionization tension. This force may be measured in terms of electrical +potential, and the table below shows the relative values for a number +of elements. + +In general, an element with a greater solution pressure tends to cause +the deposition of an element of less solution pressure when placed in +a solution of its salt, as, for instance, when a strip of zinc or +iron is placed in a solution of a copper salt, with the resulting +precipitation of metallic copper. + +Hydrogen is included in the table, and its position should be noted +with reference to the other common elements. For a more extended +discussion of this topic the student should refer to other treatises. + + POTENTIAL SERIES OF THE METALS + +__________________________________________________________________ + | | | + | POTENTIAL | | POTENTIAL + | IN VOLTS | | IN VOLTS +_____________________|___________|____________________|___________ + | | | +Sodium Na^{+} | +2.44 | Lead Pb^{++} | -0.13 +Calcium Ca^{++} | | Hydrogen H^{+} | -0.28 +Magnesium Mg^{++} | | Bismuth Bi^{+++}| +Aluminum A1^{+++} | +1.00 | Antimony | -0.75 +Manganese Mn^{++} | | Arsenic | +Zinc Zn^{++} | +0.49 | Copper Cu^{++} | -0.61 +Cadmium Cd^{++} | +0.14 | Mercury Hg^{+} | -1.03 +Iron Fe^{++} | +0.063 | Silver Ag^{+} | -1.05 +Cobalt Co^{++} | -0.045 | Platinum | +Nickel Ni^{++} | -0.049 | Gold | +Tin Sn^{++} | -0.085(?) | | +_____________________|___________|____________________|__________ + + + +THE FOLDING OF A FILTER PAPER + +If a filter paper is folded along its diameter, and again folded along +the radius at right angles to the original fold, a cone is formed on +opening, the angle of which is 60°. Funnels for analytical use are +supposed to have the same angle, but are rarely accurate. It is +possible, however, with care, to fit a filter thus folded into a +funnel in such a way as to prevent air from passing down between the +paper and the funnel to break the column of liquid in the stem, +which aids greatly, by its gentle suction, in promoting the rate of +filtration. + +Such a filter has, however, the disadvantage that there are three +thicknesses of paper back of half of its filtering surface, as a +consequence of which one half of a precipitate washes or drains more +slowly. Much time may be saved in the aggregate by learning to fold a +filter in such a way as to improve its effective filtering surface. +The directions which follow, though apparently complicated on first +reading, are easily applied and easily remembered. Use a 6-inch filter +for practice. Place four dots on the filter, two each on diameters +which are at right angles to each other. Then proceed as follows: +(1) Fold the filter evenly across one of the diameters, creasing it +carefully; (2) open the paper, turn it over, rotate it 90° to the +right, bring the edges together and crease along the other diameter; +(3) open, and rotate 45° to the right, bring edges together, and +crease evenly; (4) open, and rotate 90° to the right, and crease +evenly; (5) open, turn the filter over, rotate 22-(1/2)° to the right, +and crease evenly; (6) open, rotate 45° to the right and crease +evenly; (7) open, rotate 45° to the right and crease evenly; (8) open, +rotate 45° to the right and crease evenly; (9) open the filter, and, +starting with one of the dots between thumb and forefinger of the +right hand, fold the second crease to the left over on it, and do +the same with each of the other dots. Place it, thus folded, in the +funnel, moisten it, and fit to the side of the funnel. The filter will +then have four short segments where there are three thicknesses +and four where there is one thickness, but the latter are evenly +distributed around its circumference, thus greatly aiding the passage +of liquids through the paper and hastening both filtration and washing +of the whole contents of the filter. + + +!SAMPLE PAGES FOR LABORATORY RECORDS! + +!Page A! + +Date + +CALIBRATION OF BURETTE No. + +___________________________________________________________________________ + | | | | + BURETTE | DIFFERENCE | OBSERVED | DIFFERENCE | CALCULATED + READINGS | | WEIGHTS | | CORRECTION +_______________|______________|______________|______________|______________ + 0.02 | | 16.27 | | + 10.12 | 10.10 | 26.35 | 10.08 | -.02 + 20.09 | 9.97 | 36.26 | 9.91 | -.06 + 30.16 | 10.07 | 46.34 | 10.08 | +.01 + 40.19 | 10.03 | 56.31 | 9.97 | -.06 + 50.00 | 9.81 | 66.17 | 9.86 | +.05 +_______________|______________|______________|______________|______________ + + These data to be obtained in duplicate for each burette. + + +!Page B! + +Date + + +DETERMINATION OF COMPARATIVE STRENGTH HCl vs. NaOH + +___________________________________________________________________________ + | | + DETERMINATION | I | II +_________________________|________________________|________________________ + | | + | Corrected | Corrected +Final Reading HCl | 48.17 48.08 | 43.20 43.14 +Initial Reading HCl | 0.12 .12 | .17 .17 + | ----- ----- | ----- ----- + | 47.96 | 42.97 + | | + | Corrected | Corrected +Final Reading HCl | 46.36 46.29 | 40.51 40.37 +Initial Reading HCl | 1.75 1.75 | .50 .50 + | ----- ----- | ----- ----- + | 44.54 | 39.87 + | | + log cc. NaOH | 1.6468 | 1.6008 + colog cc. HCl | 8.3192 | 8.3668 + | ------ | ------ + | 9.9680 - 10 | 9.9676 - 10 + 1 cc. HCl | .9290 cc. NaOH | .9282 cc. NaOH + Mean | .9286 | +_________________________|________________________|________________________ + + +Signed + +!Page C! +Date + + +STANDARDIZATION OF HYDROCHLORIC ACID +===================================================================== + | | +Weight sample and tube| 9.1793 | 8.1731 + | 8.1731 | 6.9187 + | ------ | ------ + Weight sample | 1.0062 | 1.2544 + | | +Final Reading HCl | 39.97 39.83 | 49.90 49.77 +Initial Reading HCl | .00 .00 | .04 .04 + | ----- ----- | ----- ----- + | 39.83 | 49.73 + | | +Final Reading NaOH | .26 .26 | .67 .67 +Initial Reading NaOH | .12 .12 | .36 .36 + | --- --- | --- --- + | .14 | .31 + | | + | .14 | .31 +Corrected cc. HCl | 39.83 - ----- = 39.68 | 49.73 - ----- = 49.40 + | .9286 | .9286 + | | +log sample | 0.0025 | 0.0983 +colog cc | 8.4014 - 10 | 8.3063 - 10 +colog milli equivalent| 1.2757 | 1.2757 + | ------ | ------ + | 9.6796 - 10 | 9.6803 - 10 + | | +Normal value HCl | .4782 | .4789 + Mean | .4786 | + | | +===================================================================== + +Signed + + +!Page D! +Date + + +DETERMINATION OF CHLORINE IN CHLORIDE, SAMPLE No. +===================================================================== + | | +Weight sample and tube| 16.1721 | 15.9976 + | 15.9976 | 15.7117 + | ------- | ------- + Weight sample | .1745 | .2859 + | | +Weight crucible | | + + precipitate | 14.4496 | 15.6915 + Constant weights | 14.4487 | 15.6915 + | 14.4485 | + | | + Weight crucible | 14.2216 | 15.3196 + Constant weight | 14.2216 | 15.3194 + | | + Weight AgCl | .2269 | .3721 + | | + log Cl | 1.5496 | 1.5496 + log weight AgCl | 9.3558 - 10 | 9.5706 - 10 + log 100 | 2.0000 | 2.0000 + colog AgCl | 7.8438 - 10 | 7.7438 - 10 + colog sample | 0.7583 | 0.5438 + | ------- | ------- + | 1.5075 | 1.5078 + | | + Cl in sample No. | 32.18% | 32.20% + | | +===================================================================== + +Signed + + +STRENGTH OF REAGENTS + +The concentrations given in this table are those suggested for use +in the procedures described in the foregoing pages. It is obvious, +however, that an exact adherence to these quantities is not essential. + + + Approx. Approx. + Grams relation relation + per to normal to molal + liter. solution solution + +Ammonium oxalate, (NH_{4})_{2}C_{2}O_{4}.H_{2}O 40 0.5N 0.25 +Barium chloride, BaCl_{2}.2H_{2}O 25 0.2N 0.1 +Magnesium ammonium chloride (of MgCl_{2}) 71 1.5N 0.75 +Mercuric chloride, HgCl_{2} 45 0.33N 0.66 +Potassium hydroxide, KOH (sp. gr. 1.27) 480 +Potassium thiocyanate, KSCN 5 0.05N 0.55 +Silver nitrate, AgNO_{3} 21 0.125N 0.125 +Sodium hydroxide, NaOH 100 2.5N 2.5 +Sodium carbonate. Na_{2}CO_{3} 159 3N 1.5 +Sodium phosphate, Na_{2}HPO_{4}.12H_{2}O 90 0.5N or 0.75N 0.25 + +Stannous chloride, SnCl_{2}, made by saturating hydrochloric acid (sp. +gr. 1.2) with tin, diluting with an equal volume of water, and adding +a slight excess of acid from time to time. A strip of metallic tin is +kept in the bottle. + +A solution of ammonium molybdate is best prepared as follows: Stir +100 grams of molybdic acid (MoO_{3}) into 400 cc. of cold, distilled +water. Add 80 cc. of concentrated ammonium hydroxide (sp. gr. 0.90). +Filter, and pour the filtrate slowly, with constant stirring, into a +mixture of 400 cc. concentrated nitric acid (sp. gr. 1.42) and 600 +cc. of water. Add to the mixture about 0.05 gram of microcosmic salt. +Filter, after allowing the whole to stand for 24 hours. + +The following data regarding the common acids and aqueous ammonia +are based upon percentages given in the Standard Tables of the +Manufacturing Chemists' Association of the United States [!J.S.C.I.!, +24 (1905), 787-790]. All gravities are taken at 15.5°C. and compared +with water at the same temperature. + +Aqueous ammonia (sp. gr. 0.96) contains 9.91 per cent NH_{3} by +weight, and corresponds to a 5.6 N and 5.6 molal solution. + +Aqueous ammonia (sp. gr. 0.90) contains 28.52 per cent NH_{3} by +weight, and corresponds to a 5.6 N and 5.6 molal solution. + +Hydrochloric acid (sp. gr. 1.12) contains 23.81 per cent HCl by +weight, and corresponds to a 7.3 N and 7.3 molal solution. + +Hydrochloric acid (sp. gr. 1.20) contains 39.80 per cent HCl by +weight, and corresponds to a 13.1 N and 13.1 molal solution. + +Nitric acid (sp. gr. 1.20) contains 32.25 per cent HNO_{3} by weight, +and corresponds to a 6.1 N and 6.1 molal solution: + +Nitric acid (sp. gr. 1.42) contains 69.96 per cent HNO_{3} by weight, +and corresponds to a 15.8 N and 15.8 molal solution. + +Sulphuric acid (sp. gr. 1.8354) contains 93.19 per cent H_{2}SO_{4} by +weight, and corresponds to a 34.8 N or 17.4 molal solution. + +Sulphuric acid (sp. gr. 1.18) contains 24.74 per cent H_{2}SO_{4} by +weight, and corresponds to a 5.9 N or 2.95 molal solution. + +The term !normal! (N), as used above, has the same significance as +in volumetric analyses. The molal solution is assumed to contain one +molecular weight in grams in a liter of solution. + +DENSITIES AND VOLUMES OF WATER AT TEMPERATURES FROM 15-30°C. + +Temperature Density. Volume. +Centigrade. + + 4° 1.000000 1.000000 + 15° 0.999126 1.000874 + 16° 0.998970 1.001031 + 17° 0.998801 1.001200 + 18° 0.998622 1.001380 + 19° 0.998432 1.001571 + 20° 0.998230 1.001773 + 21° 0.998019 1.001985 + 22° 0.997797 1.002208 + 23° 0.997565 1.002441 + 24° 0.997323 1.002685 + 25° 0.997071 1.002938 + 26° 0.996810 1.003201 + 27° 0.996539 1.003473 + 28° 0.996259 1.003755 + 29° 0.995971 1.004046 + 30° 0.995673 1.004346 + +Authority: Landolt, Börnstein, and Meyerhoffer's !Tabellen!, third +edition. + + +CORRECTIONS FOR CHANGE OF TEMPERATURE OF STANDARD SOLUTIONS + +The values below are average values computed from data relating to a +considerable number of solutions. They are sufficiently accurate for +use in chemical analyses, except in the comparatively few cases +where the highest attainable accuracy is demanded in chemical +investigations. The expansion coefficients should then be carefully +determined for the solutions employed. For a compilation of the +existing data, consult Landolt, Börnstein, and Meyerhoffer's +!Tabellen!, third edition. + + Corrections for 1 cc. + Concentration. of solution between + 15° and 35°C. + + Normal .00029 + 0.5 Normal .00025 + 0.1 Normal or more dilute solutions .00020 + +The volume of solution used should be multiplied by the values given, +and that product multiplied by the number of degrees which the +temperature of the solution varies from the standard temperature +selected for the laboratory. The total correction thus found is +subtracted from the observed burette reading if the temperature is +higher than the standard, or added, if it is lower. Corrections are +not usually necessary for variations of temperature of 2°C. or less. + + + + INTERNATIONAL ATOMIC WEIGHTS + +========================================================== + | | | + | 1920 | | 1920 +_________________|_________|___________________|__________ + | | | +Aluminium Al | 27.1 | Molybdenum Mo | 96.0 +Antimony Sb | 120.2 | Neodymium Nd | 144.3 +Argon A | 39.9 | Neon Ne | 20.2 +Arsenic As | 74.96 | Nickel Ni | 58.68 +Barium Ba | 137.37 | Nitrogen N | 14.008 +Bismuth Bi | 208.0 | Osmium Os | 190.9 +Boron B | 11.0 | Oxygen O | 16.00 +Bromine Br | 79.92 | Palladium Pd | 106.7 +Cadmium Cd | 112.40 | Phosphorus P | 31.04 +Caesium Cs | 132.81 | Platinum Pt | 195.2 +Calcium Ca | 40.07 | Potassium K | 39.10 +Carbon C | 12.005 | Praseodymium Pr | 140.9 +Cerium Ce | 140.25 | Radium Ra | 226.0 +Chlorine Cl | 35.46 | Rhodium Rh | 102.9 +Chromium Cr | 52.0 | Rubidium Rb | 85.45 +Cobalt Co | 58.97 | Ruthenium Ru | 101.7 +Columbium Cb | 93.1 | Samarium Sm | 150.4 +Copper Cu | 63.57 | Scandium Sc | 44.1 +Dysprosium Dy | 162.5 | Selenium Se | 79.2 +Erbium Er | 167.7 | Silicon Si | 28.3 +Europium Eu | 152.0 | Silver Ag | 107.88 +Fluorine Fl | 19.0 | Sodium Na | 23.00 +Gadolinium Gd | 157.3 | Strontium Sr | 87.63 +Gallium Ga | 69.9 | Sulphur S | 32.06 +Germanium Ge | 72.5 | Tantalum Ta | 181.5 +Glucinum Gl | 9.1 | Tellurium Te | 127.5 +Gold Au | 197.2 | Terbium Tb | 159.2 +Helium He | 4.00 | Thallium Tl | 204.0 +Hydrogen H | 1.008 | Thorium Th | 232.4 +Indium In | 114.8 | Thulium Tm | 168.5 +Iodine I | 126.92 | Tin Sn | 118.7 +Iridium Ir | 193.1 | Titanium Ti | 48.1 +Iron Fe | 55.84 | Tungsten W | 184.0 +Krypton Kr | 82.92 | Uranium U | 238.2 +Lanthanum La | 139.0 | Vanadium V | 51.0 +Lead Pb | 207.2 | Xenon Xe | 130.2 +Lithium Li | 6.94 | Ytterbium Yb | 173.5 +Lutecium Lu | 175.0 | Yttrium Y | 88.7 +Magnesium Mg | 24.32 | Zinc Zn | 65.37 +Manganese Mn | 54.93 | Zirconium Zr | 90.6 +Mercury Hg | 200.6 | | +========================================================== + + + + +INDEX + +Acidimetry +Acid solutions, normal + standard +Acids, definition of +Acids, weak, action of other acids on + action of salts on +Accuracy demanded +Alkalimetry +Alkali solutions, normal + standard +Alumina, determination of in stibnite +Ammonium nitrate, acid +Analytical chemistry, subdivisions of +Antimony, determination of, in stibnite +Apatite, analysis of +Asbestos filters +Atomic weights, table of + +Balances, essential features of + use and care of +Barium sulphate, determination of sulphur in +Bases, definition of +Bichromate process for iron +Bleaching powder, analysis of +Brass, analysis of +Burette, description of + calibration of + cleaning of + reading of + +Calcium, determination of, in limestone +Calibration, definition of + of burettes + of flasks +Carbon dioxide, determination of, in limestone +Chlorimetry +Chlorine, gravimetric determination of +Chrome iron ore, analysis of +Coin, determination of silver in +Colloidal solution of precipitates +Colorimetric analyses, definition of +Copper, determination of, in brass + determination of in copper ores +Crucibles, use of +Crystalline precipitates + +Densities of water +Deposition potentials +Desiccators +Direct methods +Dissociation, degree of + +Economy of time +Electrolytic dissociation, theory of +Electrolytic separations, principles of +End-point, definition of +Equilibrium, chemical +Evaporation of liquids + +Faraday's law +Feldspar, analysis of +Ferrous ammonium sulphate, analysis of +Filters, folding of + how fitted +Filtrates, testing of +Filtration +Flasks, graduation of +Funnels +Fusions, removal of from crucibles + +General directions for gravimetric analysis + volumetric analysis +Gooch filter +Gravimetric analysis, definition of + +Hydrochloric acid, standardization of +Hydrolysis + +Ignition of precipitates +Indicators, definition of + for acidimetry + preparation of +Indirect methods +Insoluble matter, determination of in limestone +Integrity +Iodimetry +Ions, definition of +Iron, gravimetric determination of + volumetric determination of + +Jones reductor + +Lead, determination of in brass +Limestone, analysis of +Limonite, determination of iron in +Liquids, evaporation of + transfer of +Litmus +Logarithms + +Magnesium, determination of +Mass action, law of +Measuring instruments +Methyl orange +Moisture, determination of in limestone + +Neutralization methods +Normal solutions, acid and alkali + oxidizing agents + reducing agents +Notebooks, sample pages of + +Oxalic acid, determination of strength of +Oxidation processes +Oxidizing power of pyrolusite + +Permanganate process for iron +Phenolphthalein +Phosphoric anhydride, determination of +Pipette, calibration of + description of +Platinum crucibles, care of +Precipitates, colloidal + crystalline + ignition of + separation from filter + washing of +Precipitation +Precipitation methods (volumetric) +Problems +Pyrolusite, oxidizing power of + +Quantitative Analyses, subdivisions of + +Reagents, strength of +Reducing solution, normal +Reductor, Jones +Reversible reactions + +Silica, determination of, in limestone + determination of, in silicates + purification of +Silicic acid, dehydration of +Silver, determination of in coin +Soda ash, alkaline strength of +Sodium chloride, determination of chlorine in +Solubility product +Solution pressure +Solutions, normal + standard +Standardization, definition of +Standard solutions, acidimetry and alkalimetry + chlorimetry + iodimetry + oxidizing and reducing agents + thiocyanate +Starch solutions +Stibnite, determination of antimony in +Stirring rods +Stoichiometry +Strength of reagents +Suction, use of +Sulphur, determination of in ferrous ammonium sulphate + in barium sulphate + +Temperature, corrections for +Testing of washings +Theory of electrolytic dissociation +Thiocyanate process for silver +Titration, definition of +Transfer of liquids + +Volumetric analysis, definition of + general directions + +Wash-bottles +Washed filters +Washing of precipitates +Washings, testing of +Water, ionization of + densities of +Weights, care of + +Zimmermann-Reinhardt method for iron +Zinc, determination of, in brass + + + + + + + +End of the Project Gutenberg EBook of An Introductory Course of Quantitative +Chemical Analysis, by Henry P. 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You may copy it, give it away or +re-use it under the terms of the Project Gutenberg License included +with this eBook or online at www.gutenberg.org + + +Title: An Introductory Course of Quantitative Chemical Analysis + With Explanatory Notes + +Author: Henry P. Talbot + +Release Date: June 30, 2004 [EBook #12787] + +Language: English + +Character set encoding: ASCII + +*** START OF THIS PROJECT GUTENBERG EBOOK QUANTITATIVE CHEMICAL ANALYSIS *** + + + + +Produced by Kevin Handy, Dave Maddock, Josephine Paolucci and the +Online Distributed Proofreading Team. + + + + + +[Transcriber's notes: In the chemical equations, superscripts are +indicated with a ^ and subscripts are indicated with a _. The affected +item is enclosed in curly brackets {}. Examples are H^{+} for hydrogen +ion and H_{2}O for water. Since the underscore is already being used +in this project, italics are designated by an exclamation point +before and after the italicized word or phrase.] + + + +AN INTRODUCTORY COURSE + +OF + +QUANTITATIVE + +CHEMICAL ANALYSIS + +WITH + +EXPLANATORY NOTES + + +BY + +HENRY P. TALBOT + +PROFESSOR OF INORGANIC CHEMISTRY AT THE MASSACHUSETTS INSTITUTE OF +TECHNOLOGY + +SIXTH EDITION, COMPLETELY REWRITTEN + + + + +PREFACE + + +This Introductory Course of Quantitative Analysis has been prepared +to meet the needs of students who are just entering upon the subject, +after a course of qualitative analysis. It is primarily intended to +enable the student to work successfully and intelligently without the +necessity for a larger measure of personal assistance and supervision +than can reasonably be given to each member of a large class. To this +end the directions are given in such detail that there is very little +opportunity for the student to go astray; but the manual is not, the +author believes, on this account less adapted for use with small +classes, where the instructor, by greater personal influence, can +stimulate independent thought on the part of the pupil. + +The method of presentation of the subject is that suggested by +Professor A.A. Noyes' excellent manual of Qualitative Analysis. For +each analysis the procedure is given in considerable detail, and +this is accompanied by explanatory notes, which are believed to be +sufficiently expanded to enable the student to understand fully the +underlying reason for each step prescribed. The use of the book +should, nevertheless, be supplemented by classroom instruction, mainly +of the character of recitations, and the student should be taught to +consult larger works. The general directions are intended to emphasize +those matters upon which the beginner in quantitative analysis must +bestow special care, and to offer helpful suggestions. The student +can hardly be expected to appreciate the force of all the statements +contained in these directions, or, indeed, to retain them all in +the memory after a single reading; but the instructor, by frequent +reference to special paragraphs, as suitable occasion presents itself, +can soon render them familiar to the student. + +The analyses selected for practice are those comprised in the first +course of quantitative analysis at the Massachusetts Institute of +Technology, and have been chosen, after an experience of years, +as affording the best preparation for more advanced work, and as +satisfactory types of gravimetric and volumetric methods. From the +latter point of view, they also seem to furnish the best insight into +quantitative analysis for those students who can devote but a limited +time to the subject, and who may never extend their study beyond the +field covered by this manual. The author has had opportunity to test +the efficiency of the course for use with such students, and has found +the results satisfactory. + +In place of the usual custom of selecting simple salts as material for +preliminary practice, it has been found advantageous to substitute, in +most instances, approximately pure samples of appropriate minerals or +industrial products. The difficulties are not greatly enhanced, while +the student gains in practical experience. + +The analytical procedures described in the following pages have been +selected chiefly with reference to their usefulness in teaching the +subject, and with the purpose of affording as wide a variety of +processes as is practicable within an introductory course of this +character. The scope of the manual precludes any extended attempt to +indicate alternative procedures, except through general references to +larger works on analytical chemistry. The author is indebted to the +standard works for many suggestions for which it is impracticable to +make specific acknowledgment; no considerable credit is claimed by him +for originality of procedure. + +For many years, as a matter of convenience, the classes for which this +text was originally prepared were divided, one part beginning with +gravimetric processes and the other with volumetric analyses. After a +careful review of the experience thus gained the conclusion has been +reached that volumetric analysis offers the better approach to the +subject. Accordingly the arrangement of the present (the sixth) +edition of this manual has been changed to introduce volumetric +procedures first. Teachers who are familiar with earlier editions +will, however, find that the order of presentation of the material +under the various divisions is nearly the same as that previously +followed, and those who may still prefer to begin the course of +instruction with gravimetric processes will, it is believed, be able +to follow that order without difficulty. + +Procedures for the determination of sulphur in insoluble sulphates, +for the determination of copper in copper ores by iodometric methods, +for the determination of iron by permanganate in hydrochloric acid +solutions, and for the standardization of potassium permanganate +solutions using sodium oxalate as a standard, and of thiosulphate +solutions using copper as a standard, have been added. The +determination of silica in silicates decomposable by acids, as a +separate procedure, has been omitted. + +The explanatory notes have been rearranged to bring them into closer +association with the procedures to which they relate. The number of +problems has been considerably increased. + +The author wishes to renew his expressions of appreciation of the +kindly reception accorded the earlier editions of this manual. He has +received helpful suggestions from so many of his colleagues within the +Institute, and friends elsewhere, that his sense of obligation must +be expressed to them collectively. He is under special obligations +to Professor L.F. Hamilton for assistance in the preparation of the +present edition. + +HENRY P. TALBOT + +!Massachusetts Institute of Technology, September, 1921!. + + + + +CONTENTS + + +PART I. INTRODUCTION + +SUBDIVISIONS OF ANALYTICAL CHEMISTRY + +GENERAL DIRECTIONS + Accuracy and Economy of Time; Notebooks; Reagents; Wash-bottles; + Transfer of Liquids + + +PART II. VOLUMETRIC ANALYSIS + +GENERAL DISCUSSION + Subdivisions; The Analytical Balance; Weights; Burettes; + Calibration of Measuring Devices +GENERAL DIRECTIONS + Standard and Normal Solutions + +!I. Neutralization Methods! + +ALKALIMETRY AND ACIDIMETRY + Preparation and Standardization of Solutions; Indicators +STANDARDIZATION OF HYDROCHLORIC ACID +DETERMINATION OF TOTAL ALKALINE STRENGTH OF SODA ASH +DETERMINATION OF ACID STRENGTH OF OXALIC ACID + +!II. Oxidation Processes! + +GENERAL DISCUSSION +BICHROMATE PROCESS FOR THE DETERMINATION OF IRON +DETERMINATION OF IRON IN LIMONITE BY THE BICHROMATE PROCESS +DETERMINATION OF CHROMIUM IN CHROME IRON ORE +PERMANGANATE PROCESS FOR THE DETERMINATION OF IRON +DETERMINATION OF IRON IN LIMONITE BY THE PERMANGANATE PROCESS +DETERMINATION OF IRON IN LIMONITE BY THE ZIMMERMANN-REINHARDT PROCESS +DETERMINATION OF THE OXIDIZING POWER OF PYROLUSITE +IODIMETRY +DETERMINATION OF COPPER IN ORES +DETERMINATION OF ANTIMONY IN STIBNITE +CHLORIMETRY +DETERMINATION OF AVAILABLE CHLORINE IN BLEACHING POWDER + +!III. Precipitation Methods! + +DETERMINATION OF SILVER BY THE THIOCYANATE PROCESS + + +PART III. GRAVIMETRIC ANALYSIS + +GENERAL DIRECTIONS + Precipitation; Funnels and Filters; Filtration and Washing of + Precipitates; Desiccators; Crucibles and their Preparation + for Use; Ignition of Precipitates +DETERMINATION OF CHLORINE IN SODIUM CHLORIDE +DETERMINATION OF IRON AND OF SULPHUR IN FERROUS AMMONIUM SULPHATE +DETERMINATION OF SULPHUR IN BARIUM SULPHATE +DETERMINATION OF PHOSPHORIC ANHYDRIDE IN APATITE +ANALYSIS OF LIMESTONE + Determination of Moisture; Insoluble Matter and Silica; Ferric + Oxide and Alumina; Calcium; Magnesium; Carbon Dioxide +ANALYSIS OF BRASS + Electrolytic Separations; Determination of Lead, Copper, Iron + and Zinc. +DETERMINATION OF SILICA IN SILICATES + +PART IV. STOICHIOMETRY + +SOLUTIONS OF TYPICAL PROBLEMS +PROBLEMS + +APPENDIX + +ELECTROLYTIC DISSOCIATION THEORY +FOLDING OF A FILTER PAPER +SAMPLE NOTEBOOK PAGES +STRENGTH OF REAGENTS +DENSITIES AND VOLUMES OF WATER +CORRECTIONS FOR CHANGE OF TEMPERATURE OF STANDARD SOLUTIONS +ATOMIC WEIGHTS +LOGARITHM TABLES + + + + +QUANTITATIVE CHEMICAL ANALYSIS + + + + +PART I + +INTRODUCTION + +SUBDIVISIONS OF ANALYTICAL CHEMISTRY + + +A complete chemical analysis of a body of unknown composition involves +the recognition of its component parts by the methods of !qualitative +analysis!, and the determination of the proportions in which these +components are present by the processes of !quantitative analysis!. +A preliminary qualitative examination is generally indispensable, if +intelligent and proper provisions are to be made for the separation of +the various constituents under such conditions as will insure accurate +quantitative estimations. + +It is assumed that the operations of qualitative analysis are familiar +to the student, who will find that the reactions made use of in +quantitative processes are frequently the same as those employed in +qualitative analyses with respect to both precipitation and systematic +separation from interfering substances; but it should be noted that +the conditions must now be regulated with greater care, and in such +a manner as to insure the most complete separation possible. For +example, in the qualitative detection of sulphates by precipitation +as barium sulphate from acid solution it is not necessary, in most +instances, to take into account the solubility of the sulphate +in hydrochloric acid, while in the quantitative determination of +sulphates by this reaction this solubility becomes an important +consideration. The operations of qualitative analysis are, therefore, +the more accurate the nearer they are made to conform to quantitative +conditions. + +The methods of quantitative analysis are subdivided, according +to their nature, into those of !gravimetric analysis, volumetric +analysis!, and !colorimetric analysis!. In !gravimetric! processes the +constituent to be determined is sometimes isolated in elementary +form, but more commonly in the form of some compound possessing a +well-established and definite composition, which can be readily and +completely separated, and weighed either directly or after ignition. +From the weight of this substance and its known composition, the +amount of the constituent in question is determined. + +In !volumetric! analysis, instead of the final weighing of a definite +body, a well-defined reaction is caused to take place, wherein the +reagent is added from an apparatus so designed that the volume of the +solution employed to complete the reaction can be accurately measured. +The strength of this solution (and hence its value for the reaction +in question) is accurately known, and the volume employed serves, +therefore, as a measure of the substance acted upon. An example will +make clear the distinction between these two types of analysis. +The percentage of chlorine in a sample of sodium chloride may be +determined by dissolving a weighed amount of the chloride in water +and precipitating the chloride ions as silver chloride, which is +then separated by filtration, ignited, and weighed (a !gravimetric! +process); or the sodium chloride may be dissolved in water, and a +solution of silver nitrate containing an accurately known amount of +the silver salt in each cubic centimeter may be cautiously added from +a measuring device called a burette until precipitation is complete, +when the amount of chlorine may be calculated from the number of cubic +centimeters of the silver nitrate solution involved in the reaction. +This is a !volumetric! process, and is equivalent to weighing without +the use of a balance. + +Volumetric methods are generally more rapid, require less apparatus, +and are frequently capable of greater accuracy than gravimetric +methods. They are particularly useful when many determinations of the +same sort are required. + +In !colorimetric! analyses the substance to be determined is converted +into some compound which imparts to its solutions a distinct color, +the intensity of which must vary in direct proportion to the amount of +the compound in the solution. Such solutions are compared with respect +to depth of color with standard solutions containing known amounts of +the colored compound, or of other similar color-producing substance +which has been found acceptable as a color standard. Colorimetric +methods are, in general, restricted to the determinations of very +small quantities, since only in dilute solutions are accurate +comparisons of color possible. + + + + +GENERAL DIRECTIONS + + +The following paragraphs should be read carefully and thoughtfully. A +prime essential for success as an analyst is attention to details and +the avoidance of all conditions which could destroy, or even lessen, +confidence in the analyses when completed. The suggestions here given +are the outcome of much experience, and their adoption will tend to +insure permanently work of a high grade, while neglect of them will +often lead to disappointment and loss of time. + + +ACCURACY AND ECONOMY OF TIME + +The fundamental conception of quantitative analysis implies a +necessity for all possible care in guarding against loss of material +or the introduction of foreign matter. The laboratory desk, and all +apparatus, should be scrupulously neat and clean at all times. A +sponge should always be ready at hand, and desk and filter-stands +should be kept dry and in good order. Funnels should never be allowed +to drip upon the base of the stand. Glassware should always be +wiped with a clean, lintless towel just before use. All filters and +solutions should be covered to protect them from dust, just as far as +is practicable, and every drop of solution or particle of precipitate +must be regarded as invaluable for the success of the analysis. + +An economical use of laboratory hours is best secured by acquiring +a thorough knowledge of the character of the work to be done before +undertaking it, and then by so arranging the work that no time shall +be wasted during the evaporation of liquids and like time-consuming +operations. To this end the student should read thoughtfully not only +the !entire! procedure, but the explanatory notes as well, before +any step is taken in the analysis. The explanatory notes furnish, in +general, the reasons for particular steps or precautions, but they +also occasionally contain details of manipulation not incorporated, +for various reasons, in the procedure. These notes follow the +procedures at frequent intervals, and the exact points to which they +apply are indicated by references. The student should realize that a +!failure to study the notes will inevitably lead to mistakes, loss of +time, and an inadequate understanding of the subject!. + +All analyses should be made in duplicate, and in general a close +agreement of results should be expected. It should, however, be +remembered that a close concordance of results in "check analyses" is +not conclusive evidence of the accuracy of those results, although the +probability of their accuracy is, of course, considerably enhanced. +The satisfaction in obtaining "check results" in such analyses must +never be allowed to interfere with the critical examination of the +procedure employed, nor must they ever be regarded as in any measure a +substitute for absolute truth and accuracy. + +In this connection it must also be emphasized that only the operator +himself can know the whole history of an analysis, and only he can +know whether his work is worthy of full confidence. No work should be +continued for a moment after such confidence is lost, but should +be resolutely discarded as soon as a cause for distrust is fully +established. The student should, however, determine to put forth his +best efforts in each analysis; it is well not to be too ready to +condone failures and to "begin again," as much time is lost in these +fruitless attempts. Nothing less than !absolute integrity! is or can +be demanded of a quantitative analyst, and any disregard of this +principle, however slight, is as fatal to success as lack of chemical +knowledge or inaptitude in manipulation can possibly be. + + +NOTEBOOKS + +Notebooks should contain, beside the record of observations, +descriptive notes. All records of weights should be placed upon the +right-hand page, while that on the left is reserved for the notes, +calculations of factors, or the amount of reagents required. + +The neat and systematic arrangement of the records of analyses is +of the first importance, and is an evidence of careful work and an +excellent credential. Of two notebooks in which the results may be, +in fact, of equal value as legal evidence, that one which is neatly +arranged will carry with it greater weight. + +All records should be dated, and all observations should be recorded +at once in the notebook. The making of records upon loose paper is a +practice to be deprecated, as is also that of copying original entries +into a second notebook. The student should accustom himself to orderly +entries at the time of observation. Several sample pages of systematic +records are to be found in the Appendix. These are based upon +experience; but other arrangements, if clear and orderly, may prove +equally serviceable. The student is advised to follow the sample pages +until he is in a position to plan out a system of his own. + + +REAGENTS + +The habit of carefully testing reagents, including distilled water, +cannot be too early acquired or too constantly practiced; for, in +spite of all reasonable precautionary measures, inferior chemicals +will occasionally find their way into the stock room, or errors will +be made in filling reagent bottles. The student should remember that +while there may be others who share the responsibility for the purity +of materials in the laboratory of an institution, the responsibility +will later be one which he must individually assume. + +The stoppers of reagent bottles should never be laid upon the desk, +unless upon a clean watch-glass or paper. The neck and mouth of all +such bottles should be kept scrupulously clean, and care taken that no +confusion of stoppers occurs. + + +WASH-BOTTLES + +Wash-bottles for distilled water should be made from flasks of about +750 cc. capacity and be provided with gracefully bent tubes, which +should not be too long. The jet should be connected with the tube +entering the wash-bottle by a short piece of rubber tubing in such +a way as to be flexible, and should deliver a stream about one +millimeter in diameter. The neck of the flask may be wound with cord, +or covered with wash-leather, for greater comfort when hot water is +used. It is well to provide several small wash-bottles for liquids +other than distilled water, which should invariably be clearly +labeled. + + +TRANSFER OF LIQUIDS + +Liquids should never be transferred from one vessel to another, nor to +a filter, without the aid of a stirring rod held firmly against the +side or lip of the vessel. When the vessel is provided with a lip it +is not usually necessary to use other means to prevent the loss of +liquid by running down the side; whenever loss seems imminent a !very +thin! layer of vaseline, applied with the finger to the edge of the +vessel, will prevent it. The stirring rod down which the liquid runs +should never be drawn upward in such a way as to allow the solution to +collect on the under side of the rim or lip of a vessel. + +The number of transfers of liquids from one vessel to another during +an analysis should be as small as possible to avoid the risk of slight +losses. Each vessel must, of course, be completely washed to insure +the transfer of all material; but it should be remembered that this +can be accomplished better by the use of successive small portions of +wash-water (perhaps 5-10 cc.), if each wash-water is allowed to drain +away for a few seconds, than by the addition of large amounts which +unnecessarily increase the volume of the solutions, causing loss of +time in subsequent filtrations or evaporations. + +All stirring rods employed in quantitative analyses should be rounded +at the ends by holding them in the flame of a burner until they begin +to soften. If this is not done, the rods will scratch the inner +surface of beakers, causing them to crack on subsequent heating. + + +EVAPORATION OF LIQUIDS + +The greatest care must be taken to prevent loss of solutions during +processes of evaporation, either from too violent ebullition, from +evaporation to dryness and spattering, or from the evolution of gas +during the heating. In general, evaporation upon the steam bath is to +be preferred to other methods on account of the impossibility of +loss by spattering. If the steam baths are well protected from dust, +solutions should be left without covers during evaporation; but +solutions which are boiled upon the hot plate, or from which gases are +escaping, should invariably be covered. In any case a watch-glass may +be supported above the vessel by means of a glass triangle, or other +similar device, and the danger of loss of material or contamination by +dust thus be avoided. It is obvious that evaporation is promoted by +the use of vessels which admit of the exposure of a broad surface to +the air. + +Liquids which contain suspended matter (precipitates) should always +be cautiously heated, since the presence of the solid matter is +frequently the occasion of violent "bumping," with consequent risk to +apparatus and analysis. + + + + +PART II + +VOLUMETRIC ANALYSIS + + +The processes of volumetric analysis are, in general, simpler than +those of gravimetric analysis and accordingly serve best as an +introduction to the practice of quantitative analysis. For their +execution there are required, first, an accurate balance with which +to weigh the material for analysis; second, graduated instruments in +which to measure the volume of the solutions employed; third, standard +solutions, that is, solutions the value of which is accurately known; +and fourth, indicators, which will furnish accurate evidence of the +point at which the desired reaction is completed. The nature of the +indicators employed will be explained in connection with the different +analyses. + +The process whereby a !standard solution! is brought into reaction is +called !titration!, and the point at which the reaction is exactly +completed is called the !end-point!. The !indicator! should show the +!end-point! of the !titration!. The volume of the standard solution +used then furnishes the measure of the substance to be determined as +truly as if that substance had been separated and weighed. + +The processes of volumetric analysis are easily classified, according +to their character, into: + +I. NEUTRALIZATION METHODS; such, for example, as those of acidimetry +and alkalimetry. + +II. OXIDATION PROCESSES; as exemplified in the determination of +ferrous iron by its oxidation with potassium bichromate. + +III. PRECIPITATION METHODS; of which the titration for silver with +potassium thiocyanate solution is an illustration. + +From a somewhat different standpoint the methods in each case may +be subdivided into (a) DIRECT METHODS, in which the substance to be +measured is directly determined by titration to an end-point with a +standard solution; and (b) INDIRECT METHODS, in which the substance +itself is not measured, but a quantity of reagent is added which is +known to be an excess with respect to a specific reaction, and the +unused excess determined by titration. Examples of the latter class +will be pointed out as they occur in the procedures. + + +MEASURING INSTRUMENTS + + +THE ANALYTICAL BALANCE + +For a complete discussion of the physical principles underlying the +construction and use of balances, and the various methods of weighing, +the student is referred to larger manuals of Quantitative Analysis, +such as those of Fresenius, or Treadwell-Hall, and particularly to +the admirable discussion of this topic in Morse's !Exercises in +Quantitative Chemistry!. + +The statements and rules of procedure which follow are sufficient +for the intelligent use of an analytical balance in connection with +processes prescribed in this introductory manual. It is, however, +imperative that the student should make himself familiar with these +essential features of the balance, and its use. He should fully +realize that the analytical balance is a delicate instrument which +will render excellent service under careful treatment, but such +treatment is an essential condition if its accuracy is to be depended +upon. He should also understand that no set of rules, however +complete, can do away with the necessity for a sense of personal +responsibility, since by carelessness he can render inaccurate not +only his own analyses, but those of all other students using the same +balance. + +Before making any weighings the student should seat himself before a +balance and observe the following details of construction: + +1. The balance case is mounted on three brass legs, which should +preferably rest in glass cups, backed with rubber to prevent slipping. +The front legs are adjustable as to height and are used to level the +balance case; the rear leg is of permanent length. + +2. The front of the case may be raised to give access to the balance. +In some makes doors are provided also at the ends of the balance case. + +3. The balance beam is mounted upon an upright in the center of the +case on the top of which is an inlaid agate plate. To the center of +the beam there is attached a steel or agate knife-edge on which the +beam oscillates when it rests on the agate plate. + +4. The balance beam, extending to the right and left, is graduated +along its upper edge, usually on both sides, and has at its +extremities two agate or steel knife-edges from which are suspended +stirrups. Each of these stirrups has an agate plate which, when the +balance is in action, rests upon the corresponding knife-edge of the +beam. The balance pans are suspended from the stirrups. + +5. A pointer is attached to the center of the beam, and as the beam +oscillates this pointer moves in front of a scale near the base of the +post. + +6. At the base of the post, usually in the rear, is a spirit-level. + +7. Within the upright is a mechanism, controlled by a knob at the +front of the balance case, which is so arranged as to raise the entire +beam slightly above the level at which the knife-edges are in contact +with the agate plates. When the balance is not in use the beam must +be supported by this device since, otherwise, the constant jarring +to which a balance is inevitably subjected, will soon dull the +knife-edges, and lessen the sensitiveness of the balance. + +8. A small weight, or bob, is attached to the pointer (or sometimes +to the beam) by which the center of gravity of the beam and its +attachments may be regulated. The center of gravity must lie very +slightly below the level of the agate plates to secure the desired +sensitiveness of the balance. This is provided for when the balance is +set up and very rarely requires alteration. The student should never +attempt to change this adjustment. + +9. Below the balance pans are two pan-arrests operated by a button +from the front of the case. These arrests exert a very slight upward +pressure upon the pans and minimize the displacement of the beam when +objects or weights are being placed upon the pans. + +10. A movable rod, operated from one end of the balance case, extends +over the balance beam and carries a small wire weight, called a rider. +By means of this rod the rider can be placed upon any desired division +of the scale on the balance beam. Each numbered division on the beam +corresponds to one milligram, and the use of the rider obviates the +placing of very small fractional weights on the balance pan. + +If a new rider is purchased, or an old one replaced, care must be +taken that its weight corresponds to the graduations on the beam of +the balance on which it is to be used. The weight of the rider in +milligrams must be equal to the number of large divisions (5, 6, 10, +or 12) between the central knife-edge and the knife-edge at the end of +the beam. It should be noted that on some balances the last division +bears no number. Each new rider should be tested against a 5 or +10-milligram weight. + +In some of the most recent forms of the balance a chain device +replaces the smaller weights and the use of the rider as just +described. + +Before using a balance, it is always best to test its adjustment. This +is absolutely necessary if the balance is used by several workers; it +is always a wise precaution under any conditions. For this purpose, +brush off the balance pans with a soft camel's hair brush. Then note +(1) whether the balance is level; (2) that the mechanism for raising +and lowering the beams works smoothly; (3) that the pan-arrests touch +the pans when the beam is lowered; and (4) that the needle swings +equal distances on either side of the zero-point when set in motion +without any load on the pans. If the latter condition is not +fulfilled, the balance should be adjusted by means of the adjusting +screw at the end of the beam unless the variation is not more than one +division on the scale; it is often better to make a proper allowance +for this small zero error than to disturb the balance by an attempt at +correction. Unless a student thoroughly understands the construction +of a balance he should never attempt to make adjustments, but should +apply to the instructor in charge. + +The object to be weighed should be placed on the left-hand balance pan +and the weights upon the right-hand pan. Every substance which +could attack the metal of the balance pan should be weighed upon a +watch-glass, and all objects must be dry and cold. A warm body gives +rise to air currents which vitiate the accuracy of the weighing. + +The weights should be applied in the order in which they occur in the +weight-box (not at haphazard), beginning with the largest weight which +is apparently required. After a weight has been placed upon the pan +the beam should be lowered upon its knife-edges, and, if necessary, +the pan-arrests depressed. The movement of the pointer will then +indicate whether the weight applied is too great or too small. When +the weight has been ascertained, by the successive addition of small +weights, to the nearest 5 or 10 milligrams, the weighing is completed +by the use of the rider. The correct weight is that which causes the +pointer to swing an equal number of divisions to the right and left +of the zero-point, when the pointer traverses not less than five +divisions on either side. + +The balance case should always be closed during the final weighing, +while the rider is being used, to protect the pans from the effect of +air currents. + +Before the final determination of an exact weight the beam should +always be lifted from the knife-edges and again lowered into place, +as it frequently happens that the scale pans are, in spite of the +pan-arrests, slightly twisted by the impact of the weights, the beam +being thereby virtually lengthened or shortened. Lifting the beam +restores the proper alignment. + +The beam should never be set in motion by lowering it forcibly upon +the knife-edges, nor by touching the pans, but rather by lifting the +rider (unless the balance be provided with some of the newer devices +for the purpose), and the swing should be arrested only when the +needle approaches zero on the scale, otherwise the knife-edges become +dull. For the same reason the beam should never be left upon its +knife-edges, nor should weights be removed from or placed on the +pans without supporting the beam, except in the case of the small +fractional weights. + +When the process of weighing has been completed, the weight should +be recorded in the notebook by first noting the vacant spaces in the +weight-box, and then checking the weight by again noting the weights +as they are removed from the pan. This practice will often detect and +avoid errors. It is obvious that the weights should always be returned +to their proper places in the box, and be handled only with pincers. + +It should be borne in mind that if the mechanism of a balance is +deranged or if any substance is spilled upon the pans or in the +balance case, the damage should be reported at once. In many instances +serious harm can be averted by prompt action when delay might ruin the +balance. + +Samples for analysis are commonly weighed in small tubes with cork +stoppers. Since the stoppers are likely to change in weight from +the varying amounts of moisture absorbed from the atmosphere, it is +necessary to confirm the recorded weight of a tube which has been +unused for some time before weighing out a new portion of substance +from it. + + +WEIGHTS + +The sets of weights commonly used in analytical chemistry range from +20 grams to 5 milligrams. The weights from 20 grams to 1 gram are +usually of brass, lacquered or gold plated. The fractional weights +are of German silver, gold, platinum or aluminium. The rider is of +platinum or aluminium wire. + +The sets of weights purchased from reputable dealers are usually +sufficiently accurate for analytical work. It is not necessary that +such a set should be strictly exact in comparison with the absolute +standard of weight, provided they are relatively correct among +themselves, and provided the same set of weights is used in all +weighings made during a given analysis. The analyst should assure +himself that the weights in a set previously unfamiliar to him are +relatively correct by a few simple tests. For example, he should make +sure that in his set two weights of the same denomination (i.e., two +10-gram weights, or the two 100-milligram weights) are actually equal +and interchangeable, or that the 500-milligram weight is equal to +the sum of the 200, 100, 100, 50, 20, 20 and 10-milligram weights +combined, and so on. If discrepancies of more than a few tenths of a +milligram (depending upon the total weight involved) are found, the +weights should be returned for correction. The rider should also be +compared with a 5 or 10-milligram weight. + +In an instructional laboratory appreciable errors should be reported +to the instructor in charge for his consideration. + +When the highest accuracy is desired, the weights may be calibrated +and corrections applied. A calibration procedure is described in a +paper by T.W. Richards, !J. Am. Chem. Soc.!, 22, 144, and in many +large text-books. + +Weights are inevitably subject to corrosion if not properly protected +at all times, and are liable to damage unless handled with great care. +It is obvious that anything which alters the weight of a single piece +in an analytical set will introduce an error in every weighing made +in which that piece is used. This source of error is often extremely +obscure and difficult to detect. The only safeguard against such +errors is to be found in scrupulous care in handling and protection +on the part of the analyst, and an equal insistence that if several +analysts use the same set of weights, each shall realize his +responsibility for the work of others as well as his own. + + +BURETTES + +A burette is made from a glass tube which is as uniformly cylindrical +as possible, and of such a bore that the divisions which are etched +upon its surface shall correspond closely to actual contents. + +The tube is contracted at one extremity, and terminates in either a +glass stopcock and delivery-tube, or in such a manner that a piece of +rubber tubing may be firmly attached, connecting a delivery-tube of +glass. The rubber tubing is closed by means of a glass bead. Burettes +of the latter type will be referred to as "plain burettes." + +The graduations are usually numbered in cubic centimeters, and the +latter are subdivided into tenths. + +One burette of each type is desirable for the analytical procedures +which follow. + + +PREPARATION OF A BURETTE FOR USE + +The inner surface of a burette must be thoroughly cleaned in order +that the liquid as drawn out may drain away completely, without +leaving drops upon the sides. This is best accomplished by treating +the inside of the burette with a warm solution of chromic acid in +concentrated sulphuric acid, applied as follows: If the burette is of +the "plain" type, first remove the rubber tip and force the lower +end of the burette into a medium-sized cork stopper. Nearly fill the +burette with the chromic acid solution, close the upper end with a +cork stopper and tip the burette backward and forward in such a way +as to bring the solution into contact with the entire inner surface. +Remove the stopper and pour the solution into a stock bottle to be +kept for further use, and rinse out the burette with water several +times. Unless the water then runs freely from the burette without +leaving drops adhering to the sides, the process must be repeated +(Note 1). + +If the burette has a glass stopcock, this should be removed after +the cleaning and wiped, and also the inside of the ground joint. The +surface of the stopcock should then be smeared with a thin coating of +vaseline and replaced. It should be attached to the burette by means +of a wire, or elastic band, to lessen the danger of breakage. + +Fill the burettes with distilled water, and allow the water to run out +through the stopcock or rubber tip until convinced that no air +bubbles are inclosed (Note 2). Fill the burette to a point above the +zero-point and draw off the water until the meniscus is just below +that mark. It is then ready for calibration. + +[Note 1: The inner surface of the burette must be absolutely clean if +the liquid is to run off freely. Chromic acid in sulphuric acid is +usually found to be the best cleansing agent, but the mixture must be +warm and concentrated. The solution can be prepared by pouring over a +few crystals of potassium bichromate a little water and then adding +concentrated sulphuric acid.] + +[Note 2: It is always necessary to insure the absence of air bubbles +in the tips or stopcocks. The treatment described above will usually +accomplish this, but, in the case of plain burettes it is sometimes +better to allow a little of the liquid to flow out of the tip while it +is bent upwards. Any air which may be entrapped then rises with the +liquid and escapes. + +If air bubbles escape during subsequent calibration or titration, an +error is introduced which vitiates the results.] + + +READING OF A BURETTE + +All liquids when placed in a burette form what is called a meniscus at +their upper surfaces. In the case of liquids such as water or +aqueous solutions this meniscus is concave, and when the liquids are +transparent accurate readings are best obtained by observing the +position on the graduated scales of the lowest point of the meniscus. +This can best be done as follows: Wrap around the burette a piece of +colored paper, the straight, smooth edges of which are held evenly +together with the colored side next to the burette (Note 1). Hold the +paper about two small divisions below the meniscus and raise or lower +the level of the eyes until the edge of the paper at the back of the +burette is just hidden from the eye by that in front (Note 2). Note +the position of the lowest point of the curve of the meniscus, +estimating the tenths of the small divisions, thus reading its +position to hundredths of a cubic centimeter. + +[Note 1: The ends of the colored paper used as an aid to accurate +readings may be fastened together by means of a gummed label. The +paper may then remain on the burette and be ready for immediate use by +sliding it up or down, as required.] + +[Note 2: To obtain an accurate reading the eye must be very nearly on +a level with the meniscus. This is secured by the use of the paper +as described. The student should observe by trial how a reading is +affected when the meniscus is viewed from above or below. + +The eye soon becomes accustomed to estimating the tenths of the +divisions. If the paper is held as directed, two divisions below the +meniscus, one whole division is visible to correct the judgment. It is +not well to attempt to bring the meniscus exactly to a division mark +on the burette. Such readings are usually less accurate than those in +which the tenths of a division are estimated.] + + +CALIBRATION OF GLASS MEASURING DEVICES + +If accuracy of results is to be attained, the correctness of all +measuring instruments must be tested. None of the apparatus offered +for sale can be implicitly relied upon except those more expensive +instruments which are accompanied by a certificate from the !National +Bureau of Standards! at Washington, or other equally authentic source. + +The bore of burettes is subject to accidental variations, and since +the graduations are applied by machine without regard to such +variations of bore, local errors result. + +The process of testing these instruments is called !calibration!. +It is usually accomplished by comparing the actual weight of water +contained in the instrument with its apparent volume. + +There is, unfortunately, no uniform standard of volume which has been +adopted for general use in all laboratories. It has been variously +proposed to consider the volume of 1000 grams of water at 4 deg., 15.5 deg., +16 deg., 17.5 deg., and even 20 deg.C., as a liter for practical purposes, and to +consider the cubic centimeter to be one one-thousandth of that volume. +The true liter is the volume of 1000 grams of water at 4 deg.C.; but +this is obviously a lower temperature than that commonly found in +laboratories, and involves the constant use of corrections if taken as +a laboratory standard. Many laboratories use 15.5 deg.C. (60 deg. F.) as the +working standard. It is plain that any temperature which is deemed +most convenient might be chosen for a particular laboratory, but it +cannot be too strongly emphasized that all measuring instruments, +including burettes, pipettes, and flasks, should be calibrated at that +temperature in order that the contents of each burette, pipette, +etc., shall be comparable with that of every other instrument, thus +permitting general interchange and substitution. For example, it is +obvious that if it is desired to remove exactly 50 cc. from a solution +which has been diluted to 500 cc. in a graduated flask, the 50 cc. +flask or pipette used to remove the fractional portion must give +a correct reading at the same temperature as the 500 cc. flask. +Similarly, a burette used for the titration of the 50 cc. of solution +removed should be calibrated under the same conditions as the +measuring flasks or pipettes employed with it. + +The student should also keep constantly in mind the fact that all +volumetric operations, to be exact, should be carried out as nearly at +a constant temperature as is practicable. The spot selected for +such work should therefore be subject to a minimum of temperature +variations, and should have as nearly the average temperature of +the laboratory as is possible. In all work, whether of calibration, +standardization, or analysis, the temperature of the liquids employed +must be taken into account, and if the temperature of these liquids +varies more than 3 deg. or 4 deg. from the standard temperature chosen for the +laboratory, corrections must be applied for errors due to expansion or +contraction, since volumes of a liquid measured at different times are +comparable only under like conditions as to temperature. Data to be +used for this purpose are given in the Appendix. Neglect of this +correction is frequently an avoidable source of error and annoyance in +otherwise excellent work. The temperature of all solutions at the time +of standardization should be recorded to facilitate the application of +temperature corrections, if such are necessary at any later time. + + +CALIBRATION OF THE BURETTES + +Two burettes, one at least of which should have a glass stopper, are +required throughout the volumetric work. Both burettes should be +calibrated by the student to whom they are assigned. + +PROCEDURE.--Weigh a 50 cc., flat-bottomed flask (preferably a +light-weight flask), which must be dry on the outside, to the nearest +centigram. Record the weight in the notebook. (See Appendix for +suggestions as to records.) Place the flask under the burette and draw +out into it about 10 cc. of water, removing any drop on the tip by +touching it against the inside of the neck of the flask. Do not +attempt to stop exactly at the 10 cc. mark, but do not vary more than +0.1 cc. from it. Note the time, and at the expiration of three minutes +(or longer) read the burette accurately, and record the reading in the +notebook (Note 1). Meanwhile weigh the flask and water to centigrams +and record its weight (Note 2). Draw off the liquid from 10 cc. to +about 20 cc. into the same flask without emptying it; weigh, and at +the expiration of three minutes take the reading, and so on throughout +the length of the burette. When it is completed, refill the burette +and check the first calibration. + +The differences in readings represent the apparent volumes, the +differences in weights the true volumes. For example, if an apparent +volume of 10.05 cc. is found to weigh 10.03 grams, it may be assumed +with sufficient accuracy that the error in that 10 cc. amounts to +-0.02 cc., or -0.002 for each cubic centimeter (Note 3). + +In the calculation of corrections the temperature of the water must be +taken into account, if this varies more than 4 deg.C. from the laboratory +standard temperature, consulting the table of densities of water in +the Appendix. + +From the final data, plot the corrections to be applied so that they +may be easily read for each cubic centimeter throughout the burette. +The total correction at each 10 cc. may also be written on the burette +with a diamond, or etching ink, for permanence of record. + +[Note 1: A small quantity of liquid at first adheres to the side of +even a clean burette. This slowly unites with the main body of liquid, +but requires an appreciable time. Three minutes is a sufficient +interval, but not too long, and should be adopted in every instance +throughout the whole volumetric practice before final readings are +recorded.] + +[Note 2: A comparatively rough balance, capable of weighing to +centigrams, is sufficiently accurate for use in calibrations, for a +moment's reflection will show that it would be useless to weigh the +water with an accuracy greater than that of the readings taken on +the burette. The latter cannot exceed 0.01 cc. in accuracy, which +corresponds to 0.01 gram. + +The student should clearly understand that !all other weighings!, +except those for calibration, should be made accurately to 0.0001 +gram, unless special directions are given to the contrary. + +Corrections for temperature variations of less than 4 deg.C. are +negligible, as they amount to less than 0.01 gram for each 10 grams of +water withdrawn.] + +[Note 3: Should the error discovered in any interval of 10 cc. on the +burette exceed 0.10 cc., it is advisable to weigh small portions (even +1 cc.) to locate the position of the variation of bore in the +tube rather than to distribute the correction uniformly over the +corresponding 10 cc. The latter is the usual course for small +corrections, and it is convenient to calculate the correction +corresponding to each cubic centimeter and to record it in the form +of a table or calibration card, or to plot a curve representing the +values. + +Burettes may also be calibrated by drawing off the liquid in +successive portions through a 5 cc. pipette which has been accurately +calibrated, as a substitute for weighing. If many burettes are to be +tested, this is a more rapid method.] + + +PIPETTES + +A !pipette! may consist of a narrow tube, in the middle of which is +blown a bulb of a capacity a little less than that which it is desired +to measure by the pipette; or it may be a miniature burette, without +the stopcock or rubber tip at the lower extremity. In either case, the +flow of liquid is regulated by the pressure of the finger on the top, +which governs the admission of the air. + +Pipettes are usually already graduated when purchased, but they +require calibration for accurate work. + + +CALIBRATION OF PIPETTES + +PROCEDURE.--Clean the pipette. Draw distilled water into it by sucking +at the upper end until the water is well above the graduation mark. +Quickly place the forefinger over the top of the tube, thus preventing +the entrance of air and holding the water in the pipette. Cautiously +admit a little air by releasing the pressure of the finger, and allow +the level of the water to fall until the lowest point of the meniscus +is level with the graduation. Hold the water at that point by pressure +of the finger and then allow the water to run out from the pipette +into a small tared, or weighed, beaker or flask. After a definite time +interval, usually two to three minutes, touch the end of the pipette +against the side of the beaker or flask to remove any liquid adhering +to it (Note 1). The increase in weight of the flask in grams +represents the volume of the water in cubic centimeters delivered by +the pipette. Calculate the necessary correction. + +[Note 1: A definite interval must be allowed for draining, and a +definite practice adopted with respect to the removal of the liquid +which collects at the end of the tube, if the pipette is designed to +deliver a specific volume when emptied. This liquid may be removed +at the end of a definite interval either by touching the side of the +vessel or by gently blowing out the last drops. Either practice, when +adopted, must be uniformly adhered to.] + + +FLASKS + +!Graduated or measuring flasks! are similar to the ordinary +flat-bottomed flasks, but are provided with long, narrow necks in +order that slight variations in the position of the meniscus with +respect to the graduation shall represent a minimum volume of liquid. +The flasks must be of such a capacity that, when filled with the +specified volume, the liquid rises well into the neck. + + +GRADUATION OF FLASKS + +It is a general custom to purchase the flasks ungraduated and to +graduate them for use under standard conditions selected for the +laboratory in question. They may be graduated for "contents" or +"delivery." When graduated for "contents" they contain a specified +volume when filled to the graduation at a specified temperature, and +require to be washed out in order to remove all of the solution from +the flask. Flasks graduated for "delivery" will deliver the specified +volume of a liquid without rinsing. A flask may, of course, be +graduated for both contents and delivery by placing two graduation +marks upon it. + +PROCEDURE.--To calibrate a flask for !contents!, proceed as follows: +Clean the flask, using a chromic acid solution, and dry it carefully +outside and inside. Tare it accurately; pour water into the flask +until the weight of the latter counterbalances weights on the opposite +pan which equal in grams the number of cubic centimeters of water +which the flask is to contain. Remove any excess of water with the aid +of filter paper (Note 1). Take the flask from the balance, stopper +it, place it in a bath at the desired temperature, usually 15.5 deg. +or 17.5 deg.C., and after an hour mark on the neck with a diamond the +location of the lowest point of the meniscus (Note 2). The mark may +be etched upon the flask by hydrofluoric acid, or by the use of an +etching ink now commonly sold on the market. + +To graduate a flask which is designed to !deliver! a specified volume, +proceed as follows: Clean the flask as usual and wipe all moisture +from the outside. Fill it with distilled water. Pour out the water +and allow the water to drain from the flask for three minutes. +Counterbalance the flask with weights to the nearest centigram. +Add weights corresponding in grams to the volume desired, and add +distilled water to counterbalance these weights. An excess of water, +or water adhering to the neck of the flask, may be removed by means of +a strip of clean filter paper. Stopper the flask, place it in a bath +at 15.5 deg.C. or 17.5 deg.C. and, after an hour, mark the location of the +lowest point of the meniscus, as described above. + +[Note 1: The allowable error in counterbalancing the water and +weights varies with the volume of the flask. It should not exceed one +ten-thousandth of the weight of water.] + +[Note 2: Other methods are employed which involve the use of +calibrated apparatus from which the desired volume of water may be run +into the dry flask and the position of the meniscus marked directly +upon it. For a description of a procedure which is most convenient +when many flasks are to be calibrated, the student is referred to the +!Am. Chem J.!, 16, 479.] + + + + +GENERAL DIRECTIONS FOR VOLUMETRIC ANALYSES + + +It cannot be too strongly emphasized that for the success of analyses +uniformity of practice must prevail throughout all volumetric work +with respect to those factors which can influence the accuracy of the +measurement of liquids. For example, whatever conditions are imposed +during the calibration of a burette, pipette, or flask (notably the +time allowed for draining), must also prevail whenever the flask or +burette is used. + +The student should also be constantly watchful to insure parallel +conditions during both standardization and analyst with respect to the +final volume of liquid in which a titration takes place. The value +of a standard solution is only accurate under the conditions which +prevailed when it was standardized. It is plain that the standard +solutions must be scrupulously protected from concentration or +dilution, after their value has been established. Accordingly, great +care must be taken to thoroughly rinse out all burettes, flasks, etc., +with the solutions which they are to contain, in order to remove all +traces of water or other liquid which could act as a diluent. It is +best to wash out a burette at least three times with small portions of +a solution, allowing each to run out through the tip before assuming +that the burette is in a condition to be filled and used. It is, of +course, possible to dry measuring instruments in a hot closet, but +this is tedious and unnecessary. + +To the same end, all solutions should be kept stoppered and away from +direct sunlight or heat. The bottles should be shaken before use to +collect any liquid which may have distilled from the solution and +condensed on the sides. + +The student is again reminded that variations in temperature of +volumetric solutions must be carefully noted, and care should always +be taken that no source of heat is sufficiently near the solutions to +raise the temperature during use. + +Much time may be saved by estimating the approximate volume of a +standard solution which will be required for a titration (if the data +are obtainable) before beginning the operation. It is then possible to +run in rapidly approximately the required amount, after which it is +only necessary to determine the end-point slowly and with accuracy. +In such cases, however, the knowledge of the approximate amount to be +required should never be allowed to influence the judgment regarding +the actual end-point. + + +STANDARD SOLUTIONS + +The strength or value of a solution for a specific reaction is +determined by a procedure called !Standardization!, in which the +solution is brought into reaction with a definite weight of a +substance of known purity. For example, a definite weight of pure +sodium carbonate may be dissolved in water, and the volume of a +solution of hydrochloric acid necessary to exactly neutralize the +carbonate accurately determined. From these data the strength or value +of the acid is known. It is then a !standard solution!. + + +NORMAL SOLUTIONS + +Standard solutions may be made of a purely empirical strength dictated +solely by convenience of manipulation, or the concentration may +be chosen with reference to a system which is applicable to all +solutions, and based upon chemical equivalents. Such solutions are +called !Normal Solutions! and contain such an amount of the reacting +substance per liter as is equivalent in its chemical action to one +gram of hydrogen, or eight grams of oxygen. Solutions containing one +half, one tenth, or one one-hundredth of this quantity per liter are +called, respectively, half-normal, tenth-normal, or hundredth-normal +solutions. + +Since normal solutions of various reagents are all referred to a +common standard, they have an advantage not possessed by empirical +solutions, namely, that they are exactly equivalent to each other. +Thus, a liter of a normal solution of an acid will exactly neutralize +a liter of a normal alkali solution, and a liter of a normal oxidizing +solution will exactly react with a liter of a normal reducing +solution, and so on. + +Beside the advantage of uniformity, the use of normal solutions +simplifies the calculations of the results of analyses. This is +particularly true if, in connection with the normal solution, the +weight of substance for analysis is chosen with reference to the +atomic or molecular weight of the constituent to be determined. (See +problem 26.) + +The preparation of an !exactly! normal, half-normal, or tenth-normal +solution requires considerable time and care. It is usually carried +out only when a large number of analyses are to be made, or when the +analyst has some other specific purpose in view. It is, however, a +comparatively easy matter to prepare standard solutions which differ +but slightly from the normal or half-normal solution, and these have +the advantage of practical equality; that is, two approximately +half-normal solutions are more convenient to work with than two which +are widely different in strength. It is, however, true that some of +the advantage which pertains to the use of normal solutions as regards +simplicity of calculations is lost when using these approximate +solutions. + +The application of these general statements will be made clear in +connection with the use of normal solutions in the various types of +volumetric processes which follow. + + + + +I. NEUTRALIZATION METHODS + +ALKALIMETRY AND ACIDIMETRY + + + + +GENERAL DISCUSSION + + +!Standard Acid Solutions! may be prepared from either hydrochloric, +sulphuric, or oxalic acid. Hydrochloric acid has the advantage of +forming soluble compounds with the alkaline earths, but its solutions +cannot be boiled without danger of loss of strength; sulphuric acid +solutions may be boiled without loss, but the acid forms insoluble +sulphates with three of the alkaline earths; oxalic acid can be +accurately weighed for the preparation of solutions, and its solutions +may be boiled without loss, but it forms insoluble oxalates with +three of the alkaline earths and cannot be used with certain of the +indicators. + +!Standard Alkali Solutions! may be prepared from sodium or potassium +hydroxide, sodium carbonate, barium hydroxide, or ammonia. Of sodium +and potassium hydroxide, it may be said that they can be used with all +indicators, and their solutions may be boiled, but they absorb carbon +dioxide readily and attack the glass of bottles, thereby losing +strength; sodium carbonate may be weighed directly if its purity is +assured, but the presence of carbonic acid from the carbonate is a +disadvantage with many indicators; barium hydroxide solutions may +be prepared which are entirely free from carbon dioxide, and such +solutions immediately show by precipitation any contamination from +absorption, but the hydroxide is not freely soluble in water; ammonia +does not absorb carbon dioxide as readily as the caustic alkalies, +but its solutions cannot be boiled nor can they be used with all +indicators. The choice of a solution must depend upon the nature of +the work in hand. + +A !normal acid solution! should contain in one liter that quantity of +the reagent which represents 1 gram of hydrogen replaceable by a base. +For example, the normal solution of hydrochloric acid (HCl) should +contain 36.46 grams of gaseous hydrogen chloride, since that amount +furnishes the requisite 1 gram of replaceable hydrogen. On the other +hand, the normal solution of sulphuric acid (H_{2}SO_{4}) should +contain only 49.03 grams, i.e., one half of its molecular weight in +grams. + +A !normal alkali solution! should contain sufficient alkali in a liter +to replace 1 gram of hydrogen in an acid. This quantity is represented +by the molecular weight in grams (40.01) of sodium hydroxide (NaOH), +while a sodium carbonate solution (Na_{2}CO_{3}) should contain but +one half the molecular weight in grams (i.e., 53.0 grams) in a liter +of normal solution. + +Half-normal or tenth-normal solutions are employed in most analyses +(except in the case of the less soluble barium hydroxide). Solutions +of the latter strength yield more accurate results when small +percentages of acid or alkali are to be determined. + + +INDICATORS + +It has already been pointed out that the purpose of an indicator is to +mark (usually by a change of color) the point at which just enough of +the titrating solution has been added to complete the chemical change +which it is intended to bring about. In the neutralization processes +which are employed in the measurement of alkalies (!alkalimetry!) +or acids (!acidimetry!) the end-point of the reaction should, in +principle, be that of complete neutrality. Expressed in terms of ionic +reactions, it should be the point at which the H^{+} ions from an +acid[Note 1] unite with a corresponding number of OH^{-} ions from a +base to form water molecules, as in the equation + +H^{+}, Cl^{-} + Na^{+}, OH^{-} --> Na^{+}, Cl^{-} + (H_{2}O). + +It is not usually possible to realize this condition of exact +neutrality, but it is possible to approach it with sufficient +exactness for analytical purposes, since substances are known which, +in solution, undergo a sharp change of color as soon as even a minute +excess of H^{+} or OH^{-} ions are present. Some, as will be seen, +react sharply in the presence of H^{+} ions, and others with OH^{-} +ions. These substances employed as indicators are usually organic +compounds of complex structure and are closely allied to the dyestuffs +in character. + +[Note 1: A knowledge on the part of the student of the ionic theory +as applied to aqueous solutions of electrolytes is assumed. A brief +outline of the more important applications of the theory is given in +the Appendix.] + + +BEHAVIOR OF ORGANIC INDICATORS + +The indicators in most common use for acid and alkali titrations are +methyl orange, litmus, and phenolphthalein. + +In the following discussion of the principles underlying the behavior +of the indicators as a class, methyl orange and phenolphthalein will +be taken as types. It has just been pointed out that indicators are +bodies of complicated structure. In the case of the two indicators +named, the changes which they undergo have been carefully studied by +Stieglitz (!J. Am. Chem. Soc.!, 25, 1112) and others, and it appears +that the changes involved are of two sorts: First, a rearrangement +of the atoms within the molecule, such as often occurs in organic +compounds; and, second, ionic changes. The intermolecular changes +cannot appropriately be discussed here, as they involve a somewhat +detailed knowledge of the classification and general behavior of +organic compounds; they will, therefore, be merely alluded to, and +only the ionic changes followed. + +Methyl orange is a representative of the group of indicators which, +in aqueous solutions, behave as weak bases. The yellow color which it +imparts to solutions is ascribed to the presence of the undissociated +base. If an acid, such as HCl, is added to such a solution, the acid +reacts with the indicator (neutralizes it) and a salt is formed, as +indicated by the equation: + +(M.o.)^{+}, OH^{-} + H^{+}, Cl^{-} --> (M.o.)^{+} Cl^{-} + (H_{2}O). + +This salt ionizes into (M.o.)^{+} (using this abbreviation for the +positive complex) and Cl^{-}; but simultaneously with this ionization +there appears to be an internal rearrangement of the atoms which +results in the production of a cation which may be designated as +(M'.o'.)^{+}, and it is this which imparts a characteristic red color +to the solution. As these changes occur in the presence of even a +very small excess of acid (that is, of H^{+} ions), it serves as the +desired index of their presence in the solution. If, now, an alkali, +such as NaOH, is added to this reddened solution, the reverse +series of changes takes place. As soon as the free acid present is +neutralized, the slightest excess of sodium hydroxide, acting as +a strong base, sets free the weak, little-dissociated base of the +indicator, and at the moment of its formation it reverts, because of +the rearrangement of the atoms, to the yellow form: + +OH^{-} + (M'.o'.)^{+} --> [M'.o'.OH] --> [M.o.OH]. + +Phenolphthalein, on the other hand, is a very weak, little-dissociated +acid, which is colorless in neutral aqueous solution or in the +presence of free H^{+} ions. When an alkali is added to such a +solution, even in slight excess, the anion of the salt which has +formed from the acid of the indicator undergoes a rearrangement of the +atoms, and a new ion, (Ph')^{+}, is formed, which imparts a pink color +to the solution: + +H^{+}, (Ph)^{-} + Na^{+}, OH^{-} --> (H_{2}O) + Na^{+}, (Ph)^{-} +--> Na^{+}, (Ph')^{-} + +The addition of the slightest excess of an acid to this solution, on +the other hand, occasions first the reversion to the colorless ion and +then the setting free of the undissociated acid of the indicator: + +H^{+}, (Ph')^{-} --> H^{+}, (Ph)^{-} --> (HPh). + +Of the common indicators methyl orange is the most sensitive toward +alkalies and phenolphthalein toward acids; the others occupy +intermediate positions. That methyl orange should be most sensitive +toward alkalies is evident from the following considerations: Methyl +orange is a weak base and, therefore, but little dissociated. It +should, then, be formed in the undissociated condition as soon as even +a slight excess of OH^{-} ions is present in the solution, and there +should be a prompt change from red to yellow as outlined above. On the +other hand, it should be an unsatisfactory indicator for use with weak +acids (acetic acid, for example) because the salts which it forms +with such acids are, like all salts of that type, hydrolyzed to a +considerable extent. This hydrolytic change is illustrated by the +equation: + +(M.o.)^{+} C_{2}H_{3}O_{2}^{-} + H^{+}, OH^{-} --> [M.o.OH] + H^{+}, +C_{2}H_{3}O_{2}^{-}. + +Comparison of this equation with that on page 30 will make it plain +that hydrolysis is just the reverse of neutralization and must, +accordingly, interfere with it. Salts of methyl orange with weak acids +are so far hydrolyzed that the end-point is uncertain, and methyl +orange cannot be used in the titration of such acids, while with +the very weak acids, such as carbonic acid or hydrogen sulphide +(hydrosulphuric acid), the salts formed with methyl orange are, in +effect, completely hydrolyzed (i.e., no neutralization occurs), and +methyl orange is accordingly scarcely affected by these acids. This +explains its usefulness, as referred to later, for the titration of +strong acids, such as hydrochloric acid, even in the presence of +carbonates or sulphides in solution. + +Phenolphthalein, on the other hand, should be, as it is, the best of +the common indicators for use with weak acids. For, since it is +itself a weak acid, it is very little dissociated, and its nearly +undissociated, colorless molecules are promptly formed as soon as +there is any free acid (that is, free H^{+} ions) in the solution. +This indicator cannot, however, be successfully used with weak bases, +even ammonium hydroxide; for, since it is weak acid, the salts +which it forms with weak alkalies are easily hydrolyzed, and as a +consequence of this hydrolysis the change of color is not sharp. +This indicator can, however, be successfully used with strong bases, +because the salts which it forms with such bases are much less +hydrolyzed and because the excess of OH^{-} ions from these bases also +diminishes the hydrolytic action of water. + +This indicator is affected by even so weak an acid as carbonic acid, +which must be removed by boiling the solution before titration. It is +the indicator most generally employed for the titration of organic +acids. + +In general, it may be stated that when a strong acid, such as +hydrochloric, sulphuric or nitric acid, is titrated against a strong +base, such as sodium hydroxide, potassium hydroxide, or barium +hydroxide, any of these indicators may be used, since very little +hydrolysis ensues. It has been noted above that the color change does +not occur exactly at theoretical neutrality, from which it follows +that no two indicators will show exactly the same end-point when acids +and alkalis are brought together. It is plain, therefore, that the +same indicator must be employed for both standardization and analysis, +and that, if this is done, accurate results are obtainable. + +The following table (Note 1) illustrates the variations in the volume +of an alkali solution (tenth-normal sodium hydroxide) required to +produce an alkaline end-point when run into 10 cc. of tenth-normal +sulphuric acid, diluted with 50 cc. of water, using five drops of each +of the different indicator solutions. + +==================================================================== + | | | | + INDICATOR | N/10 | N/10 |COLOR IN ACID|COLOR IN ALKA- + | H_{2}SO_{4}| NaOH |SOLUTION |LINE SOLUTION +_______________|____________|__________|_____________|______________ + | cc. | cc. | cc. | +Methyl orange | 10 | 9.90 | Red | Yellow +Lacmoid | 10 | 10.00 | Red | Blue +Litmus | 10 | 10.00 | Red | Blue +Rosalic acid | 10 | 10.07 | Yellow | Pink +Phenolphthalein| 10 | 10.10 | Colorless | Pink +==================================================================== + +It should also be stated that there are occasionally secondary +changes, other than those outlined above, which depend upon the +temperature and concentration of the solutions in which the indicators +are used. These changes may influence the sensitiveness of an +indicator. It is important, therefore, to take pains to use +approximately the same volume of solution when standardizing that is +likely to be employed in analysis; and when it is necessary, as is +often the case, to titrate the solution at boiling temperature, the +standardization should take place under the same conditions. It is +also obvious that since some acid or alkali is required to react with +the indicator itself, the amount of indicator used should be uniform +and not excessive. Usually a few drops of solution will suffice. + +The foregoing statements with respect to the behavior of indicators +present the subject in its simplest terms. Many substances other than +those named may be employed, and they have been carefully studied to +determine the exact concentration of H^{+} ions at which the color +change of each occurs. It is thus possible to select an indicator +for a particular purpose with considerable accuracy. As data of this +nature do not belong in an introductory manual, reference is made to +the following papers or books in which a more extended treatment of +the subject may be found: + +Washburn, E.W., Principles of Physical Chemistry (McGraw-Hill Book +Co.), (Second Edition, 1921), pp. 380-387. + +Prideaux, E.B.R., The Theory and Use of Indicators (Constable & Co., +Ltd.), (1917). + +Salm, E., A Study of Indicators, !Z. physik. Chem.!, 57 (1906), +471-501. + +Stieglitz, J., Theories of Indicators, !J. Am. Chem. Soc.!, 25 (1903), +1112-1127. + +Noyes, A.A., Quantitative Applications of the Theory of Indicators to +Volumetric Analysis, !J. Am. Chem. Soc.!, 32 (1911), 815-861. + +Bjerrum, N., General Discussion, !Z. Anal. Chem.!, 66 (1917), 13-28 +and 81-95. + +Ostwald, W., Colloid Chemistry of Indicators, !Z. Chem. Ind. +Kolloide!, 10 (1912), 132-146. + +[Note 1: Glaser, !Indikatoren der Acidimetrie und Alkalimetrie!. +Wiesbaden, 1901.] + + +PREPARATION OF INDICATOR SOLUTIONS + +A !methyl orange solution! for use as an indicator is commonly made by +dissolving 0.05-0.1 gram of the compound (also known as Orange III) in +a few cubic centimeters of alcohol and diluting with water to 100 cc. +A good grade of material should be secured. It can be successfully +used for the titration of hydrochloric, nitric, sulphuric, phosphoric, +and sulphurous acids, and is particularly useful in the determination +of bases, such as sodium, potassium, barium, calcium, and ammonium +hydroxides, and even many of the weak organic bases. It can also be +used for the determination, by titration with a standard solution of +a strong acid, of the salts of very weak acids, such as carbonates, +sulphides, arsenites, borates, and silicates, because the weak acids +which are liberated do not affect the indicator, and the reddening of +the solution does not take place until an excess of the strong acid +is added. It should be used in cold, not too dilute, solutions. Its +sensitiveness is lessened in the presence of considerable quantities +of the salts of the alkalies. + +A !phenolphthalein solution! is prepared by dissolving 1 gram of the +pure compound in 100 cc. of 95 per cent alcohol. This indicator is +particularly valuable in the determination of weak acids, especially +organic acids. It cannot be used with weak bases, even ammonia. It +is affected by carbonic acid, which must, therefore, be removed by +boiling when other acids are to be measured. It can be used in hot +solutions. Some care is necessary to keep the volume of the solutions +to be titrated approximately uniform in standardization and in +analysis, and this volume should not in general exceed 125-150 cc. for +the best results, since the compounds formed by the indicator undergo +changes in very dilute solution which lessen its sensitiveness. + +The preparation of a !solution of litmus! which is suitable for use +as an indicator involves the separation from the commercial litmus of +azolithmine, the true coloring principle. Soluble litmus tablets are +often obtainable, but the litmus as commonly supplied to the market is +mixed with calcium carbonate or sulphate and compressed into lumps. To +prepare a solution, these are powdered and treated two or three times +with alcohol, which dissolves out certain constituents which cause a +troublesome intermediate color if not removed. The alcohol is decanted +and drained off, after which the litmus is extracted with hot water +until exhausted. The solution is allowed to settle for some time, the +clear liquid siphoned off, concentrated to one-third its volume and +acetic acid added in slight excess. It is then concentrated to a +sirup, and a large excess of 95 per cent. alcohol added to it. This +precipitates the blue coloring matter, which is filtered off, washed +with alcohol, and finally dissolved in a small volume of water and +diluted until about three drops of the solution added to 50 cc. of +water just produce a distinct color. This solution must be kept in an +unstoppered bottle. It should be protected from dust by a loose plug +of absorbent cotton. If kept in a closed bottle it soon undergoes a +reduction and loses its color, which, however, is often restored by +exposure to the air. + +Litmus can be employed successfully with the strong acids and bases, +and also with ammonium hydroxide, although the salts of the latter +influence the indicator unfavorably if present in considerable +concentration. It may be employed with some of the stronger organic +acids, but the use of phenolphthalein is to be preferred. + + +PREPARATION OF STANDARD SOLUTIONS + +!Hydrochloric Acid and Sodium Hydroxide. Approximate Strength!, 0.5 N + + +PROCEDURE.--Measure out 40 cc. of concentrated, pure hydrochloric +acid into a clean liter bottle, and dilute with distilled water to an +approximate volume of 1000 cc. Shake the solution vigorously for a +full minute to insure uniformity. Be sure that the bottle is not too +full to permit of a thorough mixing, since lack of care at this point +will be the cause of much wasted time (Note 1). + +Weigh out, upon a rough balance, 23 grams of sodium hydroxide (Note +2). Dissolve the hydroxide in water in a beaker. Pour the solution +into a liter bottle and dilute, as above, to approximately 1000 cc. +This bottle should preferably have a rubber stopper, as the hydroxide +solution attacks the glass of the ground joint of a glass stopper, and +may cement the stopper to the bottle. Shake the solution as described +above. + +[Note 1: The original solutions are prepared of a strength greater +than 0.5 N, as they are more readily diluted than strengthened if +later adjustment is desired. + +Too much care cannot be taken to insure perfect uniformity of +solutions before standardization, and thoroughness in this respect +will, as stated, often avoid much waste of time. A solution once +thoroughly mixed remains uniform.] + +[Note 2: Commercial sodium hydroxide is usually impure and always +contains more or less carbonate; an allowance is therefore made for +this impurity by placing the weight taken at 23 grams per liter. If +the hydroxide is known to be pure, a lesser amount (say 21 grams) will +suffice.] + + +COMPARISON OF ACID AND ALKALI SOLUTIONS + +PROCEDURE.--Rinse a previously calibrated burette three times with the +hydrochloric acid solution, using 10 cc. each time, and allowing the +liquid to run out through the tip to displace all water and air +from that part of the burette. Then fill the burette with the acid +solution. Carry out the same procedure with a second burette, using +the sodium hydroxide solution. + +The acid solution may be placed in a plain or in a glass-stoppered +burette as may be more convenient, but the alkaline solution should +never be allowed to remain long in a glass-stoppered burette, as it +tends to cement the stopper to the burette, rendering it useless. It +is preferable to use a plain burette for this solution. + +When the burettes are ready for use and all air bubbles displaced from +the tip (see Note 2, page 17) note the exact position of the liquid in +each, and record the readings in the notebook. (Consult page 188.) Run +out from the burette into a beaker about 40 cc. of the acid and add +two drops of a solution of methyl orange; dilute the acid to about +80 cc. and run out alkali solution from the other burette, stirring +constantly, until the pink has given place to a yellow. Wash down the +sides of the beaker with a little distilled water if the solution has +spattered upon them, return the beaker to the acid burette, and add +acid to restore the pink; continue these alternations until the point +is accurately fixed at which a single drop of either solutions served +to produce a distinct change of color. Select as the final end-point +the appearance of the faintest pink tinge which can be recognized, or +the disappearance of this tinge, leaving a pure yellow; but always +titrate to the same point (Note 1). If the titration has occupied more +than the three minutes required for draining the sides of the burette, +the final reading may be taken immediately and recorded in the +notebook. + +Refill the burettes and repeat the titration. From the records of +calibration already obtained, correct the burette readings and make +corrections for temperature, if necessary. Obtain the ratio of the +sodium hydroxide solution to that of hydrochloric acid by dividing +the number of cubic centimeters of acid used by the number of cubic +centimeters of alkali required for neutralization. The check results +of the two titrations should not vary by more than two parts in one +thousand (Note 2). If the variation in results is greater than this, +refill the burettes and repeat the titration until satisfactory values +are obtained. Use a new page in the notebook for each titration. +Inaccurate values should not be erased or discarded. They should be +retained and marked "correct" or "incorrect," as indicated by the +final outcome of the titrations. This custom should be rigidly +followed in all analytical work. + +[Note 1: The end-point should be chosen exactly at the point of +change; any darker tint is unsatisfactory, since it is impossible to +carry shades of color in the memory and to duplicate them from day to +day.] + +[Note 2: While variation of two parts in one thousand in the values +obtained by an inexperienced analyst is not excessive, the idea must +be carefully avoided that this is a standard for accurate work to be +!generally applied!. In many cases, after experience is gained, the +allowable error is less than this proportion. In a few cases a +larger variation is permissible, but these are rare and can only +be recognized by an experienced analyst. It is essential that the +beginner should acquire at least the degree of accuracy indicated if +he is to become a successful analyst.] + + + + +STANDARDIZATION OF HYDROCHLORIC ACID + +SELECTION AND PREPARATION OF STANDARD + + +The selection of the best substance to be used as a standard for acid +solutions has been the subject of much controversy. The work of Lunge +(!Ztschr. angew. Chem.! (1904), 8, 231), Ferguson (!J. Soc. Chem. +Ind.! (1905), 24, 784), and others, seems to indicate that the best +standard is sodium carbonate prepared from sodium bicarbonate by +heating the latter at temperature between 270 deg. and 300 deg.C. The +bicarbonate is easily prepared in a pure state, and at the +temperatures named the decomposition takes place according to the +equation + +2HNaCO_{3} --> Na_{2}CO_{3} + H_{2}O + CO_{2} + +and without loss of any carbon dioxide from the sodium carbonate, such +as may occur at higher temperatures. The process is carried out as +described below. + +PROCEDURE.--Place in a porcelain crucible about 6 grams (roughly +weighed) of the purest sodium bicarbonate obtainable. Rest the +crucible upon a triangle of iron or copper wire so placed within a +large crucible that there is an open air space of about three eighths +of an inch between them. The larger crucible may be of iron, nickel or +porcelain, as may be most convenient. Insert the bulb of a thermometer +reading to 350 deg.C. in the bicarbonate, supporting it with a clamp so +that the bulb does not rest on the bottom of the crucible. Heat +the outside crucible, using a rather small flame, and raise the +temperature of the bicarbonate fairly rapidly to 270 deg.C. Then regulate +the heat in such a way that the temperature rises !slowly! to 300 deg.C. +in the course of a half-hour. The bicarbonate should be frequently +stirred with a clean, dry, glass rod, and after stirring, should be +heaped up around the bulb of the thermometer in such a way as to cover +it. This will require attention during most of the heating, as the +temperature should not be permitted to rise above 310 deg.C. for any +length of time. At the end of the half-hour remove the thermometer and +transfer the porcelain crucible, which now contains sodium carbonate, +to a desiccator. When it is cold, transfer the carbonate to a +stoppered weighing tube or weighing-bottle. + + +STANDARDIZATION + +PROCEDURE.--Clean carefully the outside of a weighing-tube, or +weighing-bottle, containing the pure sodium carbonate, taking care +to handle it as little as possible after wiping. Weigh the tube +accurately to 0.0001 gram, and record the weight in the notebook. Hold +the tube over the top of a beaker (200-300 cc.) and cautiously remove +the stopper, making sure that no particles fall from it or from the +tube elsewhere than in the beaker. Pour out from the tube a portion +of the carbonate, replace the stopper and determine approximately how +much has been removed. Continue this procedure until 1.00 to 1.10 +grams has been taken from the tube. Then weigh the tube accurately +and record the weight under the first weight in the notebook. +The difference in the two weights is the weight of the carbonate +transferred to the beaker. Proceed in the same way to transfer a +second portion of the carbonate from the tube to another beaker of +about the same size as the first. The beakers should be labeled and +plainly marked to correspond with the entries in the notebook. + +Pour over the carbonate in each beaker about 80 cc. of water, stir +until solution is complete, and add two drops of methyl orange +solution. Fill the burettes with the standard acid and alkali +solutions, noting the initial readings of the burettes and temperature +of the solutions. Run in acid from the burette, stirring and avoiding +loss by effervescence, until the solution has become pink. Wash down +the sides of the beaker with a !little! water from a wash-bottle, and +then run in alkali from the other burette until the pink is replaced +by yellow; then finish the titration as described on page 37. Note the +readings of the burettes after the proper interval, and record them in +the notebook. Repeat the procedure, using the second portion of sodium +carbonate. Apply the necessary calibration corrections to the volumes +of the solutions used, and correct for temperature if necessary. + +From the data obtained, calculate the volume of the hydrochloric +acid solution which is equivalent to the volume of sodium hydroxide +solution used in this titration. Subtract this volume from the volume +of hydrochloric acid. The difference represents the volume of acid +used to react with the sodium carbonate. Divide the weight of sodium +carbonate by this volume in cubic centimeters, thus obtaining the +weight of sodium carbonate equivalent to each cubic centimeter of the +acid. + +From this weight it is possible to calculate the corresponding weight +of HCl in each cubic centimeter of the acid, and in turn the relation +of the acid to the normal. + +If, however, it is recalled that normal solutions are equivalent to +each other, it will be seen that the same result may be more readily +reached by dividing the weight in grams of sodium carbonate per cubic +centimeter just found by titration by the weight which would be +contained in the same volume of a normal solution of sodium carbonate. +A normal solution of sodium carbonate contains 53.0 grams per liter, +or 0.0530 gram per cc. (see page 29). The relation of the acid +solution to the normal is, therefore, calculated by dividing the +weight of the carbonate to which each cubic centimeter of the acid is +equivalent by 0.0530. The standardization must be repeated until the +values obtained agree within, at most, two parts in one thousand. + +When the standard of the acid solution has been determined, calculate, +from the known ratio of the two solutions, the relation of the sodium +hydroxide solution to a normal solution (Notes 1 and 2). + +[Note 1: In the foregoing procedure the acid solution is standardized +and the alkali solution referred to this standard by calculation. It +is equally possible, if preferred, to standardize the alkali solution. +The standards in a common use for this purpose are purified +oxalic acid (H_{2}C_{2}O_{4}.2H_{2}O), potassium acid oxalate +(KHC_{2}O_{4}.H_{2}O or KHC_{2}O_{4}), potassium tetroxalate +(KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O), or potassium acid tartrate +(KHC_{4}O_{6}), with the use of a suitable indicator. The oxalic acid +and the oxalates should be specially prepared to insure purity, +the main difficulty lying in the preservation of the water of +crystallization. + +It should be noted that the acid oxalate and the acid tartrate each +contain one hydrogen atom replaceable by a base, while the tetroxalate +contains three such atoms and the oxalic acid two. Each of the two +salts first named behave, therefore, as monobasic acids, and the +tetroxalate as a tribasic acid.] + +[Note 2: It is also possible to standardize a hydrochloric acid +solution by precipitating the chloride ions as silver chloride and +weighing the precipitate, as prescribed under the analysis of sodium +chloride to be described later. Sulphuric acid solutions may be +standardized by precipitation of the sulphate ions as barium sulphate +and weighing the ignited precipitate, but the results are not above +criticism on account of the difficulty in obtaining large precipitates +of barium sulphate which are uncontaminated by inclosures or are not +reduced on ignition.] + + + + +DETERMINATION OF THE TOTAL ALKALINE STRENGTH OF SODA ASH + + +Soda ash is crude sodium carbonate. If made by the ammonia process it +may contain also sodium chloride, sulphate, and hydroxide; when made +by the Le Blanc process it may contain sodium sulphide, silicate, and +aluminate, and other impurities. Some of these, notably the hydroxide, +combine with acids and contribute to the total alkaline strength, +but it is customary to calculate this strength in terms of sodium +carbonate; i.e., as though no other alkali were present. + +PROCEDURE.--In order to secure a sample which shall represent the +average value of the ash, it is well to take at least 5 grams. As this +is too large a quantity for convenient titration, an aliquot portion +of the solution is measured off, representing one fifth of the entire +quantity. This is accomplished as follows: Weigh out on an analytical +balance two samples of soda ash of about 5 grams each into beakers +of about 500 cc. capacity. (The weighings need be made to centigrams +only.) Dissolve the ash in 75 cc. of water, warming gently, and filter +off the insoluble residue; wash the filter by filling it at least +three times with distilled water, and allowing it to drain, adding the +washings to the main filtrate. Cool the filtrate to approximately the +standard temperature of the laboratory, and transfer it to a 250 cc. +measuring flask, washing out the beaker thoroughly. Add distilled +water of laboratory temperature until the lowest point of the meniscus +is level with the graduation on the neck of the flask and remove any +drops of water that may be on the neck above the graduation by means +of a strip of filter paper; make the solution thoroughly uniform by +pouring it out into a dry beaker and back into the flask several +times. Measure off 50 cc. of the solution in a measuring flask, or +pipette, either of which before use should, unless they are dry on the +inside, be rinsed out with at least two small portions of the soda ash +solution to displace any water. + +If a flask is used, fill it to the graduation with the soda ash +solution and remove any liquid from the neck above the graduation with +filter paper. Empty it into a beaker, and wash out the small flask, +unless it is graduated for !delivery!, using small quantities of +water, which are added to the liquid in the beaker. A second 50 cc. +portion from the main solution should be measured off into a second +beaker. Dilute the solutions in each beaker to 100 cc., add two drops +of a solution of methyl orange (Note 1) and titrate for the alkali +with the standard hydrochloric acid solution, using the alkali +solution to complete the titration as already prescribed. + +From the volumes of acid and alkali employed, corrected for burette +errors and temperature changes, and the data derived from the +standardization, calculate the percentage of alkali present, assuming +it all to be present as sodium carbonate (Note 2). + +[Note 1: The hydrochloric acid sets free carbonic acid which is +unstable and breaks down into water and carbon dioxide, most of which +escapes from the solution. Carbonic acid is a weak acid and, as such, +does not yield a sufficient concentration of H^{+} ions to cause the +indicator to change to a pink (see page 32). + +The chemical changes involved may be summarized as follows: + +2H^{+}, 2Cl^{-} + 2Na^{+}, CO_{3}^{--} --> 2Na^{+}, 2Cl^{-} + +[H_{2}CO_{3}] --> H_{2}O + CO_{2}] + +[Note 2: A determination of the alkali present as hydroxide in soda +ash may be determined by precipitating the carbonate by the addition +of barium chloride, removing the barium carbonate by filtration, and +titrating the alkali in the filtrate. + +The caustic alkali may also be determined by first using +phenolphthalein as an indicator, which will show by its change from +pink to colorless the point at which the caustic alkali has been +neutralized and the carbonate has been converted to bicarbonate, and +then adding methyl orange and completing the titration. The amount of +acid necessary to change the methyl orange to pink is a measure of one +half of the carbonate present. The results of the double titration +furnish the data necessary for the determination of the caustic alkali +and of the carbonate in the sample.] + + + + +DETERMINATION OF THE ACID STRENGTH OF OXALIC ACID + + +PROCEDURE.--Weigh out two portions of the acid of about 1 gram +each. Dissolve these in 50 cc. of warm water. Add two drops of +phenolphthalein solution, and run in alkali from the burette until the +solution is pink; add acid from the other burette until the pink is +just destroyed, and then add 0.3 cc. (not more) in excess. Heat the +solution to boiling for three minutes. If the pink returns during the +boiling, discharge it with acid and again add 0.3 cc. in excess and +repeat the boiling (Note 1). If the color does not then reappear, add +alkali until it does, and a !drop or two! of acid in excess and boil +again for one minute (Note 2). If no color reappears during this time, +complete the titration in the hot solution. The end-point should be +the faintest visible shade of color (or its disappearance), as the +same difficulty would exist here as with methyl orange if an attempt +were made to match shades of pink. + +From the corrected volume of alkali required to react with the +oxalic acid, calculate the percentage of the crystallized acid +(H_{2}C_{2}O_{4}.2H_{2}O) in the sample (Note 3). + +[Note 1: All commercial caustic soda such as that from which the +standard solution was made contains some sodium carbonate. This reacts +with the oxalic acid, setting free carbonic acid, which, in turn, +forms sodium bicarbonate with the remaining carbonate: + +H_{2}CO_{3} + Na_{2}CO_{3} --> 2HNaCO_{3}. + +This compound does not hydrolyze sufficiently to furnish enough OH^{-} +ions to cause phenolphthalein to remain pink; hence, the color of +the indicator is discharged in cold solutions at the point at which +bicarbonate is formed. If, however, the solution is heated to boiling, +the bicarbonate loses carbon dioxide and water, and reverts to sodium +carbonate, which causes the indicator to become again pink: + +2HNaCO_{3} --> H_{2}O + CO_{2} + Na_{2}CO_{3}. + +By adding successive portions of hydrochloric acid and boiling, the +carbonate is ultimately all brought into reaction. + +The student should make sure that the difference in behavior of the +two indicators, methyl orange and phenolphthalein, is understood.] + +[Note 2: Hydrochloric acid is volatilized from aqueous solutions, +except such as are very dilute. If the directions in the procedure +are strictly followed, no loss of acid need be feared, but the amount +added in excess should not be greater than 0.3-0.4 cc.] + +[Note 3: Attention has already been called to the fact that the color +changes in the different indicators occur at varying concentrations +of H^{+} or OH^{-} ions. They do not indicate exact theoretical +neutrality, but a particular indicator always shows its color change +at a particular concentration of H^{+} or OH^{-} ions. The results +of titration with a given indicator are, therefore, comparable. As a +matter of fact, a small error is involved in the procedure as outlined +above. The comparison of the acid and alkali solutions was made, using +methyl orange as an indicator, while the titration of the oxalic acid +is made with the use of phenolphthalein. For our present purposes the +small error may be neglected but, if time permits, the student is +recommended to standardize the alkali solution against one of the +substances named in Note 1, page 41, and also to ascertain +the comparative value of the acid and alkali solutions, using +phenolphthalein as indicator throughout, and conducting the titrations +as described above. This will insure complete accuracy.] + + + + +II. OXIDATION PROCESSES + +GENERAL DISCUSSION + + +In the oxidation processes of volumetric analysis standard solutions +of oxidizing agents and of reducing agents take the place of the acid +and alkali solutions of the neutralization processes already studied. +Just as an acid solution was the principal reagent in alkalimetry, and +the alkali solution used only to make certain of the end-point, the +solution of the oxidizing agent is the principal reagent for the +titration of substances exerting a reducing action. It is, in general, +true that oxidizable substances are determined by !direct! titration, +while oxidizing substances are determined by !indirect! titration. + +The important oxidizing agents employed in volumetric solutions are +potassium bichromate, potassium permangenate, potassium ferricyanide, +iodine, ferric chloride, and sodium hypochlorite. + +The important reducing agents which are used in the form of standard +solutions are ferrous sulphate (or ferrous ammonium sulphate), oxalic +acid, sodium thiosulphate, stannous chloride, arsenious acid, and +potassium cyanide. Other reducing agents, as sulphurous acid, +sulphureted hydrogen, and zinc (nascent hydrogen), may take part in +the processes, but not as standard solutions. + +The most important combinations among the foregoing are: Potassium +bichromate and ferrous salts; potassium permanganate and ferrous +salts; potassium permanganate and oxalic acid, or its derivatives; +iodine and sodium thiosulphate; hypochlorites and arsenious acid. + + + + +BICHROMATE PROCESS FOR THE DETERMINATION OF IRON + + +Ferrous salts may be promptly and completely oxidized to ferric salts, +even in cold solution, by the addition of potassium bichromate, +provided sufficient acid is present to hold in solution the ferric and +chromic compounds which are formed. + +The acid may be either hydrochloric or sulphuric, but the former is +usually preferred, since it is by far the best solvent for iron and +its compounds. The reaction in the presence of hydrochloric acid is as +follows: + +6FeCl_{2} + K_{2}Cr_{2}O_{7} + 14HCl --> 6FeCl_{3} + 2CrCl_{3} + 2KCl ++ 7H_{2}O. + + +NORMAL SOLUTIONS OF OXIDIZING OR REDUCING AGENTS + +It will be recalled that the system of normal solutions is based upon +the equivalence of the reagents which they contain to 8 grams of +oxygen or 1 gram of hydrogen. A normal solution of an oxidizing agent +should, therefore, contain that amount per liter which is equivalent +in oxidizing power to 8 grams of oxygen; a normal reducing solution +must be equivalent in reducing power to 1 gram of hydrogen. In order +to determine what the amount per liter will be it is necessary to know +how the reagents enter into reaction. The two solutions to be employed +in the process under consideration are those of potassium bichromate +and ferrous sulphate. The reaction between them, in the presence of an +excess of sulphuric acid, may be expressed as follows: + +6FeSO_{4} + K_{2}Cr_{2}O_{7} + 7H_{2}SO_{4} --> 3Fe_{2}(SO_{4})_{3} + +K_{2}SO_{4} + Cr_{2}(SO_{4})_{3} + 7H_{2}O. + +If the compounds of iron and chromium, with which alone we are now +concerned, be written in such a way as to show the oxides of these +elements in each, they would appear as follows: On the left-hand side +of the equation 6(FeO.SO_{3}) and K_{2}O.2CrO_{3}; on the right-hand +side, 3(Fe_{2}O_{3}.3SO_{3}) and Cr_{2}O_{3}.3SO_{3}. A careful +inspection shows that there are three less oxygen atoms associated +with chromium atoms on the right-hand side of the equation than on the +left-hand, but there are three more oxygen atoms associated with iron +atoms on the right than on the left. In other words, a molecule of +potassium bichromate has given up three atoms of oxygen for oxidation +purposes; i.e., a molecular weight in grams of the bichromate (294.2) +will furnish 3 X 16 or 48 grams of oxygen for oxidation purposes. +As this 48 grams is six times 8 grams, the basis of the system, the +normal solution of potassium bichromate should contain per liter one +sixth of 294.2 grams or 49.03 grams. + +A further inspection of the dissected compounds above shows that six +molecules of FeO.SO_{3} were required to react with the three atoms of +oxygen from the bichromate. From the two equations + +3H_{2} + 3O --> 3H_{2}O +6(FeO.SO_{3}) + 3O --> 3(Fe_{2}O_{3}.3SO_{3}) + +it is plain that one molecule of ferrous sulphate is equivalent to one +atom of hydrogen in reducing power; therefore one molecular weight in +grams of ferrous sulphate (151.9) is equivalent to 1 gram of +hydrogen. Since the ferrous sulphate crystalline form has the formula +FeSO_{4}.7H_{2}O, a normal reducing solution of this crystalline salt +should contain 277.9 grams per liter. + + +PREPARATION OF SOLUTIONS + +!Approximate Strength 0.1 N! + +It is possible to purify commercial potassium bichromate by +recrystallization from hot water. It must then be dried and cautiously +heated to fusion to expel the last traces of moisture, but not +sufficiently high to expel any oxygen. The pure salt thus prepared, +may be weighed out directly, dissolved, and the solution diluted in a +graduated flask to a definite volume. In this case no standardization +is made, as the normal value can be calculated directly. It is, +however, more generally customary to standardize a solution of +the commercial salt by comparison with some substance of definite +composition, as described below. + +PROCEDURE.--Pulverize about 5 grams of potassium bichromate of good +quality. Dissolve the bichromate in distilled water, transfer the +solution to a liter bottle, and dilute to approximately 1000 cc. Shake +thoroughly until the solution is uniform. + +To prepare the solution of the reducing agent, pulverize about 28 +grams of ferrous sulphate (FeSO_{4}.7H_{2}O) or about 40 grams of +ferrous ammonium sulphate (FeSO_{4}.(NH_{4})_{2}SO_{4}.6H_{2}O) and +dissolve in distilled water containing 5 cc. of concentrated sulphuric +acid. Transfer the solution to a liter bottle, add 5 cc. concentrated +sulphuric acid, make up to about 1000 cc. and shake vigorously to +insure uniformity. + + +INDICATOR SOLUTION + +No indicator is known which, like methyl orange, can be used within +the solution, to show when the oxidation process is complete. Instead, +an outside indicator solution is employed to which drops of the +titrated solution are transferred for testing. The reagent used is +potassium ferricyanide, which produces a blue precipitate (or color) +with ferrous compounds as long as there are unoxidized ferrous ions in +the titrated solution. Drops of the indicator solution are placed upon +a glazed porcelain tile, or upon white cardboard which has been coated +with paraffin to render it waterproof, and drops of the titrated +solution are transferred to the indicator on the end of a stirring +rod. When the oxidation is nearly completed only very small amounts +of the ferrous compounds remain unoxidized and the reaction with the +indicator is no longer instantaneous. It is necessary to allow a brief +time to elapse before determining that no blue color is formed. Thirty +seconds is a sufficient interval, and should be adopted throughout the +analytical procedure. If left too long, the combined effect of light +and dust from the air will cause a reduction of the ferric compounds +already formed and a resultant blue will appear which misleads the +observer with respect to the true end-point. + +The indicator solution must be highly diluted, otherwise its own color +interferes with accurate observation. Prepare a fresh solution, as +needed each day, by dissolving a crystal of potassium ferricyanide +about the size of a pin's head in 25 cc. of distilled water. The salt +should be carefully tested with ferric chloride for the presence of +ferrocyanides, which give a blue color with ferric salts. + +In case of need, the ferricyanide can be purified by adding to its +solution a little bromine water and recrystallizing the compound. + + +COMPARISON OF OXIDIZING AND REDUCING SOLUTIONS + +PROCEDURE.--Fill one burette with each of the solutions, observing +the general procedure with respect to cleaning and rinsing already +prescribed. The bichromate solution is preferably to be placed in a +glass-stoppered burette. + +Run out from a burette into a beaker of about 300 cc. capacity nearly +40 cc. of the ferrous solution, add 15 cc. of dilute hydrochloric acid +(sp. gr. 1.12) and 150 cc. of water and run in the bichromate +solution from another burette. Since both solutions are approximately +tenth-normal, 35 cc. of the bichromate solution may be added without +testing. Test at that point by removing a very small drop of the +iron solution on the end of a stirring rod, mixing it with a drop of +indicator on the tile (Note 1). If a blue precipitate appears at once, +0.5 cc. of the bichromate solution may be added before testing again. +The stirring rod which has touched the indicator should be dipped in +distilled water before returning it to the iron solution. As soon as +the blue appears to be less intense, add the bichromate solution in +small portions, finally a single drop at a time, until the point is +reached at which no blue color appears after the lapse of thirty +seconds from the time of mixing solution and indicator. At the close +of the titration a large drop of the iron solution should be taken for +the test. To determine the end-point beyond any question, as soon as +the thirty seconds have elapsed remove another drop of the solution +of the same size as that last taken and mix it with the indicator, +placing it beside the last previous test. If this last previous test +shows a blue tint in comparison with the fresh mixture, the end-point +has not been reached; if no difference can be noted the reaction is +complete. Should the end-point be overstepped, a little more of the +ferrous solution may be added and the end-point definitely fixed. + +From the volumes of the solutions used, after applying corrections for +burette readings, and, if need be, for the temperature of solutions, +calculate the value of the ferrous solution in terms of the oxidizing +solution. + +[Note 1: The accuracy of the work may be much impaired by the removal +of unnecessarily large quantities of solution for the tests. At the +beginning of the titration, while much ferrous iron is still present, +the end of the stirring rod need only be moist with the solution; but +at the close of the titration drops of considerable size may properly +be taken for the final tests. The stirring rod should be washed to +prevent transfer of indicator to the main solution. This cautious +removal of solution does not seriously affect the accuracy of the +determination, as it will be noted that the volume of the titrated +solution is about 200 cc. and the portions removed are very +small. Moreover, if the procedure is followed as prescribed, the +concentration of unoxidized iron decreases very rapidly as the +titration is carried out so that when the final tests are made, though +large drops may be taken, the amount of ferrous iron is not sufficient +to produce any appreciable error in results. + +If the end-point is determined as prescribed, it can be as accurately +fixed as that of other methods; and if a ferrous solution is at +hand, the titration need consume hardly more time than that of the +permanganate process to be described later on.] + + +STANDARDIZATION OF POTASSIUM BICHROMATE SOLUTIONS + +!Selection of a Standard! + +A substance which will serve satisfactorily as a standard for +oxidizing solutions must possess certain specific properties: It must +be of accurately known composition and definite in its behavior as a +reducing agent, and it must be permanent against oxidation in the air, +at least for considerable periods. Such standards may take the form of +pure crystalline salts, such as ferrous ammonium sulphate, or may be +in the form of iron wire or an iron ore of known iron content. It is +not necessary that the standard should be of 100 per cent purity, +provided the content of the active reducing agent is known and no +interfering substances are present. + +The two substances most commonly used as standards for a bichromate +solution are ferrous ammonium sulphate and iron wire. A standard wire +is to be purchased in the market which answers the purpose well, and +its iron content may be determined for each lot purchased by a number +of gravimetric determinations. It may best be preserved in jars +containing calcium chloride, but this must not be allowed to come +into contact with the wire. It should, however, even then be examined +carefully for rust before use. + +If pure ferrous ammonium sulphate is used as the standard, clear +crystals only should be selected. It is perhaps even better to +determine by gravimetric methods once for all the iron content of a +large commercial sample which has been ground and well mixed. This +salt is permanent over long periods if kept in stoppered containers. + + +STANDARDIZATION + +PROCEDURE.--Weigh out two portions of iron wire of about 0.24-0.26 +gram each, examining the wire carefully for rust. It should be handled +and wiped with filter paper (not touched by the fingers), should +be weighed on a watch-glass, and be bent in such a way as not to +interfere with the movement of the balance. + +Place 30 cc. of hydrochloric acid (sp. gr. 1.12) in each of two 300 +cc. Erlenmeyer flasks, cover them with watch-glasses, and bring the +acid just to boiling. Remove them from the flame and drop in the +portions of wire, taking great care to avoid loss of liquid during +solution. Boil for two or three minutes, keeping the flasks covered +(Note 1), then wash the sides of the flasks and the watch-glass with +a little water and add stannous chloride solution to the hot liquid +!from a dropper! until the solution is colorless, but avoid more than +a drop or two in excess (Note 2). Dilute with 150 cc. of water and +cool !completely!. When cold, add rapidly about 30 cc. of mercuric +chloride solution. Allow the solutions to stand about three minutes +and then titrate without further delay (Note 3), add about 35 cc. of +the standard solution at once and finish the titration as prescribed +above, making use of the ferrous solution if the end-point should be +passed. + +From the corrected volumes of the bichromate solution required to +oxidize the iron actually know to be present in the wire, calculate +the relation of the standard solution to the normal. + +Repeat the standardization until the results are concordant within at +least two parts in one thousand. + + +[Note 1: The hydrochloric acid is added to the ferrous solution +to insure the presence of at least sufficient free acid for the +titration, as required by the equation on page 48. + +The solution of the wire in hot acid and the short boiling insure the +removal of compounds of hydrogen and carbon which are formed from the +small amount of carbon in the iron. These might be acted upon by the +bichromate if not expelled.] + +[Note 2: It is plain that all the iron must be reduced to the ferrous +condition before the titration begins, as some oxidation may have +occurred from the oxygen of the air during solution. It is also +evident that any excess of the agent used to reduce the iron must be +removed; otherwise it will react with the bichromate added later. + +The reagents available for the reduction of iron are stannous +chloride, sulphurous acid, sulphureted hydrogen, and zinc; of these +stannous chloride acts most readily, the completion of the reaction +is most easily noted, and the excess of the reagent is most readily +removed. The latter object is accomplished by oxidation to stannic +chloride by means of mercuric chloride added in excess, as the +mercuric salts have no effect upon ferrous iron or the bichromate. The +reactions involved are: + +2FeCl_{3} + SnCl_{2} --> 2FeCl_{2} + SnCl_{4} +SnCl_{2} + 2HgCl_{2} --> SnCl_{4} + 2HgCl + +The mercurous chloride is precipitated. + +It is essential that the solution should be cold and that the stannous +chloride should not be present in great excess, otherwise a secondary +reaction takes place, resulting in the reduction of the mercurous +chloride to metallic mercury: + +SnCl_{2} + 2HgCl --> SnCl_{4} + 2Hg. + +The occurrence of this secondary reaction is indicated by the +darkening of the precipitate; and, since potassium bichromate oxidizes +this mercury slowly, solutions in which it has been precipitated are +worthless as iron determinations.] + +[Note 3: The solution should be allowed to stand about three minutes +after the addition of mercuric chloride to permit the complete +deposition of mercurous chloride. It should then be titrated without +delay to avoid possible reoxidation of the iron by the oxygen of the +air.] + + + + +DETERMINATION OF IRON IN LIMONITE + + +PROCEDURE.--Grind the mineral (Note 1) to a fine powder. Weigh out +accurately two portions of about 0.5 gram (Note 2) into porcelain +crucibles; heat these crucibles to dull redness for ten minutes, +allow them to cool, and place them, with their contents, in beakers +containing 30 cc. of dilute hydrochloric acid (sp. gr. 1.12). Heat +at a temperature just below boiling until the undissolved residue is +white or until solvent action has ceased. If the residue is white, +or known to be free from iron, it may be neglected and need not be +removed by filtration. If a dark residue remains, collect it on a +filter, wash free from hydrochloric acid, and ignite the filter in a +platinum crucible (Note 3). Mix the ash with five times its weight of +sodium carbonate and heat to fusion; cool, and disintegrate the fused +mass with boiling water in the crucible. Unite this solution and +precipitate (if any) with the acid solution, taking care to avoid loss +by effervescence. Wash out the crucible, heat the acid solution +to boiling, add stannous chloride solution until it is colorless, +avoiding a large excess (Note 4); cool, and when !cold!, add 40 cc. of +mercuric chloride solution, dilute to 200 cc., and proceed with the +titration as already described. + +From the standardization data already obtained, and the known weight +of the sample, calculate the percentage of iron (Fe) in the limonite. + +[Note 1: Limonite is selected as a representative of iron ores in +general. It is a native, hydrated oxide of iron. It frequently occurs +in or near peat beds and contains more or less organic matter which, +if brought into solution, would be acted upon by the potassium +bichromate. This organic matter is destroyed by roasting. Since a high +temperature tends to lessen the solubility of ferric oxide, the heat +should not be raised above low redness.] + +[Note 2: It is sometimes advantageous to dissolve a large portion--say +5 grams--and to take one tenth of it for titration. The sample will +then represent more closely the average value of the ore.] + +[Note 3: A platinum crucible may be used for the roasting of the +limonite and must be used for the fusion of the residue. When used, it +must not be allowed to remain in the acid solution of ferric chloride +for any length of time, since the platinum is attacked and dissolved, +and the platinic chloride is later reduced by the stannous chloride, +and in the reduced condition reacts with the bichromate, thus +introducing an error. It should also be noted that copper and antimony +interfere with the determination of iron by the bichromate process.] + +[Note 4: The quantity of stannous chloride required for the reduction +of the iron in the limonite will be much larger than that added to the +solution of iron wire, in which the iron was mainly already in the +ferrous condition. It should, however, be added from a dropper to +avoid an unnecessary excess.] + + + + +DETERMINATION OF CHROMIUM IN CHROME IRON ORE + + +PROCEDURE.--Grind the chrome iron ore (Note 1) in an agate mortar +until no grit is perceptible under the pestle. Weigh out two portions +of 0.5 gram each into iron crucibles which have been scoured inside +until bright (Note 2). Weigh out on a watch-glass (Note 3), using the +rough balances, 5 grams of dry sodium peroxide for each portion, and +pour about three quarters of the peroxide upon the ore. Mix ore and +flux by thorough stirring with a dry glass rod. Then cover the mixture +with the remainder of the peroxide. Place the crucible on a triangle +and raise the temperature !slowly! to the melting point of the flux, +using a low flame, and holding the lamp in the hand (Note 4). Maintain +the fusion for five minutes, and stir constantly with a stout iron +wire, but do not raise the temperature above moderate redness (Notes 5 +and 6). + +Allow the crucible to cool until it can be comfortably handled (Note +7) and then place it in a 300 cc. beaker, and cover it with distilled +water (Note 8). The beaker must be carefully covered to avoid loss +during the disintegration of the fused mass. When the evolution of +gas ceases, rinse off and remove the crucible; then heat the solution +!while still alkaline! to boiling for fifteen minutes. Allow the +liquid to cool for a few minutes; then acidify with dilute sulphuric +acid (1:5), adding 10 cc. in excess of the amount necessary to +dissolve the ferric hydroxide (Note 9). Dilute to 200 cc., cool, add +from a burette an excess of a standard ferrous solution, and titrate +for the excess with a standard solution of potassium bichromate, using +the outside indicator (Note 10). + +From the corrected volumes of the two standard solutions, and their +relations to normal solutions, calculate the percentage of chromium in +the ore. + +[Note 1: Chrome iron ore is essentially a ferrous chromite, or +combination of FeO and Cr_{2}O_{3}. It must be reduced to a state of +fine subdivision to ensure a prompt reaction with the flux.] + +[Note 2: The scouring of the iron crucible is rendered much easier if +it is first heated to bright redness and plunged into cold water. In +this process oily matter is burned off and adhering scale is caused to +chip off when the hot crucible contracts rapidly in the cold water.] + +[Note 3: Sodium peroxide must be kept off of balance pans and should +not be weighed out on paper, as is the usual practice in the rough +weighing of chemicals. If paper to which the peroxide is adhering is +exposed to moist air it is likely to take fire as a result of +the absorption of moisture, and consequent evolution of heat and +liberation of oxygen.] + +[Note 4: The lamp should never be allowed to remain under the +crucible, as this will raise the temperature to a point at which the +crucible itself is rapidly attacked by the flux and burned through.] + +[Note 5: The sodium peroxide acts as both a flux and an oxidizing +agent. The chromic oxide is dissolved by the flux and oxidized to +chromic anhydride (CrO_{3}) which combines with the alkali to form +sodium chromate. The iron is oxidized to ferric oxide.] + +[Note 6: The sodium peroxide cannot be used in porcelain, platinum, or +silver crucibles. It attacks iron and nickel as well; but crucibles +made from these metals may be used if care is exercised to keep the +temperature as low as possible. Preference is here given to iron +crucibles, because the resulting ferric hydroxide is more readily +brought into solution than the nickelic oxide from a nickel crucible. +The peroxide must be dry, and must be protected from any admixture of +dust, paper, or of organic matter of any kind, otherwise explosions +may ensue.] + +[Note 7: When an iron crucible is employed it is desirable to allow +the fusion to become nearly cold before it is placed in water, +otherwise scales of magnetic iron oxide may separate from the +crucible, which by slowly dissolving in acid form ferrous sulphate, +which reduces the chromate.] + +[Note 8: Upon treatment with water the chromate passes into solution, +the ferric hydroxide remains undissolved, and the excess of peroxide +is decomposed with the evolution of oxygen. The subsequent boiling +insures the complete decomposition of the peroxide. Unless this is +complete, hydrogen peroxide is formed when the solution is acidified, +and this reacts with the bichromate, reducing it and introducing a +serious error.] + +[Note 9: The addition of the sulphuric acid converts the sodium +chromate to bichromate, which behaves exactly like potassium +bichromate in acid solution.] + +[Note 10: If a standard solution of a ferrous salt is not at hand, a +weight of iron wire somewhat in excess of the amount which would be +required if the chromite were pure FeO.Cr_{2}O_{3} may be weighed out +and dissolved in sulphuric acid; after reduction of all the iron by +stannous chloride and the addition of mercuric chloride, this solution +may be poured into the chromate solution and the excess of iron +determined by titration with standard bichromate solution.] + + + + +PERMANGANATE PROCESS FOR THE DETERMINATION OF IRON + + +Potassium permanganate oxidizes ferrous salts in cold, acid solution +promptly and completely to the ferric condition, while in hot acid +solution it also enters into a definite reaction with oxalic acid, by +which the latter is oxidized to carbon dioxide and water. + +The reactions involved are these: + +10FeSO_{4} + 2KMnO_{4} + 8H_{2}S_{4} --> 5Fe_{2}(SO_{4})_{3} + +K_{2}SO_{4} + 2MnSO_{4} + 8H_{2}O + +5C_{2}H_{2}O_{4}(2H_{2}O) + 2KMnO_{4} +3H_{2}SO_{4} --> K_{2}SO_{4} + +2MnSO_{4} + 10CO_{2} + 1 H_{2}O. + +These are the fundamental reactions upon which the extensive use of +potassium permanganate depends; but besides iron and oxalic acid the +permanganate enters into reaction with antimony, tin, copper, mercury, +and manganese (the latter only in neutral solution), by which these +metals are changed from a lower to a higher state of oxidation; and it +also reacts with sulphurous acid, sulphureted hydrogen, nitrous acid, +ferrocyanides, and most soluble organic bodies. It should be noted, +however, that very few of these organic compounds react quantitatively +with the permanganate, as is the case with oxalic acid and the +oxalates. + +Potassium permanganate is acted upon by hydrochloric acid; the action +is rapid in hot or concentrated solution (particularly in the presence +of iron salts, which appear to act as catalyzers, increasing the +velocity of the reaction), but slow in cold, dilute solutions. +However, the greater solubility of iron compounds in hydrochloric acid +makes it desirable to use this acid as a solvent, and experiments made +with this end in view have shown that in cold, dilute hydrochloric +acid solution, to which considerable quantities of manganous sulphate +and an excess of phosphoric acid have been added, it is possible to +obtain satisfactory results. + +It is also possible to replace the hydrochloric acid by evaporating +the solutions with an excess of sulphuric acid until the latter fumes. +This procedure is somewhat more time-consuming, but the end-point of +the permanganate titration is more permanent. Both procedures are +described below. + +Potassium permanganate has an intense coloring power, and since the +solution resulting from the oxidation of the iron and the reduction of +the permanganate is colorless, the latter becomes its own indicator. +The slightest excess is indicated with great accuracy by the pink +color of the solution. + + +PREPARATION OF A STANDARD SOLUTION + +!Approximate Strength 0.1 N! + +A study of the reactions given above which represent the oxidation of +ferrous compounds by potassium permanganate, shows that there are 2 +molecules of KMnO_{4} and 10 molecules of FeSO_{4} on the +left-hand side, and 2 molecules of MnSO_{4} and 5 molecules of +Fe_{2}(SO_{4})_{5} on the right-hand side. Considering only these +compounds, and writing the formulas in such a way as to show the +oxides of the elements in each, the equation becomes: + +K_{2}O.Mn_{2}O_{7} + 10(FeO.SO_{3}) --> K_{2}O.SO_{3} + 2(MnO.SO_{3}) ++ 5(Fe_{2}O_{3}.3SO_{3}). + +From this it appears that two molecules of KMnO_{4} (or 316.0 grams) +have given up five atoms (or 80 grams) of oxygen to oxidize the +ferrous compound. Since 8 grams of oxygen is the basis of normal +oxidizing solutions and 80 grams of oxygen are supplied by 316.0 grams +of KMnO_{4}, the normal solution of the permanganate should contain, +per liter, 316.0/10 grams, or 31.60 grams (Note 1). + +The preparation of an approximately tenth-normal solution of the +reagent may be carried out as follows: + +PROCEDURE.--Dissolve about 3.25 grams of potassium permanganate +crystals in approximately 1000 cc. of distilled water in a large +beaker, or casserole. Heat slowly and when the crystals have +dissolved, boil the solution for 10-15 minutes. Cover the solution +with a watch-glass; allow it to stand until cool, or preferably over +night. Filter the solution through a layer of asbestos. Transfer the +filtrate to a liter bottle and mix thoroughly (Note 2). + +[Note 1: The reactions given on page 61 are those which take place in +the presence of an excess of acid. In neutral solutions the reduction +of the permanganate is less complete, and, under these conditions, +two gram-molecular weights of KMnO_{4} will furnish only 48 grams +of oxygen. A normal solution for use under these conditions should, +therefore, contain 316.0/6 grams, or 52.66 grams.] + +[Note 2: Potassium permanganate solutions are not usually stable for +long periods, and change more rapidly when first prepared than after +standing some days. This change is probably caused by interaction +with the organic matter contained in all distilled water, except that +redistilled from an alkaline permanganate solution. The solutions +should be protected from light and heat as far as possible, since both +induce decomposition with a deposition of manganese dioxide, and it +has been shown that decomposition proceeds with considerable rapidity, +with the evolution of oxygen, after the dioxide has begun to form. As +commercial samples of the permanganate are likely to be contaminated +by the dioxide, it is advisable to boil and filter solutions through +asbestos before standardization, as prescribed above. Such solutions +are relatively stable.] + + +COMPARISON OF PERMANGANATE AND FERROUS SOLUTIONS + +PROCEDURE.--Fill a glass-stoppered burette with the permanganate +solution, observing the usual precautions, and fill a second burette +with the ferrous sulphate solution prepared for use with the potassium +bichromate. The permanganate solution cannot be used in burettes with +rubber tips, as a reduction takes place upon contact with the rubber. +The solution has so deep a color that the lower line of the meniscus +cannot be detected; readings must therefore be made from the upper +edge. Run out into a beaker about 40 cc. of the ferrous solution, +dilute to about 100 cc., add 10 cc. of dilute sulphuric acid, and run +in the permanganate solution to a slight permanent pink. Repeat, until +the ratio of the two solutions is satisfactorily established. + + +STANDARDIZATION OF A POTASSIUM PERMANGANATE SOLUTION + +!Selection of a Standard! + +Commercial potassium permanganate is rarely sufficiently pure to admit +of its direct weighing as a standard. On this account, and because +of the uncertainties as to the permanence of its solutions, it is +advisable to standardize them against substances of known value. Those +in most common use are iron wire, ferrous ammonium sulphate, sodium +oxalate, oxalic acid, and some other derivatives of oxalic acid. +With the exception of sodium oxalate, these all contain water of +crystallization which may be lost on standing. They should, therefore, +be freshly prepared, and with great care. At present, sodium oxalate +is considered to be one of the most satisfactory standards. + + +!Method A! + + +!Iron Standards! + +The standardization processes employed when iron or its compounds are +selected as standards differ from those applicable in connection with +oxalate standards. The procedure which immediately follows is that in +use with iron standards. + +As in the case of the bichromate process, it is necessary to reduce +the iron completely to the ferrous condition before titration. The +reducing agents available are zinc, sulphurous acid, or sulphureted +hydrogen. Stannous chloride may also be used when the titration is +made in the presence of hydrochloric acid. Since the excess of both +the gaseous reducing agents can only be expelled by boiling, with +consequent uncertainty regarding both the removal of the excess and +the reoxidation of the iron, zinc or stannous chlorides are the most +satisfactory agents. For prompt and complete reduction it is essential +that the iron solution should be brought into ultimate contact with +the zinc. This is brought about by the use of a modified Jones +reductor, as shown in Figure 1. This reductor is a standard apparatus +and is used in other quantitative processes. + +[Illustration: Fig. 1] + +The tube A has an inside diameter of 18 mm. and is 300 mm. long; the +small tube has an inside diameter of 6 mm. and extends 100 mm. below +the stopcock. At the base of the tube A are placed some pieces of +broken glass or porcelain, covered by a plug of glass wool about 8 mm. +thick, and upon this is placed a thin layer of asbestos, such as is +used for Gooch filters, 1 mm. thick. The tube is then filled with the +amalgamated zinc (Note 1) to within 50 mm. of the top, and on the zinc +is placed a plug of glass wool. If the top of the tube is not already +shaped like the mouth of a thistle-tube (B), a 60 mm. funnel is fitted +into the tube with a rubber stopper and the reductor is connected +with a suction bottle, F. The bottle D is a safety bottle to +prevent contamination of the solution by water from the pump. After +preparation for use, or when left standing, the tube A should be +filled with water, to prevent clogging of the zinc. + +[Note 1: The use of fine zinc in the reductor is not necessary and +tends to clog the tube. Particles which will pass a 10-mesh sieve, but +are retained by one of 20 meshes to the inch, are most satisfactory. +The zinc can be amalgamated by stirring or shaking it in a mixture of +25 cc. of normal mercuric chloride solution, 25 cc. of hydrochloric +acid (sp. gr. 1.12) and 250 cc. of water for two minutes. The solution +should then be poured off and the zinc thoroughly washed. It is then +ready for bottling and preservation under water. A small quantity of +glass wool is placed in the neck of the funnel to hold back foreign +material when the reductor is in use.] + + +STANDARDIZATION + +PROCEDURE.--Weigh out into Erlenmeyer flasks two portions of iron wire +of about 0.25 gram each. Dissolve these in hot dilute sulphuric acid +(5 cc. of concentrated acid and 100 cc. of water), using a covered +flask to avoid loss by spattering. Boil the solution for two or +three minutes after the iron has dissolved to remove any volatile +hydrocarbons. Meanwhile prepare the reductor for use as follows: +Connect the vacuum bottle with the suction pump and pour into the +funnel at the top warm, dilute sulphuric acid, prepared by adding 5 +cc. of concentrated sulphuric acid to 100 cc. of distilled water. See +that the stopcock (C) is open far enough to allow the acid to run +through slowly. Continue to pour in acid until 200 cc. have passed +through, then close the stopcock !while a small quantity of liquid +is still left in the funnel!. Discard the filtrate, and again +pass through 100 cc. of the warm, dilute acid. Test this with the +permanganate solution. A single drop should color it permanently; if +it does not, repeat the washing, until assured that the zinc is not +contaminated with appreciable quantities of reducing substances. Be +sure that no air enters the reductor (Note 1). + +Pour the iron solution while hot (but not boiling) through the +reductor at a rate not exceeding 50 cc. per minute (Notes 2 and 3). +Wash out the beaker with dilute sulphuric acid, and follow the iron +solution without interruption with 175 cc. of the warm acid and +finally with 75 cc. of distilled water, leaving the funnel partially +filled. Remove the filter bottle and cool the solution quickly under +the water tap (Note 4), avoiding unnecessary exposure to the oxygen of +the air. Add 10 cc. of dilute sulphuric acid and titrate to a faint +pink with the permanganate solution, adding it directly to the +contents of the vacuum flask. Should the end-point be overstepped, the +ferrous sulphate solution may be added. + +From the volume of the solution required to oxidize the iron in +the wire, calculate the relation to the normal of the permanganate +solution. The duplicate results should be concordant within two parts +in one thousand. + +[Note 1: The funnel of the reductor must never be allowed to empty. +If it is left partially filled with water the reductor is ready for +subsequent use after a very little washing; but a preliminary test is +always necessary to safeguard against error. + +If more than a small drop of permanganate solution is required to +color 100 cc. of the dilute acid after the reductor is well washed, an +allowance must be made for the iron in the zinc. !Great care! must be +used to prevent the access of air to the reductor after it has been +washed out ready for use. If air enters, hydrogen peroxide forms, +which reacts with the permanganate, and the results are worthless.] + +[Note 2: The iron is reduced to the ferrous condition by contact with +the zinc. The active agent may be considered to be !nascent! hydrogen, +and it must be borne in mind that the visible bubbles are produced by +molecular hydrogen, which is without appreciable effect upon ferric +iron. + +The rate at which the iron solution passes through the zinc should not +exceed that prescribed, but the rate may be increased somewhat when +the wash-water is added. It is well to allow the iron solution to run +nearly, but not entirely, out of the funnel before the wash-water +is added. If it is necessary to interrupt the process, the complete +emptying of the funnel can always be avoided by closing the stopcock. + +It is also possible to reduce the iron by treatment with zinc in a +flask from which air is excluded. The zinc must be present in excess +of the quantity necessary to reduce the iron and is finally completely +dissolved. This method is, however, less convenient and more tedious +than the use of the reductor.] + +[Note 3: The dilute sulphuric acid for washing must be warmed ready +for use before the reduction of the iron begins, and it is of the +first importance that the volume of acid and of wash-water should +be measured, and the volume used should always be the same in the +standardizations and all subsequent analyses.] + +[Note 4: The end-point is more permanent in cold than hot solutions, +possibly because of a slight action of the permanganate upon the +manganous sulphate formed during titration. If the solution turns +brown, it is an evidence of insufficient acid, and more should be +immediately added. The results are likely to be less accurate in this +case, however, as a consequence of secondary reactions between the +ferrous iron and the manganese dioxide thrown down. It is wiser to +discard such results and repeat the process.] + +[Note 5: The potassium permanganate may, of course, be diluted and +brought to an exactly 0.1 N solution from the data here obtained. The +percentage of iron in the iron wire must be taken into account in all +calculations.] + + +!Method B! + +!Oxalate Standards! + +PROCEDURE.--Weigh out two portions of pure sodium oxalate of 0.25-0.3 +gram each into beakers of about 600 cc. capacity. Add about 400 cc. of +boiling water and 20 cc. of manganous sulphate solution (Note 1). +When the solution of the oxalate is complete, heat the liquid, if +necessary, until near its boiling point (70-90 deg.C.) and run in the +standard permanganate solution drop by drop from a burette, stirring +constantly until an end-point is reached (Note 2). Make a blank test +with 20 cc. of manganous sulphate solution and a volume of distilled +water equal to that of the titrated solution to determine the volume +of the permanganate solution required to produce a very slight pink. +Deduct this volume from the amount of permanganate solution used in +the titration. + +From the data obtained, calculate the relation of the permanganate +solution to the normal. The reaction involved is: + +5Na_{2}C_{2}O_{4} + 2KMnO_{4} + 8H_{2}SO_{4} --> 5Na_{2}SO_{4} + +K_{2}SO_{4} + 2MnSO_{4} + 10CO_{2} + 8H_{2}O + +[Note 1: The manganous sulphate titrating solution is made by +dissolving 20 grams of MnSO_{4} in 200 cubic centimeters of water and +adding 40 cc. of concentrated sulphuric acid (sp. gr. 1.84) and 40 cc. +or phosphoric acid (85%).] + +[Note 2: The reaction between oxalates and permanganates takes place +quantitatively only in hot acid solutions. The temperatures must not +fall below 70 deg.C.] + + + + +DETERMINATION OF IRON IN LIMONITE + + +!Method A! + +The procedures, as here prescribed, are applicable to iron ores in +general, provided these ores contain no constituents which are reduced +by zinc or stannous chloride and reoxidized by permanganates. Many +iron ores contain titanium, and this element among others does +interfere with the determination of iron by the process described. +If, however, the solutions of such ores are treated with sulphureted +hydrogen or sulphurous acid, instead of zinc or stannous chloride to +reduce the iron, and the excess reducing agent removed by boiling, an +accurate determination of the iron can be made. + +PROCEDURE.--Grind the mineral to a fine powder. Weigh out two portions +of about 0.5 gram each into small porcelain crucibles. Roast the ore +at dull redness for ten minutes (Note 1), allow the crucibles to cool, +and place them and their contents in casseroles containing 30 cc. of +dilute hydrochloric acid (sp. gr. 1.12). + +Proceed with the solution of the ore, and the treatment of the +residue, if necessary, exactly as described for the bichromate process +on page 56. When solution is complete, add 6 cc. of concentrated +sulphuric acid to each casserole, and evaporate on the steam bath +until the solution is nearly colorless (Note 2). Cover the casseroles +and heat over the flame of the burner, holding the casserole in +the hand and rotating it slowly to hasten evaporation and prevent +spattering, until the heavy white fumes of sulphuric anhydride are +freely evolved (Note 3). Cool the casseroles, add 100 cc. of water +(measured), and boil gently until the ferric sulphate is dissolved; +pour the warm solution through the reductor which has been previously +washed; proceed as described under standardization, taking pains +to use the same volume and strength of acid and the same volume of +wash-water as there prescribed, and titrate with the permanganate +solution in the reductor flask, using the ferrous sulphate solution if +the end-point should be overstepped. + +From the corrected volume of permanganate solution used, calculate the +percentage of iron (Fe) in the limonite. + +[Note 1: The preliminary roasting is usually necessary because, even +though the sulphuric acid would subsequently char the carbonaceous +matter, certain nitrogenous bodies are not thereby rendered insoluble +in the acid, and would be oxidized by the permanganate.] + +[Note 2: The temperature of the steam bath is not sufficient to +volatilize sulphuric acid. Solutions may, therefore, be left to +evaporate overnight without danger of evaporation to dryness.] + +[Note 3: The hydrochloric acid, both free and combined, is displaced +by the less volatile sulphuric acid at its boiling point. Ferric +sulphate separates at this point, since there is no water to hold +it in solution and care is required to prevent bumping. The ferric +sulphate usually has a silky appearance and is easily distinguished +from the flocculent silica which often remains undissolved.] + + +!Zimmermann-Reinhardt Procedure! + + +!Method (B)! + +PROCEDURE.--Grind the mineral to a fine powder. Weigh out two portions +of about 0.5 gram each into small porcelain crucibles. Proceed with +the solution of the ore, treat the residue, if necessary, and reduce +the iron by the addition of stannous chloride, followed by mercuric +chloride, as described for the bichromate process on page 56. Dilute +the solution to about 400 cc. with cold water, add 10 cc. of the +manganous sulphate titrating solution (Note 1, page 68) and titrate +with the standard potassium permanganate solution to a faint pink +(Note 1). + +From the standardization data already obtained calculate the +percentage of iron (Fe) in the limonite. + +[Note 1: It has already been noted that hydrochloric acid reacts +slowly in cold solutions with potassium permanganate. It is, however, +possible to obtain a satisfactory, although somewhat fugitive +end-point in the presence of manganous sulphate and phosphoric acid. +The explanation of the part played by these reagents is somewhat +obscure as yet. It is possible that an intermediate manganic compound +is formed which reacts rapidly with the ferrous compounds--thus in +effect catalyzing the oxidizing process. + +While an excess of hydrochloric acid is necessary for the successful +reduction of the iron by stannous chloride, too large an amount +should be avoided in order to lessen the chance of reduction of the +permanganate by the acid during titration.] + + + + +DETERMINATION OF THE OXIDIZING POWER OF PYROLUSITE + +INDIRECT OXIDATION + + +Pyrolusite, when pure, consists of manganese dioxide. Its value as an +oxidizing agent, and for the production of chlorine, depends upon the +percentage of MnO_{2} in the sample. This percentage is determined +by an indirect method, in which the manganese dioxide is reduced and +dissolved by an excess of ferrous sulphate or oxalic acid in the +presence of sulphuric acid, and the unused excess determined by +titration with standard permanganate solution. + +PROCEDURE.--Grind the mineral in an agate mortar until no grit +whatever can be detected under the pestle (Note 1). Transfer it to a +stoppered weighing-tube, and weigh out two portions of about 0.5 gram +into beakers (400-500 cc.) Read Note 2, and then calculate in each +case the weight of oxalic acid (H_{2}C_{2}O_{4}.2H_{2}O) required to +react with the weights of pyrolusite taken. The reaction involved is + +MnO_{2} + H_{2}C_{2}O_{4}(2H_{2}O) + H_{2}SO_{4} --> MnSO_{4} + +2CO_{2} + 4H_{2}O. + +Weigh out about 0.2 gram in excess of this quantity of !pure! oxalic +acid into the corresponding beakers, weighing the acid accurately and +recording the weight in the notebook. Pour into each beaker 25 cc. of +water and 50 cc. of dilute sulphuric acid (1:5), cover and warm the +beaker and its contents gently until the evolution of carbon dioxide +ceases (Note 3). If a residue remains which is sufficiently colored to +obscure the end-reaction of the permanganate, it must be removed by +filtration. + +Finally, dilute the solution to 200-300 cc., heat the solution to a +temperature just below boiling, add 15 cc. of a manganese sulphate +solution and while hot, titrate for the excess of the oxalic acid with +standard permanganate solution (Notes 4 and 5). + +From the corrected volume of the solution required, calculate the +amount of oxalic acid undecomposed by the pyrolusite; subtract this +from the total quantity of acid used, and calculate the weight of +manganese dioxide which would react with the balance of the acid, and +from this the percentage in the sample. + +[Note 1: The success of the analysis is largely dependent upon the +fineness of the powdered mineral. If properly ground, solution should +be complete in fifteen minutes or less.] + +[Note 2: A moderate excess of oxalic acid above that required to react +with the pyrolusite is necessary to promote solution; otherwise the +residual quantity of oxalic acid would be so small that the last +particles of the mineral would scarcely dissolve. It is also desirable +that a sufficient excess of the acid should be present to react with a +considerable volume of the permanganate solution during the titration, +thus increasing the accuracy of the process. On the other hand, the +excess of oxalic acid should not be so large as to react with more of +the permanganate solution than is contained in a 50 cc. burette. If +the pyrolusite under examination is known to be of high grade, say 80 +per cent pure, or above the calculation of the oxalic acid needed may +be based upon an assumption that the mineral is all MnO_{2}. If the +quality of the mineral is unknown, it is better to weigh out three +portions instead of two and to add to one of these the amount of +oxalic prescribed, assuming complete purity of the mineral. Then run +in the permanganate solution from a pipette or burette to determine +roughly the amount required. If the volume exceeds the contents of a +burette, the amount of oxalic acid added to the other two portions is +reduced accordingly.] + +[Note 3: Care should be taken that the sides of the beaker are not +overheated, as oxalic acid would be decomposed by heat alone if +crystallization should occur on the sides of the vessel. Strong +sulphuric acid also decomposes the oxalic acid. The dilute acid +should, therefore, be prepared before it is poured into the beaker.] + +[Note 4: Ferrous ammonium sulphate, ferrous sulphate, or iron wire +may be substituted for the oxalic acid. The reaction is then the +following: + +2 FeSO_{4} + MnO_{2} + 2H_{2}SO_{4} --> Fe_{2}(SO_{4})_{3} + 2H_{2}O + +The excess of ferrous iron may also be determined by titration with +potassium bichromate, if desired. Care is required to prevent the +oxidation of the iron by the air, if ferrous salts are employed.] + +[Note 5: The oxidizing power of pyrolusite may be determined by other +volumetric processes, one of which is outlined in the following +reactions: + +MnO_{2} + 4HCl --> MnCl_{2} + Cl_{2} + 2H_{2}O +Cl_{2} + 2KI --> I_{2} + 2KCl +I_{2} + 2Na_{2}S_{2}O_{3} --> Na_{2}S_{4}O_{6} + 2NaI. + +The chlorine generated by the pyrolusite is passed into a solution of +potassium iodide. The liberated iodine is then determined by titration +with sodium thiosulphate, as described on page 78. This is a direct +process, although it involves three steps.] + + + + +IODIMETRY + + +The titration of iodine against sodium thiosulphate, with starch as an +indicator, may perhaps be regarded as the most accurate of volumetric +processes. The thiosulphate solution may be used in both acid and +neutral solutions to measure free iodine and the latter may, in turn, +serve as a measure of any substance capable of liberating iodine from +potassium iodide under suitable conditions for titration, as, for +example, in the process outlined in Note 5 on page 74. + +The fundamental reaction upon which iodometric processes are based is +the following: + +I_{2} + 2 Na_{2}S_{2}O_{3} --> 2 NaI + Na_{2}S_{4}O_{6}. + +This reaction between iodine and sodium thiosulphate, resulting in +the formation of the compound Na_{2}S_{4}O_{6}, called sodium +tetrathionate, is quantitatively exact, and differs in that +respect from the action of chlorine or bromine, which oxidize the +thiosulphate, but not quantitatively. + +NORMAL SOLUTIONS OF IODINE AND SODIUM THIOSULPHATE + +If the formulas of sodium thiosulphate and sodium tetrathionate are +written in a manner to show the atoms of oxygen associated +with sulphur atoms in each, thus, 2(Na_{2}).S_{2}O_{2} and +Na_{2}O.S_{4}O_{5}, it is plain that in the tetrathionate there are +five atoms of oxygen associated with sulphur, instead of the four +in the two molecules of the thiosulphate taken together. Although, +therefore, the iodine contains no oxygen, the two atoms of iodine +have, in effect, brought about the addition of one oxygen atoms to the +sulphur atoms. That is the same thing as saying that 253.84 grams of +iodine (I_{2}) are equivalent to 16 grams of oxygen; hence, since 8 +grams of oxygen is the basis of normal solutions, 253.84/2 or 126.97 +grams of iodine should be contained in one liter of normal iodine +solution. By a similar course of reasoning the conclusion is reached +that the normal solution of sodium thiosulphate should contain, +per liter, its molecular weight in grams. As the thiosulphate in +crystalline form has the formula Na_{2}S_{2}O_{3}.5H_{2}O, this weight +is 248.12 grams. Tenth-normal or hundredth-normal solutions are +generally used. + + +PREPARATION OF STANDARD SOLUTIONS + +!Approximate Strength, 0.1 N! + +PROCEDURE.--Weigh out on the rough balances 13 grams of commercial +iodine. Place it in a mortar with 18 grams of potassium iodide and +triturate with small portions of water until all is dissolved. Dilute +the solution to 1000 cc. and transfer to a liter bottle and mix +thoroughly (Note 1).[1] + +[Footnote 1: It will be found more economical to have a considerable +quantity of the solution prepared by a laboratory attendant, and to +have all unused solutions returned to the common stock.] + +Weigh out 25 grams of sodium thiosulphate, dissolve it in water which +has been previously boiled and cooled, and dilute to 1000 cc., also +with boiled water. Transfer the solution to a liter bottle and mix +thoroughly (Note 2). + +[Note 1: Iodine solutions react with water to form hydriodic acid +under the influence of the sunlight, and even at low room temperatures +the iodine tends to volatilize from solution. They should, therefore, +be protected from light and heat. Iodine solutions are not stable for +long periods under the best of conditions. They cannot be used in +burettes with rubber tips, since they attack the rubber.] + +[Note 2: Sodium thiosulphate (Na_{2}S_{2}O_{3}.5H_{2}O) is +rarely wholly pure as sold commercially, but may be purified by +recrystallization. The carbon dioxide absorbed from the air by +distilled water decomposes the salt, with the separation of sulphur. +Boiled water which has been cooled out of contact with the air should +be used in preparing solutions.] + + +INDICATOR SOLUTION + +The starch solution for use as an indicator must be freshly prepared. +A soluble starch is obtainable which serves well, and a solution of +0.5 gram of this starch in 25 cc. of boiling water is sufficient. The +solution should be filtered while hot and is ready for use when cold. + +If soluble starch is not at hand, potato starch may be used. Mix about +1 gram with 5 cc. of cold water to a smooth paste, pour 150 cc. of +!boiling! water over it, warm for a moment on the hot plate, and put +it aside to settle. Decant the supernatant liquid through a filter +and use the clear filtrate; 5 cc. of this solution are needed for a +titration. + +The solution of potato starch is less stable than the soluble starch. +The solid particles of the starch, if not removed by filtration, +become so colored by the iodine that they are not readily decolorized +by the thiosulphate (Note 1). + +[Note 1: The blue color which results when free iodine and starch +are brought together is probably not due to the formation of a true +chemical compound. It is regarded as a "solid solution" of iodine in +starch. Although it is unstable, and easily destroyed by heat, it +serves as an indicator for the presence of free iodine of remarkable +sensitiveness, and makes the iodometric processes the most +satisfactory of any in the field of volumetric analysis.] + + +COMPARISON OF IODINE AND THIOSULPHATE SOLUTIONS + +PROCEDURE.--Place the solutions in burettes (the iodine in a +glass-stoppered burette), observing the usual precautions. Run out 40 +cc. of the thiosulphate solution into a beaker, dilute with 150 cc. of +water, add 1 cc. to 2 cc. of the soluble starch solution, and titrate +with the iodine to the appearance of the blue of the iodo-starch. +Repeat until the ratio of the two solutions is established, +remembering all necessary corrections for burettes and for temperature +changes. + + +STANDARDIZATION OF SOLUTIONS + +Commercial iodine is usually not sufficiently pure to permit of its +use as a standard for thiosulphate solutions or the direct preparation +of a standard solution of iodine. It is likely to contain, beside +moisture, some iodine chloride, if chlorine was used to liberate the +iodine when it was prepared. It may be purified by sublimation after +mixing it with a little potassium iodide, which reacts with the iodine +chloride, forming potassium chloride and setting free the iodine. The +sublimed iodine is then dried by placing it in a closed container over +concentrated sulphuric acid. It may then be weighed in a stoppered +weighing-tube and dissolved in a solution of potassium iodide in a +stoppered flask to prevent loss of iodine by volatilization. About 18 +grams of the iodide and twelve grams of iodine per liter are required +for an approximately tenth-normal solution. + +An iodine solution made from commercial iodine may also be +standardized against arsenious oxide (As_{4}O_{6}). This substance +also usually requires purification by sublimation before use. + +The substances usually employed for the standardization of a +thiosulphate solution are potassium bromate and metallic copper. The +former is obtainable in pure condition or may be easily purified by +re-crystallization. Copper wire of high grade is sufficiently pure +to serve as a standard. Both potassium bromate and cupric salts in +solution will liberate iodine from an iodide, which is then titrated +with the thiosulphate solution. + +The reactions involved are the following: + +(a) KBrO_{3} + 6KI + 3H_{2}SO_{4} --> KBr + 3I_{2} + 3K_{2}SO_{4} + 3H_{2}O, + +(b) 3Cu + 8HNO_{3} --> 3Cu(NO_{3})_{2} + 2NO + 4H_{2}O, + 2Cu(NO_{3})_{2} + 4KI --> 2CuI + 4KNO_{3} + I_{2}. + +Two methods for the direct standardization of the sodium thiosulphate +solution are here described, and one for the direct standardization of +the iodine solution. + + +!Method A! + +PROCEDURE.--Weigh out into 500 cc. beakers two portions of about +0.150-0.175 gram of potassium bromate. Dissolve each of these in 50 +cc. of water, and add 10 cc. of a potassium iodide solution containing +3 grams of the salt in that volume (Note 1). Add to the mixture 10 cc. +of dilute sulphuric acid (1 volume of sulphuric acid with 5 volumes of +water), allow the solution to stand for three minutes, and dilute to +150 cc. (Note 2). Run in thiosulphate solution from a burette until +the color of the liberated iodine is nearly destroyed, and then add 1 +cc. or 2 cc. of starch solution, titrate to the disappearance of the +iodo-starch blue, and finally add iodine solution until the color +is just restored. Make a blank test for the amount of thiosulphate +solution required to react with the iodine liberated by the iodate +which is generally present in the potassium iodide solution, and +deduct this from the total volume used in the titration. + +From the data obtained, calculate the relation of the thiosulphate +solution to a normal solution, and subsequently calculate the similar +value for the iodine solution. + +[Note 1:--Potassium iodide usually contains small amounts of potassium +iodate as impurity which, when the iodide is brought into an acid +solution, liberates iodine, just as does the potassium bromate used as +a standard. It is necessary to determine the amount of thiosulphate +which reacts with the iodine thus liberated by making a "blank test" +with the iodide and acid alone. As the iodate is not always uniformly +distributed throughout the iodide, it is better to make up a +sufficient volume of a solution of the iodide for the purposes of the +work in hand, and to make the blank test by using the same volume of +the iodide solution as is added in the standardizing process. The +iodide solution should contain about 3 grams of the salt in 10 cc.] + +[Note 2: The color of the iodo-starch is somewhat less satisfactory in +concentrated solutions of the alkali salts, notably the iodides. The +dilution prescribed obviates this difficulty.] + + +!Method B! + +PROCEDURE.--Weigh out two portions of 0.25-0.27 gram of clean copper +wire into 250 cc. Erlenmeyer flasks (Note 1). Add to each 5 cc. of +concentrated nitric acid (sp. gr. 1.42) and 25 cc. of water, cover, +and warm until solution is complete. Add 5 cc. of bromine water and +boil until the excess of bromine is expelled. Cool, and add strong +ammonia (sp. gr. 0.90) drop by drop until a deep blue color indicates +the presence of an excess. Boil the solution until the deep blue is +replaced by a light bluish green, or a brown stain appears on the +sides of the flask (Note 2). Add 10 cc. of strong acetic acid (sp. +gr. 1.04), cool under the water tap, and add a solution of potassium +iodide (Note 3) containing about 3 grams of the salt, and titrate +with thiosulphate solution until the color of the liberated iodine +is nearly destroyed. Then add 1-2 cc. of freshly prepared starch +solution, and add thiosulphate solution, drop by drop, until the blue +color is discharged. + +From the data obtained, including the "blank test" of the iodide, +calculate the relation of the thiosulphate solution to the normal. + +[Note 1: While copper wire of commerce is not absolutely pure, the +requirements for its use as a conductor of electricity are such that +the impurities constitute only a few hundredths of one per cent and +are negligible for analytical purposes.] + +[Note 2: Ammonia neutralizes the free nitric acid. It should be added +in slight excess only, since the excess must be removed by boiling, +which is tedious. If too much ammonia is present when acetic acid is +added, the resulting ammonium acetate is hydrolyzed, and the ammonium +hydroxide reacts with the iodine set free.] + +[Note 3: A considerable excess of potassium iodide is necessary for +the prompt liberation of iodine. While a large excess will do no harm, +the cost of this reagent is so great that waste should be avoided.] + + +!Method C! + +PROCEDURE.--Weigh out into 500 cc. beakers two portions of 0.175-0.200 +gram each of pure arsenious oxide. Dissolve each of these in 10 cc. of +sodium hydroxide solution, with stirring. Dilute the solutions to 150 +cc. and add dilute hydrochloric acid until the solutions contain a few +drops in excess, and finally add to each a concentrated solution of +5 grams of pure sodium bicarbonate (NaHCO_{3}) in water. Cover the +beakers before adding the bicarbonate, to avoid loss. Add the starch +solution and titrate with the iodine to the appearance of the blue of +the iodo-starch, taking care not to pass the end-point by more than a +few drops (Note 1). + +From the corrected volume of the iodine solution used to oxidize the +arsenious oxide, calculate its relation to the normal. From the +ratio between the solutions, calculate the similar value for the +thiosulphate solution. + +[Note 1: Arsenious oxide dissolves more readily in caustic alkali than +in a bicarbonate solution, but the presence of caustic alkali during +the titration is not admissible. It is therefore destroyed by the +addition of acid, and the solution is then made neutral with the +solution of bicarbonate, part of which reacts with the acid, the +excess remaining in solution. + +The reaction during titration is the following: + +Na_{3}AsO_{3} + I_{2} + 2NaHCO_{3} --> Na_{3}AsO_{4} + 2NaI + 2CO_{2} ++ H_{2}O + +As the reaction between sodium thiosulphate and iodine is not always +free from secondary reactions in the presence of even the weakly +alkaline bicarbonate, it is best to avoid the addition of any +considerable excess of iodine. Should the end-point be passed by a few +drops, the thiosulphate may be used to correct it.] + + + + +DETERMINATION OF COPPER IN ORES + + +Copper ores vary widely in composition from the nearly pure copper +minerals, such as malachite and copper sulphide, to very low grade +materials which contain such impurities as silica, lead, iron, silver, +sulphur, arsenic, and antimony. In nearly all varieties there will be +found a siliceous residue insoluble in acids. The method here given, +which is a modification of that described by A.H. Low (!J. Am. Chem. +Soc.! (1902), 24, 1082), provides for the extraction of the copper +from commonly occurring ores, and for the presence of their common +impurities. For practice analyses it is advisable to select an ore of +a fair degree of purity. + +PROCEDURE.-- Weigh out two portions of about 0.5 gram each of the +ore (which should be ground until no grit is detected) into 250 cc. +Erlenmeyer flasks or small beakers. Add 10 cc. of concentrated nitric +acid (sp. gr. 1.42) and heat very gently until the ore is decomposed +and the acid evaporated nearly to dryness (Note 1). Add 5 cc. of +concentrated hydrochloric acid (sp. gr. 1.2) and warm gently. Then +add about 7 cc. of concentrated sulphuric acid (sp. gr. 1.84) and +evaporate over a free flame until the sulphuric acid fumes freely +(Note 2). It has then displaced nitric and hydrochloric acid from +their compounds. + +Cool the flask or beaker, add 25 cc. of water, heat the solution +to boiling, and boil for two minutes. Filter to remove insoluble +sulphates, silica and any silver that may have been precipitated as +silver chloride, and receive the filtrate in a small beaker, washing +the precipitate and filter paper with warm water until the filtrate +and washings amount to 75 cc. Bend a strip of aluminium foil (5 cm. x +12 cm.) into triangular form and place it on edge in the beaker. Cover +the beaker and boil the solution (being careful to avoid loss of +liquid by spattering) for ten minutes, but do not evaporate to small +volume. + +Wash the cover glass and sides of the beaker. The copper should now be +in the form of a precipitate at the bottom of the beaker or adhering +loosely to the aluminium sheet. Remove the sheet, wash it carefully +with hydrogen sulphide water and place it in a small beaker. Decant +the solution through a filter, wash the precipitated copper twice by +decantation with hydrogen sulphide water, and finally transfer the +copper to the filter paper, where it is again washed thoroughly, being +careful at all times to keep the precipitated copper covered with the +wash water. Remove and discard the filtrate and place an Erlenmeyer +flask under the funnel. Pour 15 cc. of dilute nitric acid (sp. gr. +1.20) over the aluminium foil in the beaker, thus dissolving any +adhering copper. Wash the foil with hot water and remove it. Warm this +nitric acid solution and pour it slowly through the filter paper, +thereby dissolving the copper on the paper, receiving the acid +solution in the Erlenmeyer flask. Before washing the paper, pour 5 cc. +of saturated bromine water (Note 3) through it and finally wash the +paper carefully with hot water and transfer any particles of copper +which may be left on it to the Erlenmeyer flask. Boil to expel the +bromine. Add concentrated ammonia drop by drop until the appearance of +a deep blue coloration indicates an excess. Boil until the deep blue +is displaced by a light bluish green coloration, or until brown stains +form on the sides of the flask. Add 10 cc. of strong acetic acid (Note +4) and cool under the water tap. Add a solution containing about 3 +grams of potassium iodide, as in the standardization, and titrate with +thiosulphate solution until the yellow of the liberated iodine is +nearly discharged. Add 1-2 cc. of freshly prepared starch solution and +titrate to the disappearance of the blue color. + +From the data obtained, calculate the percentage of copper (Cu) in the +ore. + +[Note 1: Nitric acid, because of its oxidizing power, is used as a +solvent for the sulphide ores. As a strong acid it will also dissolve +the copper from carbonate ores. The hydrochloric acid is added to +dissolve oxides of iron and to precipitate silver and lead. The +sulphuric acid displaces the other acids, leaving a solution +containing sulphates only. It also, by its dehydrating action, renders +silica from silicates insoluble.] + +[Note 2: Unless proper precautions are taken to insure the correct +concentrations of acid the copper will not precipitate quantitatively +on the aluminium foil; hence care must be taken to follow directions +carefully at this point. Lead and silver have been almost completely +removed as sulphate and chloride respectively, or they too would +be precipitated on the aluminium. Bismuth, though precipitated on +aluminium, has no effect on the analysis. Arsenic and antimony +precipitate on aluminium and would interfere with the titration if +allowed to remain in the lower state of oxidation.] + +[Note 3: Bromine is added to oxidize arsenious and antimonious +compounds from the original sample, and to oxidize nitrous acid formed +by the action of nitric acid on copper and copper sulphide.] + +[Note 4: This reaction can be carried out in the presence of sulphuric +and hydrochloric acids as well as acetic acid, but in the presence +of these strong acids arsenic and antimonic acids may react with the +hydriodic acid produced with the liberation of free iodine, thereby +reversing the process and introducing an error.] + + + + +DETERMINATION OF ANTIMONY IN STIBNITE + + +Stibnite is native antimony sulphide. Nearly pure samples of this +mineral are easily obtainable and should be used for practice, since +many impurities, notably iron, seriously interfere with the accurate +determination of the antimony by iodometric methods. It is, moreover, +essential that the directions with respect to amounts of reagents +employed and concentration of solutions should be followed closely. + +PROCEDURE.--Grind the mineral with great care, and weigh out two +portions of 0.35-0.40 gram into small, dry beakers (100 cc.). +Cover the beakers and pour over the stibnite 5 cc. of concentrated +hydrochloric acid (sp. gr. 1.20) and warm gently on the water bath +(Note 1). When the residue is white, add to each beaker 2 grams of +powdered tartaric acid (Note 2). Warm the solution on the water bath +for ten minutes longer, dilute the solution very cautiously by adding +water in portions of 5 cc., stopping if the solution turns red. It +is possible that no coloration will appear, in which case cautiously +continue the dilution to 125 cc. If a red precipitate or coloration +does appear, warm the solution until it is colorless, and again dilute +cautiously to a total volume of 125 cc. and boil for a minute (Note +3). + +If a white precipitate of the oxychloride separates during dilution +(which should not occur if the directions are followed), it is best to +discard the determination and to start anew. + +Carefully neutralize most of the acid with ammonium hydroxide solution +(sp. gr. 0.96), but leave it distinctly acid (Note 4). Dissolve 3 +grams of sodium bicarbonate in 200 cc. of water in a 500 cc. beaker, +and pour the cold solution of the antimony chloride into this, +avoiding loss by effervescence. Make sure that the solution contains +an excess of the bicarbonate, and then add 1 cc. or 2 cc. of starch +solution and titrate with iodine solution to the appearance of the +blue, avoiding excess (Notes 5 and 6). + +From the corrected volume of the iodine solution required to oxidize +the antimony, calculate the percentage of antimony (Sb) in the +stibnite. + +[Note 1: Antimony chloride is volatile with steam from its +concentrated solutions; hence these solutions must not be boiled until +they have been diluted.] + +[Note 2: Antimony salts, such as the chloride, are readily hydrolyzed, +and compounds such as SbOCl are formed which are often relatively +insoluble; but in the presence of tartaric acid compounds with complex +ions are formed, and these are soluble. An excess of hydrochloric acid +also prevents precipitation of the oxychloride because the H^{+} ions +from the acid lessen the dissociation of the water and thus prevent +any considerable hydrolysis.] + +[Note 3: The action of hydrochloric acid upon the sulphide sets free +sulphureted hydrogen, a part of which is held in solution by the acid. +This is usually expelled by the heating upon the water bath; but if it +is not wholly driven out, a point is reached during dilution at which +the antimony sulphide, being no longer held in solution by the acid, +separates. If the dilution is immediately stopped and the solution +warmed, this sulphide is again brought into solution and at the same +time more of the sulphureted hydrogen is expelled. This procedure must +be continued until the sulphureted hydrogen is all removed, since it +reacts with iodine. If no precipitation of the sulphide occurs, it +is an indication that the sulphureted hydrogen was all expelled on +solution of the stibnite.] + +[Note 4: Ammonium hydroxide is added to neutralize most of the acid, +thus lessening the amount of sodium bicarbonate to be added. The +ammonia should not neutralize all of the acid.] + +[Note 5: The reaction which takes place during titration may be +expressed thus: + +Na_{3}SbO_{3} + 2NaHCO_{3} + I_{2} --> Na_{3}SbO_{4} + 2NaI + H_{2}O + +2CO_{2}.] + +[Note 6: If the end-point is not permanent, that is, if the blue of +the iodo-starch is discharged after standing a few moments, the cause +may be an insufficient quantity of sodium bicarbonate, leaving the +solution slightly acid, or a very slight precipitation of an antimony +compound which is slowly acted upon by the iodine when the latter is +momentarily present in excess. In either case it is better to discard +the analysis and to repeat the process, using greater care in the +amounts of reagents employed.] + + + + +CHLORIMETRY + + +The processes included under the term !chlorimetry! comprise +those employed to determine chlorine, hypochlorites, bromine, and +hypobromites. The reagent employed is sodium arsenite in the presence +of sodium bicarbonate. The reaction in the case of the hypochlorites +is + +NaClO + Na_{3}AsO_{3} --> Na_{3}AsO_{4} + NaCl. + +The sodium arsenite may be prepared from pure arsenious oxide, +as described below, and is stable for considerable periods; but +commercial oxide requires resublimation to remove arsenic sulphide, +which may be present in small quantity. To prepare the solution, +dissolve about 5 grams of the powdered oxide, accurately weighed, +in 10 cc. of a concentrated sodium hydroxide solution, dilute the +solution to 300 cc., and make it faintly acid with dilute hydrochloric +acid. Add 30 grams of sodium bicarbonate dissolved in a little water, +and dilute the solution to exactly 1000 cc. in a measuring flask. +Transfer the solution to a dry liter bottle and mix thoroughly. + +It is possible to dissolve the arsenious oxide directly in a solution +of sodium bicarbonate, with gentle warming, but solution in sodium +hydroxide takes place much more rapidly, and the excess of the +hydroxide is readily neutralized by hydrochloric acid, with subsequent +addition of the bicarbonate to maintain neutrality during the +titration. + +The indicator required for this process is made by dipping strips of +filter paper in a starch solution prepared as described on page 76, +to which 1 gram of potassium iodide has been added. These strips are +allowed to drain and spread upon a watch-glass until dry. When touched +by a drop of the solution the paper turns blue until the hypochlorite +has all been reduced and an excess of the arsenite has been added. + + + + +DETERMINATION OF THE AVAILABLE CHLORINE IN BLEACHING POWDER + + +Bleaching powder consists mainly of a calcium compound which is a +derivative of both hydrochloric and hypochlorous acids. Its formula is +CaClOCl. Its use as a bleaching or disinfecting agent, or as a source +of chlorine, depends upon the amount of hypochlorous acid which it +yields when treated with a stronger acid. It is customary to express +the value of bleaching powder in terms of "available chlorine," by +which is meant the chlorine present as hypochlorite, but not the +chlorine present as chloride. + +PROCEDURE.--Weigh out from a stoppered test tube into a porcelain +mortar about 3.5 grams of bleaching powder (Note 1). Triturate the +powder in the mortar with successive portions of water until it is +well ground and wash the contents into a 500 cc. measuring flask +(Note 2). Fill the flask to the mark with water and shake thoroughly. +Measure off 25 cc. of this semi-solution in a measuring flask, or +pipette, observing the precaution that the liquid removed shall +contain approximately its proportion of suspended matter. + +Empty the flask or pipette into a beaker and wash it out. Run in the +arsenite solution from a burette until no further reaction takes place +on the starch-iodide paper when touched by a drop of the solution of +bleaching powder. Repeat the titration, using a second 25 cc. portion. + +From the volume of solution required to react with the bleaching +powder, calculate the percentage of available chlorine in the latter, +assuming the titration reaction to be that between chlorine and +arsenious oxide: + +As_{4}O_{6} + 4Cl_{2} + 4H_{2}O --> 2As_{2}O_{5} + 8HCl + +Note that only one twentieth of the original weight of bleaching +powder enters into the reaction. + +[Note 1: The powder must be triturated until it is fine, otherwise the +lumps will inclose calcium hypochlorite, which will fail to react with +the arsenious acid. The clear supernatant liquid gives percentages +which are below, and the sediment percentages which are above, the +average. The liquid measured off should, therefore, carry with it its +proper proportion of the sediment, so far as that can be brought about +by shaking the solution just before removal of the aliquot part for +titration.] + +[Note 2: Bleaching powder is easily acted upon by the carbonic acid in +the air, which liberates the weak hypochlorous acid. This, of course, +results in a loss of available chlorine. The original material for +analysis should be kept in a closed container and protected form the +air as far as possible. It is difficult to obtain analytical samples +which are accurately representative of a large quantity of the +bleaching powder. The procedure, as outlined, will yield results which +are sufficiently exact for technical purposes.] + + + + +III. PRECIPITATION METHODS + + + + +DETERMINATION OF SILVER BY THE THIOCYANATE PROCESS + + +The addition of a solution of potassium or ammonium thiocyanate to one +of silver in nitric acid causes a deposition of silver thiocyanate as +a white, curdy precipitate. If ferric nitrate is also present, the +slightest excess of the thiocyanate over that required to combine with +the silver is indicated by the deep red which is characteristic of the +thiocyanate test for iron. + +The reactions involved are: + +AgNO_{3} + KSCN --> AgSCN + KNO_{3}, +3KSCN + Fe(NO_{3})_{3} --> Fe(SCN)_{3} + 3KNO_{3}. + +The ferric thiocyanate differs from the great majority of salts in +that it is but very little dissociated in aqueous solutions, and the +characteristic color appears to be occasioned by the formation of the +un-ionized ferric salt. + +The normal solution of potassium thiocyanate should contain an amount +of the salt per liter of solution which would yield sufficient +(CNS)^{-} to combine with one gram of hydrogen to form HCNS, i.e., +a gram-molecular weight of the salt or 97.17 grams. If the ammonium +thiocyanate is used, the amount is 76.08 grams. To prepare the +solution for this determination, which should be approximately 0.05 +N, dissolve about 5 grams of potassium thiocyanate, or 4 grams of +ammonium thiocyanate, in a small amount of water; dilute this solution +to 1000 cc. in a liter bottle and mix as usual. + +Prepare 20 cc. of a saturated solution of ferric alum and add 5 cc. of +dilute nitric acid (sp. gr. 1.20). About 5 cc. of this solution should +be used as an indicator. + + +STANDARDIZATION + +PROCEDURE.--Crush a small quantity of silver nitrate crystals in a +mortar (Note 1). Transfer them to a watch-glass and dry them for an +hour at 110 deg.C., protecting them from dust or other organic matter +(Note 2). Weigh out two portions of about 0.5 gram each and dissolve +them in 50 cc. of water. Add 10 cc. of dilute nitric acid which has +been recently boiled to expel the lower oxides of nitrogen, if any, +and then add 5 cc. of the indicator solution. Run in the thiocyanate +solution from a burette, with constant stirring, allowing the +precipitate to settle occasionally to obtain an exact recognition +of the end-point, until a faint red tinge can be detected in the +solution. + +From the data obtained, calculate the relation of the thiocyanate +solution to the normal. + +[Note 1: The thiocyanate cannot be accurately weighed; its solutions +must, therefore, be standardized against silver nitrate (or pure +silver), either in the form of a standard solution or in small, +weighed portions.] + +[Note 2: The crystals of silver nitrate sometimes inclose water which +is expelled on drying. If the nitrate has come into contact with +organic bodies it suffers a reduction and blackens during the heating. + +It is plain that a standard solution of silver nitrate (made by +weighing out the crystals) is convenient or necessary if many +titrations of this nature are to be made. In the absence of such a +solution the liability of passing the end-point is lessened by setting +aside a small fraction of the silver solution, to be added near the +close of the titration.] + + +DETERMINATION OF SILVER IN COIN + +PROCEDURE.-- Weigh out two portions of the coin of about 0.5 gram +each. Dissolve them in 15 cc. of dilute nitric acid (sp. gr. 1.2) and +boil until all the nitrous compounds are expelled (Note 1). Cool the +solution, dilute to 50 cc., and add 5 cc. of the indicator solution, +and titrate with the thiocyanate to the appearance of the faint red +coloration (Note 2). + +From the corrected volume of the thiocyanate solution required, +calculate the percentage of silver in the coin. + +[Note 1: The reaction with silver may be carried out in nitric acid +solutions and in the presence of copper, if the latter does not exceed +70 per cent. Above that percentage it is necessary to add silver in +known quantity to the solution. The liquid must be cold at the time of +titration and entirely free from nitrous compounds, as these sometimes +cause a reddening of the indicator solution. All utensils, distilled +water, the nitric acid and the beakers must be free from chlorides, +as the presence of these will cause precipitation of silver chloride, +thereby introducing an error.] + +[Note 2: The solution containing the silver precipitate, as well as +those from the standardization, should be placed in the receptacle for +"silver residues" as a matter of economy.] + + + + +PART III + +GRAVIMETRIC ANALYSIS + + + + +GENERAL DIRECTIONS + + +Gravimetric analyses involve the following principal steps: first, the +weighing of the sample; second, the solution of the sample; third, the +separation of some substance from solution containing, or bearing a +definite relation to, the constituent to be measured, under conditions +which render this separation as complete as possible; and finally, +the segregation of that substance, commonly by filtration, and the +determination of its weight, or that of some stable product formed +from it on ignition. For example, the gravimetric determination of +aluminium is accomplished by solution of the sample, by precipitation +in the form of hydroxide, collection of the hydroxide upon a filter, +complete removal by washing of all foreign soluble matter, and the +burning of the filter and ignition of the precipitate to aluminium +oxide, in which condition it is weighed. + +Among the operations which are common to nearly all gravimetric +analyses are precipitation, washing of precipitates, ignition of +precipitates, and the use of desiccators. In order to avoid burdensome +repetitions in the descriptions of the various gravimetric procedures +which follow, certain general instructions are introduced at this +point. These instructions must, therefore, be considered to be as much +a part of all subsequent procedures as the description of apparatus, +reagents, or manipulations. + +The analytical balance, the fundamentally important instrument in +gravimetric analysis, has already been described on pages 11 to 15. + + +PRECIPITATION + +For successful quantitative precipitations those substances are +selected which are least soluble under conditions which can be easily +established, and which separate from solution in such a state that +they can be filtered readily and washed free from admixed material. +In general, the substances selected are the same as those already +familiar to the student of Qualitative Analysis. + +When possible, substances are selected which separate in crystalline +form, since such substances are less likely to clog the pores of +filter paper and can be most quickly washed. In order to increase the +size of the crystals, which further promotes filtration and washing, +it is often desirable to allow a precipitate to remain for some time +in contact with the solution from which it has separated. The solution +is often kept warm during this period of "digestion." The small +crystals gradually disappear and the larger crystals increase in size, +probably as the result of the force known as surface tension, which +tends to reduce the surface of a given mass of material to a minimum, +combined with a very slightly greater solubility of small crystals as +compared with the larger ones. + +Amorphous substances, such as ferric hydroxide, aluminium hydroxide, +or silicic acid, separate in a gelatinous form and are relatively +difficult to filter and wash. Substances of this class also exhibit +a tendency to form, with pure water, what are known as colloidal +solutions. To prevent this as far as possible, they are washed with +solutions of volatile salts, as will be described in some of the +following procedures. + +In all precipitations the reagent should be added slowly, with +constant stirring, and should be hot when circumstances permit. +The slow addition is less likely to occasion contamination of the +precipitate by the inclosure of other substances which may be in the +solution, or of the reagent itself. + + +FUNNELS AND FILTERS + +Filtration in analytical processes is most commonly effected through +paper filters. In special cases these may be advantageously replaced +by an asbestos filter in a perforated porcelain or platinum crucible, +commonly known, from its originator, as a "Gooch filter." The +operation and use of a filter of this type is described on page 103. +Porous crucibles of a material known as alundum may also be employed +to advantage in special cases. + +The glass funnels selected for use with paper filters should have an +angle as near 60 deg. as possible, and a narrow stem about six inches in +length. The filters employed should be washed filters, i.e., those +which have been treated with hydrochloric and hydrofluoric acids, and +which on incineration leave a very small and definitely known weight +of ash, generally about .00003 gram. Such filters are readily +obtainable on the market. + +The filter should be carefully folded to fit the funnel according to +either of the two well-established methods described in the Appendix. +It should always be placed so that the upper edge of the paper +is about one fourth inch below the top of the funnel. Under no +circumstances should the filter extend above the edge of the funnel, +as it is then utterly impossible to effect complete washing. + +To test the efficiency of the filter, fill it with distilled water. +This water should soon fill the stem completely, forming a continuous +column of liquid which, by its hydrostatic pressure, produces a gentle +suction, thus materially promoting the rapidity of filtration. Unless +the filter allows free passage of water under these conditions, it is +likely to give much trouble when a precipitate is placed upon it. + +The use of a suction pump to promote filtration is rarely altogether +advantageous in quantitative analysis, if paper filters are employed. +The tendency of the filter to break, unless the point of the filter +paper is supported by a perforated porcelain cone or a small "hardened +filter" of parchment, and the tendency of the precipitates to pass +through the pores of the filter, more than compensate for the possible +gain in time. On the other hand, filtration by suction may be useful +in the case of precipitates which do not require ignition before +weighing, or in the case of precipitates which are to be discarded +without weighing. This is best accomplished with the aid of the +special apparatus called a Gooch filter referred to above. + + +FILTRATION AND WASHING OF PRECIPITATES + +Solutions should be filtered while hot, as far as possible, since +the passage of a liquid through the pores of a filter is retarded by +friction, and this, for water at 100 deg.C., is less than one sixth of the +resistance at 0 deg.C. + +When the filtrate is received in a beaker, the stem of the funnel +should touch the side of the receiving vessel to avoid loss by +spattering. Neglect of this precaution is a frequent source of error. + +The vessels which contain the initial filtrate should !always! be +replaced by clean ones, properly labeled, before the washing of a +precipitate begins. In many instances a finely divided precipitate +which shows no tendency to pass through the filter at first, while the +solution is relatively dense, appears at once in the washings. Under +such conditions the advantages accruing from the removal of the first +filtrate are obvious, both as regards the diminished volume requiring +refiltration, and also the smaller number of washings subsequently +required. + +Much time may often be saved by washing precipitates by decantation, +i.e., by pouring over them, while still in the original vessel, +considerable volumes of wash-water and allowing them to settle. The +supernatant, clear wash-water is then decanted through the filter, +so far as practicable without disturbing the precipitate, and a new +portion of wash-water is added. This procedure can be employed to +special advantage with gelatinous precipitates, which fill up the +pores of the filter paper. As the medium from which the precipitate +is to settle becomes less dense it subsides less readily, and it +ultimately becomes necessary to transfer it to the filter and complete +the washing there. + +A precipitate should never completely fill a filter. The wash-water +should be applied at the top of the filter, above the precipitate. +It may be shown mathematically that the washing is most !rapidly! +accomplished by filling the filter well to the top with wash-water +each time, and allowing it to drain completely after each addition; +but that when a precipitate is to be washed with the !least possible +volume! of liquid the latter should be applied in repeated !small! +quantities. + +Gelatinous precipitates should not be allowed to dry before complete +removal of foreign matter is effected. They are likely to shrink and +crack, and subsequent additions of wash-water pass through these +channels only. + +All filtrates and wash-waters without exception must be properly +tested. !This lies at the foundation of accurate work!, and the +student should clearly understand that it is only by the invariable +application of this rule that assurance of ultimate reliability can +be secured. Every original filtrate must be tested to prove complete +precipitation of the compound to be separated, and the wash-waters +must also be tested to assure complete removal of foreign material. In +testing the latter, the amount first taken should be but a few +drops if the filtrate contains material which is to be subsequently +determined. When, however, the washing of the filter and precipitate +is nearly completed the amount should be increased, and for the final +test not less than 3 cc. should be used. + +It is impossible to trust to one's judgment with regard to the washing +of precipitates; the washings from !each precipitate! of a series +simultaneously treated must be tested, since the rate of washing will +often differ materially under apparently similar conditions, !No +exception can ever be made to this rule!. + +The habit of placing a clean common filter paper under the receiving +beaker during filtration is one to be commended. On this paper a +record of the number of washings can very well be made as the portions +of wash-water are added. + +It is an excellent practice, when possible, to retain filtrates and +precipitates until the completion of an analysis, in order that, in +case of question, they may be examined to discover sources of error. + +For the complete removal of precipitates from containing vessels, it +is often necessary to rub the sides of these vessels to loosen the +adhering particles. This can best be done by slipping over the end of +a stirring rod a soft rubber device sometimes called a "policeman." + + +DESICCATORS + +Desiccators should be filled with fused, anhydrous calcium chloride, +over which is placed a clay triangle, or an iron triangle covered with +silica tubes, to support the crucible or other utensils. The cover of +the desiccator should be made air-tight by the use of a thin coating +of vaseline. + +Pumice moistened with concentrated sulphuric acid may be used in place +of the calcium chloride, and is essential in special cases; but for +most purposes the calcium chloride, if renewed occasionally and not +allowed to cake together, is practically efficient and does not slop +about when the desiccator is moved. + +Desiccators should never remain uncovered for any length of time. The +dehydrating agents rapidly lose their efficiency on exposure to the +air. + + +CRUCIBLES + +It is often necessary in quantitative analysis to employ fluxes to +bring into solution substances which are not dissolved by acids. The +fluxes in most common use are sodium carbonate and sodium or potassium +acid sulphate. In gravimetric analysis it is usually necessary to +ignite the separated substance after filtration and washing, in order +to remove moisture, or to convert it through physical or chemical +changes into some definite and stable form for weighing. Crucibles +to be used in fusion processes must be made of materials which will +withstand the action of the fluxes employed, and crucibles to be used +for ignitions must be made of material which will not undergo any +permanent change during the ignition, since the initial weight of the +crucible must be deducted from the final weight of the crucible and +product to obtain the weight of the ignited substance. The three +materials which satisfy these conditions, in general, are platinum, +porcelain, and silica. + +Platinum crucibles have the advantage that they can be employed at +high temperatures, but, on the other hand, these crucibles can never +be used when there is a possibility of the reduction to the metallic +state of metals like lead, copper, silver, or gold, which would alloy +with and ruin the crucible. When platinum crucibles are used with +compounds of arsenic or phosphorus, special precautions are necessary +to prevent damage. This statement applies to both fusions and +ignitions. + +Fusions with sodium carbonate can be made only in platinum, since +porcelain or silica crucibles are attacked by this reagent. Acid +sulphate fusions, which require comparatively low temperatures, can +sometimes be made in platinum, although platinum is slightly attacked +by the flux. Porcelain or silica crucibles may be used with acid +fluxes. + +Silica crucibles are less likely to crack on heating than porcelain +crucibles on account of their smaller coefficient of expansion. +Ignition of substances not requiring too high a temperature may be +made in porcelain or silica crucibles. + +Iron, nickel or silver crucibles are used in special cases. + +In general, platinum crucibles should be used whenever such use is +practicable, and this is the custom in private, research or commercial +laboratories. Platinum has, however, become so valuable that it is +liable to theft unless constantly under the protection of the user. As +constant protection is often difficult in instructional laboratories, +it is advisable, in order to avoid serious monetary losses, to use +porcelain or silica crucibles whenever these will give satisfactory +service. When platinum utensils are used the danger of theft should +always be kept in mind. + + +PREPARATION OF CRUCIBLES FOR USE + +All crucibles, of whatever material, must always be cleaned, ignited +and allowed to cool in a desiccator before weighing, since all bodies +exposed to the air condense on their surfaces a layer of moisture +which increases their weight. The amount and weight of this moisture +varies with the humidity of the atmosphere, and the latter may change +from hour to hour. The air in the desiccator (see above) is kept at +a constant and low humidity by the drying agent which it contains. +Bodies which remain in a desiccator for a sufficient time (usually +20-30 minutes) retain, therefore, on their surfaces a constant weight +of moisture which is the same day after day, thus insuring constant +conditions. + +Hot objects, such as ignited crucibles, should be allowed to cool in +the air until, when held near the skin, but little heat is noticeable. +If this precaution is not taken, the air within the desiccator is +strongly heated and expands before the desiccator is covered. As the +temperature falls, the air contracts, causing a reduction of air +pressure within the covered vessel. When the cover is removed (which +is often rendered difficult) the inrush of air from the outside may +sweep light particles out of a crucible, thus ruining an entire +analysis. + +Constant heating of platinum causes a slight crystallization of the +surface which, if not removed, penetrates into the crucible. Gentle +polishing of the surface destroys the crystalline structure and +prevents further damage. If sea sand is used for this purpose, great +care is necessary to keep it from the desk, since beakers are easily +scratched by it, and subsequently crack on heating. + +Platinum crucibles stained in use may often be cleaned by the fusion +in them of potassium or sodium acid sulphate, or by heating with +ammonium chloride. If the former is used, care should be taken not +to heat so strongly as to expel all of the sulphuric acid, since the +normal sulphates sometimes expand so rapidly on cooling as to split +the crucible. The fused material should be poured out, while hot, on +to a !dry! tile or iron surface. + + +IGNITION OF PRECIPITATES + +Most precipitates may, if proper precautions are taken, be ignited +without previous drying. If, however, such precipitates can be dried +without loss of time to the analyst (as, for example, over night), it +is well to submit them to this process. It should, nevertheless, be +remembered that a partially dried precipitate often requires more care +during ignition than a thoroughly moist one. + +The details of the ignition of precipitates vary so much with the +character of the precipitate, its moisture content, and temperature to +which it is to be heated, that these details will be given under the +various procedures which follow. + + + + +DETERMINATION OF CHLORINE IN SODIUM CHLORIDE + + +!Method A. With the Use of a Gooch Filter! + +PROCEDURE.--Carefully clean a weighing-tube containing the sodium +chloride, handling it as little as possible with the moist fingers, +and weigh it accurately to 0.0001 gram, recording the weight at once +in the notebook (see Appendix). Hold the tube over the top of a beaker +(200-300 cc.), and cautiously remove the stopper, noting carefully +that no particles fall from it, or from the tube, elsewhere than into +the beaker. Pour out a small portion of the chloride, replace the +stopper, and determine by approximate weighing how much has been +removed. Continue this procedure until 0.25-0.30 gram has been taken +from the tube, then weigh accurately and record the weight beneath the +first in the notebook. The difference of the two weights represents +the weight of the chloride taken for analysis. Again weigh a second +portion of 0.25-0.30 gram into a second beaker of the same size as the +first. The beakers should be plainly marked to correspond with the +entries in the notebook. Dissolve each portion of the chloride in 150 +cc. of distilled water and add about ten drops of dilute nitric acid +(sp. gr. 1.20) (Note 2). Calculate the volume of silver nitrate +solution required to effect complete precipitation in each case, +and add slowly about 5 cc. in excess of that amount, with constant +stirring. Heat the solutions cautiously to boiling, stirring +occasionally, and continue the heating and stirring until the +precipitates settle promptly, leaving a nearly clear supernatant +liquid (Note 3). This heating should not take place in direct sunlight +(Note 4). The beaker should be covered with a watch-glass, and both +boiling and stirring so regulated as to preclude any possibility of +loss of material. Add to the clear liquid one or two drops of silver +nitrate solution, to make sure that an excess of the reagent is +present. If a precipitate, or cloudiness, appears as the drops fall +into the solution, heat again, and stir until the whole precipitate +has coagulated. The solution is then ready for filtration. + +Prepare a Gooch filter as follows: Fold over the top of a Gooch funnel +(Fig. 2) a piece of rubber-band tubing, such as is known as "bill-tie" +tubing, and fit into the mouth of the funnel a perforated porcelain +crucible (Gooch crucible), making sure that when the crucible is +gently forced into the mouth of the funnel an airtight joint results. +(A small 1 or 1-1/4-inch glass funnel may be used, in which case the +rubber tubing is stretched over the top of the funnel and then drawn +up over the side of the crucible until an air-tight joint is secured.) + +[ILLUSTRATION: FIG. 2] + +Fit the funnel into the stopper of a filter bottle, and connect the +filter bottle with the suction pump. Suspend some finely divided +asbestos, which has been washed with acid, in 20 to 30 cc. of water +(Note 1); allow this to settle, pour off the very fine particles, and +then pour some of the mixture cautiously into the crucible until an +even felt of asbestos, not over 1/32 inch in thickness, is formed. A +gentle suction must be applied while preparing this felt. Wash the +felt thoroughly by passing through it distilled water until all fine +or loose particles are removed, increasing the suction at the last +until no more water can be drawn out of it; place on top of the felt +the small, perforated porcelain disc and hold it in place by pouring a +very thin layer of asbestos over it, washing the whole carefully; +then place the crucible in a small beaker, and place both in a drying +closet at 100-110 deg.C. for thirty to forty minutes. Cool the crucible +in a desiccator, and weigh. Heat again for twenty to thirty minutes, +cool, and again weigh, repeating this until the weight is constant +within 0.0003 gram. The filter is then ready for use. + +Place the crucible in the funnel, and apply a gentle suction, !after +which! the solution to be filtered may be poured in without disturbing +the asbestos felt. When pouring liquid onto a Gooch filter hold the +stirring-rod at first well down in the crucible, so that the liquid +does not fall with any force upon the asbestos, and afterward keep the +crucible will filled with the solution. + +Pour the liquid above the silver chloride slowly onto the filter, +leaving the precipitate in the beaker as far as possible. Wash the +precipitate twice by decantation with warm water; then transfer it +to the filter with the aid of a stirring-rod with a rubber tip and a +stream from the wash-bottle. + +Examine the first portions of the filtrate which pass through the +filter with great care for asbestos fibers, which are most likely to +be lost at this point. Refilter the liquid if any fibers are visible. +Finally, wash the precipitate thoroughly with warm water until free +from soluble silver salts. To test the washings, disconnect the +suction at the flask and remove the funnel or filter tube from the +suction flask. Hold the end of the tube over the mouth of a small test +tube and add from a wash-bottle 2-3 cc. of water. Allow the water to +drip through into the test tube and add a drop of dilute hydrochloric +acid. No precipitate or cloud should form in the wash-water (Note 16). +Dry the filter and contents at 100-110 deg.C. until the weight is constant +within 0.0003 gram, as described for the preparation of the filter. +Deduct the weight of the dry crucible from the final weight, and from +the weight of silver chloride thus obtained calculate the percentage +of chlorine in the sample of sodium chloride. + +[Note 1: The washed asbestos for this type of filter is prepared by +digesting in concentrated hydrochloric acid, long-fibered asbestos +which has been cut in pieces of about 0.5 cm. in length. After +digestion, the asbestos is filtered off on a filter plate and washed +with hot, distilled water until free from chlorides. A small portion +of the asbestos is shaken with water, forming a thin suspension, which +is bottled and kept for use.] + +[Note 2: The nitric acid is added before precipitation to lessen the +tendency of the silver chloride to carry down with it other substances +which might be precipitated from a neutral solution. A large excess of +the acid would exert a slight solvent action upon the chloride.] + +[Note 3: The solution should not be boiled after the addition of the +nitric acid before the presence of an excess of silver nitrate is +assured, since a slight interaction between the nitric acid and the +sodium chloride is possible, by which a loss of chlorine, either as +such or as hydrochloric acid, might ensue. The presence of an excess +of the precipitant can usually be recognized at the time of its +addition, by the increased readiness with which the precipitate +coagulates and settles.] + +[Note 4: The precipitate should not be exposed to strong sunlight, +since under those conditions a reduction of the silver chloride ensues +which is accompanied by a loss of chlorine. The superficial alteration +which the chloride undergoes in diffused daylight is not sufficient +to materially affect the accuracy of the determination. It should be +noted, however, that a slight error does result from the effect of +light upon the silver chloride precipitate and in cases in which the +greatest obtainable accuracy is required, the procedure described +under "Method B" should be followed, in which this slight reduction of +the silver chloride is corrected by subsequent treatment with nitric +and hydrochloric acids.] + +[Note 5: The asbestos used in the Gooch filter should be of the finest +quality and capable of division into minute fibrous particles. A +coarse felt is not satisfactory.] + +[Note 6: The precipitate must be washed with warm water until it is +absolutely free from silver and sodium nitrates. It may be assumed +that the sodium salt is completely removed when the wash-water shows +no evidence of silver. It must be borne in mind that silver chloride +is somewhat soluble in hydrochloric acid, and only a single drop +should be added. The washing should be continued until no cloudiness +whatever can be detected in 3 cc. of the washings. + +Silver chloride is but slightly soluble in water. The solubility +varies with its physical condition within small limits, and is +about 0.0018 gram per liter at 18 deg.C. for the curdy variety usually +precipitated. The chloride is also somewhat soluble in solutions of +many chlorides, in solutions of silver nitrate, and in concentrated +nitric acid. + +As a matter of economy, the filtrate, which contains whatever silver +nitrate was added in excess, may be set aside. The silver can be +precipitated as chloride and later converted into silver nitrate.] + +[Note 7: The use of the Gooch filter commends itself strongly when a +considerable number of halogen determinations are to be made, since +successive portions of the silver halides may be filtered on the same +filter, without the removal of the preceding portions, until the +crucible is about two thirds filled. If the felt is properly prepared, +filtration and washing are rapidly accomplished on this filter, and +this, combined with the possibility of collecting several precipitates +on the same filter, is a strong argument in favor of its use with any +but gelatinous precipitates.] + + +!Method B. With the Use of a Paper Filter! + +PROCEDURE.--Weigh out two portions of sodium chloride of about +0.25-0.3 gram each and proceed with the precipitation of the silver +chloride as described under Method A above. When the chloride is ready +for filtration prepare two 9 cm. washed paper filters (see Appendix). +Pour the liquid above the precipitates through the filters, wash twice +by decantation and transfer the precipitates to the filters, finally +washing them until free from silver solution as described. The funnel +should then be covered with a moistened filter paper by stretching it +over the top and edges, to which it will adhere on drying. It should +be properly labeled with the student's name and desk number, and then +placed in a drying closet, at a temperature of about 100-110 deg.C., until +completely dry. + +The perfectly dry filter is then opened over a circular piece of +clean, smooth, glazed paper about six inches in diameter, placed upon +a larger piece about twelve inches in diameter. The precipitate is +removed from the filter as completely as possible by rubbing the sides +gently together, or by scraping them cautiously with a feather which +has been cut close to the quill and is slightly stiff (Note 1). In +either case, care must be taken not to rub off any considerable +quantity of the paper, nor to lose silver chloride in the form of +dust. Cover the precipitate on the glazed paper with a watch-glass to +prevent loss of fine particles and to protect it from dust from the +air. Fold the filter paper carefully, roll it into a small cone, and +wind loosely around !the top! a piece of small platinum wire (Note 2). +Hold the filter by the wire over a small porcelain crucible (which has +been cleaned, ignited, cooled in a desiccator, and weighed), ignite +it, and allow the ash to fall into the crucible. Place the crucible +upon a clean clay triangle, on its side, and ignite, with a low +flame well at its base, until all the carbon of the filter has been +consumed. Allow the crucible to cool, add two drops of concentrated +nitric acid and one drop of concentrated hydrochloric acid, and heat +!very cautiously!, to avoid spattering, until the acids have been +expelled; then transfer the main portion of the precipitate from the +glazed paper to the cooled crucible, placing the latter on the larger +piece of glazed paper and brushing the precipitate from the +smaller piece into it, sweeping off all particles belonging to the +determination. + +Moisten the precipitate with two drops of concentrated nitric acid and +one drop of concentrated hydrochloric acid, and again heat with great +caution until the acids are expelled and the precipitate is white, +when the temperature is slowly raised until the silver chloride just +begins to fuse at the edges (Note 3). The crucible is then cooled in a +desiccator and weighed, after which the heating (without the addition +of acids) is repeated, and it is again weighed. This must be continued +until the weight is constant within 0.0003 gram in two consecutive +weighings. Deduct the weight of the crucible, and calculate the +percentage of chlorine in the sample of sodium chloride taken for +analysis. + +[Note 1: The separation of the silver chloride from the filter is +essential, since the burning carbon of the paper would reduce a +considerable quantity of the precipitate to metallic silver, and its +complete reconversion to the chloride within the crucible, by means of +acids, would be accompanied by some difficulty. The small amount of +silver reduced from the chloride adhering to the filter paper after +separating the bulk of the precipitate, and igniting the paper +as prescribed, can be dissolved in nitric acid, and completely +reconverted to chloride by hydrochloric acid. The subsequent addition +of the two acids to the main portion of the precipitate restores the +chlorine to any chloride which may have been partially reduced by the +sunlight. The excess of the acids is volatilized by heating.] + +[Note 2: The platinum wire is wrapped around the top of the filter +during its incineration to avoid contact with any reduced silver from +the reduction of the precipitate. If the wire were placed nearer the +apex, such contact could hardly be avoided.] + +[Note 3: Silver chloride should not be heated to complete fusion, +since a slight loss by volatilization is possible at high +temperatures. The temperature of fusion is not always sufficient +to destroy filter shreds; hence these should not be allowed to +contaminate the precipitate.] + + + + +DETERMINATION OF IRON AND OF SULPHUR IN FERROUS AMMONIUM SULPHATE, + +FESO_{4}.(NH_{4})_{2}SO_{4}.6H_{2}O + + +DETERMINATION OF IRON + +PROCEDURE.--Weigh out into beakers (200-250 cc.) two portions of the +sample (Note 1) of about 1 gram each and dissolve these in 50 cc. of +water, to which 1 cc. of dilute hydrochloric acid (sp. gr. 1.12) has +been added (Note 2). Heat the solution to boiling, and while at the +boiling point add concentrated nitric acid (sp. gr. 1.42), !drop by +drop! (noting the volume used), until the brown coloration, which +appears after the addition of a part of the nitric acid, gives place +to a yellow or red (Note 3). Avoid a large excess of nitric acid, but +be sure that the action is complete. Pour this solution cautiously +into about 200 cc. of water, containing a slight excess of ammonia. +Calculate for this purpose the amount of aqueous ammonia required to +neutralize the hydrochloric and nitric acids added (see Appendix for +data), and also to precipitate the iron as ferric hydroxide from the +weight of the ferrous ammonium sulphate taken for analysis, assuming +it to be pure (Note 4). The volume thus calculated will be in excess +of that actually required for precipitation, since the acids are in +part consumed in the oxidation process, or are volatilized. Heat the +solution to boiling, and allow the precipitated ferric hydroxide to +settle. Decant the clear liquid through a washed filter (9 cm.), +keeping as much of the precipitate in the beaker as possible. Wash +twice by decantation with 100 cc. of hot water. Reserve the filtrate. +Dissolve the iron from the filter with hot, dilute hydrochloric acid +(sp. gr. 1.12), adding it in small portions, using as little as +possible and noting the volume used. Collect the solution in the +beaker in which precipitation took place. Add 1 cc. of nitric acid +(sp. gr. 1.42), boil for a few moments, and again pour into a +calculated excess of ammonia. + +Wash the precipitate twice by decantation, and finally transfer it to +the original filter. Wash continuously with hot water until finally +3 cc. of the washings, acidified with nitric acid (Note 5), show +no evidences of the presence of chlorides when tested with silver +nitrate. The filtrate and washings are combined with those from the +first precipitation and treated for the determination of sulphur, as +prescribed on page 112. + +[Note 1: If a selection of pure material for analysis is to be made, +crystals which are cloudy are to be avoided on account of loss of +water of crystallization; and also those which are red, indicating +the presence of ferric iron. If, on the other hand, the value of an +average sample of material is desired, it is preferable to grind the +whole together, mix thoroughly, and take a sample from the mixture for +analysis.] + +[Note 2: When aqueous solutions of ferrous compounds are heated in the +air, oxidation of the Fe^{++} ions to Fe^{+++} ions readily occurs in +the absence of free acid. The H^{+} and OH^{-} ions from water are +involved in the oxidation process and the result is, in effect, the +formation of some ferric hydroxide which tends to separate. Moreover, +at the boiling temperature, the ferric sulphate produced by the +oxidation hydrolyzes in part with the formation of a basic ferric +sulphate, which also tends to separate from solution. The addition of +the hydrochloric acid prevents the formation of ferric hydroxide, and +so far reduces the ionization of the water that the hydrolysis of the +ferric sulphate is also prevented, and no precipitation occurs on +heating.] + +[Note 3: The nitric acid, after attaining a moderate strength, +oxidizes the Fe^{++} ions to Fe^{+++} ions with the formation of an +intermediate nitroso-compound similar in character to that formed in +the "ring-test" for nitrates. The nitric oxide is driven out by heat, +and the solution then shows by its color the presence of ferric +compounds. A drop of the oxidized solution should be tested on +a watch-glass with potassium ferricyanide, to insure a complete +oxidation. This oxidation of the iron is necessary, since Fe^{++} ions +are not completely precipitated by ammonia. + +The ionic changes which are involved in this oxidation are perhaps +most simply expressed by the equation + +3Fe^{++} + NO_{3}^{-}+ 4H^{+} --> 3Fe^{+++} + 2H_{2}O + NO, + +the H^{+} ions coming from the acid in the solution, in this case +either the nitric or the hydrochloric acid. The full equation on which +this is based may be written thus: + +6FeSO_{4} + 2HNO_{3} + 6HCl --> 2Fe_{2}(SO_{4})_{3} + 2FeCl_{3} + 2NO ++ 4H_{2}O, + +assuming that only enough nitric acid is added to complete the +oxidation.] + +[Note 4: The ferric hydroxide precipitate tends to carry down some +sulphuric acid in the form of basic ferric sulphate. This tendency is +lessened if the solution of the iron is added to an excess of OH^{-} +ions from the ammonium hydroxide, since under these conditions +immediate and complete precipitation of the ferric hydroxide ensues. +A gradual neutralization with ammonia would result in the local +formation of a neutral solution within the liquid, and subsequent +deposition of a basic sulphate as a consequence of a local deficiency +of OH^{-} ions from the NH_{4}OH and a partial hydrolysis of the +ferric salt. Even with this precaution the entire absence of sulphates +from the first iron precipitate is not assured. It is, therefore, +redissolved and again thrown down by ammonia. The organic matter of +the filter paper may occasion a partial reduction of the iron during +solution, with consequent possibility of incomplete subsequent +precipitation with ammonia. The nitric acid is added to reoxidize this +iron. + +To avoid errors arising from the solvent action of ammoniacal +liquids upon glass, the iron precipitate should be filtered without +unnecessary delay.] + +[Note 5: The washings from the ferric hydroxide are acidified with +nitric acid, before testing with silver nitrate, to destroy the +ammonia which is a solvent of silver chloride. + +The use of suction to promote filtration and washing is permissible, +though not prescribed. The precipitate should not be allowed to dry +during the washing.] + + +!Ignition of the Iron Precipitate! + +Heat a platinum or porcelain crucible, cool it in a desiccator and +weigh, repeating until a constant weight is obtained. + +Fold the top of the filter paper over the moist precipitate of ferric +hydroxide and transfer it cautiously to the crucible. Wipe the inside +of the funnel with a small fragment of washed filter paper, if +necessary, and place the paper in the crucible. + +Incline the crucible on its side, on a triangle supported on a +ring-stand, and stand the cover on edge at the mouth of the crucible. +Place a burner below the front edge of the crucible, using a low flame +and protecting it from drafts of air by means of a chimney. The heat +from the burner is thus reflected into the crucible and dries +the precipitate without danger of loss as the result of a sudden +generation of steam within the mass of ferric hydroxide. As the drying +progresses the burner may be gradually moved toward the base of the +crucible and the flame increased until the paper of the filter begins +to char and finally to smoke, as the volatile matter is expelled. This +is known as "smoking off" a filter, and the temperature should not be +raised sufficiently high during this process to cause the paper to +ignite, as the air currents produced by the flame of the blazing paper +may carry away particles of the precipitate. + +When the paper is fully charred, move the burner to the base of the +crucible and raise the temperature to the full heat of the burner for +fifteen minutes, with the crucible still inclined on its side, but +without the cover (Note 1). Finally set the crucible upright in the +triangle, cover it, and heat at the full temperature of a blast lamp +or other high temperature burner. Cool and weigh in the usual manner +(Note 2). Repeat the strong heating until the weight is constant +within 0.0003 gram. + +From the weight of ferric oxide (Fe_{2}O_{3}) calculate the percentage +of iron (Fe) in the sample (Note 3). + +[Note 1: These directions for the ignition of the precipitate must be +closely followed. A ready access of atmospheric oxygen is of special +importance to insure the reoxidation to ferric oxide of any iron which +may be reduced to magnetic oxide (Fe_{3}O_{4}) during the combustion +of the filter. The final heating over the blast lamp is essential +for the complete expulsion of the last traces of water from the +hydroxide.] + +[Note 2: Ignited ferric oxide is somewhat hygroscopic. On this account +the weighings must be promptly completed after removal from the +desiccator. In all weighings after the first it is well to place the +weights upon the balance-pan before removing the crucible from the +desiccator. It is then only necessary to move the rider to obtain the +weight.] + +[Note 3: The gravimetric determination of aluminium or chromium is +comparable with that of iron just described, with the additional +precaution that the solution must be boiled until it contains but a +very slight excess of ammonia, since the hydroxides of aluminium and +chromium are more soluble than ferric hydroxide. + +The most important properties of these hydroxides, from a quantitative +standpoint, other than those mentioned, are the following: All are +precipitable by the hydroxides of sodium and potassium, but always +inclose some of the precipitant, and should be reprecipitated with +ammonium hydroxide before ignition to oxides. Chromium and aluminium +hydroxides dissolve in an excess of the caustic alkalies and form +anions, probably of the formula AlO_2^{-} and CrO_{2}^{-}. Chromium +hydroxide is reprecipitated from this solution on boiling. When first +precipitated the hydroxides are all readily soluble in acids, but +aluminium hydroxide dissolves with considerable difficulty after +standing or boiling for some time. The precipitation of the hydroxides +is promoted by the presence of ammonium chloride, but is partially +or entirely prevented by the presence of tartaric or citric acids, +glycerine, sugars, and some other forms of soluble organic matter. +The hydroxides yield on ignition an oxide suitable for weighing +(Al_{2}O_{3}, Cr_{2}O_{3}, Fe_{2}O_{3}).] + + + + +DETERMINATION OF SULPHUR + + +PROCEDURE.--Add to the combined filtrates from the ferric hydroxide +about 0.6 gram of anhydrous sodium carbonate; cover the beaker, and +then add dilute hydrochloric acid (sp. gr. 1.12) in moderate excess +and evaporate to dryness on the water bath. Add 10 cc. of concentrated +hydrochloric acid (sp. gr. 1.20) to the residue, and again evaporate +to dryness on the bath. Dissolve the residue in water, filter if not +clear, transfer to a 700 cc. beaker, dilute to about 400 cc., and +cautiously add hydrochloric acid until the solution shows a distinctly +acid reaction (Note 1). Heat the solution to boiling, and add !very +slowly! and with constant stirring, 20 cc. in excess of the calculated +amount of a hot barium chloride solution, containing about 20 grams +BaCl_{2}.2H_{2}O per liter (Notes 2 and 3). Continue the boiling for +about two minutes, allow the precipitate to settle, and decant the +liquid at the end of half an hour (Note 4). Replace the beaker +containing the original filtrate by a clean beaker, wash the +precipitated sulphate by decantation with hot water, and subsequently +upon the filter until it is freed from chlorides, testing the washings +as described in the determination of iron. The filter is then +transferred to a platinum or porcelain crucible and ignited, as +described above, until the weight is constant (Note 5). From the +weight of barium sulphate (BaSO_{4}) obtained, calculate the +percentage of sulphur (S) in the sample. + +[Note 1: Barium sulphate is slightly soluble in hydrochloric acid, +even dilute, probably as a result of the reduction in the degree of +dissociation of sulphuric acid in the presence of the H^{+} ions of +the hydrochloric acid, and possibly because of the formation of a +complex anion made up of barium and chlorine; hence only the smallest +excess should be added over the amount required to acidify the +solution.] + +[Note 2: The ionic changes involved in the precipitation of barium +sulphate are very simple: + +Ba^{++} + SO_{4}^{--} --> [BaSO_{4}] + +This case affords one of the best illustrations of the effect of an +excess of a precipitant in decreasing the solubility of a precipitate. +If the conditions are considered which exist at the moment when just +enough of the Ba^{++} ions have been added to correspond to the +SO_{4}^{--} ions in the solution, it will be seen that nearly all of +the barium sulphate has been precipitated, and that the small amount +which then remains in the solution which is in contact with the +precipitate must represent a saturated solution for the existing +temperature, and that this solution is comparable with a solution of +sugar to which more sugar has been added than will dissolve. It +should be borne in mind that the quantity of barium sulphate in +this !saturated solution is a constant quantity! for the existing +conditions. The dissolved barium sulphate, like any electrolyte, is +dissociated, and the equilibrium conditions may be expressed thus: + +(!Conc'n Ba^{++} x Conc'n SO_{4}^{--})/(Conc'n BaSO_{4}) = Const.!, + +and since !Conc'n BaSO_{4}! for the saturated solution has a constant +value (which is very small), it may be eliminated, when the expression +becomes !Conc'n Ba^{++} x Conc'n SO_{4}^{--} = Const.!, which is +the "solubility product" of BaSO_{4}. If, now, an excess of the +precipitant, a soluble barium salt, is added in the form of a +relatively concentrated solution (the slight change of volume of a few +cubic centimeters may be disregarded for the present discussion) +the concentration of the Ba^{++} ions is much increased, and as a +consequence the !Conc'n SO_{4}! must decrease in proportion if the +value of the expression is to remain constant, which is a requisite +condition if the law of mass action upon which our argument depends +holds true. In other words, SO_{4}^{--} ions must combine with some +of the added Ba^{++} ions to form [BaSO_{4}]; but it will be recalled +that the solution is already saturated with BaSO_{4}, and this freshly +formed quantity must, therefore, separate and add itself to the +precipitate. This is exactly what is desired in order to insure +more complete precipitation and greater accuracy, and leads to the +conclusion that the larger the excess of the precipitant added the +more successful the analysis; but a practical limit is placed upon +the quantity of the precipitant which may be properly added by other +conditions, as stated in the following note.] + +[Note 3: Barium sulphate, in a larger measure than most compounds, +tends to carry down other substances which are present in the solution +from which it separates, even when these other substances are +relatively soluble, and including the barium chloride used as the +precipitant. This is also notably true in the case of nitrates and +chlorates of the alkalies, and of ferric compounds; and, since in this +analysis ammonium nitrate has resulted from the neutralization of the +excess of the nitric acid added to oxidize the iron, it is essential +that this should be destroyed by repeated evaporation with a +relatively large quantity of hydrochloric acid. During evaporation a +mutual decomposition of the two acids takes place, and the nitric acid +is finally decomposed and expelled by the excess of hydrochloric acid. + +Iron is usually found in the precipitate of barium sulphate when +thrown down from hot solutions in the presence of ferric salts. This, +according to Kuster and Thiel (!Zeit. anorg. Chem.!, 22, 424), is due +to the formation of a complex ion (Fe(SO_{4})_{2}) which precipitates +with the Ba^{++} ion, while Richards (!Zeit. anorg. Chem.!, 23, 383) +ascribes it to hydrolytic action, which causes the formation of a +basic ferric complex which is occluded in the barium precipitate. +Whatever the character of the compound may be, it has been shown that +it loses sulphuric anhydride upon ignition, causing low results, even +though the precipitate contains iron. + +The contamination of the barium sulphate by iron is much less in the +presence of ferrous than ferric salts. If, therefore, the sulphur +alone were to be determined in the ferrous ammonium sulphate, the +precipitation by barium might be made directly from an aqueous +solution of the salt, which had been made slightly acid with +hydrochloric acid.] + +[Note 4: The precipitation of the barium sulphate is probably complete +at the end of a half-hour, and the solution may safely be filtered at +the expiration of that time if it is desired to hasten the analysis. + +As already noted, many precipitates of the general character of this +sulphate tend to grow more coarsely granular if digested for some time +with the liquid from which they have separated. It is therefore well +to allow the precipitate to stand in a warm place for several hours, +if practicable, to promote ease of filtration. The filtrate and +washings should always be carefully examined for minute quantities of +the sulphate which may pass through the pores of the filter. This is +best accomplished by imparting to the filtrate a gentle rotary motion, +when the sulphate, if present, will collect at the center of the +bottom of the beaker.] + +[Note 5: A reduction of barium sulphate to the sulphide may very +readily be caused by the reducing action of the burning carbon of the +filter, and much care should be taken to prevent any considerable +reduction from this cause. Subsequent ignition, with ready access +of air, reconverts the sulphide to sulphate unless a considerable +reduction has occurred. In the latter case it is expedient to add one +or two drops of sulphuric acid and to heat cautiously until the excess +of acid is expelled.] + +[Note 6: Barium sulphate requires about 400,000 parts of water for +its solution. It is not decomposed at a red heat but suffers loss, +probably of sulphur trioxide, at a temperature above 900 deg.C.] + + + + +DETERMINATION OF SULPHUR IN BARIUM SULPHATE + + +PROCEDURE.--Weigh out, into platinum crucibles, two portions of about +0.5 gram of the sulphate. Mix each in the crucible with five to six +times its weight of anhydrous sodium carbonate. This can best be done +by placing the crucible on a piece of glazed paper and stirring the +mixture with a clean, dry stirring-rod, which may finally be wiped off +with a small fragment of filter paper, the latter being placed in the +crucible. Cover the crucible and heat until a quiet, liquid fusion +ensues. Remove the burner, and tip the crucible until the fused mass +flows nearly to its mouth. Hold it in that position until the mass has +solidified. When cold, the material may usually be detached in a lump +by tapping the crucible or gently pressing it near its upper edge. If +it still adheres, a cubic centimeter or so of water may be placed in +the cold crucible and cautiously brought to boiling, when the cake +will become loosened and may be removed and placed in about 250 cc. +of hot, distilled water to dissolve. Clean the crucible completely, +rubbing the sides with a rubber-covered stirring-rod, if need be. + +When the fused mass has completely disintegrated and nothing further +will dissolve, decant the solution from the residue of barium +carbonate (Note 1). Pour over the residue 20 cc. of a solution of +sodium carbonate and 10 cc. of water and heat to gentle boiling for +about three minutes (Note 2). Filter off the carbonate and wash it +with hot water, testing the slightly acidified washings for sulphate +and preserving any precipitates which appear in these tests. Acidify +the filtrate with hydrochloric acid until just acid, bring to boiling, +and slowly add hot barium chloride solution, as in the preceding +determination. Add also any tests from the washings in which +precipitates have appeared. Filter, wash, ignite, and weigh. + +From the weight of barium sulphate, calculate the percentage of +sulphur (S) in the sample. + +[Note 1: This alkaline fusion is much employed to disintegrate +substances ordinarily insoluble in acids into two components, one +of which is water soluble and the other acid soluble. The reaction +involved is: + +BaSO_{4} + Na_{2}CO_{3}, --> BaCO_{3}, + Na_{2}SO_{4}. + +As the sodium sulphate is soluble in water, and the barium carbonate +insoluble, a separation between them is possible and the sulphur can +be determined in the water-soluble portion. + +It should be noted that this method can be applied to the purification +of a precipitate of barium sulphate if contaminated by most of the +substances mentioned in Note 3 on page 114. The impurities pass into +the water solution together with the sodium sulphate, but, being +present in such minute amounts, do not again precipitate with the +barium sulphate.] + +[Note 2: The barium carbonate is boiled with sodium carbonate solution +before filtration because the reaction above is reversible; and it is +only by keeping the sodium carbonate present in excess until nearly +all of the sodium sulphate solution has been removed by filtration +that the reversion of some of the barium carbonate to barium sulphate +is prevented. This is an application of the principle of mass action, +in which the concentration of the reagent (the carbonate ion) is +kept as high as practicable and that of the sulphate ion as low as +possible, in order to force the reaction in the desired direction (see +Appendix).] + + + + +DETERMINATION OF PHOSPHORIC ANHYDRIDE IN APATITE + + +The mineral apatite is composed of calcium phosphate, associated with +calcium chloride, or fluoride. Specimens are easily obtainable which +are nearly pure and leave on treatment with acid only a slight +siliceous residue. + +For the purpose of gravimetric determination, phosphoric acid is +usually precipitated from ammoniacal solutions in the form of +magnesium ammonium phosphate which, on ignition, is converted into +magnesium pyrophosphate. Since the calcium phosphate of the apatite +is also insoluble in ammoniacal solutions, this procedure cannot be +applied directly. The separation of the phosphoric acid from the +calcium must first be accomplished by precipitation in the form of +ammonium phosphomolybdate in nitric acid solution, using ammonium +molybdate as the precipitant. The "yellow precipitate," as it is often +called, is not always of a definite composition, and therefore not +suitable for direct weighing, but may be dissolved in ammonia, and the +phosphoric acid thrown out as magnesium ammonium phosphate from the +solution. + +Of the substances likely to occur in apatite, silicic acid alone +interferes with the precipitation of the phosphoric acid in nitric +acid solution. + + +PRECIPITATION OF AMMONIUM PHOSPHOMOLYBDATE + +PROCEDURE.--Grind the mineral in an agate mortar until no grit is +perceptible. Transfer the substance to a weighing-tube, and weigh out +two portions, not exceeding 0.20 gram each (Note 1) into two beakers +of about 200 cc. capacity. Pour over them 20 cc. of dilute nitric acid +(sp. gr. 1.2) and warm gently until solvent action has apparently +ceased. Evaporate the solution cautiously to dryness, heat the residue +for about an hour at 100-110 deg.C., and treat it again with nitric acid +as described above; separate the residue of silica by filtration on +a small filter (7 cm.) and wash with warm water, using as little as +possible (Note 2). Receive the filtrate in a beaker (200-500 cc.). +Test the washings with ammonia for calcium phosphate, but add all such +tests in which a precipitate appears to the original nitrate (Note 3). +The filtrate and washings must be kept as small as possible and should +not exceed 100 cc. in volume. Add aqueous ammonia (sp. gr. 0.96) until +the precipitate of calcium phosphate first produced just fails to +redissolve, and then add a few drops of nitric acid until this is +again brought into solution (Note 4). Warm the solution until it +cannot be comfortably held in the hand (about 60 deg.C.) and, after +removal of the burner, add 75 cc. of ammonium molybdate solution which +has been !gently! warmed, but which must be perfectly clear. Allow +the mixture to stand at a temperature of about 50 or 60 deg.C. for twelve +hours (Notes 5 and 6). Filter off the yellow precipitate on a 9 cm. +filter, and wash by decantation with a solution of ammonium nitrate +made acid with nitric acid.[1] Allow the precipitate to remain in the +beaker as far as possible. Test the washings for calcium with ammonia +and ammonium oxalate (Note 3). + +[Footnote 1: This solution is prepared as follows: Mix 100 cc. of +ammonia solution (sp. gr. 0.96) with 325 cc. of nitric acid (sp. gr. +1.2) and dilute with 100 cc. of water.] + +Add 10 cc. of molybdate solution to the nitrate, and leave it for +a few hours. It should then be carefully examined for a !yellow! +precipitate; a white precipitate may be neglected. + +[Note 1: Magnesium ammonium phosphate, as noted below, is slightly +soluble under the conditions of operation. Consequently the +unavoidable errors of analysis are greater in this determination than +in those which have preceded it, and some divergence may be expected +in duplicate analyses. It is obvious that the larger the amount of +substance taken for analysis the less will be the relative loss or +gain due to unavoidable experimental errors; but, in this instance, a +check is placed upon the amount of material which may be taken both by +the bulk of the resulting precipitate of ammonium phosphomolybdate +and by the excessive amount of ammonium molybdate required to effect +complete separation of the phosphoric acid, since a liberal excess +above the theoretical quantity is demanded. Molybdic acid is one of +the more expensive reagents.] + +[Note 2: Soluble silicic acid would, if present, partially separate +with the phosphomolybdate, although not in combination with +molybdenum. Its previous removal by dehydration is therefore +necessary.] + +[Note 3: When washing the siliceous residue the filtrate may be tested +for calcium by adding ammonia, since that reagent neutralizes the +acid which holds the calcium phosphate in solution and causes +precipitation; but after the removal of the phosphoric acid in +combination with the molybdenum, the addition of an oxalate is +required to show the presence of calcium.] + +[Note 4: An excess of nitric acid exerts a slight solvent +action, while ammonium nitrate lessens the solubility; hence the +neutralization of the former by ammonia.] + +[Note 5: The precipitation of the phosphomolybdate takes place more +promptly in warm than in cold solutions, but the temperature should +not exceed 60 deg.C. during precipitation; a higher temperature tends to +separate molybdic acid from the solution. This acid is nearly white, +and its deposition in the filtrate on long standing should not be +mistaken for a second precipitation of the yellow precipitate. The +addition of 75 cc. of ammonium molybdate solution insures the presence +of a liberal excess of the reagent, but the filtrate should be tested +as in all quantitative procedures. + +The precipitation is probably complete in many cases in less than +twelve hours; but it is better, when practicable, to allow the +solution to stand for this length of time. Vigorous shaking or +stirring promotes the separation of the precipitate.] + +[Note 6: The composition of the "yellow precipitate" undoubtedly +varies slightly with varying conditions at the time of its formation. +Its composition may probably fairly be represented by the formula, +(NH_{4})_{3}PO_{4}.12MoO_{3}.H_{2}O, when precipitated under the +conditions prescribed in the procedure. Whatever other variations may +occur in its composition, the ratio of 12 MoO_{3}:1 P seems to +hold, and this fact is utilized in volumetric processes for the +determination of phosphorus, in which the molybdenum is reduced to +a lower oxide and reoxidized by a standard solution of potassium +permanganate. In principle, the procedure is comparable with that +described for the determination of iron by permanganate.] + + +PRECIPITATION OF MAGNESIUM AMMONIUM PHOSPHATE + +PROCEDURE.--Dissolve the precipitate of phosphomolybdate upon the +filter by pouring through it dilute aqueous ammonia (one volume of +dilute ammonia (sp. gr. 0.96) and three volumes of water, which +should be carefully measured), and receive the solution in the beaker +containing the bulk of the precipitate. The total volume of nitrate +and washings should not much exceed 100 cc. Acidify the solution with +dilute hydrochloric acid, and heat it nearly to boiling. Calculate the +volume of magnesium ammonium chloride solution ("magnesia mixture") +required to precipitate the phosphoric acid, assuming 40 per cent +P_{2}O_{5} in the apatite. Measure out about 5 cc. in excess of this +amount, and pour it into the acid solution. Then add slowly dilute +ammonium hydroxide (1 volume of strong ammonia (sp. gr. 0.90) and 9 +volumes of water), stirring constantly until a precipitate forms. Then +add a volume of filtered, concentrated ammonia (sp. gr. 0.90) equal to +one third of the volume of liquid in the beaker (Note 1). Allow the +whole to cool. The precipitated magnesium ammonium phosphate should +then be definitely crystalline in appearance (Note 2). (If it is +desired to hasten the precipitation, the solution may be cooled, first +in cold and then in ice-water, and stirred !constantly! for half an +hour, when precipitation will usually be complete.) + +Decant the clear liquid through a filter, and transfer the precipitate +to the filter, using as wash-water a mixture of one volume of +concentrated ammonia and three volumes of water. It is not necessary +to clean the beaker completely or to wash the precipitate thoroughly +at this point, as it is necessary to purify it by reprecipitation. + +[Note 1: Magnesium ammonium phosphate is not a wholly insoluble +substance, even under the most favorable analytical conditions. It +is least soluble in a liquid containing one fourth of its volume of +concentrated aqueous ammonia (sp. gr. 0.90) and this proportion should +be carefully maintained as prescribed in the procedure. On account of +this slight solubility the volume of solutions should be kept as small +as possible and the amount of wash-water limited to that absolutely +required. + +A large excess of the magnesium solution tends both to throw out +magnesium hydroxide (shown by a persistently flocculent precipitate) +and to cause the phosphate to carry down molybdic acid. The tendency +of the magnesium precipitate to carry down molybdic acid is also +increased if the solution is too concentrated. The volume should not +be less than 90 cc., nor more than 125 cc., at the time of the first +precipitation with the magnesia mixture.] + +[Note 2: The magnesium ammonium phosphate should be perfectly +crystalline, and will be so if the directions are followed. The slow +addition of the reagent is essential, and the stirring not less so. +Stirring promotes the separation of the precipitate and the formation +of larger crystals, and may therefore be substituted for digestion in +the cold. The stirring-rod must not be allowed to scratch the glass, +as the crystals adhere to such scratches and are removed with +difficulty.] + + +REPRECIPITATION AND IGNITION OF MAGNESIUM AMMONIUM PHOSPHATE + +A single precipitation of the magnesium compound in the presence of +molybdenum compounds rarely yields a pure product. The molybdenum can +be removed by solution of the precipitate in acid and precipitation +of the molybdenum by sulphureted hydrogen, after which the magnesium +precipitate may be again thrown down. It is usually more satisfactory +to dissolve the magnesium precipitate and reprecipitate the phosphate +as magnesium ammonium phosphate as described below. + +PROCEDURE.--Dissolve the precipitate from the filter in a little +dilute hydrochloric acid (sp. gr. 1.12), allowing the acid solution to +run into the beaker in which the original precipitation was made (Note +1). Wash the filter with water until the wash-water shows no test for +chlorides, but avoid an unnecessary amount of wash-water. Add to +the solution 2 cc. (not more) of magnesia mixture, and then dilute +ammonium hydroxide solution (sp. gr. 0.96), drop by drop, with +constant stirring, until the liquid smells distinctly of ammonia. Stir +for a few moments and then add a volume of strong ammonia (sp. gr. +0.90), equal to one third of the volume of the solution. Allow the +solution to stand for some hours, and then filter off the magnesium +ammonium phosphate, which should be distinctly crystalline in +character. Wash the precipitate with dilute ammonia water, as +prescribed above, until, finally, 3 cc. of the washings, after +acidifying with nitric acid, show no evidence of chlorides. Test both +filtrates for complete precipitation by adding a few cubic centimeters +of magnesia mixture and allowing them to stand for some time. + +Transfer the moist precipitate to a weighed porcelain or platinum +crucible and ignite, using great care to raise the temperature slowly +while drying the filter in the crucible, and to insure the ready +access of oxygen during the combustion of the filter paper, thus +guarding against a possible reduction of the phosphate, which would +result in disastrous consequences both to the crucible, if of +platinum, and the analysis. Do not raise the temperature above +moderate redness until the precipitate is white. (Keep this precaution +well in mind.) Ignite finally at the highest temperature of the +Tirrill burner, and repeat the heating until the weight is constant. +If the ignited precipitate is persistently discolored by particles of +unburned carbon, moisten the mass with a drop or two of concentrated +nitric acid and heat cautiously, finally igniting strongly. The +acid will dissolve magnesium pyrophosphate from the surface of the +particles of carbon, which will then burn away. Nitric acid also aids +as an oxidizing agent in supplying oxygen for the combustion of the +carbon. + +From the weight of magnesium pyrophosphate (Mg_{2}P_{2}O_{7}) +obtained, calculate the phosphoric anhydride (P_{2}O_{5}) in the +sample of apatite. + +[Note 1: The ionic change involved in the precipitation of the +magnesium compound is + +PO_{4}^{---} + NH_{4}^{+} + Mg^{++} --> [MgNH_{4}PO_{4}]. + +The magnesium ammonium phosphate is readily dissolved by acids, even +those which are no stronger than acetic acid. This is accounted for +by the fact that two of the ions into which phosphoric acid may +dissociate, the HPO_{4}^{--} or H_{2}PO_{4}^{-} ions, exhibit the +characteristics of very weak acids, in that they show almost no +tendency to dissociate further into H^{+} and PO_{4}^{--} ions. +Consequently the ionic changes which occur when the magnesium ammonium +phosphate is brought into contact with an acid may be typified by the +reaction: + +H^{+} + Mg^{++} + NH_{4}^{+} + PO_{4}^{---} --> Mg^{++} + NH_{4}^{+} + +HPO_{4}^{--}; + +that is, the PO_{4}^{--} ions and the H^{+} ions lose their identity +in the formation of the new ion, HPO_{4}^{--}, and this continues +until the magnesium ammonium phosphate is entirely dissolved.] + +[Note 2: During ignition the magnesium ammonium phosphate loses +ammonia and water and is converted into magnesium pyrophosphate: + +2MgNH_{4}PO_{4} --> Mg_{2}P_{2}O_{7} + 2NH_{3} + H_{2}O. + +The precautions mentioned on pages 111 and 123 must be observed with +great care during the ignition of this precipitate. The danger here +lies in a possible reduction of the phosphate by the carbon of the +filter paper, or by the ammonia evolved, which may act as a reducing +agent. The phosphorus then attacks and injures a platinum crucible, +and the determination is valueless.] + + + + +ANALYSIS OF LIMESTONE + + +Limestones vary widely in composition from a nearly pure marble +through the dolomitic limestones, containing varying amounts of +magnesium, to the impure varieties, which contain also ferrous and +manganous carbonates and siliceous compounds in variable proportions. +Many other minerals may be inclosed in limestones in small quantities, +and an exact qualitative analysis will often show the presence of +sulphides or sulphates, phosphates, and titanates, and the alkali or +even the heavy metals. No attempt is made in the following procedures +to provide a complete quantitative scheme which would take into +account all of these constituents. Such a scheme for a complete +analysis of a limestone may be found in Bulletin No. 700 of the United +States Geological Survey. It is assumed that, for these practice +determinations, a limestone is selected which contains only the more +common constituents first enumerated above. + + +DETERMINATION OF MOISTURE + +The determination of the amount of moisture in minerals or ores is +often of great importance. Ores which have been exposed to the weather +during shipment may have absorbed enough moisture to appreciably +affect the results of analysis. Since it is essential that the seller +and buyer should make their analyses upon comparable material, it is +customary for each analyst to determine the moisture in the sample +examined, and then to calculate the percentages of the various +constituents with reference to a sample dried in the air, or at a +temperature a little above 100 deg.C., which, unless the ore has undergone +chemical change because of the wetting, should be the same before and +after shipment. + +PROCEDURE.--Spread 25 grams of the powdered sample on a weighed +watch-glass; weigh to the nearest 10 milligrams only and heat at +105 deg.C.; weigh at intervals of an hour, after cooling in a desiccator, +until the loss of weight after an hour's heating does not exceed +10 milligrams. It should be noted that a variation in weight of 10 +milligrams in a total weight of 25 grams is no greater relatively than +a variation of 0.1 milligram when the sample taken weighs 0.25 gram + +DETERMINATION OF THE INSOLUBLE MATTER AND SILICA + +PROCEDURE.--Weigh out two portions of the original powdered sample +(not the dried sample), of about 5 grams each, into 250 cc. +casseroles, and cover each with a watch-glass (Note 1). Pour over the +powder 25 cc. of water, and then add 50 cc. of dilute hydrochloric +acid (sp. gr. 1.12) in small portions, warming gently, until nothing +further appears to dissolve (Note 2). Evaporate to dryness on the +water bath. Pour over the residue a mixture of 5 cc. of water and +5 cc. of concentrated hydrochloric acid (sp. gr. 1.2) and again +evaporate to dryness, and finally heat for at least an hour at +a temperature of 110 deg.C. Pour over this residue 50 cc. of dilute +hydrochloric acid (one volume acid (sp. gr. 1.12) to five volumes +water), and boil for about five minutes; then filter and wash twice +with the dilute hydrochloric acid, and then with hot water until +free from chlorides. Transfer the filter and contents to a porcelain +crucible, dry carefully over a low flame, and ignite to constant +weight. The residue represents the insoluble matter and the silica +from any soluble silicates (Note 3). + +Calculate the combined percentage of these in the limestone. + +[Note 1: The relatively large weight (5 grams) taken for analysis +insures greater accuracy in the determination of the ingredients which +are present in small proportions, and is also more likely to be a +representative sample of the material analyzed.] + +[Note 2: It is plain that the amount of the insoluble residue and also +its character will often depend upon the strength of acid used for +solution of the limestone. It cannot, therefore, be regarded as +representing any well-defined constituent, and its determination is +essentially empirical.] + +[Note 3: It is probable that some of the silicates present are wholly +or partly decomposed by the acid, and the soluble silicic acid must +be converted by evaporation to dryness, and heating, into white, +insoluble silica. This change is not complete after one evaporation. +The heating at a temperature somewhat higher than that of the water +bath for a short time tends to leave the silica in the form of a +powder, which promotes subsequent filtration. The siliceous residue +is washed first with dilute acid to prevent hydrolytic changes, which +would result in the formation of appreciable quantities of insoluble +basic iron or aluminium salts on the filter when washing with hot +water. + +If it is desired to determine the percentage of silica separately, the +ignited residue should be mixed in a platinum crucible with about six +times its weight of anhydrous sodium carbonate, and the procedure +given on page 151 should be followed. The filtrate from the silica is +then added to the main filtrate from the insoluble residue.] + + + + +DETERMINATION OF FERRIC OXIDE AND ALUMINIUM OXIDE (WITH MANGANESE) + + +PROCEDURE.--To the filtrate from the insoluble residue add ammonium +hydroxide until the solution just smells distinctly of ammonia, but do +not add an excess. Then add 5 cc. of saturated bromine water (Note 1), +and boil for five minutes. If the smell of ammonia has disappeared, +again add ammonium hydroxide in slight excess, and 3 cc. of bromine +water, and heat again for a few minutes. Finally add 10 cc. of +ammonium chloride solution and keep the solution warm until it barely +smells of ammonia; then filter promptly (Note 2). Wash the filter +twice with hot water, then (after replacing the receiving beaker) pour +through it 25 cc. of hot, dilute hydrochloric acid (one volume dilute +HCl [sp. gr. 1.12] to five volumes water). A brown residue insoluble +in the acid may be allowed to remain on the filter. Wash the filter +five times with hot water, add to the filtrate ammonium hydroxide +and bromine water as described above, and repeat the precipitation. +Collect the precipitate on the filter already used, wash it free from +chlorides with hot water, and ignite and weigh as described for ferric +hydroxide on page 110. The residue after ignition consists of ferric +oxide, alumina, and mangano-manganic oxide (Mn_{3}O_{4}), if manganese +is present. These are commonly determined together (Note 3). + +Calculate the percentage of the combined oxides in the limestone. + +[Note 1: The addition of bromine water to the ammoniacal solutions +serves to oxidize any ferrous hydroxide to ferric hydroxide and to +precipitate manganese as MnO(OH)_{2}. The solution must contain not +more than a bare excess of hydroxyl ions (ammonium hydroxide) when it +is filtered, on account of the tendency of the aluminium hydroxide to +redissolve. + +The solution should not be strongly ammoniacal when the bromine is +added, as strong ammonia reacts with the bromine, with the evolution +of nitrogen.] + +[Note 2: The precipitate produced by ammonium hydroxide and bromine +should be filtered off promptly, since the alkaline solution absorbs +carbon dioxide from the air, with consequent partial precipitation +of the calcium as carbonate. This is possible even under the most +favorable conditions, and for this reason the iron precipitate is +redissolved and again precipitated to free it from calcium. When the +precipitate is small, this reprecipitation may be omitted.] + +[Note 3: In the absence of significant amounts of manganese the iron +and aluminium may be separately determined by fusion of the mixed +ignited precipitate, after weighing, with about ten times its weight +of acid potassium sulphate, solution of the cold fused mass in water, +and volumetric determination of the iron, as described on page 66. +The aluminium is then determined by difference, after subtracting the +weight of ferric oxide corresponding to the amount of iron found. + +If a separate determination of the iron, aluminium, and manganese +is desired, the mixed precipitate may be dissolved in acid before +ignition, and the separation effected by special methods (see, for +example, Fay, !Quantitative Analyses!, First Edition, pp. 15-19 and +23-27).] + + + + +DETERMINATION OF CALCIUM + + +PROCEDURE.--To the combined filtrates from the double precipitation of +the hydroxides just described, add 5 cc. of dilute ammonium hydroxide +(sp. gr. 0.96), and transfer the liquid to a 500 cc. graduated flask, +washing out the beaker carefully. Cool to laboratory temperature, and +fill the flask with distilled water until the lowest point of the +meniscus is exactly level with the mark on the neck of the flask. +Carefully remove any drops of water which are on the inside of the +neck of the flask above the graduation by means of a strip of filter +paper, make the solution uniform by pouring it out into a dry beaker +and back into the flask several times. Measure off one fifth of this +solution as follows (Note 1): Pour into a 100 cc. graduated flask +about 10 cc. of the solution, shake the liquid thoroughly over the +inner surface of the small flask, and pour it out. Repeat the same +operation. Fill the 100 cc. flask until the lowest point of the +meniscus is exactly level with the mark on its neck, remove any drops +of solution from the upper part of the neck with filter paper, and +pour the solution into a beaker (400-500 cc.). Wash out the flask with +small quantities of water until it is clean, adding these to the 100 +cc. of solution. When the duplicate portion of 100 cc. is measured out +from the solution, remember that the flask must be rinsed out twice +with that solution, as prescribed above, before the measurement is +made. (A 100 cc. pipette may be used to measure out the aliquot +portions, if preferred.) + +Dilute each of the measured portions to 250 cc. with distilled water, +heat the whole to boiling, and add ammonium oxalate solution slowly +in moderate excess, stirring well. Boil for two minutes; allow the +precipitated calcium oxalate to settle for a half-hour, and decant +through a filter. Test the filtrate for complete precipitation by +adding a few cubic centimeters of the precipitant, allowing it to +stand for fifteen minutes. If no precipitate forms, make the solution +slightly acid with hydrochloric acid (Note 2); see that it is properly +labeled and reserve it to be combined with the filtrate from the +second calcium oxalate precipitation (Notes 3 and 4). + +Redissolve the calcium oxalate in the beaker with warm hydrochloric +acid, pouring the acid through the filter. Wash the filter five times +with water, and finally pour through it aqueous ammonia. Dilute the +solution to 250 cc., bring to boiling, and add 1 cc. ammonium oxalate +solution (Note 5) and ammonia in slight excess; boil for two minutes, +and set aside for a half-hour. Filter off the calcium oxalate upon the +filter first used, and wash free from chlorides. The filtrate should +be made barely acid with hydrochloric acid and combined with the +filtrate from the first precipitation. Begin at once the evaporation +of the solutions for the determination of magnesium as described +below. + +The precipitate of calcium oxalate may be converted into calcium oxide +by ignition without previous drying. After burning the filter, it may +be ignited for three quarters of an hour in a platinum crucible at +the highest heat of the Bunsen or Tirrill burner, and finally for ten +minutes at the blast lamp (Note 6). Repeat the heating over the blast +lamp until the weight is constant. As the calcium oxide absorbs +moisture from the air, it must (after cooling) be weighed as rapidly +as possible. + +The precipitate may, if preferred, be placed in a weighted porcelain +crucible. After burning off the filter and heating for ten minutes the +calcium precipitate may be converted into calcium sulphate by placing +2 cc. of dilute sulphuric acid in the crucible (cold), heating the +covered crucible very cautiously over a low flame to drive off the +excess of acid, and finally at redness to constant weight (Note 7). + +From the weight of the oxide or sulphate, calculate the percentage of +the calcium (Ca) in the limestone, remembering that only one fifth of +the total solution is used for this determination. + +[Note 1: If the calcium were precipitated from the entire solution, +the quantity of the precipitate would be greater than could be +properly treated. The solution is, therefore, diluted to a definite +volume (500 cc.), and exactly one fifth (100 cc.) is measured off in a +graduated flask or by means of a pipette.] + +[Note 2: The filtrate from the calcium oxalate should be made slightly +acid immediately after filtration, in order to avoid the solvent +action of the alkaline liquid upon the glass.] + +[Note 3: The accurate quantitative separation of calcium and magnesium +as oxalates requires considerable care. The calcium precipitate +usually carries down with it some magnesium, and this can best +be removed by redissolving the precipitate after filtration, and +reprecipitation in the presence of only the small amount of magnesium +which was included in the first precipitate. When, however, the +proportion of magnesium is not very large, the second precipitation of +the calcium can usually be avoided by precipitating it from a rather +dilute solution (800 cc. or so) and in the presence of a considerable +excess of the precipitant, that is, rather more than enough to convert +both the magnesium and calcium into oxalates.] + +[Note 4: The ionic changes involved in the precipitation of calcium +as oxalate are exceedingly simple, and the principles discussed in +connection with the barium sulphate precipitation on page 113 also +apply here. The reaction is + +C_{2}O_{4}^{--} + Ca^{++} --> [CaC_{2}O_{4}]. + +Calcium oxalate is nearly insoluble in water, and only very slightly +soluble in acetic acid, but is readily dissolved by the strong mineral +acids. This behavior with acids is explained by the fact that oxalic +acid is a stronger acid than acetic acid; when, therefore, the oxalate +is brought into contact with the latter there is almost no tendency to +diminish the concentration of C_{2}O_{4}^{--} ions by the formation of +an acid less dissociated than the acetic acid itself, and practically +no solvent action ensues. When a strong mineral acid is present, +however, the ionization of the oxalic acid is much reduced by the high +concentration of the H^{+} ions from the strong acid, the formation +of the undissociated acid lessens the concentration of the +C_{2}O_{4}^{--} ions in solution, more of the oxalate passes into +solution to re-establish equilibrium, and this process repeats itself +until all is dissolved. + +The oxalate is immediately reprecipitated from such a solution on the +addition of OH^{-} ions, which, by uniting with the H^{+} ions of the +acids (both the mineral acid and the oxalic acid) to form water, leave +the Ca^{++} and C_{2}O_{4}^{--} ions in the solution to recombine to +form [CaC_{2}O_{4}], which is precipitated in the absence of the +H^{+} ions. It is well at this point to add a small excess of +C_{2}O_{4}^{--} ions in the form of ammonium oxalate to decrease the +solubility of the precipitate. + +The oxalate precipitate consists mainly of CaC_{2}O_{4}.H_{2}O when +thrown down.] + +[Note 5: The small quantity of ammonium oxalate solution is added +before the second precipitation of the calcium oxalate to insure +the presence of a slight excess of the reagent, which promotes the +separation of the calcium compound.] + +[Note 6: On ignition the calcium oxalate loses carbon dioxide and +carbon monoxide, leaving calcium oxide: + +CaC_{2}O_{4}.H_{2}O --> CaO + CO_{2} + CO + H_{2}O. + +For small weights of the oxalate (0.6 gram or less) this reaction may +be brought about in a platinum crucible at the highest temperature of +a Tirrill burner, but it is well to ignite larger quantities than this +over the blast lamp until the weight is constant.] + +[Note 7: The heat required to burn the filter, and that subsequently +applied as described, will convert most of the calcium oxalate to +calcium carbonate, which is changed to sulphate by the sulphuric acid. +The reactions involved are + +CaC_{2}O_{4} --> CaCO_{3} + CO, +CaCO_{3} + H_{2}SO_{4} --> CaSO_{4} + H_{2}O + CO_{2}. + +If a porcelain crucible is employed for ignition, this conversion to +sulphate is to be preferred, as a complete conversion to oxide is +difficult to accomplish.] + +[Note 8: The determination of the calcium may be completed +volumetrically by washing the calcium oxalate precipitate from +the filter into dilute sulphuric acid, warming, and titrating +the liberated oxalic acid with a standard solution of potassium +permanganate as described on page 72. When a considerable number of +analyses are to be made, this procedure will save much of the time +otherwise required for ignition and weighing.] + + + + +DETERMINATION OF MAGNESIUM + + +PROCEDURE.--Evaporate the acidified filtrates from the calcium +precipitates until the salts begin to crystallize, but do !not! +evaporate to dryness (Note 1). Dilute the solution cautiously until +the salts are brought into solution, adding a little acid if the +solution has evaporated to very small volume. The solution should be +carefully examined at this point and must be filtered if a precipitate +has appeared. Heat the clear solution to boiling; remove the burner +and add 25 cc. of a solution of disodium phosphate. Then add slowly +dilute ammonia (1 volume strong ammonia (sp. gr. 0.90) and 9 volumes +water) as long as a precipitate continues to form. Finally, add a +volume of concentrated ammonia (sp. gr. 0.90) equal to one third of +the volume of the solution, and allow the whole to stand for about +twelve hours. + +Decant the solution through a filter, wash it with dilute ammonia +water, proceeding as prescribed for the determination of phosphoric +anhydride on page 122, including; the reprecipitation (Note 2), +except that 3 cc. of disodium phosphate solution are added before the +reprecipitation of the magnesium ammonium phosphate instead of +the magnesia mixture there prescribed. From the weight of the +pyrophosphate, calculate the percentage of magnesium oxide (MgO) in +the sample of limestone. Remember that the pyrophosphate finally +obtained is from one fifth of the original sample. + +[Note 1: The precipitation of the magnesium should be made in as small +volume as possible, and the ratio of ammonia to the total volume of +solution should be carefully provided for, on account of the relative +solubility of the magnesium ammonium phosphate. This matter has +been fully discussed in connection with the phosphoric anhydride +determination.] + +[Note 2: The first magnesium ammonium phosphate precipitate is rarely +wholly crystalline, as it should be, and is not always of the proper +composition when precipitated in the presence of such large amounts of +ammonium salts. The difficulty can best be remedied by filtering the +precipitate and (without washing it) redissolving in a small quantity +of hydrochloric acid, from which it may be again thrown down by +ammonia after adding a little disodium phosphate solution. If the +flocculent character was occasioned by the presence of magnesium +hydroxide, the second precipitation, in a smaller volume containing +fewer salts, will often result more favorably. + +The removal of iron or alumina from a contaminated precipitate is +a matter involving a long procedure, and a redetermination of the +magnesium from a new sample, with additional precautions, is usually +to be preferred.] + + + + +DETERMINATION OF CARBON DIOXIDE + + +!Absorption Apparatus! + +[Illustration: Fig. 3] + +The apparatus required for the determination of the carbon dioxide +should be arranged as shown in the cut (Fig. 3). The flask (A) is +an ordinary wash bottle, which should be nearly filled with dilute +hydrochloric acid (100 cc. acid (sp. gr. 1.12) and 200 cc. of water). +The flask is connected by rubber tubing (a) with the glass tube (b) +leading nearly to the bottom of the evolution flask (B) and having its +lower end bent upward and drawn out to small bore, so that the carbon +dioxide evolved from the limestone cannot bubble back into (b). The +evolution flask should preferably be a wide-mouthed Soxhlet extraction +flask of about 150 cc. capacity because of the ease with which tubes +and stoppers may be fitted into the neck of a flask of this type. The +flask should be fitted with a two-hole rubber stopper. The condenser +(C) may consist of a tube with two or three large bulbs blown in +it, for use as an air-cooled condenser, or it may be a small +water-jacketed condenser. The latter is to be preferred if a number of +determinations are to be made in succession. + +A glass delivery tube (c) leads from the condenser to the small U-tube +(D) containing some glass beads or small pieces of glass rod and 3 cc. +of a saturated solution of silver sulphate, with 3 cc. of concentrated +sulphuric acid (sp. gr. 1.84). The short rubber tubing (d) connects +the first U-tube to a second U-tube (E) which is filled with small +dust-free lumps of dry calcium chloride, with a small, loose plug of +cotton at the top of each arm. Both tubes should be closed by cork +stoppers, the tops of which are cut off level with, or preferably +forced a little below, the top of the U-tube, and then neatly sealed +with sealing wax. + +The carbon dioxide may be absorbed in a tube containing soda lime +(F) or in a Geissler bulb (F') containing a concentrated solution +of potassium hydroxide (Note 2). The tube (F) is a glass-stoppered +side-arm U-tube in which the side toward the evolution flask and one +half of the other side are filled with small, dust-free lumps of soda +lime of good quality (Note 3). Since soda lime contains considerable +moisture, the other half of the right side of the tube is filled with +small lumps of dry, dust-free calcium chloride to retain the moisture +from the soda lime. Loose plugs of cotton are placed at the top of +each arm and between the soda lime and the calcium chloride. + +The Geissler bulb (F'), if used, should be filled with potassium +hydroxide solution (1 part of solid potassium hydroxide dissolved in +two parts of water) until each small bulb is about two thirds full +(Note 4). A small tube containing calcium chloride is connected with +the Geissler bulb proper by a ground joint and should be wired to the +bulb for safety. This is designed to retain any moisture from the +hydroxide solution. A piece of clean, fine copper wire is so attached +to the bulb that it can be hung from the hook above a balance pan, or +other support. + +The small bottle (G) with concentrated sulphuric acid (sp. gr. 1.84) +is so arranged that the tube (f) barely dips below the surface. This +will prevent the absorption of water vapor by (F) or (F') and serves +as an aid in regulating the flow of air through the apparatus. (H) is +an aspirator bottle of about four liters capacity, filled with water; +(k) is a safety tube and a means of refilling (H); (h) is a screw +clamp, and (K) a U-tube filled with soda lime. + +[Note 1: The air current, which is subsequently drawn through the +apparatus, to sweep all of the carbon dioxide into the absorption +apparatus, is likely to carry with it some hydrochloric acid from +the evolution flask. This acid is retained by the silver sulphate +solution. The addition of concentrated sulphuric acid to this solution +reduces its vapor pressure so far that very little water is carried on +by the air current, and this slight amount is absorbed by the calcium +chloride in (E). As the calcium chloride frequently contains a small +amount of a basic material which would absorb carbon dioxide, it is +necessary to pass carbon dioxide through (E) for a short time and then +drive all the gas out with a dry air current for thirty minutes before +use.] + +[Note 2: Soda-lime absorption tubes are to be preferred if a +satisfactory quality of soda lime is available and the number of +determinations to be made successively is small. The potash bulbs will +usually permit of a larger number of successive determinations without +refilling, but they require greater care in handling and in the +analytical procedure.] + +[Note 3: Soda lime is a mixture of sodium and calcium hydroxides. Both +combine with carbon dioxide to form carbonates, with the evolution +of water. Considerable heat is generated by the reaction, and the +temperature of the tube during absorption serves as a rough index of +the progress of the reaction through the mass of soda lime. + +It is essential that soda lime of good quality for analytical purposes +should be used. The tube should not contain dust, as this is likely to +be swept away.] + +[Note 4: The solution of the hydroxide for use in the Geissler bulb +must be highly concentrated to insure complete absorption of the +carbon dioxide and also to reduce the vapor pressure of the solution, +thus lessening the danger of loss of water with the air which passes +through the bulbs. The small quantity of moisture which is then +carried out of the bulbs is held by the calcium chloride in the +prolong tube. The best form of absorption bulb is that to which the +prolong tube is attached by a ground glass joint. + +After the potassium hydroxide is approximately half consumed in the +first bulb of the absorption apparatus, potassium bicarbonate is +formed, and as it is much less soluble than the carbonate, it often +precipitates. Its formation is a warning that the absorbing power of +the hydroxide is much diminished.] + + +!The Analysis! + +PROCEDURE.-- Weigh out into the flask (B) about 1 gram of limestone. +Cover it with 15 cc. of water. Weigh the absorption apparatus (F) +or (F') accurately after allowing it to stand for 30 minutes in the +balance case, and wiping it carefully with a lintless cloth, taking +care to handle it as little as possible after wiping (Note 1). Connect +the absorption apparatus with (e) and (f). If a soda-lime tube is +used, be sure that the arm containing the soda lime is next the tube +(E) and that the glass stopcocks are open. + +To be sure that the whole apparatus is airtight, disconnect the rubber +tube from the flask (A), making sure that the tubes (a) and (b) do not +contain any hydrochloric acid, close the pinchcocks (a) and (k) and +open (h). No bubbles should pass through (D) or (G) after a few +seconds. When assured that the fittings are tight, close (h) and open +(a) cautiously to admit air to restore atmospheric pressure. This +precaution is essential, as a sudden inrush of air will project liquid +from (D) or (F'). Reconnect the rubber tube with the flask (A). +Open the pinchcocks (a) and (k) and blow over about 10 cc. of the +hydrochloric acid from (A) into (B). When the action of the acid +slackens, blow over (slowly) another 10 cc. + +The rate of gas evolution should not exceed for more than a few +seconds that at which about two bubbles per second pass through (G) +(Note 2). Repeat the addition of acid in small portions until the +action upon the limestone seems to be at an end, taking care to close +(a) after each addition of acid (Note 3). Disconnect (A) and connect +the rubber tubing with the soda-lime tube (K) and open (a). Then close +(k) and open (h), regulating the flow of water from (H) in such a way +that about two bubbles per second pass through (G). Place a small +flame under (B) and !slowly! raise the contents to boiling and boil +for three minutes. Then remove the burner from under (B) and continue +to draw air through the apparatus for 20-30 minutes, or until (H) +is emptied (Note 4). Remove the absorption apparatus, closing the +stopcocks on (F) or stoppering the open ends of (F'), leave the +apparatus in the balance case for at least thirty minutes, wipe it +carefully and weigh, after opening the stopcocks (or removing plugs). +The increase in weight is due to absorption of CO_{2}, from which its +percentage in the sample may be calculated. + +After cleaning (B) and refilling (H), the apparatus is ready for the +duplicate analysis. + +[Note 1: The absorption tubes or bulbs have large surfaces on which +moisture may collect. By allowing them to remain in the balance case +for some time before weighing, the amount of moisture absorbed on the +surface is as nearly constant as practicable during two weighings, and +a uniform temperature is also assured. The stopcocks of the U-tube +should be opened, or the plugs used to close the openings of the +Geissler bulb should be removed before weighing in order that the air +contents shall always be at atmospheric pressure.] + +[Note 2: If the gas passes too rapidly into the absorption apparatus, +some carbon dioxide may be carried through, not being completely +retained by the absorbents.] + +[Note 3: The essential ionic changes involved in this procedure are +the following: It is assumed that the limestone, which is typified by +calcium carbonate, is very slightly soluble in water, and the ions +resulting are Ca^{++} and CO_{3}^{--}. In the presence of H^{+} ions +of the mineral acid, the CO_{3}^{--} ions form [H_{2}CO_{3}]. This +is not only a weak acid which, by its formation, diminishes the +concentration of the CO_{3}^{--} ions, thus causing more of the +carbonate to dissolve to re-establish equilibrium, but it is also an +unstable compound and breaks down into carbon dioxide and water.] + +[Note 4: Carbon dioxide is dissolved by cold water, but the gas is +expelled by boiling, and, together with that which is distributed +through the apparatus, is swept out into the absorption bulb by the +current of air. This air is purified by drawing it through the tube +(K) containing soda lime, which removes any carbon dioxide which may +be in it.] + + + + +DETERMINATION OF LEAD, COPPER, IRON, AND ZINC IN BRASS + +ELECTROLYTIC SEPARATIONS + + +!General Discussion! + +When a direct current of electricity passes from one electrode to +another through solutions of electrolytes, the individual ions present +in these solutions tend to move toward the electrode of opposite +electrical charge to that which each ion bears, and to be discharged +by that electrode. Whether or not such discharge actually occurs in +the case of any particular ion depends upon the potential (voltage) of +the current which is passing through the solution, since for each ion +there is, under definite conditions, a minimum potential below which +the discharge of the ion cannot be effected. By taking advantage +of differences in discharge-potentials, it is possible to effect +separations of a number of the metallic ions by electrolysis, and at +the same time to deposit the metals in forms which admit of direct +weighing. In this way the slower procedures of precipitation and +filtration may frequently be avoided. The following paragraphs present +a brief statement of the fundamental principles and conditions +underlying electro-analysis. + +The total energy of an electric current as it passes through a +solution is distributed among three factors, first, its potential, +which is measured in volts, and corresponds to what is called "head" +in a stream of water; second, current strength, which is measured +in amperes, and corresponds to the volume of water passing a +cross-section of a stream in a given time interval; and third, the +resistance of the conducting medium, which is measured in ohms. The +relation between these three factors is expressed by Ohm's law, +namely, that !I = E/R!, when I is current strength, E potential, and R +resistance. It is plain that, for a constant resistance, the +strength of the current and its potential are mutually and directly +interdependent. + +As already stated, the applied electrical potential determines whether +or not deposition of a metal upon an electrode actually occurs. The +current strength determines the rate of deposition and the physical +characteristics of the deposit. The resistance of the solution is +generally so small as to fall out of practical consideration. + +Approximate deposition-potentials have been determined for a number +of the metallic elements, and also for hydrogen and some of the +acid-forming radicals. The values given below are those required +for deposition from normal solutions at ordinary temperatures +with reference to a hydrogen electrode. They must be regarded as +approximate, since several disturbing factors and some secondary +reactions render difficult their exact application under the +conditions of analysis. They are: + + Zn Cd Fe Ni Pb H Cu Sb Hg Ag SO_{4} ++0.77 +0.42 +0.34 +0.33 +0.13 0 -0.34 -0.67 -0.76 -0.79 +1.90 + +From these data it is evident that in order to deposit copper from a +normal solution of copper sulphate a minimum potential equal to the +algebraic sum of the deposition-potentials of copper ions and sulphate +ions must be applied, that is, +1.56 volts. The deposition of zinc +from a solution of zinc sulphate would require +2.67 volts, but, since +the deposition of hydrogen from sulphuric acid solution requires only ++1.90 volts, the quantitative deposition of zinc by electrolysis from +a sulphuric acid solution of a zinc salt is not practicable. On the +other hand, silver, if present in a solution of copper sulphate, would +deposit with the copper. + +The foregoing examples suffice to illustrate the application of the +principle of deposition potentials, but it must further be noted +that the values stated apply to normal solutions of the compounds in +question, that is, to solutions of considerable concentrations. As the +concentration of the ions diminishes, and hence fewer ions approach +the electrodes, somewhat higher voltages are required to attract and +discharge them. From this it follows that the concentrations should be +kept as high as possible to effect complete deposition in the least +practicable time, or else the potentials applied must be progressively +increased as deposition proceeds. In practice, the desired result is +obtained by starting with small volumes of solution, using as large an +electrode surface as possible, and by stirring the solution to bring +the ions in contact with the electrodes. This is, in general, a more +convenient procedure than that of increasing the potential of the +current during electrolysis, although that method is also used. + +As already stated, those ions in a solution of electrolytes will first +be discharged which have the lowest deposition potentials, and so +long as these ions are present around the electrode in considerable +concentration they, almost alone, are discharged, but, as their +concentration diminishes, other ions whose deposition potentials are +higher but still within that of the current applied, will also begin +to separate. For example, from a nitric acid solution of copper +nitrate, the copper ions will first be discharged at the cathode, but +as they diminish in concentration hydrogen ions from the acid (or +water) will be also discharged. Since the hydrogen thus liberated is a +reducing agent, the nitric acid in the solution is slowly reduced to +ammonia, and it may happen that if the current is passed through for a +long time, such a solution will become alkaline. Oxygen is liberated +at the anode, but, since there is no oxidizable substance present +around that electrode, it escapes as oxygen gas. It should be noted +that, in general, the changes occurring at the cathode are reductions, +while those at the anode are oxidations. + +For analytical purposes, solutions of nitrates or sulphates of the +metals are preferable to those of the chlorides, since liberated +chlorine attacks the electrodes. In some cases, as for example, that +of silver, solution of salts forming complex ions, like that of +the double cyanide of silver and potassium, yield better metallic +deposits. + +Most metals are deposited as such upon the cathode; a few, notably +lead and manganese, separate in the form of dioxides upon the anode. +It is evidently important that the deposited material should be so +firmly adherent that it can be washed, dried, and weighed without +loss in handling. To secure these conditions it is essential that the +current density (that is, the amount of current per unit of area of +the electrodes) shall not be too high. In prescribing analytical +conditions it is customary to state the current strength in "normal +densities" expressed in amperes per 100 sq. cm. of electrode surface, +as, for example, "N.D_{100} = 2 amps." + +If deposition occurs too rapidly, the deposit is likely to be spongy +or loosely adherent and falls off on subsequent treatment. This places +a practical limit to the current density to be employed, for a given +electrode surface. The cause of the unsatisfactory character of +the deposit is apparently sometimes to be found in the coincident +liberation of considerable hydrogen and sometimes in the failure of +the rapidly deposited material to form a continuous adherent surface. +The effect of rotating electrodes upon the character of the deposit is +referred to below. + +The negative ions of an electrolyte are attracted to the anode and are +discharged on contact with it. Anions such as the chloride ion yield +chlorine atoms, from which gaseous chlorine molecules are formed +and escape. The radicals which compose such ions as NO_{3}^{-} or +SO_{4}^{--} are not capable of independent existence after discharge, +and break down into oxygen and N_{2}O_{5} and SO_{3} respectively. The +oxygen escapes and the anhydrides, reacting with water, re-form nitric +and sulphuric acids. + +The law of Faraday expresses the relation between current strength and +the quantities of the decomposition products which, under constant +conditions, appear at the electrodes, namely, that a given quantity +of electricity, acting for a given time, causes the separation of +chemically equivalent quantities of the various elements or radicals. +For example, since 107.94 grams of silver is equivalent to 1.008 grams +of hydrogen, and that in turn to 8 grams of oxygen, or 31.78 grams of +copper, the quantity of electricity which will cause the deposit of +107.94 grams of silver in a given time will also separate the weights +just indicated of the other substances. Experiments show that a +current of one ampere passing for one second, i.e., a coulomb of +electricity, causes the deposition of 0.001118 gram of silver from a +normal solution of a silver salt. The number of coulombs required to +deposit 107.94 grams is 107.94/0.001118 or 96,550 and the same number +of coulombs will also cause the separation of 1.008 grams of hydrogen, +8 grams of oxygen or 31.78 grams of copper. While it might at first +appear that Faraday's law could thus be used as a basis for the +calculation of the time required for the deposition of a given +quantity of an electrolyte from solution, it must be remembered that +the law expresses what occurs when the concentration of the ions in +the solution is kept constant, as, for example, when the anode in +a silver salt solution is a plate of metallic silver. Under the +conditions of electro-analysis the concentration of the ions is +constantly diminishing as deposition proceeds and the time actually +required for complete deposition of a given weight of material by +a current of constant strength is, therefore, greater than that +calculated on the basis of the law as stated above. + +The electrodes employed in electro-analysis are almost exclusively +of platinum, since that metal alone satisfactorily resists chemical +action of the electrolytes, and can be dried and weighed without +change in composition. The platinum electrodes may be used in the +form of dishes, foil or gauze. The last, on account of the ease of +circulation of the electrolyte, its relatively large surface in +proportion to its weight and the readiness with which it can be washed +and dried, is generally preferred. + +Many devices have been described by the use of which the electrode +upon which deposition occurs can be mechanically rotated. This has an +effect parallel to that of greatly increasing the electrode surface +and also provides a most efficient means of stirring the solution. +With such an apparatus the amperage may be increased to 5 or even 10 +amperes and depositions completed with great rapidity and accuracy. It +is desirable, whenever practicable, to provide a rotating or stirring +device, since, for example, the time consumed in the deposition of the +amount of copper usually found in analysis may be reduced from the +20 to 24 hours required with stationary electrodes, and unstirred +solutions, to about one half hour. + + + + +DETERMINATION OF COPPER AND LEAD + + +PROCEDURE.--Weigh out two portions of about 0.5 gram each (Note 1) +into tall, slender lipless beakers of about 100 cc. capacity. Dissolve +the metal in a solution of 5 cc. of dilute nitric acid (sp. gr. 1.20) +and 5 cc. of water, heating gently, and keeping the beaker covered. +When the sample has all dissolved (Note 2), wash down the sides of the +beaker and the bottom of the watch-glass with water and dilute the +solution to about 50 cc. Carefully heat to boiling and boil for a +minute or two to expel nitrous fumes. + +Meanwhile, four platinum electrodes, two anodes and two cathodes, +should be cleaned by dipping in dilute nitric acid, washing with water +and finally with 95 per cent alcohol (Note 3). The alcohol may be +ignited and burned off. The electrodes are then cooled in a desiccator +and weighed. Connect the electrodes with the binding posts (or other +device for connection with the electric circuit) in such a way that +the copper will be deposited upon the electrode with the larger +surface, which is made the cathode. The beaker containing the solution +should then be raised into place from below the electrodes until the +latter reach nearly to the bottom of the beaker. The support for the +beaker must be so arranged that it can be easily raised or lowered. + +If the electrolytic apparatus is provided with a mechanism for the +rotation of the electrode or stirring of the electrolyte, proceed as +follows: Arrange the resistance in the circuit to provide a direct +current of about one ampere. Pass this current through the solution +to be electrolyzed, and start the rotating mechanism. Keep the beaker +covered as completely as possible, using a split watch-glass (or other +device) to avoid loss by spattering. When the solution is colorless, +which is usually the case after about 35 minutes, rinse off the cover +glass, wash down the sides of the beaker, add about 0.30 gram of urea +and continue the electrolysis for another five minutes (Notes 4 and +5). + +If stationary electrodes are employed, the current strength should be +about 0.1 ampere, which may, after 12 to 15 hours, be increased to 0.2 +ampere. The time required for complete deposition is usually from 20 +to 24 hours. It is advisable to add 5 cc. of nitric acid (sp. gr. 1.2) +if the electrolysis extends over this length of time. No urea is added +in this case. + +When the deposition of the copper appears to be complete, stop the +rotating mechanism and slowly lower the beaker with the left hand, +directing at the same time a stream of water from a wash bottle on +both electrodes. Remove the beaker, shut off the current, and, if +necessary, complete the washing of the electrodes (Note 6). Rinse the +electrodes cautiously with alcohol and heat them in a hot closet until +the alcohol has just evaporated, but no longer, since the copper is +likely to oxidize at the higher temperature. (The alcohol may be +removed by ignition if care is taken to keep the electrodes in motion +in the air so that the copper deposit is not too strongly heated at +any one point.) + +Test the solution in the beaker for copper as follows, remembering +that it is to be used for subsequent determinations of iron and zinc: +Remove about 5 cc. and add a slight excess of ammonia. Compare the +mixture with some distilled water, holding both above a white surface. +The solution should not show any tinge of blue. If the presence of +copper is indicated, add the test portion to the main solution, +evaporate the whole to a volume of about 100 cc., and again +electrolyze with clean electrodes (Note 7). + +After cooling the electrodes in a desiccator, weigh them and from the +weight of copper on the cathode and of lead dioxide (PbO_{2}) on the +anode, calculate the percentage of copper (Cu) and of lead (Pb) in the +brass. + +[Note 1: It is obvious that the brass taken for analysis should be +untarnished, which can be easily assured, when wire is used, by +scouring with emery. If chips or borings are used, they should be well +mixed, and the sample for analysis taken from different parts of the +mixture.] + +[Note 2: If a white residue remains upon treatment of the alloy with +nitric acid, it indicates the presence of tin. The material is not, +therefore, a true brass. This may be treated as follows: Evaporate the +solution to dryness, moisten the residue with 5 cc. of dilute nitric +acid (sp. gr. 1.2) and add 50 cc. of hot water. Filter off the +meta-stannic acid, wash, ignite in porcelain and weigh as SnO_{2}. +This oxide is never wholly free from copper and must be purified for +an exact determination. If it does not exceed 2 per cent of the alloy, +the quantity of copper which it contains may usually be neglected.] + +[Note 3: The electrodes should be freed from all greasy matter before +using, and those portions upon which the metal will deposit should not +be touched with the fingers after cleaning.] + +[Note 4: Of the ions in solution, the H^{+}, Cu^{++}, Zn^{++}, and +Fe^{+++} ions tend to move toward the cathode. The NO_{3}^{-} ions and +the lead, probably in the form of PbO_{2}^{--} ions, move toward the +anode. At the cathode the Cu^{++} ions are discharged and plate out as +metallic copper. This alone occurs while the solution is relatively +concentrated. Later on, H^{+} ions are also discharged. In the +presence of considerable quantities of H^{+} ions, as in this acid +solution, no Zn^{++} or Fe^{+++} ions are discharged because of their +greater deposition potentials. At the anode the lead is deposited as +PbO_{2} and oxygen is evolved. + +For the reasons stated on page 141 care must be taken that the +solution does not become alkaline if the electrolysis is long +continued.] + +[Note 5: Urea reacts with nitrous acid, which may be formed in the +solution as a result of the reducing action of the liberated hydrogen. +Its removal promotes the complete precipitation of the copper. The +reaction is + +CO(NH_{2})_{2} + 2HNO_{2} --> CO_{2} + 2N_{2} + 3H_{2}O.] + +[Note 6: The electrodes must be washed nearly or quite free from +the nitric acid solution before the circuit is broken to prevent +re-solution of the copper. + +If several solutions are connected in the same circuit it is obvious +that some device must be used to close the circuit as soon as the +beaker is removed.] + +[Note 7: The electrodes upon which the copper has been deposited +may be cleaned by immersion in warm nitric acid. To remove the lead +dioxide, add a few crystals of oxalic acid to the nitric acid.] + + + + +DETERMINATION OF IRON + + +Most brasses contain small percentages of iron (usually not over 0.1 +per cent) which, unless removed, is precipitated as phosphate and +weighed with the zinc. + +PROCEDURE.--To the solution from the precipitation of copper and +lead by electrolysis, add dilute ammonia (sp. gr. 0.96) until the +precipitate of zinc hydroxide which first forms re-dissolves, leaving +only a slight red precipitate of ferric hydroxide. Filter off the +iron precipitate, using a washed filter, and wash five times with hot +water. Test a portion of the last washing with a dilute solution of +ammonium sulphide to assure complete removal of the zinc. + +The precipitate may then be ignited and weighed as ferric oxide, as +described on page 110. + +Calculate the percentage of iron (Fe) in the brass. + + + + +DETERMINATION OF ZINC + + +PROCEDURE.--Acidify the filtrate from the iron determination with +dilute nitric acid. Concentrate it to 150 cc. Add to the cold solution +dilute ammonia (sp. gr. 0.96) cautiously until it barely smells of +ammonia; then add !one drop! of a dilute solution of litmus (Note 1), +and drop in, with the aid of a dropper, dilute nitric acid until the +blue of the litmus just changes to red. It is important that this +point should not be overstepped. Heat the solution nearly to boiling +and pour into it slowly a filtered solution of di-ammonium hydrogen +phosphate[1] containing a weight of the phosphate about equal +to twelve times that of the zinc to be precipitated. (For this +calculation the approximate percentage of zinc is that found by +subtracting the sum of the percentages of the copper, lead and iron +from 100 per cent.) Keep the solution just below boiling for fifteen +minutes, stirring frequently (Note 2). If at the end of this time the +amorphous precipitate has become crystalline, allow the solution to +cool for about four hours, although a longer time does no harm (Note +3), and filter upon an asbestos filter in a porcelain Gooch crucible. +The filter is prepared as described on page 103, and should be dried +to constant weight at 105 deg.C. + +[Footnote 1: The ammonium phosphate which is commonly obtainable +contains some mono-ammonium salt, and this is not satisfactory as a +precipitant. It is advisable, therefore, to weigh out the amount of +the salt required, dissolve it in a small volume of water, add a drop +of phenolphthalein solution, and finally add dilute ammonium hydroxide +solution cautiously until the solution just becomes pink, but do not +add an excess.] + +Wash the precipitate until free from sulphates with a warm 1 per cent +solution of the di-ammonium phosphate, and then five times with 50 per +cent alcohol (Note 4). Dry the crucible and precipitate for an hour at +105 deg.C., and finally to constant weight (Note 5). The filtrate should +be made alkaline with ammonia and tested for zinc with a few drops of +ammonium sulphide, allowing it to stand (Notes 6, 7 and 8). + +From the weight of the zinc ammonium phosphate (ZnNH_{4}PO_{4}) +calculate the percentage of the zinc (Zn) in the brass. + +[Note 1: The zinc ammonium phosphate is soluble both in acids and in +ammonia. It is, therefore, necessary to precipitate the zinc in a +nearly neutral solution, which is more accurately obtained by adding +a drop of a litmus solution to the liquid than by the use of litmus +paper.] + +[Note 2: The precipitate which first forms is amorphous, and may have +a variable composition. On standing it becomes crystalline and then +has the composition ZnNH_{4}PO_{4}. The precipitate then settles +rapidly and is apt to occasion "bumping" if the solution is heated to +boiling. Stirring promotes the crystallization.] + +[Note 3: In a carefully neutralized solution containing a considerable +excess of the precipitant, and also ammonium salts, the separation +of the zinc is complete after standing four hours. The ionic changes +connected with the precipitation of the zinc as zinc ammonium +phosphate are similar to those described for magnesium ammonium +phosphate, except that the zinc precipitate is soluble in an excess of +ammonium hydroxide, probably as a result of the formation of complex +ions of the general character Zn(NH_{3})_{4}^{++}.] + +[Note 4: The precipitate is washed first with a dilute solution of the +phosphate to prevent a slight decomposition of the precipitate (as a +result of hydrolysis) if hot water alone is used. The alcohol is added +to the final wash-water to promote the subsequent drying.] + +[Note 5: The precipitate may be ignited and weighed as +Zn_{2}P_{2}O_{7}, by cautiously heating the porcelain Gooch crucible +within a nickel or iron crucible, used as a radiator. The heating +must be very slow at first, as the escaping ammonia may reduce the +precipitate if it is heated too quickly.] + +[Note 6: If the ammonium sulphide produced a distinct precipitate, +this should be collected on a small filter, dissolved in a few cubic +centimeters of dilute nitric acid, and the zinc reprecipitated as +phosphate, filtered off, dried, and weighed, and the weight added to +that of the main precipitate.] + +[Note 7: It has been found that some samples of asbestos are acted +upon by the phosphate solution and lose weight. An error from this +source may be avoided by determining the weight of the crucible +and filter after weighing the precipitate. For this purpose the +precipitate may be dissolved in dilute nitric acid, the asbestos +washed thoroughly, and the crucible reweighed.] + +[Note 8. The details of this method of precipitation of zinc are fully +discussed in an article by Dakin, !Ztschr. Anal. Chem.!, 39 (1900), +273.] + + + + +DETERMINATION OF SILICA IN SILICATES + + +Of the natural silicates, or artificial silicates such as slags and +some of the cements, a comparatively few can be completely decomposed +by treatment with acids, but by far the larger number require fusion +with an alkaline flux to effect decomposition and solution +for analysis. The procedure given below applies to silicates +undecomposable by acids, of which the mineral feldspar is taken as a +typical example. Modifications of the procedure, which are applicable +to silicates which are completely or partially decomposable by acids, +are given in the Notes on page 155. + + +PREPARATION OF THE SAMPLE + +Grind about 3 grams of the mineral in an agate mortar (Note 1) until +no grittiness is to be detected, or, better, until it will entirely +pass through a sieve made of fine silk bolting cloth. The sieve may be +made by placing a piece of the bolting cloth over the top of a small +beaker in which the ground mineral is placed, holding the cloth in +place by means of a rubber band below the lip of the beaker. By +inverting the beaker over clean paper and gently tapping it, the fine +particles pass through the sieve, leaving the coarser particles within +the beaker. These must be returned to the mortar and ground, and the +process of sifting and grinding repeated until the entire sample +passes through the sieve. + +[Note 1: If the sample of feldspar for analysis is in the massive or +crystalline form, it should be crushed in an iron mortar until the +pieces are about half the size of a pea, and then transferred to a +steel mortar, in which they are reduced to a coarse powder. A wooden +mallet should always be used to strike the pestle of the steel mortar, +and the blows should not be sharp. + +It is plain that final grinding in an agate mortar must be continued +until the whole of the portion of the mineral originally taken has +been ground so that it will pass the bolting cloth, otherwise the +sifted portion does not represent an average sample, the softer +ingredients, if foreign matter is present, being first reduced to +powder. For this reason it is best to start with not more than the +quantity of the feldspar needed for analysis. The mineral must be +thoroughly mixed after the grinding.] + + +FUSION AND SOLUTION + +PROCEDURE.--Weigh into platinum crucibles two portions of the ground +feldspar of about 0.8 gram each. Weigh on rough balances two portions +of anhydrous sodium carbonate, each amounting to about six times the +weight of the feldspar taken for analysis (Note 1). Pour about three +fourths of the sodium carbonate into the crucible, place the latter on +a piece of clean, glazed paper, and thoroughly mix the substance and +the flux by carefully stirring for several minutes with a dry glass +rod, the end of which has been recently heated and rounded in a flame +and slowly cooled. The rod may be wiped off with a small fragment of +filter paper, which may be placed in the crucible. Place the remaining +fourth of the carbonate on the top of the mixture. Cover the crucible, +heat it to dull redness for five minutes, and then gradually increase +the heat to the full capacity of a Bunsen or Tirrill burner for +twenty minutes, or until a quiet, liquid fusion is obtained (Note 2). +Finally, heat the sides and cover strongly until any material which +may have collected upon them is also brought to fusion. + +Allow the crucible to cool, and remove the fused mass as directed on +page 116. Disintegrate the mass by placing it in a previously prepared +mixture of 100 cc. of water and 50 cc. of dilute hydrochloric acid +(sp. gr. 1.12) in a covered casserole (Note 3). Clean the crucible and +lid by means of a little hydrochloric acid, adding this acid to the +main solution (Notes 4 and 5). + +[Note 1: Quartz, and minerals containing very high percentages of +silica, may require eight or ten parts by weight of the flux to insure +a satisfactory decomposition.] + +[Note 2: During the fusion the feldspar, which, when pure, is a +silicate of aluminium and either sodium or potassium, but usually +contains some iron, calcium, and magnesium, is decomposed by the +alkaline flux. The sodium of the latter combines with the silicic acid +of the silicate, with the evolution of carbon dioxide, while about two +thirds of the aluminium forms sodium aluminate and the remainder +is converted into basic carbonate, or the oxide. The calcium and +magnesium, if present, are changed to carbonates or oxides. + +The heat is applied gently to prevent a too violent reaction when +fusion first takes place.] + +[Note 3: The solution of a silicate by a strong acid is the result of +the combination of the H^{+} ions of the acid and the silicate ions +of the silicate to form a slightly ionized silicic acid. As a +consequence, the concentration of the silicate ions in the solution is +reduced nearly to zero, and more silicate dissolves to re-establish +the disturbed equilibrium. This process repeats itself until all of +the silicate is brought into solution. + +Whether the resulting solution of the silicate contains ortho-silicic +acid (H_{4}SiO_{4}) or whether it is a colloidal solution of some +other less hydrated acid, such as meta-silicic acid (H_{2}SiO_{3}), +is a matter that is still debatable. It is certain, however, that the +gelatinous material which readily separates from such solutions is of +the nature of a hydrogel, that is, a colloid which is insoluble in +water. This substance when heated to 100 deg.C., or higher, is completely +dehydrated, leaving only the anhydride, SiO_{2}. The changes may be +represented by the equation: + +SiO_{3}^{--} + 2H^{+} --> [H_{2}SiO_{3}] --> H_{2}O + SiO_{2}.] + +[Note 4: A portion of the fused mass is usually projected upward by +the escaping carbon dioxide during the fusion. The crucible must +therefore be kept covered as much as possible and the lid carefully +cleaned.] + +[Note 5: A gritty residue remaining after the disintegration of +the fused mass by acid indicates that the substance has been but +imperfectly decomposed. Such a residue should be filtered, washed, +dried, ignited, and again fused with the alkaline flux; or, if the +quantity of material at hand will permit, it is better to reject the +analysis, and to use increased care in grinding the mineral and in +mixing it with the flux.] + + +DEHYDRATION AND FILTRATION + +PROCEDURE.--Evaporate the solution of the fusion to dryness, stirring +frequently until the residue is a dry powder. Moisten the residue with +5 cc. of strong hydrochloric acid (sp. gr. 1.20) and evaporate again +to dryness. Heat the residue for at least one hour at a temperature +of 110 deg.C. (Note 1). Again moisten the residue with concentrated +hydrochloric acid, warm gently, making sure that the acid comes into +contact with the whole of the residue, dilute to about 200 cc. and +bring to boiling. Filter off the silica without much delay (Note 2), +and wash five times with warm dilute hydrochloric acid (one part +dilute acid (1.12 sp. gr.) to three parts of water). Allow the filter +to drain for a few moments, then place a clean beaker below the funnel +and wash with water until free from chlorides, discarding these +washings. Evaporate the original filtrate to dryness, dehydrate at +110 deg.C. for one hour (Note 3), and proceed as before, using a second +filter to collect the silica after the second dehydration. Wash this +filter with warm, dilute hydrochloric acid (Note 4), and finally with +hot water until free from chlorides. + +[Note 1: The silicic acid must be freed from its combination with +a base (sodium, in this instance) before it can be dehydrated. +The excess of hydrochloric acid accomplishes this liberation. By +disintegrating the fused mass with a considerable volume of dilute +acid the silicic acid is at first held in solution to a large extent. +Immediate treatment of the fused mass with strong acid is likely +to cause a semi-gelatinous silicic acid to separate at once and to +inclose alkali salts or alumina. + +A flocculent residue will often remain after the decomposition of the +fused mass is effected. This is usually partially dehydrated silicic +acid and does not require further treatment at this point. The +progress of the dehydration is indicated by the behavior of the +solution, which as evaporation proceeds usually gelatinizes. On this +account it is necessary to allow the solution to evaporate on a steam +bath, or to stir it vigorously, to avoid loss by spattering.] + +[Note 2: To obtain an approximately pure silica, the residue after +evaporation must be thoroughly extracted by warming with hydrochloric +acid, and the solution freely diluted to prevent, as far as possible, +the inclosure of the residual salts in the particles of silica. The +filtration should take place without delay, as the dehydrated silica +slowly dissolves in hydrochloric acid on standing.] + +[Note 3: It has been shown by Hillebrand that silicic acid cannot be +completely dehydrated by a single evaporation and heating, nor by +several such treatments, unless an intermediate filtration of the +silica occurs. If, however, the silica is removed and the filtrates +are again evaporated and the residue heated, the amount of silica +remaining in solution is usually negligible, although several +evaporations and filtrations are required with some silicates to +insure absolute accuracy. + +It is probable that temperatures above 100 deg.C. are not absolutely +necessary to dehydrate the silica; but it is recommended, as tending +to leave the silica in a better condition for filtration than when +the lower temperature of the water bath is used. This, and many other +points in the analysis of silicates, are fully discussed by Dr. +Hillebrand in the admirable monograph on "The Analysis of Silicate and +Carbonate Rocks," Bulletin No. 700 of the United States Geological +Survey. + +The double evaporation and filtration spoken of above are essential +because of the relatively large amount of alkali salts (sodium +chloride) present after evaporation. For the highest accuracy in the +determination of silica, or of iron and alumina, it is also necessary +to examine for silica the precipitate produced in the filtrate by +ammonium hydroxide by fusing it with acid potassium sulphate and +solution of the fused mass in water. The insoluble silica is filtered, +washed, and weighed, and the weight added to the weight of silica +previously obtained.] + +[Note 4: Aluminium and iron are likely to be thrown down as basic +salts from hot, very dilute solutions of their chlorides, as a result +of hydrolysis. If the silica were washed only with hot water, the +solution of these chlorides remaining in the filter after the passage +of the original filtrate would gradually become so dilute as to throw +down basic salts within the pores of the filter, which would remain +with the silica. To avoid this, an acid wash-water is used until the +aluminium and iron are practically removed. The acid is then removed +by water.] + + +IGNITION AND TESTING OF SILICA + +PROCEDURE.--Transfer the two washed filters belonging to each +determination to a platinum crucible, which need not be previously +weighed, and burn off the filter (Note 1). Ignite for thirty minutes +over the blast lamp with the cover on the crucible, and then for +periods of ten minutes, until the weight is constant. + +When a constant weight has been obtained, pour into the crucible about +3 cc. of water, and then 3 cc. of hydrofluoric acid. !This must be +done in a hood with a good draft and great care must be taken not to +come into contact with the acid or to inhale its fumes (Note 2!). + +If the precipitate has dissolved in this quantity of acid, add two +drops of concentrated sulphuric acid, and heat very slowly (always +under the hood) until all the liquid has evaporated, finally igniting +to redness. Cool in a desiccator, and weigh the crucible and residue. +Deduct this weight from the previous weight of crucible and impure +silica, and from the difference calculate the percentage of silica in +the sample (Note 3). + +[Note 1: The silica undergoes no change during the ignition beyond the +removal of all traces of water; but Hillebrand (!loc. cit.!) has shown +that the silica holds moisture so tenaciously that prolonged ignition +over the blast lamp is necessary to remove it entirely. This finely +divided, ignited silica tends to absorb moisture, and should be +weighed quickly.] + +[Note 2: Notwithstanding all precautions, the ignited precipitate of +silica is rarely wholly pure. It is tested by volatilisation of the +silica as silicon fluoride after solution in hydrofluoric acid, and, +if the analysis has been properly conducted, the residue, after +treatment with the acids and ignition, should not exceed 1 mg. + +The acid produces ulceration if brought into contact with the skin, +and its fumes are excessively harmful if inhaled.] + +[Note 3: The impurities are probably weighed with the original +precipitate in the form of oxides. The addition of the sulphuric +acid displaces the hydrofluoric acid, and it may be assumed that the +resulting sulphates (usually of iron or aluminium) are converted to +oxides by the final ignition. + +It is obvious that unless the sulphuric and hydrofluoric acids used +are known to leave no residue on evaporation, a quantity equal to that +employed in the analysis must be evaporated and a correction applied +for any residue found.] + +[Note 4: If the silicate to be analyzed is shown by a previous +qualitative examination to be completely decomposable, it may be +directly treated with hydrochloric acid, the solution evaporated to +dryness, and the silica dehydrated and further treated as described in +the case of the feldspar after fusion. + +A silicate which gelatinizes on treatment with acids should be mixed +first with a little water, and the strong acid added in small portions +with stirring, otherwise the gelatinous silicic acid incloses +particles of the original silicate and prevents decomposition. The +water, by separating the particles and slightly lessening the rapidity +of action, prevents this difficulty. This procedure is one which +applies in general to the solution of fine mineral powders in acids. + +If a small residue remains undecomposed by the treatment of the +silicate with acid, this may be filtered, washed, ignited and fused +with sodium carbonate and a solution of the fused mass added to the +original acid solution. This double procedure has an advantage, in +that it avoids adding so large a quantity of sodium salts as is +required for disintegration of the whole of the silicate by the fusion +method.] + + + + +PART IV + +STOICHIOMETRY + + +The problems with which the analytical chemist has to deal are not, as +a matter of actual fact, difficult either to solve or to understand. +That they appear difficult to many students is due to the fact that, +instead of understanding the principles which underlie each of the +small number of types into which these problems may be grouped, each +problem is approached as an individual puzzle, unrelated to others +already solved or explained. This attitude of mind should be carefully +avoided. + +It is obvious that ability to make the calculations necessary for +the interpretation of analytical data is no less important than the +manipulative skill required to obtain them, and that a moderate time +spent in the careful study of the solutions of the typical problems +which follow may save much later embarrassment. + +1. It is often necessary to calculate what is known as a "chemical +factor," or its equivalent logarithmic value called a "log factor," +for the conversion of the weight of a given chemical substance into an +equivalent weight of another substance. This is, in reality, a very +simple problem in proportion, making use of the atomic or molecular +weights of the substances in question which are chemically equivalent +to each other. One of the simplest cases of this sort is the +following: What is the factor for the conversion of a given weight of +barium sulphate (BaSO_{4}) into an equivalent weight of sulphur (S)? +The molecular weight of BaSO_{4} is 233.5. There is one atom of S in +the molecule and the atomic weight of S is 32.1. The chemical factor +is, therefore, 32.1/233.5, or 0.1375 and the weight of S corresponding +to a given weight of BaSO_{4} is found by multiplying the weight of +BaSO_{4} by this factor. If the problem takes the form, "What is +the factor for the conversion of a given weight of ferric oxide +(Fe_{2}O_{3}) into ferrous oxide (FeO), or of a given weight of +mangano-manganic oxide (Mn_{3}O_{4}) into manganese (Mn)?" the +principle involved is the same, but it must then be noted that, in the +first instance, each molecule of Fe_{2}O_{3} will be equivalent to two +molecules of FeO, and in the second instance that each molecule of +Mn_{3}O_{4} is equivalent to three atoms of Mn. The respective factors +then become + +(2FeO/Fe_{2}O_{3}) or (143.6/159.6) and (3Mn/Mn_{3}O_{4}) or +(164.7/228.7). + +It is obvious that the arithmetical processes involved in this type +of problem are extremely simple. It is only necessary to observe +carefully the chemical equivalents. It is plainly incorrect to express +the ratio of ferrous to ferric oxide as (FeO/Fe_{2}O_{3}), since each +molecule of the ferric oxide will yield two molecules of the ferrous +oxide. Mistakes of this sort are easily made and constitute one of the +most frequent sources of error. + +2. A type of problem which is slightly more complicated in appearance, +but exactly comparable in principle, is the following: "What is the +factor for the conversion of a given weight of ferrous sulphate +(FeSO_{4}), used as a reducing agent against potassium permanganate, +into the equivalent weight of sodium oxalate (Na_{2}C_{2}O_{4})?" To +determine the chemical equivalents in such an instance it is necessary +to inspect the chemical reactions involved. These are: + +10FeSO_{4} + 2KMnO_{4} + 8H_{2}SO_{4} --> 5Fe_{2}(SO_{4})_{3} + +K_{2}SO_{4} + 2MnSO_{4} + 8H_{2}O, + +5Na_{2}C_{2}O_{4} + 2KMnO_{4} + 8H_{2}SO_{4} --> 5Na_{2}SO_{4} + +10CO_{2} + K_{2}SO_{4} + 2MnSO_{4} + 8H_{2}O. + +It is evident that 10FeSO_{4} in the one case, and 5Na_{2}C_{2}O_{4} +in the other, each react with 2KMnO_{4}. These molecular +quantities are therefore equivalent, and the factor becomes +(10FeSO_{4}/5Na_{2}C_{2}O_{4}) or (2FeSO_{4}/Na_{2}C_{2}O_{4}) or +(303.8/134). + +Again, let it be assumed that it is desired to determine the +factor required for the conversion of a given weight of potassium +permanganate (KMnO_{4}) into an equivalent weight of potassium +bichromate (K_{2}Cr_{2}O_{7}), each acting as an oxidizing agent +against ferrous sulphate. The reactions involved are: + +10FeSO_{4} + 2KMnO_{4} + 8H_{2}SO_{4} --> 5Fe_{2}(SO_{4})_{3} + +K_{2}SO_{4} + 2MnSO_{4} + 8H_{2}O, + +6FeSO_{4} + K_{2}Cr_{2}O_{7} + 7H_{2}SO_{4} --> 3Fe_{2}(SO_{3})_{3} + +K_{2}SO_{4} + Cr_{2}(SO_{4})_{3} + 7H_{2}O. + +An inspection of these equations shows that 2KMO_{4} react with +10FeSO_{4}, while K_{2}Cr_{2}O_{7} reacts with 6FeSO_{4}. These are +not equivalent, but if the first equation is multiplied by 3 and the +second by 5 the number of molecules of FeSO_{4} is then the same in +both, and the number of molecules of KMnO_{4} and K_{2}Cr_{2}O_{7} +reacting with these 30 molecules become 6 and 5 respectively. These +are obviously chemically equivalent and the desired factor is +expressed by the fraction (6KMnO_{4}/5K_{2}Cr_{2}O_{7}) or +(948.0/1471.0). + +3. It is sometimes necessary to calculate the value of solutions +according to the principles just explained, when several successive +reactions are involved. Such problems may be solved by a series of +proportions, but it is usually possible to eliminate the common +factors and solve but a single one. For example, the amount of MnO_{2} +in a sample of the mineral pyrolusite may be determined by dissolving +the mineral in hydrochloric acid, absorbing the evolved chlorine in a +solution of potassium iodide, and measuring the liberated iodine +by titration with a standard solution of sodium thiosulphate. The +reactions involved are: + +MnO_{2} + 4HCl --> MnCl_{2} + 2H_{2}O + Cl_{2} +Cl_{2} + 2KI --> I_{2} + 2KCl +I_{2} + 2Na_{2}S_{2}O_{3} --> 2NaI + Na_{2}S_{4}O_{6} + +Assuming that the weight of thiosulphate corresponding to the +volume of sodium thiosulphate solution used is known, what is the +corresponding weight of manganese dioxide? From the reactions given +above, the following proportions may be stated: + +2Na_{2}S_{2}O_{3}:I_{2} = 316.4:253.9, + +I_{2}:Cl_{2} = 253.9:71, + +Cl_{2}:MnO_{2} = 71:86.9. + +After canceling the common factors, there remains +2Na_{2}S_{2}O_{3}:MnO_{2} = 316.4:86.9, and the factor for the +conversion of thiosulphate into an equivalent of manganese dioxide is +86.9/316.4. + +4. To calculate the volume of a reagent required for a specific +operation, it is necessary to know the exact reaction which is to be +brought about, and, as with the calculation of factors, to keep in +mind the molecular relations between the reagent and the substance +reacted upon. For example, to estimate the weight of barium chloride +necessary to precipitate the sulphur from 0.1 gram of pure pyrite +(FeS_{2}), the proportion should read + + 488. 120.0 + 2(BaCl_{2}.2H_{2}O):FeS_{2} = x:0.1, + +where !x! represents the weight of the chloride required. Each of the +two atoms of sulphur will form upon oxidation a molecule of sulphuric +acid or a sulphate, which, in turn, will require a molecule of the +barium chloride for precipitation. To determine the quantity of the +barium chloride required, it is necessary to include in its molecular +weight the water of crystallization, since this is inseparable from +the chloride when it is weighed. This applies equally to other similar +instances. + +If the strength of an acid is expressed in percentage by weight, due +regard must be paid to its specific gravity. For example, hydrochloric +acid (sp. gr. 1.12) contains 23.8 per cent HCl !by weight!; that is, +0.2666 gram HCl in each cubic centimeter. + +5. It is sometimes desirable to avoid the manipulation required for +the separation of the constituents of a mixture of substances by +making what is called an "indirect analysis." For example, in the +analysis of silicate rocks, the sodium and potassium present may be +obtained in the form of their chlorides and weighed together. If the +weight of such a mixture is known, and also the percentage of chlorine +present, it is possible to calculate the amount of each chloride in +the mixture. Let it be assumed that the weight of the mixed chlorides +is 0.15 gram, and that it contains 53 per cent of chlorine. + +The simplest solution of such a problem is reached through algebraic +methods. The weight of chlorine is evidently 0.15 x 0.53, or 0.0795 +gram. Let x represent the weight of sodium chloride present and y +that of potassium chloride. The molecular weight of NaCl is 58.5 and +that of KCl is 74.6. The atomic weight of chlorine is 35.5. Then + +x + y = 0.15 +(35.5/58.5)x + (35.5/74.6)y = 0.00795 + +Solving these equations for x shows the weight of NaCl to be 0.0625 +gram. The weight of KCl is found by subtracting this from 0.15. + +The above is one of the most common types of indirect analyses. Others +are more complex but they can be reduced to algebraic expressions and +solved by their aid. It should, however, be noted that the results +obtained by these indirect methods cannot be depended upon for high +accuracy, since slight errors in the determination of the common +constituent, as chlorine in the above mixture, will cause considerable +variations in the values found for the components. They should not be +employed when direct methods are applicable, if accuracy is essential. + + + + +PROBLEMS + + +(The reactions necessary for the solution of these problems are either +stated with the problem or may be found in the earlier text. In the +calculations from which the answers are derived, the atomic weights +given on page 195 have been employed, using, however, only the first +decimal but increasing this by 1 when the second decimal is 5 or +above. Thus, 39.1 has been taken as the atomic weight of potassium, +32.1 for sulphur, etc. This has been done merely to secure uniformity +of treatment, and the student should remember that it is always well +to take into account the degree of accuracy desired in a particular +instance in determining the number of decimal places to retain. +Four-place logarithms were employed in the calculations. Where four +figures are given in the answer, the last figure may vary by one or +(rarely) by two units, according to the method by which the problem is +solved.) + + +VOLUMETRIC ANALYSIS + +1. How many grams of pure potassium hydroxide are required for exactly +1 liter of normal alkali solution? + +!Answer!: 56.1 grams. + +2. Calculate the equivalent in grams (a) of sulphuric acid as an acid; +(b) of hydrochloric acid as an acid; (c) of oxalic acid as an acid; +(d) of nitric acid as an acid. + +!Answers!: (a) 49.05; (b) 36.5; (c) 63; (d) 63. + +3. Calculate the equivalent in grams of (a) potassium hydroxide; +(b) of sodium carbonate; (c) of barium hydroxide; (d) of sodium +bicarbonate when titrated with an acid. + +!Answers!: (a) 56.1; (b) 53.8; (c) 85.7; (d) 84. + +4. What is the equivalent in grams of Na_{2}HPO_{4} (a) as a +phosphate; (b) as a sodium salt? + +!Answers!: (a) 47.33; (b) 71.0. + +5. A sample of aqueous hydrochloric acid has a specific gravity +of 1.12 and contains 23.81 per cent hydrochloric acid by weight. +Calculate the grams and the milliequivalents of hydrochloric acid +(HCl) in each cubic centimeter of the aqueous acid. + +!Answers!: 0.2667 gram; 7.307 milliequivalents. + +6. How many cubic centimeters of hydrochloric acid (sp. gr. 1.20 +containing 39.80 per cent HCl by weight) are required to furnish 36.45 +grams of the gaseous compound? + +!Answer!: 76.33 cc. + +7. A given solution contains 0.1063 equivalents of hydrochloric acid +in 976 cc. What is its normal value? + +!Answer!: 0.1089 N. + +8. In standardizing a hydrochloric acid solution it is found that +47.26 cc. of hydrochloric acid are exactly equivalent to 1.216 grams +of pure sodium carbonate, using methyl orange as an indicator. What is +the normal value of the hydrochloric acid? + +!Answer!: 0.4855 N. + +9. Convert 42.75 cc. of 0.5162 normal hydrochloric acid to the +equivalent volume of normal hydrochloric acid. + +!Answer!: 22.07 cc. + +10. A solution containing 25.27 cc. of 0.1065 normal hydrochloric acid +is added to one containing 92.21 cc. of 0.5431 normal sulphuric acid +and 50 cc. of exactly normal potassium hydroxide added from a pipette. +Is the solution acid or alkaline? How many cubic centimeters of +0.1 normal acid or alkali must be added to exactly neutralize the +solution? + +!Answer!: 27.6 cc. alkali (solution is acid). + +11. By experiment the normal value of a sulphuric acid solution is +found to be 0.5172. Of this acid 39.65 cc. are exactly equivalent to +21.74 cc. of a standard alkali solution. What is the normal value of +the alkali? + +!Answer!: 0.9432 N. + +12. A solution of sulphuric acid is standardized against a sample of +calcium carbonate which has been previously accurately analyzed and +found to contain 92.44% CaCO_{3} and no other basic material. The +sample weighing 0.7423 gram was titrated by adding an excess of acid +(42.42 cc.) and titrating the excess with sodium hydroxide solution +(11.22 cc.). 1 cc. of acid is equivalent to 0.9976 cc. of sodium +hydroxide. Calculate the normal value of each. + +!Answers!: Acid 0.4398 N; alkali 0.4409 N. + +13. Given five 10 cc. portions of 0.1 normal hydrochloric acid, (a) +how many grams of silver chloride will be precipitated by a portion +when an excess of silver nitrate is added? (b) how many grams of pure +anhydrous sodium carbonate (Na_{2}CO_{3}) will be neutralized by a +portion of it? (c) how many grams of silver will there be in the +silver chloride formed when an excess of silver nitrate is added to a +portion? (d) how many grams of iron will be dissolved to FeCl_{2} by a +portion of it? (e) how many grams of magnesium chloride will be formed +and how many grams of carbon dioxide liberated when an excess of +magnesium carbonate is treated with a portion of the acid? + +!Answers!: (a) 0.1434; (b) 0.053; (c) 0.1079; (d) 0.0279; (e) 0.04765, +and 0.022. + +14. If 30.00 grams of potassium tetroxalate +(KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O) are dissolved and the solution +diluted to exactly 1 liter, and 40 cc. are neutralized with 20 cc. +of a potassium carbonate solution, what is the normal value of the +carbonate solution? + +!Answer!: 0.7084 N. + +15. How many cubic centimeters of 0.3 normal sulphuric acid will be +required to neutralize (a) 30 cc. of 0.5 normal potassium hydroxide; +(b) to neutralize 30 cc. of 0.5 normal barium hydroxide; (c) to +neutralize 20 cc. of a solution containing 10.02 grams of potassium +bicarbonate per 100 cc.; (d) to give a precipitate of barium sulphate +weighing 0.4320 gram? + +!Answers!: (a) 50 cc.; (b) 50 cc.; (c) 66.73 cc.; (d) 12.33 cc. + +16. It is desired to dilute a solution of sulphuric acid of which 1 +cc. is equivalent to 0.1027 gram of pure sodium carbonate to make it +exactly 1.250 normal. 700 cc. of the solution are available. To what +volume must it be diluted? + +!Answer!: 1084 cc. + +17. Given the following data: 1 cc. of NaOH = 1.117 cc. HCl. The HCl +is 0.4876 N. How much water must be added to 100 cc. of the alkali to +make it exactly 0.5 N.? + +!Answer!: 9.0 cc. + +18. What is the normal value of a sulphuric acid solution which has a +specific gravity of 1.839 and contains 95% H_{2}SO_{4} by weight? + +!Answer!: 35.61 N. + +19. A sample of Rochelle Salt (KNaC_{4}H_{4}O_{6}.4H_{2}O), after +ignition in platinum to convert it to the double carbonate, is +titrated with sulphuric acid, using methyl orange as an indicator. +From the following data calculate the percentage purity of the sample: + +Wt. sample = 0.9500 gram +H_{2}SO_{4} used = 43.65 cc. +NaOH used = 1.72 cc. +1 cc. H_{2}SO_{4} = 1.064 cc. NaOH +Normal value NaOH = 0.1321 N. + +!Answer!: 87.72 cc. + +20. One gram of a mixture of 50% sodium carbonate and 50% potassium +carbonate is dissolved in water, and 17.36 cc. of 1.075 N acid is +added. Is the resulting solution acid or alkaline? How many cubic +centimeters of 1.075 N acid or alkali will have to be added to make +the solution exactly neutral? + +!Answers!: Acid; 1.86 cc. alkali. + +21. In preparing an alkaline solution for use in volumetric work, an +analyst, because of shortage of chemicals, mixed exactly 46.32 grams +of pure KOH and 27.64 grams of pure NaOH, and after dissolving in +water, diluted the solution to exactly one liter. How many cubic +centimeters of 1.022 N hydrochloric acid are necessary to neutralize +50 cc. of the basic solution? + +!Answer!: 74.18 cc. + +22. One gram of crude ammonium salt is treated with strong potassium +hydroxide solution. The ammonia liberated is distilled and collected +in 50 cc. of 0.5 N acid and the excess titrated with 1.55 cc. of 0.5 N +sodium hydroxide. Calculate the percentage of NH_{3} in the sample. + +!Answer!: 41.17%. + + +23. In titrating solutions of alkali carbonates in the presence of +phenolphthalein, the color change takes place when the carbonate has +been converted to bicarbonate. In the presence of methyl orange, the +color change takes place only when the carbonate has been completely +neutralized. From the following data, calculate the percentages of +Na_{2}CO_{3} and NaOH in an impure mixture. Weight of sample, 1.500 +grams; HCl (0.5 N) required for phenolphthalein end-point, 28.85 cc.; +HCl (0.5 N) required to complete the titration after adding methyl +orange, 23.85 cc. + +!Answers!: 6.67% NaOH; 84.28% Na_{2}CO_{3}. + +24. A sample of sodium carbonate containing sodium hydroxide weighs +1.179 grams. It is titrated with 0.30 N hydrochloric acid, using +phenolphthalein in cold solution as an indicator and becomes colorless +after the addition of 48.16 cc. Methyl orange is added and 24.08 cc. +are needed for complete neutralization. What is the percentage of NaOH +and Na_{2}CO_{3}? + +!Answers!: 24.50% NaOH; 64.92% Na_{2}CO_{3}. + +25. From the following data, calculate the percentages of Na_{2}CO_{3} +and NaHCO_{3} in an impure mixture. Weight of sample 1.000 gram; +volume of 0.25 N hydrochloric acid required for phenolphthalein +end-point, 26.40 cc.; after adding an excess of acid and boiling out +the carbon dioxide, the total volume of 0.25 N hydrochloric acid +required for phenolphthalein end-point, 67.10 cc. + +!Answer!: 69.95% Na_{2}CO_{3}; 30.02% NaHCO_{3}. + +26. In the analysis of a one-gram sample of soda ash, what must be the +normality of the acid in order that the number of cubic centimeters of +acid used shall represent the percentage of carbon dioxide present? + +!Answer!: 0.4544 gram. + +27. What weight of pearl ash must be taken for analysis in order that +the number of cubic centimeters of 0.5 N acid used may be equal to one +third the percentage of K_{2}CO_{3}? + +!Answer!: 1.152 grams. + +28. What weight of cream of tartar must have been taken for analysis +in order to have obtained 97.60% KHC_{4}H_{4}O_{6} in an analysis +involving the following data: NaOH used = 30.06 cc.; H_{2}SO_{4} +solution used = 0.50 cc.; 1 cc. H_{2}SO_{4} sol. = 0.0255 gram +CaCO_{3}; 1 cc. H_{2}SO_{4} sol. = 1.02 cc. NaOH sol.? + +!Answer!: 2.846 grams. + +29. Calculate the percentage of potassium oxide in an impure sample of +potassium carbonate from the following data: Weight of sample = 1.00 +gram; HCl sol. used = 55.90 cc.; NaOH sol. used = 0.42 cc.; 1 cc. NaOH +sol. = 0.008473 gram of KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O; 2 cc. +HCl sol. = 5 cc. NaOH sol. + +!Answer!: 65.68%. + +30. Calculate the percentage purity of a sample of calcite +(CaCO_{3}) from the following data: (Standardization); Weight of +H_{2}C_{2}O_{4}.2H_{2}O = 0.2460 gram; NaOH solution used = 41.03 +cc.; HCl solution used = 0.63; 1 cc. NaOH solution = 1.190 cc. HCl +solution. (Analysis); Weight of sample 0.1200 gram; HCl used = 36.38 +cc.; NaOH used = 6.20 cc. + +!Answer!: 97.97%. + +31. It is desired to dilute a solution of hydrochloric acid to exactly +0.05 N. The following data are given: 44.97 cc. of the hydrochloric +acid are equivalent to 43.76 cc. of the NaOH solution. The NaOH +is standardized against a pure potassium tetroxalate +(KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O) weighing 0.2162 gram and +requires 49.14 cc. How many cc. of water must be added to 1000 cc. of +the aqueous hydrochloric acid? + +!Answer!: 11 cc. + +32. How many cubic centimeters of 3 N phosphoric acid must be added to +300 cc. of 0.4 N phosphoric acid in order that the resulting solution +may be 0.6 N? + +!Answer!: 25 cc. + +33. To oxidize the iron in 1 gram of +FeSO_{4}(NH_{4})_{2}SO_{4}.6H_{2}O (mol. wgt. 392) requires 3 cc. of +a given solution of HNO_{3}. What is the normality of the nitric +acid when used as an acid? 6FeSO_{4} + 2HNO_{3} + 2H_{2}SO_{4} = +3Fe_{2}(SO_{4})_{3} + 2NO + 4H_{2}O. + +!Answer!: 0.2835 N. + +34. The same volume of carbon dioxide at the same temperature and the +same pressure is liberated from a 1 gram sample of dolomite, by adding +an excess of hydrochloric acid, as can be liberated by the addition of +35 cc. of 0.5 N hydrochloric acid to an excess of any pure or impure +carbonate. Calculate the percentage of CO_{2} in the dolomite. + +!Answer!: 38.5%. + +35. How many cubic centimeters of sulphuric acid (sp. gr. 1.84, +containing 96% H_{2}SO_{4} by weight) will be required to displace the +chloride in the calcium chloride formed by the action of 100 cc. of +0.1072 N hydrochloric acid on an excess of calcium carbonate, and how +many grams of CaSO_{4} will be formed? + +!Answers!: 0.298 cc.; 0.7300 gram. + +36. Potassium hydroxide which has been exposed to the air is found on +analysis to contain 7.62% water, 2.38% K_{2}CO_{3}. and 90% KOH. What +weight of residue will be obtained if one gram of this sample is added +to 46 cc. of normal hydrochloric acid and the resulting solution, +after exact neutralization with 1.070 N potassium hydroxide solution, +is evaporated to dryness? + +!Answer!: 3.47 grams. + +37. A chemist received four different solutions, with the statement +that they contained either pure NaOH; pure Na_{2}CO_{3}; pure +NaHCO_{3}, or mixtures of these substances. From the following data +identify them: + +Sample I. On adding phenolphthalein to a solution of the substance, it +gave no color to the solution. + +Sample II. On titrating with standard acid, it required 15.26 cc. for +a change in color, using phenolphthalein, and 17.90 cc. additional, +using methyl orange as an indicator. + +Sample III. The sample was titrated with hydrochloric acid until the +pink of phenolphthalein disappeared, and on the addition of methyl +orange the solution was colored pink. + +Sample IV. On titrating with hydrochloric acid, using phenolphthalein, +15.00 cc. were required. A new sample of the same weight required +exactly 30 cc. of the same acid for neutralization, using methyl +orange. + +!Answers!: (a) NaHCO_{3}; (b) NaHCO_{3}+Na_{2}CO_{3}; (c)NaOH; (d) +Na_{2}CO_{3}. + +38. In the analysis of a sample of KHC_{4}H_{4}O_{6} the following +data are obtained: Weight sample = 0.4732 gram. NaOH solution used = +24.97 cc. 3.00 cc. NaOH = 1 cc. of H_{3}PO_{4} solution of which 1 +cc. will precipitate 0.01227 gram of magnesium as MgNH_{4}PO_{4}. +Calculate the percentage of KHC_{4}H_{4}O_{6}. + +!Answer!: 88.67%. + +39. A one-gram sample of sodium hydroxide which has been exposed to +the air for some time, is dissolved in water and diluted to exactly +500 cc. One hundred cubic centimeters of the solution, when titrated +with 0.1062 N hydrochloric acid, using methyl orange as an indicator, +requires 38.60 cc. for complete neutralization. Barium chloride in +excess is added to a second portion of 100 cc. of the solution, which +is diluted to exactly 250 cc., allowed to stand and filtered. Two +hundred cubic centimeters of this filtrate require 29.62 cc. of 0.1062 +N hydrochloric acid for neutralization, using phenolphthalein as an +indicator. Calculate percentage of NaOH, Na_{2}CO_{3}, and H_{2}O. + +!Answers!: 78.63% NaOH; 4.45% Na_{2}CO_{3}; 16.92% H_{2}O. + +40. A sodium hydroxide solution (made from solid NaOH which has been +exposed to the air) was titrated against a standard acid using methyl +orange as an indicator, and was found to be exactly 0.1 N. This +solution was used in the analysis of a material sold at 2 cents per +pound per cent of an acid constituent A, and always mixed so that +it was supposed to contain 15% of A, on the basis of the analyst's +report. Owing to the carelessness of the analyst's assistant, the +sodium hydroxide solution was used with phenolphthalein as an +indicator in cold solution in making the analyses. The concern +manufacturing this material sells 600 tons per year, and when the +mistake was discovered it was estimated that at the end of a year +the error in the use of indicators would either cost them or their +customers $6000. Who would lose and why? Assuming the impure NaOH used +originally in making the titrating solution consisted of NaOH and +Na_{2}CO_{3} only, what per cent of each was present? + +!Answers!: Customer lost; 3.94% Na_{2}CO_{3}; 96.06% NaOH. + +41. In the standardization of a K_{2}Cr_{2}O_{7} solution against iron +wire, 99.85% pure, 42.42 cc. of the solution were added. The weight of +the wire used was 0.22 gram. 3.27 cc. of a ferrous sulphate solution +having a normal value as a reducing agent of 0.1011 were added +to complete the titration. Calculate the normal value of the +K_{2}Cr_{2}O_{7}. + +!Answer!: 0.1006 N. + +42. What weight of iron ore containing 56.2% Fe should be taken to +standardize an approximately 0.1 N oxidizing solution, if not more +than 47 cc. are to be used? + +!Answer!: 0.4667 gram. + +43. One tenth gram of iron wire, 99.78% pure, is dissolved in +hydrochloric acid and the iron oxidized completely with bromine water. +How many grams of stannous chloride are there in a liter of solution +if it requires 9.47 cc. to just reduce the iron in the above? What +is the normal value of the stannous chloride solution as a reducing +agent? + +!Answer!: 17.92 grams; 0.1888 N. + +44. One gram of an oxide of iron is fused with potassium acid sulphate +and the fusion dissolved in acid. The iron is reduced with stannous +chloride, mercuric chloride is added, and the iron titrated with a +normal K_{2}Cr_{2}O_{7} solution. 12.94 cc. were used. What is the +formula of the oxide, FeO, Fe_{2}O_{3}, or Fe_{3}O_{4}? + +!Answer!: Fe_{3}O_{4}. + +45. If an element has 98 for its atomic weight, and after reduction +with stannous chloride could be oxidized by bichromate to a state +corresponding to an XO_{4}^{-} anion, compute the oxide, or valence, +corresponding to the reduced state from the following data: 0.3266 +gram of the pure element, after being dissolved, was reduced with +stannous chloride and oxidized by 40 cc. of K_{2}Cr_{2}O_{7}, of which +one cc. = 0.1960 gram of FeSO_{4}(NH_{4})_{2}SO_{4}.6H_{2}O. + +!Answer!: Monovalent. + +46. Determine the percentage of iron in a sample of limonite from the +following data: Sample = 0.5000 gram. KMnO_{4} used = 50 cc. 1 cc. +KMnO_{4} = 0.005317 gram Fe. FeSO_{4} used = 6 cc. 1 cc. FeSO_{4} = +0.009200 gram FeO. + +!Answer!: 44.60%. + +47. If 1 gram of a silicate yields 0.5000 gram of Fe_{2}O_{3} and +Al_{2}O_{3} and the iron present requires 25 cc. of 0.2 N KMnO_{4}, +calculate the percentage of FeO and Al_{2}O_{3} in the sample. + +!Answer!: 35.89% FeO; 10.03% Al_{2}O_{3}. + +48. A sample of magnesia limestone has the following composition: +Silica, 3.00%; ferric oxide and alumina, 0.20%; calcium oxide, 33.10%; +magnesium oxide, 20.70%; carbon dioxide, 43.00%. In manufacturing lime +from the above the carbon dioxide is reduced to 3.00%. How many cubic +centimeters of normal KMnO_{4} will be required to determine the +calcium oxide volumetrically in a 1 gram sample of the lime? + +!Answer!: 20.08 cc. + +49. If 100 cc. of potassium bichromate solution (10 gram +K_{2}Cr_{2}O_{7} per liter), 5 cc. of 6 N sulphuric acid, and 75 cc. +of ferrous sulphate solution (80 grams FeSO_{4}.7H_{2}O per liter) are +mixed, and the resulting solution titrated with 0.2121 N KMnO_{4}, how +many cubic centimeters of the KMnO_{4} solution will be required to +oxidize the iron? + +!Answer!: 5.70 cc. + +50. If a 0.5000 gram sample of limonite containing 59.50 per cent +Fe_{2}O_{3} requires 40 cc. of KMnO_{4} to oxidize the iron, what +is the value of 1 cc. of the permanganate in terms of (a) Fe, (b) +H_{2}C_{2}O_{4}.2H_{2}O? + +!Answers!: (a) 0.005189 gram; (b) 0.005859 gram. + +51. A sample of pyrolusite weighing 0.6000 gram is treated with 0.9000 +gram of oxalic acid. The excess oxalic acid requires 23.95 cc. of +permanganate (1 cc. = 0.03038 gram FeSO_{4}.7H_{2}O). What is the +percentage of MnO_{2}, in the sample? + +!Answer!: 84.47%. + +52. A solution contains 50 grams of +KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O per liter. What is the normal +value of the solution (a) as an acid, and (b) as a reducing agent? + +!Answers!: (a) 0.5903 N; (b) 0.7872 N. + +53. In the analysis of an iron ore containing 60% Fe_{2}O_{3}, a +sample weighing 0.5000 gram is taken and the iron is reduced with +sulphurous acid. On account of failure to boil out all the excess +SO_{2}, 38.60 cubic centimeters of 0.1 N KMnO_{4} were required to +titrate the solution. What was the error, percentage error, and what +weight of sulphur dioxide was in the solution? + +!Answers!: (a) 1.60%; (b) 2.67%; (c) 0.00322 gram. + +54. From the following data, calculate the ratio of the nitric acid as +an oxidizing agent to the tetroxalate solution as a reducing agent: +1 cc. HNO_{3} = 1.246 cc. NaOH solution; 1 cc. NaOH = 1.743 cc. +KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O solution; Normal value NaOH = +0.12. + +!Answer!: 4.885. + +55. Given the following data: 25 cc. of a hydrochloric acid, when +standardized gravimetrically as silver chloride, yields a precipitate +weighing 0.5465 gram. 24.35 cc. of the hydrochloric acid are exactly +equivalent to 30.17 cc. of KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O +solution. How much water must be added to a liter of the oxalate +solution to make it exactly 0.025 N as a reducing agent? + +!Answer!: 5564 cc. + +56. Ten grams of a mixture of pure potassium tetroxalate +(KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O) and pure oxalic acid +(H_{2}C_{2}O_{4}.2H_{2}O) are dissolved in water and diluted to +exactly 1000 cc. The normal value of the oxalate solution when used as +an acid is 0.1315. Calculate the ratio of tetroxalate to oxalate used +in making up the solution and the normal value of the solution as a +reducing agent. + +!Answers!: 2:1; 0.1577 N. + +57. A student standardized a solution of NaOH and one of KMnO_{4} +against pure KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O and found the former +to be 0.07500 N as an alkali and the latter exactly 0.1 N as an +oxidizing agent. By coincidence, exactly 47.26 cc. were used in each +standardization. Find the ratio of the oxalate used in the +NaOH standardization to the oxalate used in the permanganate +standardization. + +!Answer!: 1:1. + +58. A sample of apatite weighing 0.60 gram is analyzed for its +phosphoric anhydride content. If the phosphate is precipitated as +(NH_{4})_{3}PO_{4}.12MoO_{3}, and the precipitate (after solution and +reduction of the MoO_{3} to Mo_{24}O_{37}), requires 100 cc. of normal +KMnO_{4} to oxidize it back to MoO_{3}, what is the percentage of +P_{2}O_{5}? + +!Answer!: 33.81%. + +59. In the analysis of a sample of steel weighing 1.881 grams the +phosphorus was precipitated with ammonium molybdate and the yellow +precipitate was dissolved, reduced and titrated with KMnO_{4}. If the +sample contained 0.025 per cent P and 6.01 cc. of KMnO_{4} were used, +to what oxide was the molybdenum reduced? 1 cc. KMnO_{4} = 0.007188 +gram Na_{2}C_{2}O_{4}. + +!Answer!: Mo_{4}O_{5}. + +60. What is the value of 1 cc. of an iodine solution (1 cc. equivalent +to 0.0300 gram Na_{2}S_{2}O_{3}) in terms of As_{2}O_{3}? + +!Answer!: 0.009385 gram. + +61. 48 cc. of a solution of sodium thiosulphate are required to +titrate the iodine liberated from an excess of potassium iodide +solution by 0.3000 gram of pure KIO_{3}. (KIO_{3} + 5KI + 3H_{2}SO_{4} += 3K_{2}SO_{4} + 3I_{2} + 3H_{2}O.) What is the normal strength of the +sodium thiosulphate and the value of 1 cc. of it in terms of iodine? + +!Answers!: 0.1753 N; 0.02224 gram. + +62. One thousand cubic centimeters of 0.1079 N sodium thiosulphate +solution is allowed to stand. One per cent by weight of the +thiosulphate is decomposed by the carbonic acid present in the +solution. To what volume must the solution be diluted to make it +exactly 0.1 N as a reducing agent? (Na_{2}S_{2}O_{3} + 2H_{2}CO_{3} = +H_{2}SO_{3} + 2NaHCO_{3} + S.) + +!Answer!: 1090 cc. + +63. An analyzed sample of stibnite containing 70.05% Sb is given for +analysis. A student titrates it with a solution of iodine of which 1 +cc. is equivalent to 0.004950 gram of As_{2}O_{3}. Due to an error on +his part in standardization, the student's analysis shows the sample +to contain 70.32% Sb. Calculate the true normal value of the iodine +solution, and the percentage error in the analysis. + +!Answers!: 0.1000 N; 0.39%. + +64. A sample of pyrolusite weighing 0.5000 gram is treated with an +excess of hydrochloric acid, the liberated chlorine is passed into +potassium iodide and the liberated iodine is titrated with sodium +thiosulphate solution (49.66 grams of pure Na_{2}S_{2}O_{3}.5H_{2}O +per liter). If 38.72 cc. are required, what volume of 0.25 normal +permanganate solution will be required in an indirect determination +in which a similar sample is reduced with 0.9012 gram +H_{2}C_{2}O_{4}.2H_{2}O and the excess oxalic acid titrated? + +!Answer!: 26.22 cc. + +65. In the determination of sulphur in steel by evolving the sulphur +as hydrogen sulphide, precipitating cadmium sulphide by passing the +liberated hydrogen sulphide through ammoniacal cadmium chloride +solution, and decomposing the CdS with acid in the presence of a +measured amount of standard iodine, the following data are obtained: +Sample, 5.027 grams; cc. Na_{2}S_{2}O_{3} sol. = 12.68; cc. Iodine +sol. = 15.59; 1 cc. Iodine sol. = 1.086 cc. Na_{2}S_{2}O_{3} sol.; 1 +cc. Na_{2}S_{2}O_{3}= 0.005044 gram Cu. Calculate the percentage of +sulphur. (H_{2}S + I_{2} = 2HI + S.) + +!Answer!: 0.107%. + +66. Given the following data, calculate the percentage of iron in +a sample of crude ferric chloride weighing 1.000 gram. The iodine +liberated by the reaction 2FeCl_{3}+ 2HI = 2HCl + 2FeCl_{2} + I_{2} is +reduced by the addition of 50 cc. of sodium thiosulphate solution and +the excess thiosulphate is titrated with standard iodine and requires +7.85 cc. 45 cc. I_{2} solution = 45.95 cc. Na_{2}S_{2}O_{3} solution; +45 cc. As_{2}O_{3} solution = 45.27 cc. I_{2} solution. 1 cc. arsenite +solution = 0.005160 gram As_{2}O_{3}. + +!Answer!: 23.77%. + +67. Sulphide sulphur was determined in a sample of reduced barium +sulphate by the evolution method, in which the sulphur was evolved as +hydrogen sulphide and was passed into CdCl_{2} solution, the acidified +precipitate being titrated with iodine and thiosulphate. Sample, 5.076 +grams; cc. I_{2} = 20.83; cc. Na_{2}S_{2}O_{3} = 12.37; 43.45 cc. +Na_{2}S_{2}O_{3} = 43.42 cc. I_{2}; 8.06 cc. KMnO_{4} = 44.66 cc. +Na_{2}S_{2}O_{3}; 28.87 cc. KMnO_{4} = 0.2004 gram Na_{2}C_{2}O_{4}. +Calculate the percentage of sulphide sulphur in the sample. + +!Answer!: 0.050%. + +68. What weight of pyrolusite containing 89.21% MnO_{2} will oxidize +the same amount of oxalic acid as 37.12 cc. of a permanganate +solution, of which 1 cc. will liberate 0.0175 gram of I_{2} from KI? + +!Answer!: 0.2493 gram. + +69. A sample of pyrolusite weighs 0.2400 gram and is 92.50% pure +MnO_{2}. The iodine liberated from KI by the manganese dioxide is +sufficient to react with 46.24 cc. of Na_{2}S_{2}O_{3} sol. What is +the normal value of the thiosulphate? + +!Answer!:: 0.1105 N. + +70. In the volumetric analysis of silver coin (90% Ag), using a +0.5000 gram sample, what is the least normal value that a potassium +thiocyanate solution may have and not require more than 50 cc. of +solution in the analysis? + +!Answer!: 0.08339 N. + +71. A mixture of pure lithium chloride and barium bromide weighing +0.6 gram is treated with 45.15 cubic centimeters of 0.2017 N silver +nitrate, and the excess titrated with 25 cc. of 0.1 N KSCN solution, +using ferric alum as an indicator. Calculate the percentage of bromine +in the sample. + +!Answer!: 40.11%. + +72. A mixture of the chlorides of sodium and potassium from 0.5000 +gram of a feldspar weighs 0.1500 gram, and after solution in water +requires 22.71 cc. of 0.1012 N silver nitrate for the precipitation of +the chloride ions. What are the percentages of Na_{2}O and K_{2}O in +the feldspar? + +!Answer!: 8.24% Na_{2}O; 9.14% K_{2}O. + + +GRAVIMETRIC ANALYSIS + +73. Calculate (a) the grams of silver in one gram of silver chloride; +(b) the grams of carbon dioxide liberated by the addition of an excess +of acid to one gram of calcium carbonate; (c) the grams of MgCl_{2} +necessary to precipitate 1 gram of MgNH_{4}PO_{4}. + +!Answers!: (a) 0.7526; (b) 0.4397; (c) 0.6940. + +74. Calculate the chemical factor for (a) Sn in SnO_{2}; (b) MgO +in Mg_{2}P_{2}O_{7}; (c) P_{2}O_{5} in Mg_{2}P_{2}O_{7}; (d) Fe in +Fe_{2}O_{3}; (e) SO_{4} in BaSO_{4}. + +!Answers!: (a) 0.7879; (b) 0.3620; (c) 0.6378; (d) 0.6990; (e) 0.4115. + +75. Calculate the log factor for (a) Pb in PbCrO_{4}; (b) Cr_{2}O_{3} +in PbCrO_{4}; (c) Pb in PbO_{2} and (d) CaO in CaC_{2}O_{4}. + +!Answers!: (a) 9.8069-10, (b) 9.3713-10; (c) 9.9376-10; (d) 9.6415-10. + +76. How many grams of Mn_{3}O_{4} can be obtained from 1 gram of +MnO_{2}? + +!Answer!: 0.8774 gram. + +77. If a sample of silver coin weighing 0.2500 gram gives a +precipitate of AgCl weighing 0.2991 gram, what weight of AgI could +have been obtained from the same weight of sample, and what is the +percentage of silver in the coin? + +!Answers!: 0.4898 gr.; 90.05%. + +78. How many cubic centimeters of hydrochloric acid (sp. gr. 1.13 +containing 25.75% HCl by weight) are required to exactly neutralize +25 cc. of ammonium hydroxide (sp. gr. .90 containing 28.33% NH_{3} by +weight)? + +!Answer!: 47.03 cc. + +79. How many cubic centimeters of ammonium hydroxide solution (sp. gr. +0.96 containing 9.91% NH_{3} by weight) are required to precipitate +the aluminium as aluminium hydroxide from a two-gram sample of alum +(KAl(SO_{4})_{2}.12H_{2}O)? What will be the weight of the ignited +precipitate? + +!Answers!: 2.26 cc.; 0.2154 gram. + +80. What volume of nitric acid (sp. gr. 1.05 containing 9.0% +HNO_{3} by weight) is required to oxidize the iron in one gram of +FeSO_{4}.7H_{2}O in the presence of sulphuric acid? 6FeSO_{4} + +2HNO_{3} + 3H_{2}SO_{4} = 3Fe_{2}(SO_{4})_{3} + 2NO + 4H_{2}O. + +!Answer!: 0.80 cc. + +81. If 0.7530 gram of ferric nitrate (Fe(NO_{3})_{3}.9H_{2}O) is +dissolved in water and 1.37 cc. of HCl (sp. gr. 1.11 containing 21.92% +HCl by weight) is added, how many cubic centimeters of ammonia (sp. +gr. 0.96 containing 9.91% NH_{3} by weight) are required to neutralize +the acid and precipitate the iron as ferric hydroxide? + +!Answer!: 2.63 cc. + +82. To a suspension of 0.3100 gram of Al(OH)_{3} in water are added +13.00 cc. of aqueous ammonia (sp. gr. 0.90 containing 28.4% NH_{3} by +weight). How many cubic centimeters of sulphuric acid (sp. gr. 1.18 +containing 24.7% H_{2}SO_{4} by weight) must be added to the mixture +in order to bring the aluminium into solution? + +!Answer!: 34.8 cc. + +83. How many cubic centimeters of sulphurous acid (sp. gr. 1.04 +containing 75 grams SO_{2} per liter) are required to reduce the +iron in 1 gram of ferric alum (KFe(SO_{4})_{2}.12H_{2}O)? +Fe_{2}(SO_{4})_{3} + SO_{2} + 2H_{2}O = 2FeSO_{4} + 2H_{2}SO_{4}. + +!Answer!: 0.85 cc. + +84. How many cubic centimeters of a solution of potassium bichromate +containing 26.30 grams of K_{2}Cr_{2}O_{7} per liter must be taken +in order to yield 0.6033 gram of Cr_{2}O_{3} after reduction and +precipitation of the chromium? + +K_{2}Cr_{2}O_{7} + 3SO_{2} + H_{2}SO_{4} = K_{2}SO_{4} + +Cr_{2}(SO_{4})_{3} + H_{2}O. + +!Answer!: 44.39 cc. + +85. How many cubic centimeters of ammonium hydroxide (sp. gr. 0.946 +containing 13.88% NH_{3} by weight) are required to precipitate +the iron as Fe(OH)_{3} from a sample of pure +FeSO_{4}.(NH_{4})_{2}SO_{4}.6H_{2}O, which requires 0.34 cc. of nitric +acid (sp. gr. 1.350 containing 55.79% HNO_{3} by weight) for oxidation +of the iron? (See problem No. 80 for reaction.) + +!Answer!: 4.74 cc. + +86. In the analysis of an iron ore by solution, oxidation and +precipitation of the iron as Fe(OH)_{3}, what weight of sample must be +taken for analysis so that each one hundredth of a gram of the ignited +precipitate of Fe_{2}O_{3} shall represent one tenth of one per cent +of iron? + +!Answer!: 6.99 grams. + +87. What weight in grams of impure ferrous ammonium sulphate should +be taken for analysis so that the number of centigrams of BaSO_{4} +obtained will represent five times the percentage of sulphur in the +sample? + +!Answer!: 0.6870 gram. + +88. What weight of magnetite must be taken for analysis in order that, +after precipitating and igniting all the iron to Fe_{2}O_{3}, the +percentage of Fe_{2}O_{4} in the sample may be found by multiplying +the weight in grams of the ignited precipitate by 100? + +!Answer!: 0.9665 gram. + +89. After oxidizing the arsenic in 0.5000 gram of pure As_{2}S_{3} to +arsenic acid, it is precipitated with "magnesia mixture" (MgCl_{2} + +2NH_{4}Cl). If exactly 12.6 cc. of the mixture are required, how many +grams of MgCl_{2} per liter does the solution contain? H_{3}AsO_{4} + +MgCl_{2} + 3NH_{4}OH = MgNH_{4}AsO_{4} + 2NH_{4}Cl + 3H_{2}O. + +!Answer!: 30.71 grams. + +90. A sample is prepared for student analysis by mixing pure apatite +(Ca_{3}(PO_{4})_{2}.CaCl_{2}) with an inert material. If 1 gram of +the sample gives 0.4013 gram of Mg_{2}P_{2}O_{7}, how many cubic +centimeters of ammonium oxalate solution (containing 40 grams of +(NH_{4})_{2}C_{2}O_{4}.H_{2}O per liter) would be required to +precipitate the calcium from the same weight of sample? + +!Answer!: 25.60 cc. + +91. If 0.6742 gram of a mixture of pure magnesium carbonate and pure +calcium carbonate, when treated with an excess of hydrochloric acid, +yields 0.3117 gram of carbon dioxide, calculate the percentage of +magnesium oxide and of calcium oxide in the sample. + +!Answers!: 13.22% MgO; 40.54% CaO. 92. The calcium in a sample of +dolomite weighing 0.9380 gram is precipitated as calcium oxalate and +ignited to calcium oxide. What volume of gas, measured over water +at 20 deg.C. and 765 mm. pressure, is given off during ignition, if the +resulting oxide weighs 0.2606 gram? (G.M.V. = 22.4 liters; V.P. water +at 20 deg.C. = 17.4 mm.) + +!Answer!: 227 cc. + +93. A limestone is found to contain 93.05% CaCO_{3}, and 5.16 % +MgCO_{3}. Calculate the weight of CaO obtainable from 3 tons of the +limestone, assuming complete conversion to oxide. What weight of +Mg_{2}P_{2}O_{7} could be obtained from a 3-gram sample of the +limestone? + +!Answers!: 1.565 tons; 0.2044 gram. + +94. A sample of dolomite is analyzed for calcium by precipitating +as the oxalate and igniting the precipitate. The ignited product is +assumed to be CaO and the analyst reports 29.50% Ca in the sample. +Owing to insufficient ignition, the product actually contained 8% of +its weight of CaCO_{3}. What is the correct percentage of calcium in +the sample, and what is the percentage error? + +!Answers!: 28.46%; 3.65% error. + +95. What weight of impure calcite (CaCO_{3}) should be taken for +analysis so that the volume in cubic centimeters of CO_{2} obtained by +treating with acid, measured dry at 18 deg.C. and 763 mm., shall equal the +percentage of CaO in the sample? + +!Answer!: 0.2359 gram. + +96. How many cubic centimeters of HNO_{3} (sp. gr. 1.13 containing +21.0% HNO_{3} by weight) are required to dissolve 5 grams of brass, +containing 0.61% Pb, 24.39% Zn, and 75% Cu, assuming reduction of the +nitric acid to NO by each constituent? What fraction of this volume of +acid is used for oxidation? + +!Answers!: 55.06 cc.; 25%. + +97. What weight of metallic copper will be deposited from a cupric +salt solution by a current of 1.5 amperes during a period of 45 +minutes, assuming 100% current efficiency? (1 Faraday = 96,500 +coulombs.) + +!Answer!: 1.335 grams. + +98. In the electrolysis of a 0.8000 gram sample of brass, there is +obtained 0.0030 gram of PbO_{2}, and a deposit of metallic copper +exactly equal in weight to the ignited precipitate of Zn_{2}P_{2}O_{7} +subsequently obtained from the solution. What is the percentage +composition of the brass? + +!Answers!: 69.75% Cu; 29.92% Zn; 0.33% Pb. + +99. A sample of brass (68.90% Cu; 1.10% Pb and 30.00% Zn) weighing +0.9400 gram is dissolved in nitric acid. The lead is determined by +weighing as PbSO_{4}, the copper by electrolysis and the zinc by +precipitation with (NH_{4})_{2}HPO_{4} in a neutral solution. + +(a) Calculate the cubic centimeters of nitric acid (sp. gr. 1.42 +containing 69.90% HNO_{3} by weight) required to just dissolve the +brass, assuming reduction to NO. + +!Answer!: 2.48 cc. + +(b) Calculate the cubic centimeters of sulphuric acid (sp. gr. 1.84 +containing 94% H_{2}SO_{4} by weight) to displace the nitric acid. + +!Answer!: 0.83 cc. + +(c) Calculate the weight of PbSO_{4}. + +!Answer!: 0.0152 gram. + +(d) The clean electrode weighs 10.9640 grams. Calculate the weight +after the copper has been deposited. + +!Answer!: 11.6116 grams. + +(e) Calculate the grams of (NH_{4})_{2}HPO_{4} required to precipitate +the zinc as ZnNH_{4}PO_{4}. + +!Answer!: 0.5705 gram. + +(f) Calculate the weight of ignited Zn_{2}P_{2}O_{7}. + +!Answer!: 0.6573 gram. + +100. If in the analysis of a brass containing 28.00% zinc an error is +made in weighing a 2.5 gram portion by which 0.001 gram too much is +weighed out, what percentage error in the zinc determination would +result? What volume of a solution of sodium hydrogen phosphate, +containing 90 grams of Na_{2}HPO_{4}.12H_{2}O per liter, would be +required to precipitate the zinc as ZnNH_{4}PO_{4} and what weight of +precipitate would be obtained? + +!Answers!: (a) 0.04% error; (b) 39.97 cc.; (c) 1.909 grams. + +101. A sample of magnesium carbonate, contaminated with SiO_{2} as its +only impurity, weighs 0.5000 gram and loses 0.1000 gram on ignition. +What volume of disodium phosphate solution (containing 90 grams +Na_{2}HPO_{4}.12H_{2}O per liter) will be required to precipitate the +magnesium as magnesium ammonium phosphate? + +!Answer!: 9.07 cc. + +102. 2.62 cubic centimeters of nitric acid (sp. gr. 1.42 containing +69.80% HNO_{2} by weight) are required to just dissolve a sample +of brass containing 69.27% Cu; 0.05% Pb; 0.07% Fe; and 30.61% Zn. +Assuming the acid used as oxidizing agent was reduced to NO in every +case, calculate the weight of the brass and the cubic centimeters of +acid used as acid. + +!Answer!: 0.992 gram; 1.97 cc. + +103. One gram of a mixture of silver chloride and silver bromide is +found to contain 0.6635 gram of silver. What is the percentage of +bromine? + +!Answer!: 21.30%. + +104. A precipitate of silver chloride and silver bromide weighs 0.8132 +gram. On heating in a current of chlorine, the silver bromide is +converted to silver chloride, and the mixture loses 0.1450 gram +in weight. Calculate the percentage of chlorine in the original +precipitate. + +!Answer!: 6.13%. + +105. A sample of feldspar weighing 1.000 gram is fused and the silica +determined. The weight of silica is 0.6460 gram. This is fused with 4 +grams of sodium carbonate. How many grams of the carbonate actually +combined with the silica in fusion, and what was the loss in weight +due to carbon dioxide during the fusion? + +!Answers!: 1.135 grams; 0.4715 gram. + +106. A mixture of barium oxide and calcium oxide weighing 2.2120 grams +is transformed into mixed sulphates, weighing 5.023 grams. Calculate +the grams of calcium oxide and barium oxide in the mixture. + +!Answers!: 1.824 grams CaO; 0.3877 gram BaO. + + + + +APPENDIX + + +ELECTROLYTIC DISSOCIATION THEORY + +The following brief statements concerning the ionic theory and a few +of its applications are intended for reference in connection with the +explanations which are given in the Notes accompanying the various +procedures. The reader who desires a more extended discussion of the +fundamental theory and its uses is referred to such books as Talbot +and Blanchard's !Electrolytic Dissociation Theory! (Macmillan +Company), or Alexander Smith's !Introduction to General Inorganic +Chemistry! (Century Company). + +The !electrolytic dissociation theory!, as propounded by Arrhenius in +1887, assumes that acids, bases, and salts (that is, electrolytes) +in aqueous solution are dissociated to a greater or less extent into +!ions!. These ions are assumed to be electrically charged atoms or +groups of atoms, as, for example, H^{+} and Br^{-} from hydrobromic +acid, Na^{+} and OH^{-} from sodium hydroxide, 2NH_{4}^{+} and +SO_{4}^{--} from ammonium sulphate. The unit charge is that which is +dissociated with a hydrogen ion. Those upon other ions vary in sign +and number according to the chemical character and valence of the +atoms or radicals of which the ions are composed. In any solution the +aggregate of the positive charges upon the positive ions (!cations!) +must always balance the aggregate negative charges upon the negative +ions (!anions!). + +It is assumed that the Na^{+} ion, for example, differs from the +sodium atom in behavior because of the very considerable electrical +charge which it carries and which, as just stated, must, in an +electrically neutral solution, be balanced by a corresponding negative +charge on some other ion. When an electric current is passed through a +solution of an electrolyte the ions move with and convey the current, +and when the cations come into contact with the negatively charged +cathode they lose their charges, and the resulting electrically +neutral atoms (or radicals) are liberated as such, or else enter at +once into chemical reaction with the components of the solution. + +Two ions of identically the same composition but with different +electrical charges may exhibit widely different properties. For +example, the ion MnO_{4}^{-} from permanganates yields a purple-red +solution and differs in its chemical behavior from the ion +MnO_{4}^{--} from manganates, the solutions of which are green. + +The chemical changes upon which the procedures of analytical chemistry +depend are almost exclusively those in which the reacting substances +are electrolytes, and analytical chemistry is, therefore, essentially +the chemistry of the ions. The percentage dissociation of the same +electrolyte tends to increase with increasing dilution of its +solution, although not in direct proportion. The percentage +dissociation of different electrolytes in solutions of equivalent +concentrations (such, for example, as normal solutions) varies widely, +as is indicated in the following tables, in which approximate figures +are given for tenth-normal solutions at a temperature of about 18 deg.C. + + ACIDS +========================================================================= + | + SUBSTANCE | PERCENTAGE DISSOCIATION IN + | 0.1 EQUIVALENT SOLUTION +_____________________________________________|___________________________ + | +HCl, HBr, HI, HNO_{3} | 90 + | +HClO_{3}, HClO_{4}, HMnO_{4} | 90 + | +H_{2}SO_{4} <--> H^{+} + HSO_{4}^{-} | 90 + | +H_{2}C_{2}O_{4} <--> H^{+} + HC_{2}O_{4}^{-} | 50 + | +H_{2}SO_{3} <--> H^{+} + HSO{_}3^{-} | 20 + | +H_{3}PO_{4} <--> H^{+} + H_{2}PO_{4}^{-} | 27 + | +H_{2}PO_{4}^{-} <--> H^{+} + HPO_{4}^{--} | 0.2 + | +H_{3}AsO_{4} <--> H^{+} + H_{2}AsO_{4}^{-} | 20 + | +HF | 9 + | +HC_{2}H_{3}O_{2} | 1.4 + | +H_{2}CO_{3} <--> H^{+} + HCO_{3}^{-} | 0.12 + | +H_{2}S <--> H^{+} + HS^{-} | 0.05 + | +HCN | 0.01 + | +========================================================================= + + + BASES +========================================================================= + | + SUBSTANCE | PERCENTAGE DISSOCIATION IN + | 0.1 EQUIVALENT SOLUTION +_____________________________________________|___________________________ + | +KOH, NaOH | 86 + | +Ba(OH)_{2} | 75 + | +NH_{4}OH | 1.4 + | +========================================================================= + + + SALTS +========================================================================= + | + TYPE OF SALT | PERCENTAGE DISSOCIATION IN + | 0.1 EQUIVALENT SOLUTION +_____________________________________________|___________________________ + | +R^{+}R^{-} | 86 + | +R^{++}(R^{-})_{2} | 72 + | +(R^{+})_{2}R^{--} | 72 + | +R^{++}R^{--} | 45 + | +========================================================================= + +The percentage dissociation is determined by studying the electrical +conductivity of the solutions and by other physico-chemical methods, +and the following general statements summarize the results: + +!Salts!, as a class, are largely dissociated in aqueous solution. + +!Acids! yield H^{+} ions in water solution, and the comparative +!strength!, that is, the activity, of acids is proportional to the +concentration of the H^{+} ions and is measured by the percentage +dissociation in solutions of equivalent concentration. The common +mineral acids are largely dissociated and therefore give a relatively +high concentration of H^{+} ions, and are commonly known as "strong +acids." The organic acids, on the other hand, belong generally to the +group of "weak acids." + +!Bases! yield OH^{-} ions in water solution, and the comparative +strength of the bases is measured by their relative dissociation in +solutions of equivalent concentration. Ammonium hydroxide is a weak +base, as shown in the table above, while the hydroxides of sodium and +potassium exhibit strongly basic properties. + +Ionic reactions are all, to a greater or less degree, !reversible +reactions!. A typical example of an easily reversible reaction is that +representing the changes in ionization which an electrolyte such as +acetic acid undergoes on dilution or concentration of its solutions, +!i.e.!, HC_{2}H_{3}O_{2} <--> H^{+} + C_{2}H_{3}O_{2}^{-}. As was +stated above, the ionization increases with dilution, the reaction +then proceeding from left to right, while concentration of the +solution occasions a partial reassociation of the ions, and the +reaction proceeds from right to left. To understand the principle +underlying these changes it is necessary to consider first the +conditions which prevail when a solution of acetic acid, which has +been stirred until it is of uniform concentration throughout, has come +to a constant temperature. A careful study of such solutions has shown +that there is a definite state of equilibrium between the constituents +of the solution; that is, there is a definite relation between the +undissociated acetic acid and its ions, which is characteristic for +the prevailing conditions. It is not, however, assumed that this is a +condition of static equilibrium, but rather that there is continual +dissociation and association, as represented by the opposing +reactions, the apparent condition of rest resulting from the fact that +the amount of change in one direction during a given time is exactly +equal to that in the opposite direction. A quantitative study of +the amount of undissociated acid, and of H^{+} ions and +C_{2}H_{3}O_{2}^{-} ions actually to be found in a large number of +solutions of acetic acid of varying dilution (assuming them to be in +a condition of equilibrium at a common temperature), has shown that +there is always a definite relation between these three quantities +which may be expressed thus: + +(!Conc'n H^{+} x Conc'n C_{2}H_{3}O_{2}^{-})/Conc'n HC_{2}H_{3}O_{2} = +Constant!. + +In other words, there is always a definite and constant ratio between +the product of the concentrations of the ions and the concentration of +the undissociated acid when conditions of equilibrium prevail. + +It has been found, further, that a similar statement may be made +regarding all reversible reactions, which may be expressed in general +terms thus: The rate of chemical change is proportional to the product +of the concentrations of the substances taking part in the reaction; +or, if conditions of equilibrium are considered in which, as stated, +the rate of change in opposite directions is assumed to be equal, then +the product of the concentrations of the substances entering into +the reaction stands in a constant ratio to the product of the +concentrations of the resulting substances, as given in the expression +above for the solutions of acetic acid. This principle is called the +!Law of Mass Action!. + +It should be borne in mind that the expression above for acetic acid +applies to a wide range of dilutions, provided the temperature remains +constant. If the temperature changes the value of the constant changes +somewhat, but is again uniform for different dilutions at that +temperature. The following data are given for temperatures of about +18 deg.C.[1] + +========================================================================== + | | | | + MOLAL | FRACTION | MOLAL CONCENTRA- | MOLAL CONCENTRA- | VALUE OF +CONCENTRATION | IONIZED | TION OF H^{+} AND| TION OF UNDIS- | CONSTANT + CONSTANT | | ACETATE^{-} IONS | SOCIATED ACID | +______________|__________|__________________|__________________|__________ + | | | | + 1.0 | .004 | .004 | .996 | .0000161 + | | | | + 0.1 | .013 | .0013 | .0987 | .0000171 + | | | | + 0.01 | .0407 | .000407 | .009593 | .0000172 + | | | | +=========================================================================== + +[Footnote 1: Alexander Smith, !General Inorganic Chemistry!, p. 579.] + +The molal concentrations given in the table refer to fractions of a +gram-molecule per liter of the undissociated acid, and to fractions of +the corresponding quantities of H^{+} and C_{2}H_{3}O_{2}^{-} ions +per liter which would result from the complete dissociation of a +gram-molecule of acetic acid. The values calculated for the constant +are subject to some variation on account of experimental errors in +determining the percentage ionized in each case, but the approximate +agreement between the values found for molal and centimolal (one +hundredfold dilution) is significant. + +The figures given also illustrate the general principle, that the +!relative! ionization of an electrolyte increases with the dilution of +its solution. If we consider what happens during the (usually) brief +period of dilution of the solution from molal to 0.1 molal, for +example, it will be seen that on the addition of water the conditions +of concentration which led to equality in the rate of change, and +hence to equilibrium in the molal solution, cease to exist; and since +the dissociating tendency increases with dilution, as just stated, +it is true at the first instant after the addition of water that the +concentration of the undissociated acid is too great to be +permanent under the new conditions of dilution, and the reaction, +HC_{2}H_{3}O_{2} <--> H^{+} + C_{2}H_{3}O_{2}^{-}, will proceed from +left to right with great rapidity until the respective concentrations +adjust themselves to the new conditions. + +That which is true of this reaction is also true of all reversible +reactions, namely, that any change of conditions which occasions +an increase or a decrease in concentration of one or more of the +components causes the reaction to proceed in one direction or the +other until a new state of equilibrium is established. This principle +is constantly applied throughout the discussion of the applications +of the ionic theory in analytical chemistry, and it should be clearly +understood that whenever an existing state of equilibrium is disturbed +as a result of changes of dilution or temperature, or as a consequence +of chemical changes which bring into action any of the constituents of +the solution, thus altering their concentrations, there is always a +tendency to re-establish this equilibrium in accordance with the law. +Thus, if a base is added to the solution of acetic acid the H^{+} ions +then unite with the OH^{-} ions from the base to form undissociated +water. The concentration of the H^{+} ions is thus diminished, and +more of the acid dissociates in an attempt to restore equilbrium, +until finally practically all the acid is dissociated and neutralized. + +Similar conditions prevail when, for example, silver ions react with +chloride ions, or barium ions react with sulphate ions. In the former +case the dissociation reaction of the silver nitrate is AgNO_{3} <--> +Ag^{+} + NO_{3}^{-}, and as soon as the Ag^{+} ions unite with the +Cl^{-} ions the concentration of the former is diminished, more of the +AgNO_{3} dissociates, and this process goes on until the Ag^{+} ions +are practically all removed from the solution, if the Cl^{-} ions are +present in sufficient quantity. + +For the sake of accuracy it should be stated that the mass law cannot +be rigidly applied to solutions of those electrolytes which are +largely dissociated. While the explanation of the deviation from +quantitative exactness in these cases is not known, the law is still +of marked service in developing analytical methods along more logical +lines than was formerly practicable. It has not seemed wise to qualify +each statement made in the Notes to indicate this lack of quantitative +exactness. The student should recognize its existence, however, and +will realize its significance better as his knowledge of physical +chemistry increases. + +If we apply the mass law to the case of a substance of small +solubility, such as the compounds usually precipitated in quantitative +analysis, we derive what is known as the !solubility product!, as +follows: Taking silver chloride as an example, and remembering that it +is not absolutely insoluble in water, the equilibrium expression for +its solution is: + +(!Conc'n Ag^{+} x Conc'n Cl^{-})/Conc'n AgCl = Constant!. + +But such a solution of silver chloride which is in contact with the +solid precipitate must be saturated for the existing temperature, and +the quantity of undissociated AgCl in the solution is definite and +constant for that temperature. Since it is a constant, it may be +eliminated, and the expression becomes !Conc'n Ag^{+} x Conc'n +Cl^{-} = Constant!, and this is known as the solubility product. No +precipitation of a specific substance will occur until the product of +the concentrations of its ions in a solution exceeds the solubility +product for that substance; whenever that product is exceeded +precipitation must follow. + +It will readily be seen that if a substance which yields an ion in +common with the precipitated compound is added to such a solution as +has just been described, the concentration of that ion is +increased, and as a result the concentration of the other ion must +proportionately decrease, which can only occur through the formation +of some of the undissociated compound which must separate from the +already saturated solution. This explains why the addition of an +excess of the precipitant is often advantageous in quantitative +procedures. Such a case is discussed at length in Note 2 on page 113. + +Similarly, the ionization of a specific substance in solution tends to +diminish on the addition of another substance with a common ion, as, +for instance, the addition of hydrochloric acid to a solution +of hydrogen sulphide. Hydrogen sulphide is a weak acid, and the +concentration of the hydrogen ions in its aqueous solutions is very +small. The equilibrium in such a solution may be represented as: + +(!(Conc'n H^{+})^{2} x Conc'n S^{--})/Conc'n H_{2}S = Constant!, and a +marked increase in the concentration of the H^{+} ions, such as would +result from the addition of even a small amount of the highly ionized +hydrochloric acid, displaces the point of equilibrium and some of the +S^{--} ions unite with H^{+} ions to form undissociated H_{2}S. This +is of much importance in studying the reactions in which hydrogen +sulphide is employed, as in qualitative analysis. By a parallel course +of reasoning it will be seen that the addition of a salt of a weak +acid or base to solutions of that acid or base make it, in effect, +still weaker because they decrease its percentage ionization. + +To understand the changes which occur when solids are dissolved where +chemical action is involved, it should be remembered that no substance +is completely insoluble in water, and that those products of a +chemical change which are least dissociated will first form. Consider, +for example, the action of hydrochloric acid upon magnesium hydroxide. +The minute quantity of dissolved hydroxide dissociates thus: +Mg(OH)_{2} <--> Mg^{++} + 2OH^{-}. When the acid is introduced, +the H^{+} ions of the acid unite with the OH^{-} ions to form +undissociated water. The concentration of the OH^{-} ions is thus +diminished, more Mg(OH)_{2} dissociates, the solution is no longer +saturated with the undissociated compound, and more of the solid +dissolves. This process repeats itself with great rapidity until, if +sufficient acid is present, the solid passes completely into solution. + +Exactly the same sort of process takes place if calcium oxalate, for +example, is dissolved in hydrochloric acid. The C_{2}O_{4}^{--} ions +unite with the H^{+} ions to form undissociated oxalic acid, the acid +being less dissociated than normally in the presence of the H^{+} ions +from the hydrochloric acid (see statements regarding hydrogen sulphide +above). As the undissociated oxalic acid forms, the concentration of +the C_{2}O_{4}^{--} ions lessens and more CaC_{2}O_{4} dissolves, +as described for the Mg(OH)_{2} above. Numerous instances of the +applications of these principles are given in the Notes. + +Water itself is slightly dissociated, and although the resulting H^{+} +and OH^{-} ions are present only in minute concentrations (1 mol. of +dissociated water in 10^{7} liters), yet under some conditions they +may give rise to important consequences. The term !hydrolysis! is +applied to the changes which result from the reaction of these ions. +Any salt which is derived from a weak base or a weak acid (or both) +is subject to hydrolytic action. Potassium cyanide, for example, when +dissolved in water gives an alkaline solution because some of the +H^{+} ions from the water unite with CN^{-} ions to form (HCN), which +is a very weak acid, and is but very slightly dissociated. Potassium +hydroxide, which might form from the OH^{-} ions, is so largely +dissociated that the OH^{-} ions remain as such in the solution. The +union of the H^{+} ions with the CN^{-} ions to form the undissociated +HCN diminishes the concentration of the H^{+} ions, and more water +dissociates (H_{2}O <--> H^{+} + OH^{-}) to restore the equilibrium. +It is clear, however, that there must be a gradual accumulation of +OH^{-} ions in the solution as a result of these changes, causing the +solution to exhibit an alkaline reaction, and also that ultimately the +further dissociation of the water will be checked by the presence of +these ions, just as the dissociation of the H_{2}S was lessened by the +addition of HCl. + +An exactly opposite result follows the solution of such a salt as +Al_{2}(SO_{4})_{3} in water. In this case the acid is strong and the +base weak, and the OH^{-} ions form the little dissociated Al(OH)_{3}, +while the H^{+} ions remain as such in the solution, sulphuric acid +being extensively dissociated. The solution exhibits an acid reaction. + +Such hydrolytic processes as the above are of great importance in +analytical chemistry, especially in the understanding of the action of +indicators in volumetric analysis. (See page 32.) + +The impelling force which causes an element to pass from the atomic +to the ionic condition is termed !electrolytic solution pressure!, or +ionization tension. This force may be measured in terms of electrical +potential, and the table below shows the relative values for a number +of elements. + +In general, an element with a greater solution pressure tends to cause +the deposition of an element of less solution pressure when placed in +a solution of its salt, as, for instance, when a strip of zinc or +iron is placed in a solution of a copper salt, with the resulting +precipitation of metallic copper. + +Hydrogen is included in the table, and its position should be noted +with reference to the other common elements. For a more extended +discussion of this topic the student should refer to other treatises. + + POTENTIAL SERIES OF THE METALS + +__________________________________________________________________ + | | | + | POTENTIAL | | POTENTIAL + | IN VOLTS | | IN VOLTS +_____________________|___________|____________________|___________ + | | | +Sodium Na^{+} | +2.44 | Lead Pb^{++} | -0.13 +Calcium Ca^{++} | | Hydrogen H^{+} | -0.28 +Magnesium Mg^{++} | | Bismuth Bi^{+++}| +Aluminum A1^{+++} | +1.00 | Antimony | -0.75 +Manganese Mn^{++} | | Arsenic | +Zinc Zn^{++} | +0.49 | Copper Cu^{++} | -0.61 +Cadmium Cd^{++} | +0.14 | Mercury Hg^{+} | -1.03 +Iron Fe^{++} | +0.063 | Silver Ag^{+} | -1.05 +Cobalt Co^{++} | -0.045 | Platinum | +Nickel Ni^{++} | -0.049 | Gold | +Tin Sn^{++} | -0.085(?) | | +_____________________|___________|____________________|__________ + + + +THE FOLDING OF A FILTER PAPER + +If a filter paper is folded along its diameter, and again folded along +the radius at right angles to the original fold, a cone is formed on +opening, the angle of which is 60 deg.. Funnels for analytical use are +supposed to have the same angle, but are rarely accurate. It is +possible, however, with care, to fit a filter thus folded into a +funnel in such a way as to prevent air from passing down between the +paper and the funnel to break the column of liquid in the stem, +which aids greatly, by its gentle suction, in promoting the rate of +filtration. + +Such a filter has, however, the disadvantage that there are three +thicknesses of paper back of half of its filtering surface, as a +consequence of which one half of a precipitate washes or drains more +slowly. Much time may be saved in the aggregate by learning to fold a +filter in such a way as to improve its effective filtering surface. +The directions which follow, though apparently complicated on first +reading, are easily applied and easily remembered. Use a 6-inch filter +for practice. Place four dots on the filter, two each on diameters +which are at right angles to each other. Then proceed as follows: +(1) Fold the filter evenly across one of the diameters, creasing it +carefully; (2) open the paper, turn it over, rotate it 90 deg. to the +right, bring the edges together and crease along the other diameter; +(3) open, and rotate 45 deg. to the right, bring edges together, and +crease evenly; (4) open, and rotate 90 deg. to the right, and crease +evenly; (5) open, turn the filter over, rotate 22-(1/2) deg. to the right, +and crease evenly; (6) open, rotate 45 deg. to the right and crease +evenly; (7) open, rotate 45 deg. to the right and crease evenly; (8) open, +rotate 45 deg. to the right and crease evenly; (9) open the filter, and, +starting with one of the dots between thumb and forefinger of the +right hand, fold the second crease to the left over on it, and do +the same with each of the other dots. Place it, thus folded, in the +funnel, moisten it, and fit to the side of the funnel. The filter will +then have four short segments where there are three thicknesses +and four where there is one thickness, but the latter are evenly +distributed around its circumference, thus greatly aiding the passage +of liquids through the paper and hastening both filtration and washing +of the whole contents of the filter. + + +!SAMPLE PAGES FOR LABORATORY RECORDS! + +!Page A! + +Date + +CALIBRATION OF BURETTE No. + +___________________________________________________________________________ + | | | | + BURETTE | DIFFERENCE | OBSERVED | DIFFERENCE | CALCULATED + READINGS | | WEIGHTS | | CORRECTION +_______________|______________|______________|______________|______________ + 0.02 | | 16.27 | | + 10.12 | 10.10 | 26.35 | 10.08 | -.02 + 20.09 | 9.97 | 36.26 | 9.91 | -.06 + 30.16 | 10.07 | 46.34 | 10.08 | +.01 + 40.19 | 10.03 | 56.31 | 9.97 | -.06 + 50.00 | 9.81 | 66.17 | 9.86 | +.05 +_______________|______________|______________|______________|______________ + + These data to be obtained in duplicate for each burette. + + +!Page B! + +Date + + +DETERMINATION OF COMPARATIVE STRENGTH HCl vs. NaOH + +___________________________________________________________________________ + | | + DETERMINATION | I | II +_________________________|________________________|________________________ + | | + | Corrected | Corrected +Final Reading HCl | 48.17 48.08 | 43.20 43.14 +Initial Reading HCl | 0.12 .12 | .17 .17 + | ----- ----- | ----- ----- + | 47.96 | 42.97 + | | + | Corrected | Corrected +Final Reading HCl | 46.36 46.29 | 40.51 40.37 +Initial Reading HCl | 1.75 1.75 | .50 .50 + | ----- ----- | ----- ----- + | 44.54 | 39.87 + | | + log cc. NaOH | 1.6468 | 1.6008 + colog cc. HCl | 8.3192 | 8.3668 + | ------ | ------ + | 9.9680 - 10 | 9.9676 - 10 + 1 cc. HCl | .9290 cc. NaOH | .9282 cc. NaOH + Mean | .9286 | +_________________________|________________________|________________________ + + +Signed + +!Page C! +Date + + +STANDARDIZATION OF HYDROCHLORIC ACID +===================================================================== + | | +Weight sample and tube| 9.1793 | 8.1731 + | 8.1731 | 6.9187 + | ------ | ------ + Weight sample | 1.0062 | 1.2544 + | | +Final Reading HCl | 39.97 39.83 | 49.90 49.77 +Initial Reading HCl | .00 .00 | .04 .04 + | ----- ----- | ----- ----- + | 39.83 | 49.73 + | | +Final Reading NaOH | .26 .26 | .67 .67 +Initial Reading NaOH | .12 .12 | .36 .36 + | --- --- | --- --- + | .14 | .31 + | | + | .14 | .31 +Corrected cc. HCl | 39.83 - ----- = 39.68 | 49.73 - ----- = 49.40 + | .9286 | .9286 + | | +log sample | 0.0025 | 0.0983 +colog cc | 8.4014 - 10 | 8.3063 - 10 +colog milli equivalent| 1.2757 | 1.2757 + | ------ | ------ + | 9.6796 - 10 | 9.6803 - 10 + | | +Normal value HCl | .4782 | .4789 + Mean | .4786 | + | | +===================================================================== + +Signed + + +!Page D! +Date + + +DETERMINATION OF CHLORINE IN CHLORIDE, SAMPLE No. +===================================================================== + | | +Weight sample and tube| 16.1721 | 15.9976 + | 15.9976 | 15.7117 + | ------- | ------- + Weight sample | .1745 | .2859 + | | +Weight crucible | | + + precipitate | 14.4496 | 15.6915 + Constant weights | 14.4487 | 15.6915 + | 14.4485 | + | | + Weight crucible | 14.2216 | 15.3196 + Constant weight | 14.2216 | 15.3194 + | | + Weight AgCl | .2269 | .3721 + | | + log Cl | 1.5496 | 1.5496 + log weight AgCl | 9.3558 - 10 | 9.5706 - 10 + log 100 | 2.0000 | 2.0000 + colog AgCl | 7.8438 - 10 | 7.7438 - 10 + colog sample | 0.7583 | 0.5438 + | ------- | ------- + | 1.5075 | 1.5078 + | | + Cl in sample No. | 32.18% | 32.20% + | | +===================================================================== + +Signed + + +STRENGTH OF REAGENTS + +The concentrations given in this table are those suggested for use +in the procedures described in the foregoing pages. It is obvious, +however, that an exact adherence to these quantities is not essential. + + + Approx. Approx. + Grams relation relation + per to normal to molal + liter. solution solution + +Ammonium oxalate, (NH_{4})_{2}C_{2}O_{4}.H_{2}O 40 0.5N 0.25 +Barium chloride, BaCl_{2}.2H_{2}O 25 0.2N 0.1 +Magnesium ammonium chloride (of MgCl_{2}) 71 1.5N 0.75 +Mercuric chloride, HgCl_{2} 45 0.33N 0.66 +Potassium hydroxide, KOH (sp. gr. 1.27) 480 +Potassium thiocyanate, KSCN 5 0.05N 0.55 +Silver nitrate, AgNO_{3} 21 0.125N 0.125 +Sodium hydroxide, NaOH 100 2.5N 2.5 +Sodium carbonate. Na_{2}CO_{3} 159 3N 1.5 +Sodium phosphate, Na_{2}HPO_{4}.12H_{2}O 90 0.5N or 0.75N 0.25 + +Stannous chloride, SnCl_{2}, made by saturating hydrochloric acid (sp. +gr. 1.2) with tin, diluting with an equal volume of water, and adding +a slight excess of acid from time to time. A strip of metallic tin is +kept in the bottle. + +A solution of ammonium molybdate is best prepared as follows: Stir +100 grams of molybdic acid (MoO_{3}) into 400 cc. of cold, distilled +water. Add 80 cc. of concentrated ammonium hydroxide (sp. gr. 0.90). +Filter, and pour the filtrate slowly, with constant stirring, into a +mixture of 400 cc. concentrated nitric acid (sp. gr. 1.42) and 600 +cc. of water. Add to the mixture about 0.05 gram of microcosmic salt. +Filter, after allowing the whole to stand for 24 hours. + +The following data regarding the common acids and aqueous ammonia +are based upon percentages given in the Standard Tables of the +Manufacturing Chemists' Association of the United States [!J.S.C.I.!, +24 (1905), 787-790]. All gravities are taken at 15.5 deg.C. and compared +with water at the same temperature. + +Aqueous ammonia (sp. gr. 0.96) contains 9.91 per cent NH_{3} by +weight, and corresponds to a 5.6 N and 5.6 molal solution. + +Aqueous ammonia (sp. gr. 0.90) contains 28.52 per cent NH_{3} by +weight, and corresponds to a 5.6 N and 5.6 molal solution. + +Hydrochloric acid (sp. gr. 1.12) contains 23.81 per cent HCl by +weight, and corresponds to a 7.3 N and 7.3 molal solution. + +Hydrochloric acid (sp. gr. 1.20) contains 39.80 per cent HCl by +weight, and corresponds to a 13.1 N and 13.1 molal solution. + +Nitric acid (sp. gr. 1.20) contains 32.25 per cent HNO_{3} by weight, +and corresponds to a 6.1 N and 6.1 molal solution: + +Nitric acid (sp. gr. 1.42) contains 69.96 per cent HNO_{3} by weight, +and corresponds to a 15.8 N and 15.8 molal solution. + +Sulphuric acid (sp. gr. 1.8354) contains 93.19 per cent H_{2}SO_{4} by +weight, and corresponds to a 34.8 N or 17.4 molal solution. + +Sulphuric acid (sp. gr. 1.18) contains 24.74 per cent H_{2}SO_{4} by +weight, and corresponds to a 5.9 N or 2.95 molal solution. + +The term !normal! (N), as used above, has the same significance as +in volumetric analyses. The molal solution is assumed to contain one +molecular weight in grams in a liter of solution. + +DENSITIES AND VOLUMES OF WATER AT TEMPERATURES FROM 15-30 deg.C. + +Temperature Density. Volume. +Centigrade. + + 4 deg. 1.000000 1.000000 + 15 deg. 0.999126 1.000874 + 16 deg. 0.998970 1.001031 + 17 deg. 0.998801 1.001200 + 18 deg. 0.998622 1.001380 + 19 deg. 0.998432 1.001571 + 20 deg. 0.998230 1.001773 + 21 deg. 0.998019 1.001985 + 22 deg. 0.997797 1.002208 + 23 deg. 0.997565 1.002441 + 24 deg. 0.997323 1.002685 + 25 deg. 0.997071 1.002938 + 26 deg. 0.996810 1.003201 + 27 deg. 0.996539 1.003473 + 28 deg. 0.996259 1.003755 + 29 deg. 0.995971 1.004046 + 30 deg. 0.995673 1.004346 + +Authority: Landolt, Boernstein, and Meyerhoffer's !Tabellen!, third +edition. + + +CORRECTIONS FOR CHANGE OF TEMPERATURE OF STANDARD SOLUTIONS + +The values below are average values computed from data relating to a +considerable number of solutions. They are sufficiently accurate for +use in chemical analyses, except in the comparatively few cases +where the highest attainable accuracy is demanded in chemical +investigations. The expansion coefficients should then be carefully +determined for the solutions employed. For a compilation of the +existing data, consult Landolt, Boernstein, and Meyerhoffer's +!Tabellen!, third edition. + + Corrections for 1 cc. + Concentration. of solution between + 15 deg. and 35 deg.C. + + Normal .00029 + 0.5 Normal .00025 + 0.1 Normal or more dilute solutions .00020 + +The volume of solution used should be multiplied by the values given, +and that product multiplied by the number of degrees which the +temperature of the solution varies from the standard temperature +selected for the laboratory. The total correction thus found is +subtracted from the observed burette reading if the temperature is +higher than the standard, or added, if it is lower. Corrections are +not usually necessary for variations of temperature of 2 deg.C. or less. + + + + INTERNATIONAL ATOMIC WEIGHTS + +========================================================== + | | | + | 1920 | | 1920 +_________________|_________|___________________|__________ + | | | +Aluminium Al | 27.1 | Molybdenum Mo | 96.0 +Antimony Sb | 120.2 | Neodymium Nd | 144.3 +Argon A | 39.9 | Neon Ne | 20.2 +Arsenic As | 74.96 | Nickel Ni | 58.68 +Barium Ba | 137.37 | Nitrogen N | 14.008 +Bismuth Bi | 208.0 | Osmium Os | 190.9 +Boron B | 11.0 | Oxygen O | 16.00 +Bromine Br | 79.92 | Palladium Pd | 106.7 +Cadmium Cd | 112.40 | Phosphorus P | 31.04 +Caesium Cs | 132.81 | Platinum Pt | 195.2 +Calcium Ca | 40.07 | Potassium K | 39.10 +Carbon C | 12.005 | Praseodymium Pr | 140.9 +Cerium Ce | 140.25 | Radium Ra | 226.0 +Chlorine Cl | 35.46 | Rhodium Rh | 102.9 +Chromium Cr | 52.0 | Rubidium Rb | 85.45 +Cobalt Co | 58.97 | Ruthenium Ru | 101.7 +Columbium Cb | 93.1 | Samarium Sm | 150.4 +Copper Cu | 63.57 | Scandium Sc | 44.1 +Dysprosium Dy | 162.5 | Selenium Se | 79.2 +Erbium Er | 167.7 | Silicon Si | 28.3 +Europium Eu | 152.0 | Silver Ag | 107.88 +Fluorine Fl | 19.0 | Sodium Na | 23.00 +Gadolinium Gd | 157.3 | Strontium Sr | 87.63 +Gallium Ga | 69.9 | Sulphur S | 32.06 +Germanium Ge | 72.5 | Tantalum Ta | 181.5 +Glucinum Gl | 9.1 | Tellurium Te | 127.5 +Gold Au | 197.2 | Terbium Tb | 159.2 +Helium He | 4.00 | Thallium Tl | 204.0 +Hydrogen H | 1.008 | Thorium Th | 232.4 +Indium In | 114.8 | Thulium Tm | 168.5 +Iodine I | 126.92 | Tin Sn | 118.7 +Iridium Ir | 193.1 | Titanium Ti | 48.1 +Iron Fe | 55.84 | Tungsten W | 184.0 +Krypton Kr | 82.92 | Uranium U | 238.2 +Lanthanum La | 139.0 | Vanadium V | 51.0 +Lead Pb | 207.2 | Xenon Xe | 130.2 +Lithium Li | 6.94 | Ytterbium Yb | 173.5 +Lutecium Lu | 175.0 | Yttrium Y | 88.7 +Magnesium Mg | 24.32 | Zinc Zn | 65.37 +Manganese Mn | 54.93 | Zirconium Zr | 90.6 +Mercury Hg | 200.6 | | +========================================================== + + + + +INDEX + +Acidimetry +Acid solutions, normal + standard +Acids, definition of +Acids, weak, action of other acids on + action of salts on +Accuracy demanded +Alkalimetry +Alkali solutions, normal + standard +Alumina, determination of in stibnite +Ammonium nitrate, acid +Analytical chemistry, subdivisions of +Antimony, determination of, in stibnite +Apatite, analysis of +Asbestos filters +Atomic weights, table of + +Balances, essential features of + use and care of +Barium sulphate, determination of sulphur in +Bases, definition of +Bichromate process for iron +Bleaching powder, analysis of +Brass, analysis of +Burette, description of + calibration of + cleaning of + reading of + +Calcium, determination of, in limestone +Calibration, definition of + of burettes + of flasks +Carbon dioxide, determination of, in limestone +Chlorimetry +Chlorine, gravimetric determination of +Chrome iron ore, analysis of +Coin, determination of silver in +Colloidal solution of precipitates +Colorimetric analyses, definition of +Copper, determination of, in brass + determination of in copper ores +Crucibles, use of +Crystalline precipitates + +Densities of water +Deposition potentials +Desiccators +Direct methods +Dissociation, degree of + +Economy of time +Electrolytic dissociation, theory of +Electrolytic separations, principles of +End-point, definition of +Equilibrium, chemical +Evaporation of liquids + +Faraday's law +Feldspar, analysis of +Ferrous ammonium sulphate, analysis of +Filters, folding of + how fitted +Filtrates, testing of +Filtration +Flasks, graduation of +Funnels +Fusions, removal of from crucibles + +General directions for gravimetric analysis + volumetric analysis +Gooch filter +Gravimetric analysis, definition of + +Hydrochloric acid, standardization of +Hydrolysis + +Ignition of precipitates +Indicators, definition of + for acidimetry + preparation of +Indirect methods +Insoluble matter, determination of in limestone +Integrity +Iodimetry +Ions, definition of +Iron, gravimetric determination of + volumetric determination of + +Jones reductor + +Lead, determination of in brass +Limestone, analysis of +Limonite, determination of iron in +Liquids, evaporation of + transfer of +Litmus +Logarithms + +Magnesium, determination of +Mass action, law of +Measuring instruments +Methyl orange +Moisture, determination of in limestone + +Neutralization methods +Normal solutions, acid and alkali + oxidizing agents + reducing agents +Notebooks, sample pages of + +Oxalic acid, determination of strength of +Oxidation processes +Oxidizing power of pyrolusite + +Permanganate process for iron +Phenolphthalein +Phosphoric anhydride, determination of +Pipette, calibration of + description of +Platinum crucibles, care of +Precipitates, colloidal + crystalline + ignition of + separation from filter + washing of +Precipitation +Precipitation methods (volumetric) +Problems +Pyrolusite, oxidizing power of + +Quantitative Analyses, subdivisions of + +Reagents, strength of +Reducing solution, normal +Reductor, Jones +Reversible reactions + +Silica, determination of, in limestone + determination of, in silicates + purification of +Silicic acid, dehydration of +Silver, determination of in coin +Soda ash, alkaline strength of +Sodium chloride, determination of chlorine in +Solubility product +Solution pressure +Solutions, normal + standard +Standardization, definition of +Standard solutions, acidimetry and alkalimetry + chlorimetry + iodimetry + oxidizing and reducing agents + thiocyanate +Starch solutions +Stibnite, determination of antimony in +Stirring rods +Stoichiometry +Strength of reagents +Suction, use of +Sulphur, determination of in ferrous ammonium sulphate + in barium sulphate + +Temperature, corrections for +Testing of washings +Theory of electrolytic dissociation +Thiocyanate process for silver +Titration, definition of +Transfer of liquids + +Volumetric analysis, definition of + general directions + +Wash-bottles +Washed filters +Washing of precipitates +Washings, testing of +Water, ionization of + densities of +Weights, care of + +Zimmermann-Reinhardt method for iron +Zinc, determination of, in brass + + + + + + + +End of the Project Gutenberg EBook of An Introductory Course of Quantitative +Chemical Analysis, by Henry P. 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