[Transcriber's notes: In the chemical equations, superscripts areindicated with a ^ and subscripts are indicated with a _. The affecteditem is enclosed in curly brackets {}. Examples are H^{+} for hydrogenion and H_{2}O for water. Since the underscore is already being usedin this project, italics are designated by an exclamation pointbefore 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 OFTECHNOLOGY

SIXTH EDITION, COMPLETELY REWRITTEN

PREFACE

This Introductory Course of Quantitative Analysis has been preparedto meet the needs of students who are just entering upon the subject, after a course of qualitative analysis. It is primarily intended toenable the student to work successfully and intelligently without thenecessity for a larger measure of personal assistance and supervisionthan can reasonably be given to each member of a large class. To thisend the directions are given in such detail that there is very littleopportunity for the student to go astray; but the manual is not, theauthor believes, on this account less adapted for use with smallclasses, where the instructor, by greater personal influence, canstimulate independent thought on the part of the pupil. 

The method of presentation of the subject is that suggested byProfessor A. A. Noyes' excellent manual of Qualitative Analysis. Foreach analysis the procedure is given in considerable detail, andthis is accompanied by explanatory notes, which are believed to besufficiently expanded to enable the student to understand fully theunderlying reason for each step prescribed. The use of the bookshould, nevertheless, be supplemented by classroom instruction, mainlyof the character of recitations, and the student should be taught toconsult larger works. The general directions are intended to emphasizethose matters upon which the beginner in quantitative analysis mustbestow special care, and to offer helpful suggestions. The studentcan hardly be expected to appreciate the force of all the statementscontained in these directions, or, indeed, to retain them all inthe memory after a single reading; but the instructor, by frequentreference 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 firstcourse of quantitative analysis at the Massachusetts Institute ofTechnology, and have been chosen, after an experience of years, as affording the best preparation for more advanced work, and assatisfactory types of gravimetric and volumetric methods. From thelatter point of view, they also seem to furnish the best insight intoquantitative analysis for those students who can devote but a limitedtime to the subject, and who may never extend their study beyond thefield covered by this manual. The author has had opportunity to testthe efficiency of the course for use with such students, and has foundthe results satisfactory. 

In place of the usual custom of selecting simple salts as material forpreliminary practice, it has been found advantageous to substitute, inmost instances, approximately pure samples of appropriate minerals orindustrial products. The difficulties are not greatly enhanced, whilethe student gains in practical experience. 

The analytical procedures described in the following pages have beenselected chiefly with reference to their usefulness in teaching thesubject, and with the purpose of affording as wide a variety ofprocesses as is practicable within an introductory course of thischaracter. The scope of the manual precludes any extended attempt toindicate alternative procedures, except through general references tolarger works on analytical chemistry. The author is indebted to thestandard works for many suggestions for which it is impracticable tomake specific acknowledgment; no considerable credit is claimed by himfor originality of procedure. 

For many years, as a matter of convenience, the classes for which thistext was originally prepared were divided, one part beginning withgravimetric processes and the other with volumetric analyses. After acareful review of the experience thus gained the conclusion has beenreached that volumetric analysis offers the better approach to thesubject. Accordingly the arrangement of the present (the sixth)edition of this manual has been changed to introduce volumetricprocedures first. Teachers who are familiar with earlier editionswill, however, find that the order of presentation of the materialunder the various divisions is nearly the same as that previouslyfollowed, and those who may still prefer to begin the course ofinstruction with gravimetric processes will, it is believed, be ableto 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 acidsolutions, and for the standardization of potassium permanganatesolutions using sodium oxalate as a standard, and of thiosulphatesolutions using copper as a standard, have been added. Thedetermination of silica in silicates decomposable by acids, as aseparate procedure, has been omitted. 

The explanatory notes have been rearranged to bring them into closerassociation with the procedures to which they relate. The number ofproblems has been considerably increased. 

The author wishes to renew his expressions of appreciation of thekindly reception accorded the earlier editions of this manual. He hasreceived helpful suggestions from so many of his colleagues within theInstitute, and friends elsewhere, that his sense of obligation mustbe expressed to them collectively. He is under special obligationsto Professor L. F. Hamilton for assistance in the preparation of thepresent 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 DevicesGENERAL DIRECTIONS Standard and Normal Solutions

!I. Neutralization Methods!

ALKALIMETRY AND ACIDIMETRY Preparation and Standardization of Solutions; IndicatorsSTANDARDIZATION OF HYDROCHLORIC ACIDDETERMINATION OF TOTAL ALKALINE STRENGTH OF SODA ASHDETERMINATION OF ACID STRENGTH OF OXALIC ACID

!II. Oxidation Processes!

GENERAL DISCUSSIONBICHROMATE PROCESS FOR THE DETERMINATION OF IRONDETERMINATION OF IRON IN LIMONITE BY THE BICHROMATE PROCESSDETERMINATION OF CHROMIUM IN CHROME IRON OREPERMANGANATE PROCESS FOR THE DETERMINATION OF IRONDETERMINATION OF IRON IN LIMONITE BY THE PERMANGANATE PROCESSDETERMINATION OF IRON IN LIMONITE BY THE ZIMMERMANN-REINHARDT PROCESSDETERMINATION OF THE OXIDIZING POWER OF PYROLUSITEIODIMETRYDETERMINATION OF COPPER IN ORESDETERMINATION OF ANTIMONY IN STIBNITECHLORIMETRYDETERMINATION 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 PrecipitatesDETERMINATION OF CHLORINE IN SODIUM CHLORIDEDETERMINATION OF IRON AND OF SULPHUR IN FERROUS AMMONIUM SULPHATEDETERMINATION OF SULPHUR IN BARIUM SULPHATEDETERMINATION OF PHOSPHORIC ANHYDRIDE IN APATITEANALYSIS OF LIMESTONE Determination of Moisture; Insoluble Matter and Silica; Ferric Oxide and Alumina; Calcium; Magnesium; Carbon DioxideANALYSIS OF BRASS Electrolytic Separations; Determination of Lead, Copper, Iron and Zinc. DETERMINATION OF SILICA IN SILICATES

PART IV. STOICHIOMETRY

SOLUTIONS OF TYPICAL PROBLEMSPROBLEMS

APPENDIX

ELECTROLYTIC DISSOCIATION THEORYFOLDING OF A FILTER PAPERSAMPLE NOTEBOOK PAGESSTRENGTH OF REAGENTSDENSITIES AND VOLUMES OF WATERCORRECTIONS FOR CHANGE OF TEMPERATURE OF STANDARD SOLUTIONSATOMIC WEIGHTSLOGARITHM TABLES

QUANTITATIVE CHEMICAL ANALYSIS

PART I

INTRODUCTION

SUBDIVISIONS OF ANALYTICAL CHEMISTRY

A complete chemical analysis of a body of unknown composition involvesthe recognition of its component parts by the methods of !qualitativeanalysis!, and the determination of the proportions in which thesecomponents are present by the processes of !quantitative analysis!. A preliminary qualitative examination is generally indispensable, ifintelligent and proper provisions are to be made for the separation ofthe various constituents under such conditions as will insure accuratequantitative estimations. 

It is assumed that the operations of qualitative analysis are familiarto the student, who will find that the reactions made use of inquantitative processes are frequently the same as those employed inqualitative analyses with respect to both precipitation and systematicseparation from interfering substances; but it should be noted thatthe conditions must now be regulated with greater care, and in sucha manner as to insure the most complete separation possible. Forexample, in the qualitative detection of sulphates by precipitationas barium sulphate from acid solution it is not necessary, in mostinstances, to take into account the solubility of the sulphatein hydrochloric acid, while in the quantitative determination ofsulphates by this reaction this solubility becomes an importantconsideration. The operations of qualitative analysis are, therefore, the more accurate the nearer they are made to conform to quantitativeconditions. 

The methods of quantitative analysis are subdivided, accordingto their nature, into those of !gravimetric analysis, volumetricanalysis!, and !colorimetric analysis!. In !gravimetric! processes theconstituent to be determined is sometimes isolated in elementaryform, but more commonly in the form of some compound possessing awell-established and definite composition, which can be readily andcompletely separated, and weighed either directly or after ignition. From the weight of this substance and its known composition, theamount of the constituent in question is determined. 

In !volumetric! analysis, instead of the final weighing of a definitebody, a well-defined reaction is caused to take place, wherein thereagent is added from an apparatus so designed that the volume of thesolution employed to complete the reaction can be accurately measured. The strength of this solution (and hence its value for the reactionin question) is accurately known, and the volume employed serves, therefore, as a measure of the substance acted upon. An example willmake clear the distinction between these two types of analysis. The percentage of chlorine in a sample of sodium chloride may bedetermined by dissolving a weighed amount of the chloride in waterand precipitating the chloride ions as silver chloride, which isthen separated by filtration, ignited, and weighed (a !gravimetric!process); or the sodium chloride may be dissolved in water, and asolution of silver nitrate containing an accurately known amount ofthe silver salt in each cubic centimeter may be cautiously added froma measuring device called a burette until precipitation is complete, when the amount of chlorine may be calculated from the number of cubiccentimeters of the silver nitrate solution involved in the reaction. This is a !volumetric! process, and is equivalent to weighing withoutthe use of a balance. 

Volumetric methods are generally more rapid, require less apparatus, and are frequently capable of greater accuracy than gravimetricmethods. They are particularly useful when many determinations of thesame sort are required. 

In !colorimetric! analyses the substance to be determined is convertedinto some compound which imparts to its solutions a distinct color, the intensity of which must vary in direct proportion to the amount ofthe compound in the solution. Such solutions are compared with respectto depth of color with standard solutions containing known amounts ofthe colored compound, or of other similar color-producing substancewhich has been found acceptable as a color standard. Colorimetricmethods are, in general, restricted to the determinations of verysmall quantities, since only in dilute solutions are accuratecomparisons of color possible. 

GENERAL DIRECTIONS

The following paragraphs should be read carefully and thoughtfully. Aprime essential for success as an analyst is attention to details andthe avoidance of all conditions which could destroy, or even lessen, confidence in the analyses when completed. The suggestions here givenare the outcome of much experience, and their adoption will tend toinsure permanently work of a high grade, while neglect of them willoften lead to disappointment and loss of time. 

ACCURACY AND ECONOMY OF TIME

The fundamental conception of quantitative analysis implies anecessity for all possible care in guarding against loss of materialor the introduction of foreign matter. The laboratory desk, and allapparatus, should be scrupulously neat and clean at all times. Asponge should always be ready at hand, and desk and filter-standsshould be kept dry and in good order. Funnels should never be allowedto drip upon the base of the stand. Glassware should always bewiped with a clean, lintless towel just before use. All filters andsolutions should be covered to protect them from dust, just as far asis practicable, and every drop of solution or particle of precipitatemust be regarded as invaluable for the success of the analysis. 

An economical use of laboratory hours is best secured by acquiringa thorough knowledge of the character of the work to be done beforeundertaking it, and then by so arranging the work that no time shallbe wasted during the evaporation of liquids and like time-consumingoperations. To this end the student should read thoughtfully not onlythe !entire! procedure, but the explanatory notes as well, beforeany step is taken in the analysis. The explanatory notes furnish, ingeneral, the reasons for particular steps or precautions, but theyalso occasionally contain details of manipulation not incorporated, for various reasons, in the procedure. These notes follow theprocedures at frequent intervals, and the exact points to which theyapply are indicated by references. The student should realize that a!failure to study the notes will inevitably lead to mistakes, loss oftime, and an inadequate understanding of the subject!. 

All analyses should be made in duplicate, and in general a closeagreement of results should be expected. It should, however, beremembered that a close concordance of results in "check analyses" isnot conclusive evidence of the accuracy of those results, although theprobability of their accuracy is, of course, considerably enhanced. The satisfaction in obtaining "check results" in such analyses mustnever be allowed to interfere with the critical examination of theprocedure employed, nor must they ever be regarded as in any measure asubstitute for absolute truth and accuracy. 

In this connection it must also be emphasized that only the operatorhimself can know the whole history of an analysis, and only he canknow whether his work is worthy of full confidence. No work should becontinued for a moment after such confidence is lost, but shouldbe resolutely discarded as soon as a cause for distrust is fullyestablished. The student should, however, determine to put forth hisbest efforts in each analysis; it is well not to be too ready tocondone failures and to "begin again, " as much time is lost in thesefruitless attempts. Nothing less than !absolute integrity! is or canbe demanded of a quantitative analyst, and any disregard of thisprinciple, however slight, is as fatal to success as lack of chemicalknowledge 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 theright-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 isof the first importance, and is an evidence of careful work and anexcellent credential. Of two notebooks in which the results may be, in fact, of equal value as legal evidence, that one which is neatlyarranged will carry with it greater weight. 

All records should be dated, and all observations should be recordedat once in the notebook. The making of records upon loose paper is apractice to be deprecated, as is also that of copying original entriesinto a second notebook. The student should accustom himself to orderlyentries at the time of observation. Several sample pages of systematicrecords are to be found in the Appendix. These are based uponexperience; but other arrangements, if clear and orderly, may proveequally serviceable. The student is advised to follow the sample pagesuntil 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, inspite of all reasonable precautionary measures, inferior chemicalswill occasionally find their way into the stock room, or errors willbe made in filling reagent bottles. The student should remember thatwhile there may be others who share the responsibility for the purityof materials in the laboratory of an institution, the responsibilitywill 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 allsuch bottles should be kept scrupulously clean, and care taken that noconfusion of stoppers occurs. 

WASH-BOTTLES

Wash-bottles for distilled water should be made from flasks of about750 cc. Capacity and be provided with gracefully bent tubes, whichshould not be too long. The jet should be connected with the tubeentering the wash-bottle by a short piece of rubber tubing in sucha way as to be flexible, and should deliver a stream about onemillimeter in diameter. The neck of the flask may be wound with cord, or covered with wash-leather, for greater comfort when hot water isused. It is well to provide several small wash-bottles for liquidsother than distilled water, which should invariably be clearlylabeled. 

TRANSFER OF LIQUIDS

Liquids should never be transferred from one vessel to another, nor toa filter, without the aid of a stirring rod held firmly against theside or lip of the vessel. When the vessel is provided with a lip itis not usually necessary to use other means to prevent the loss ofliquid by running down the side; whenever loss seems imminent a !verythin! layer of vaseline, applied with the finger to the edge of thevessel, will prevent it. The stirring rod down which the liquid runsshould never be drawn upward in such a way as to allow the solution tocollect on the under side of the rim or lip of a vessel. 

The number of transfers of liquids from one vessel to another duringan analysis should be as small as possible to avoid the risk of slightlosses. Each vessel must, of course, be completely washed to insurethe transfer of all material; but it should be remembered that thiscan be accomplished better by the use of successive small portions ofwash-water (perhaps 5-10 cc. ), if each wash-water is allowed to drainaway for a few seconds, than by the addition of large amounts whichunnecessarily increase the volume of the solutions, causing loss oftime in subsequent filtrations or evaporations. 

All stirring rods employed in quantitative analyses should be roundedat the ends by holding them in the flame of a burner until they beginto soften. If this is not done, the rods will scratch the innersurface of beakers, causing them to crack on subsequent heating. 

EVAPORATION OF LIQUIDS

The greatest care must be taken to prevent loss of solutions duringprocesses of evaporation, either from too violent ebullition, fromevaporation to dryness and spattering, or from the evolution of gasduring the heating. In general, evaporation upon the steam bath is tobe preferred to other methods on account of the impossibility ofloss by spattering. If the steam baths are well protected from dust, solutions should be left without covers during evaporation; butsolutions which are boiled upon the hot plate, or from which gases areescaping, should invariably be covered. In any case a watch-glass maybe supported above the vessel by means of a glass triangle, or othersimilar device, and the danger of loss of material or contamination bydust thus be avoided. It is obvious that evaporation is promoted bythe use of vessels which admit of the exposure of a broad surface tothe air. 

Liquids which contain suspended matter (precipitates) should alwaysbe cautiously heated, since the presence of the solid matter isfrequently the occasion of violent "bumping, " with consequent risk toapparatus and analysis. 

PART II

VOLUMETRIC ANALYSIS

The processes of volumetric analysis are, in general, simpler thanthose of gravimetric analysis and accordingly serve best as anintroduction to the practice of quantitative analysis. For theirexecution there are required, first, an accurate balance with whichto weigh the material for analysis; second, graduated instruments inwhich to measure the volume of the solutions employed; third, standardsolutions, that is, solutions the value of which is accurately known;and fourth, indicators, which will furnish accurate evidence of thepoint at which the desired reaction is completed. The nature of theindicators employed will be explained in connection with the differentanalyses. 

The process whereby a !standard solution! is brought into reaction iscalled !titration!, and the point at which the reaction is exactlycompleted is called the !end-point!. The !indicator! should show the!end-point! of the !titration!. The volume of the standard solutionused then furnishes the measure of the substance to be determined astruly as if that substance had been separated and weighed. 

The processes of volumetric analysis are easily classified, accordingto their character, into:

I. NEUTRALIZATION METHODS; such, for example, as those of acidimetryand alkalimetry. 

II. OXIDATION PROCESSES; as exemplified in the determination offerrous iron by its oxidation with potassium bichromate. 

III. PRECIPITATION METHODS; of which the titration for silver withpotassium thiocyanate solution is an illustration. 

From a somewhat different standpoint the methods in each case maybe subdivided into (a) DIRECT METHODS, in which the substance to bemeasured is directly determined by titration to an end-point with astandard solution; and (b) INDIRECT METHODS, in which the substanceitself is not measured, but a quantity of reagent is added which isknown to be an excess with respect to a specific reaction, and theunused excess determined by titration. Examples of the latter classwill be pointed out as they occur in the procedures. 

MEASURING INSTRUMENTS

THE ANALYTICAL BALANCE

For a complete discussion of the physical principles underlying theconstruction 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 tothe admirable discussion of this topic in Morse's !Exercises inQuantitative Chemistry!. 

The statements and rules of procedure which follow are sufficientfor the intelligent use of an analytical balance in connection withprocesses prescribed in this introductory manual. It is, however, imperative that the student should make himself familiar with theseessential features of the balance, and its use. He should fullyrealize that the analytical balance is a delicate instrument whichwill render excellent service under careful treatment, but suchtreatment is an essential condition if its accuracy is to be dependedupon. He should also understand that no set of rules, howevercomplete, can do away with the necessity for a sense of personalresponsibility, since by carelessness he can render inaccurate notonly his own analyses, but those of all other students using the samebalance. 

Before making any weighings the student should seat himself before abalance and observe the following details of construction:

1. The balance case is mounted on three brass legs, which shouldpreferably rest in glass cups, backed with rubber to prevent slipping. The front legs are adjustable as to height and are used to level thebalance 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 thecase on the top of which is an inlaid agate plate. To the center ofthe beam there is attached a steel or agate knife-edge on which thebeam oscillates when it rests on the agate plate. 

4. The balance beam, extending to the right and left, is graduatedalong its upper edge, usually on both sides, and has at itsextremities two agate or steel knife-edges from which are suspendedstirrups. Each of these stirrups has an agate plate which, when thebalance is in action, rests upon the corresponding knife-edge of thebeam. The balance pans are suspended from the stirrups. 

5. A pointer is attached to the center of the beam, and as the beamoscillates this pointer moves in front of a scale near the base of thepost. 

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 thefront of the balance case, which is so arranged as to raise the entirebeam slightly above the level at which the knife-edges are in contactwith the agate plates. When the balance is not in use the beam mustbe supported by this device since, otherwise, the constant jarringto which a balance is inevitably subjected, will soon dull theknife-edges, and lessen the sensitiveness of the balance. 

8. A small weight, or bob, is attached to the pointer (or sometimesto the beam) by which the center of gravity of the beam and itsattachments may be regulated. The center of gravity must lie veryslightly below the level of the agate plates to secure the desiredsensitiveness of the balance. This is provided for when the balance isset up and very rarely requires alteration. The student should neverattempt to change this adjustment. 

9. Below the balance pans are two pan-arrests operated by a buttonfrom the front of the case. These arrests exert a very slight upwardpressure upon the pans and minimize the displacement of the beam whenobjects or weights are being placed upon the pans. 

10. A movable rod, operated from one end of the balance case, extendsover 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 divisionof the scale on the balance beam. Each numbered division on the beamcorresponds to one milligram, and the use of the rider obviates theplacing of very small fractional weights on the balance pan. 

If a new rider is purchased, or an old one replaced, care must betaken that its weight corresponds to the graduations on the beam ofthe balance on which it is to be used. The weight of the rider inmilligrams 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 ofthe beam. It should be noted that on some balances the last divisionbears no number. Each new rider should be tested against a 5 or10-milligram weight. 

In some of the most recent forms of the balance a chain devicereplaces the smaller weights and the use of the rider as justdescribed. 

Before using a balance, it is always best to test its adjustment. Thisis absolutely necessary if the balance is used by several workers; itis 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 raisingand lowering the beams works smoothly; (3) that the pan-arrests touchthe pans when the beam is lowered; and (4) that the needle swingsequal distances on either side of the zero-point when set in motionwithout any load on the pans. If the latter condition is notfulfilled, the balance should be adjusted by means of the adjustingscrew at the end of the beam unless the variation is not more than onedivision on the scale; it is often better to make a proper allowancefor this small zero error than to disturb the balance by an attempt atcorrection. Unless a student thoroughly understands the constructionof a balance he should never attempt to make adjustments, but shouldapply to the instructor in charge. 

The object to be weighed should be placed on the left-hand balance panand the weights upon the right-hand pan. Every substance whichcould attack the metal of the balance pan should be weighed upon awatch-glass, and all objects must be dry and cold. A warm body givesrise to air currents which vitiate the accuracy of the weighing. 

The weights should be applied in the order in which they occur in theweight-box (not at haphazard), beginning with the largest weight whichis apparently required. After a weight has been placed upon the panthe beam should be lowered upon its knife-edges, and, if necessary, the pan-arrests depressed. The movement of the pointer will thenindicate whether the weight applied is too great or too small. Whenthe weight has been ascertained, by the successive addition of smallweights, to the nearest 5 or 10 milligrams, the weighing is completedby the use of the rider. The correct weight is that which causes thepointer to swing an equal number of divisions to the right and leftof the zero-point, when the pointer traverses not less than fivedivisions 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 ofair currents. 

Before the final determination of an exact weight the beam shouldalways be lifted from the knife-edges and again lowered into place, as it frequently happens that the scale pans are, in spite of thepan-arrests, slightly twisted by the impact of the weights, the beambeing thereby virtually lengthened or shortened. Lifting the beamrestores the proper alignment. 

The beam should never be set in motion by lowering it forcibly uponthe knife-edges, nor by touching the pans, but rather by lifting therider (unless the balance be provided with some of the newer devicesfor the purpose), and the swing should be arrested only when theneedle approaches zero on the scale, otherwise the knife-edges becomedull. For the same reason the beam should never be left upon itsknife-edges, nor should weights be removed from or placed on thepans without supporting the beam, except in the case of the smallfractional weights. 

When the process of weighing has been completed, the weight shouldbe recorded in the notebook by first noting the vacant spaces in theweight-box, and then checking the weight by again noting the weightsas they are removed from the pan. This practice will often detect andavoid errors. It is obvious that the weights should always be returnedto 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 isderanged or if any substance is spilled upon the pans or in thebalance case, the damage should be reported at once. In many instancesserious harm can be averted by prompt action when delay might ruin thebalance. 

Samples for analysis are commonly weighed in small tubes with corkstoppers. Since the stoppers are likely to change in weight fromthe varying amounts of moisture absorbed from the atmosphere, it isnecessary to confirm the recorded weight of a tube which has beenunused for some time before weighing out a new portion of substancefrom it. 

WEIGHTS

The sets of weights commonly used in analytical chemistry range from20 grams to 5 milligrams. The weights from 20 grams to 1 gram areusually of brass, lacquered or gold plated. The fractional weightsare of German silver, gold, platinum or aluminium. The rider is ofplatinum or aluminium wire. 

The sets of weights purchased from reputable dealers are usuallysufficiently accurate for analytical work. It is not necessary thatsuch a set should be strictly exact in comparison with the absolutestandard of weight, provided they are relatively correct amongthemselves, and provided the same set of weights is used in allweighings made during a given analysis. The analyst should assurehimself that the weights in a set previously unfamiliar to him arerelatively correct by a few simple tests. For example, he should makesure that in his set two weights of the same denomination (i. E. , two10-gram weights, or the two 100-milligram weights) are actually equaland interchangeable, or that the 500-milligram weight is equal tothe sum of the 200, 100, 100, 50, 20, 20 and 10-milligram weightscombined, and so on. If discrepancies of more than a few tenths of amilligram (depending upon the total weight involved) are found, theweights should be returned for correction. The rider should also becompared with a 5 or 10-milligram weight. 

In an instructional laboratory appreciable errors should be reportedto the instructor in charge for his consideration. 

When the highest accuracy is desired, the weights may be calibratedand corrections applied. A calibration procedure is described in apaper by T. W. Richards, !J. Am. Chem. Soc. !, 22, 144, and in manylarge text-books. 

Weights are inevitably subject to corrosion if not properly protectedat all times, and are liable to damage unless handled with great care. It is obvious that anything which alters the weight of a single piecein an analytical set will introduce an error in every weighing madein which that piece is used. This source of error is often extremelyobscure and difficult to detect. The only safeguard against sucherrors is to be found in scrupulous care in handling and protectionon the part of the analyst, and an equal insistence that if severalanalysts use the same set of weights, each shall realize hisresponsibility for the work of others as well as his own. 

BURETTES

A burette is made from a glass tube which is as uniformly cylindricalas possible, and of such a bore that the divisions which are etchedupon its surface shall correspond closely to actual contents. 

The tube is contracted at one extremity, and terminates in either aglass stopcock and delivery-tube, or in such a manner that a piece ofrubber tubing may be firmly attached, connecting a delivery-tube ofglass. The rubber tubing is closed by means of a glass bead. Burettesof the latter type will be referred to as "plain burettes. "

The graduations are usually numbered in cubic centimeters, and thelatter are subdivided into tenths. 

One burette of each type is desirable for the analytical procedureswhich follow. 

PREPARATION OF A BURETTE FOR USE

The inner surface of a burette must be thoroughly cleaned in orderthat the liquid as drawn out may drain away completely, withoutleaving drops upon the sides. This is best accomplished by treatingthe inside of the burette with a warm solution of chromic acid inconcentrated sulphuric acid, applied as follows: If the burette is ofthe "plain" type, first remove the rubber tip and force the lowerend of the burette into a medium-sized cork stopper. Nearly fill theburette with the chromic acid solution, close the upper end with acork stopper and tip the burette backward and forward in such a wayas to bring the solution into contact with the entire inner surface. Remove the stopper and pour the solution into a stock bottle to bekept for further use, and rinse out the burette with water severaltimes. Unless the water then runs freely from the burette withoutleaving drops adhering to the sides, the process must be repeated(Note 1). 

If the burette has a glass stopcock, this should be removed afterthe cleaning and wiped, and also the inside of the ground joint. Thesurface of the stopcock should then be smeared with a thin coating ofvaseline and replaced. It should be attached to the burette by meansof a wire, or elastic band, to lessen the danger of breakage. 

Fill the burettes with distilled water, and allow the water to run outthrough the stopcock or rubber tip until convinced that no airbubbles are inclosed (Note 2). Fill the burette to a point above thezero-point and draw off the water until the meniscus is just belowthat mark. It is then ready for calibration. 

[Note 1: The inner surface of the burette must be absolutely clean ifthe liquid is to run off freely. Chromic acid in sulphuric acid isusually found to be the best cleansing agent, but the mixture must bewarm and concentrated. The solution can be prepared by pouring over afew crystals of potassium bichromate a little water and then addingconcentrated sulphuric acid. ]

[Note 2: It is always necessary to insure the absence of air bubblesin the tips or stopcocks. The treatment described above will usuallyaccomplish this, but, in the case of plain burettes it is sometimesbetter to allow a little of the liquid to flow out of the tip while itis bent upwards. Any air which may be entrapped then rises with theliquid and escapes. 

If air bubbles escape during subsequent calibration or titration, anerror is introduced which vitiates the results. ]

READING OF A BURETTE

All liquids when placed in a burette form what is called a meniscus attheir upper surfaces. In the case of liquids such as water oraqueous solutions this meniscus is concave, and when the liquids aretransparent accurate readings are best obtained by observing theposition on the graduated scales of the lowest point of the meniscus. This can best be done as follows: Wrap around the burette a piece ofcolored paper, the straight, smooth edges of which are held evenlytogether with the colored side next to the burette (Note 1). Hold thepaper about two small divisions below the meniscus and raise or lowerthe level of the eyes until the edge of the paper at the back of theburette is just hidden from the eye by that in front (Note 2). Notethe position of the lowest point of the curve of the meniscus, estimating the tenths of the small divisions, thus reading itsposition to hundredths of a cubic centimeter. 

[Note 1: The ends of the colored paper used as an aid to accuratereadings may be fastened together by means of a gummed label. Thepaper may then remain on the burette and be ready for immediate use bysliding it up or down, as required. ]

[Note 2: To obtain an accurate reading the eye must be very nearly ona level with the meniscus. This is secured by the use of the paperas described. The student should observe by trial how a reading isaffected when the meniscus is viewed from above or below. 

The eye soon becomes accustomed to estimating the tenths of thedivisions. If the paper is held as directed, two divisions below themeniscus, one whole division is visible to correct the judgment. It isnot well to attempt to bring the meniscus exactly to a division markon the burette. Such readings are usually less accurate than those inwhich the tenths of a division are estimated. ]

CALIBRATION OF GLASS MEASURING DEVICES

If accuracy of results is to be attained, the correctness of allmeasuring instruments must be tested. None of the apparatus offeredfor sale can be implicitly relied upon except those more expensiveinstruments which are accompanied by a certificate from the !NationalBureau of Standards! at Washington, or other equally authentic source. 

The bore of burettes is subject to accidental variations, and sincethe graduations are applied by machine without regard to suchvariations of bore, local errors result. 

The process of testing these instruments is called !calibration!. It is usually accomplished by comparing the actual weight of watercontained in the instrument with its apparent volume. 

There is, unfortunately, no uniform standard of volume which has beenadopted for general use in all laboratories. It has been variouslyproposed 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 toconsider 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. ; butthis is obviously a lower temperature than that commonly found inlaboratories, and involves the constant use of corrections if taken asa laboratory standard. Many laboratories use 15. 5°C. (60° F. ) as theworking standard. It is plain that any temperature which is deemedmost convenient might be chosen for a particular laboratory, but itcannot be too strongly emphasized that all measuring instruments, including burettes, pipettes, and flasks, should be calibrated at thattemperature in order that the contents of each burette, pipette, etc. , shall be comparable with that of every other instrument, thuspermitting general interchange and substitution. For example, it isobvious that if it is desired to remove exactly 50 cc. From a solutionwhich has been diluted to 500 cc. In a graduated flask, the 50 cc. Flask or pipette used to remove the fractional portion must givea correct reading at the same temperature as the 500 cc. Flask. Similarly, a burette used for the titration of the 50 cc. Of solutionremoved should be calibrated under the same conditions as themeasuring flasks or pipettes employed with it. 

The student should also keep constantly in mind the fact that allvolumetric operations, to be exact, should be carried out as nearly ata constant temperature as is practicable. The spot selected forsuch work should therefore be subject to a minimum of temperaturevariations, and should have as nearly the average temperature ofthe laboratory as is possible. In all work, whether of calibration, standardization, or analysis, the temperature of the liquids employedmust be taken into account, and if the temperature of these liquidsvaries more than 3° or 4° from the standard temperature chosen for thelaboratory, corrections must be applied for errors due to expansion orcontraction, since volumes of a liquid measured at different times arecomparable only under like conditions as to temperature. Data to beused for this purpose are given in the Appendix. Neglect of thiscorrection is frequently an avoidable source of error and annoyance inotherwise excellent work. The temperature of all solutions at the timeof standardization should be recorded to facilitate the application oftemperature 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, arerequired throughout the volumetric work. Both burettes should becalibrated by the student to whom they are assigned. 

PROCEDURE. --Weigh a 50 cc. , flat-bottomed flask (preferably alight-weight flask), which must be dry on the outside, to the nearestcentigram. Record the weight in the notebook. (See Appendix forsuggestions as to records. ) Place the flask under the burette and drawout into it about 10 cc. Of water, removing any drop on the tip bytouching it against the inside of the neck of the flask. Do notattempt to stop exactly at the 10 cc. Mark, but do not vary more than0. 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 thenotebook (Note 1). Meanwhile weigh the flask and water to centigramsand record its weight (Note 2). Draw off the liquid from 10 cc. Toabout 20 cc. Into the same flask without emptying it; weigh, and atthe expiration of three minutes take the reading, and so on throughoutthe length of the burette. When it is completed, refill the buretteand check the first calibration. 

The differences in readings represent the apparent volumes, thedifferences in weights the true volumes. For example, if an apparentvolume of 10. 05 cc. Is found to weigh 10. 03 grams, it may be assumedwith 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 betaken into account, if this varies more than 4°C. From the laboratorystandard temperature, consulting the table of densities of water inthe Appendix. 

From the final data, plot the corrections to be applied so that theymay be easily read for each cubic centimeter throughout the burette. The total correction at each 10 cc. May also be written on the burettewith a diamond, or etching ink, for permanence of record. 

[Note 1: A small quantity of liquid at first adheres to the side ofeven a clean burette. This slowly unites with the main body of liquid, but requires an appreciable time. Three minutes is a sufficientinterval, but not too long, and should be adopted in every instancethroughout the whole volumetric practice before final readings arerecorded. ]

[Note 2: A comparatively rough balance, capable of weighing tocentigrams, is sufficiently accurate for use in calibrations, for amoment's reflection will show that it would be useless to weigh thewater with an accuracy greater than that of the readings taken onthe burette. The latter cannot exceed 0. 01 cc. In accuracy, whichcorresponds to 0. 01 gram. 

The student should clearly understand that !all other weighings!, except those for calibration, should be made accurately to 0. 0001gram, unless special directions are given to the contrary. 

Corrections for temperature variations of less than 4°C. Arenegligible, as they amount to less than 0. 01 gram for each 10 grams ofwater withdrawn. ]

[Note 3: Should the error discovered in any interval of 10 cc. On theburette exceed 0. 10 cc. , it is advisable to weigh small portions (even1 cc. ) to locate the position of the variation of bore in thetube rather than to distribute the correction uniformly over thecorresponding 10 cc. The latter is the usual course for smallcorrections, and it is convenient to calculate the correctioncorresponding to each cubic centimeter and to record it in the formof a table or calibration card, or to plot a curve representing thevalues. 

Burettes may also be calibrated by drawing off the liquid insuccessive portions through a 5 cc. Pipette which has been accuratelycalibrated, as a substitute for weighing. If many burettes are to betested, this is a more rapid method. ]

PIPETTES

A !pipette! may consist of a narrow tube, in the middle of which isblown a bulb of a capacity a little less than that which it is desiredto measure by the pipette; or it may be a miniature burette, withoutthe stopcock or rubber tip at the lower extremity. In either case, theflow 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 theyrequire calibration for accurate work. 

CALIBRATION OF PIPETTES

PROCEDURE. --Clean the pipette. Draw distilled water into it by suckingat the upper end until the water is well above the graduation mark. Quickly place the forefinger over the top of the tube, thus preventingthe entrance of air and holding the water in the pipette. Cautiouslyadmit a little air by releasing the pressure of the finger, and allowthe level of the water to fall until the lowest point of the meniscusis level with the graduation. Hold the water at that point by pressureof the finger and then allow the water to run out from the pipetteinto a small tared, or weighed, beaker or flask. After a definite timeinterval, usually two to three minutes, touch the end of the pipetteagainst the side of the beaker or flask to remove any liquid adheringto it (Note 1). The increase in weight of the flask in gramsrepresents the volume of the water in cubic centimeters delivered bythe pipette. Calculate the necessary correction. 

[Note 1: A definite interval must be allowed for draining, and adefinite practice adopted with respect to the removal of the liquidwhich collects at the end of the tube, if the pipette is designed todeliver a specific volume when emptied. This liquid may be removedat the end of a definite interval either by touching the side of thevessel or by gently blowing out the last drops. Either practice, whenadopted, must be uniformly adhered to. ]

FLASKS

!Graduated or measuring flasks! are similar to the ordinaryflat-bottomed flasks, but are provided with long, narrow necks inorder that slight variations in the position of the meniscus withrespect to the graduation shall represent a minimum volume of liquid. The flasks must be of such a capacity that, when filled with thespecified volume, the liquid rises well into the neck. 

GRADUATION OF FLASKS

It is a general custom to purchase the flasks ungraduated and tograduate them for use under standard conditions selected for thelaboratory in question. They may be graduated for "contents" or"delivery. " When graduated for "contents" they contain a specifiedvolume when filled to the graduation at a specified temperature, andrequire to be washed out in order to remove all of the solution fromthe flask. Flasks graduated for "delivery" will deliver the specifiedvolume of a liquid without rinsing. A flask may, of course, begraduated for both contents and delivery by placing two graduationmarks upon it. 

PROCEDURE. --To calibrate a flask for !contents!, proceed as follows:Clean the flask, using a chromic acid solution, and dry it carefullyoutside and inside. Tare it accurately; pour water into the flaskuntil the weight of the latter counterbalances weights on the oppositepan which equal in grams the number of cubic centimeters of waterwhich the flask is to contain. Remove any excess of water with the aidof filter paper (Note 1). Take the flask from the balance, stopperit, 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 thelocation of the lowest point of the meniscus (Note 2). The mark maybe etched upon the flask by hydrofluoric acid, or by the use of anetching 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 moisturefrom the outside. Fill it with distilled water. Pour out the waterand 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 adddistilled water to counterbalance these weights. An excess of water, or water adhering to the neck of the flask, may be removed by means ofa strip of clean filter paper. Stopper the flask, place it in a bathat 15. 5°C. Or 17. 5°C. And, after an hour, mark the location of thelowest point of the meniscus, as described above. 

[Note 1: The allowable error in counterbalancing the water andweights varies with the volume of the flask. It should not exceed oneten-thousandth of the weight of water. ]

[Note 2: Other methods are employed which involve the use ofcalibrated apparatus from which the desired volume of water may be runinto the dry flask and the position of the meniscus marked directlyupon it. For a description of a procedure which is most convenientwhen 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 analysesuniformity of practice must prevail throughout all volumetric workwith respect to those factors which can influence the accuracy of themeasurement of liquids. For example, whatever conditions are imposedduring the calibration of a burette, pipette, or flask (notably thetime allowed for draining), must also prevail whenever the flask orburette is used. 

The student should also be constantly watchful to insure parallelconditions during both standardization and analyst with respect to thefinal volume of liquid in which a titration takes place. The valueof a standard solution is only accurate under the conditions whichprevailed when it was standardized. It is plain that the standardsolutions must be scrupulously protected from concentration ordilution, after their value has been established. Accordingly, greatcare must be taken to thoroughly rinse out all burettes, flasks, etc. , with the solutions which they are to contain, in order to remove alltraces of water or other liquid which could act as a diluent. It isbest to wash out a burette at least three times with small portions ofa solution, allowing each to run out through the tip before assumingthat the burette is in a condition to be filled and used. It is, ofcourse, possible to dry measuring instruments in a hot closet, butthis is tedious and unnecessary. 

To the same end, all solutions should be kept stoppered and away fromdirect sunlight or heat. The bottles should be shaken before use tocollect any liquid which may have distilled from the solution andcondensed on the sides. 

The student is again reminded that variations in temperature ofvolumetric solutions must be carefully noted, and care should alwaysbe taken that no source of heat is sufficiently near the solutions toraise the temperature during use. 

Much time may be saved by estimating the approximate volume of astandard solution which will be required for a titration (if the dataare obtainable) before beginning the operation. It is then possible torun in rapidly approximately the required amount, after which it isonly necessary to determine the end-point slowly and with accuracy. In such cases, however, the knowledge of the approximate amount to berequired should never be allowed to influence the judgment regardingthe actual end-point. 

STANDARD SOLUTIONS

The strength or value of a solution for a specific reaction isdetermined by a procedure called !Standardization!, in which thesolution is brought into reaction with a definite weight of asubstance of known purity. For example, a definite weight of puresodium carbonate may be dissolved in water, and the volume of asolution of hydrochloric acid necessary to exactly neutralize thecarbonate accurately determined. From these data the strength or valueof the acid is known. It is then a !standard solution!. 

NORMAL SOLUTIONS

Standard solutions may be made of a purely empirical strength dictatedsolely by convenience of manipulation, or the concentration maybe chosen with reference to a system which is applicable to allsolutions, and based upon chemical equivalents. Such solutions arecalled !Normal Solutions! and contain such an amount of the reactingsubstance per liter as is equivalent in its chemical action to onegram of hydrogen, or eight grams of oxygen. Solutions containing onehalf, one tenth, or one one-hundredth of this quantity per liter arecalled, respectively, half-normal, tenth-normal, or hundredth-normalsolutions. 

Since normal solutions of various reagents are all referred to acommon standard, they have an advantage not possessed by empiricalsolutions, namely, that they are exactly equivalent to each other. Thus, a liter of a normal solution of an acid will exactly neutralizea liter of a normal alkali solution, and a liter of a normal oxidizingsolution will exactly react with a liter of a normal reducingsolution, and so on. 

Beside the advantage of uniformity, the use of normal solutionssimplifies the calculations of the results of analyses. This isparticularly true if, in connection with the normal solution, theweight of substance for analysis is chosen with reference to theatomic or molecular weight of the constituent to be determined. (Seeproblem 26. )

The preparation of an !exactly! normal, half-normal, or tenth-normalsolution requires considerable time and care. It is usually carriedout only when a large number of analyses are to be made, or when theanalyst has some other specific purpose in view. It is, however, acomparatively easy matter to prepare standard solutions which differbut slightly from the normal or half-normal solution, and these havethe advantage of practical equality; that is, two approximatelyhalf-normal solutions are more convenient to work with than two whichare widely different in strength. It is, however, true that some ofthe advantage which pertains to the use of normal solutions as regardssimplicity of calculations is lost when using these approximatesolutions. 

The application of these general statements will be made clear inconnection with the use of normal solutions in the various types ofvolumetric 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 offorming soluble compounds with the alkaline earths, but its solutionscannot be boiled without danger of loss of strength; sulphuric acidsolutions may be boiled without loss, but the acid forms insolublesulphates with three of the alkaline earths; oxalic acid can beaccurately weighed for the preparation of solutions, and its solutionsmay be boiled without loss, but it forms insoluble oxalates withthree of the alkaline earths and cannot be used with certain of theindicators. 

!Standard Alkali Solutions! may be prepared from sodium or potassiumhydroxide, sodium carbonate, barium hydroxide, or ammonia. Of sodiumand potassium hydroxide, it may be said that they can be used with allindicators, and their solutions may be boiled, but they absorb carbondioxide readily and attack the glass of bottles, thereby losingstrength; sodium carbonate may be weighed directly if its purity isassured, but the presence of carbonic acid from the carbonate is adisadvantage with many indicators; barium hydroxide solutions maybe prepared which are entirely free from carbon dioxide, and suchsolutions immediately show by precipitation any contamination fromabsorption, but the hydroxide is not freely soluble in water; ammoniadoes not absorb carbon dioxide as readily as the caustic alkalies, but its solutions cannot be boiled nor can they be used with allindicators. The choice of a solution must depend upon the nature ofthe work in hand. 

A !normal acid solution! should contain in one liter that quantity ofthe reagent which represents 1 gram of hydrogen replaceable by a base. For example, the normal solution of hydrochloric acid (HCl) shouldcontain 36. 46 grams of gaseous hydrogen chloride, since that amountfurnishes the requisite 1 gram of replaceable hydrogen. On the otherhand, the normal solution of sulphuric acid (H_{2}SO_{4}) shouldcontain only 49. 03 grams, i. E. , one half of its molecular weight ingrams. 

A !normal alkali solution! should contain sufficient alkali in a literto replace 1 gram of hydrogen in an acid. This quantity is representedby the molecular weight in grams (40. 01) of sodium hydroxide (NaOH), while a sodium carbonate solution (Na_{2}CO_{3}) should contain butone half the molecular weight in grams (i. E. , 53. 0 grams) in a literof normal solution. 

Half-normal or tenth-normal solutions are employed in most analyses(except in the case of the less soluble barium hydroxide). Solutionsof the latter strength yield more accurate results when smallpercentages of acid or alkali are to be determined. 

INDICATORS

It has already been pointed out that the purpose of an indicator is tomark (usually by a change of color) the point at which just enough ofthe titrating solution has been added to complete the chemical changewhich it is intended to bring about. In the neutralization processeswhich are employed in the measurement of alkalies (!alkalimetry!)or acids (!acidimetry!) the end-point of the reaction should, inprinciple, be that of complete neutrality. Expressed in terms of ionicreactions, it should be the point at which the H^{+} ions from anacid[Note 1] unite with a corresponding number of OH^{-} ions from abase 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 exactneutrality, but it is possible to approach it with sufficientexactness for analytical purposes, since substances are known which, in solution, undergo a sharp change of color as soon as even a minuteexcess 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 organiccompounds of complex structure and are closely allied to the dyestuffsin character. 

[Note 1: A knowledge on the part of the student of the ionic theoryas applied to aqueous solutions of electrolytes is assumed. A briefoutline of the more important applications of the theory is given inthe Appendix. ]

BEHAVIOR OF ORGANIC INDICATORS

The indicators in most common use for acid and alkali titrations aremethyl orange, litmus, and phenolphthalein. 

In the following discussion of the principles underlying the behaviorof the indicators as a class, methyl orange and phenolphthalein willbe taken as types. It has just been pointed out that indicators arebodies of complicated structure. In the case of the two indicatorsnamed, the changes which they undergo have been carefully studied byStieglitz (!J. Am. Chem. Soc. !, 25, 1112) and others, and it appearsthat the changes involved are of two sorts: First, a rearrangementof the atoms within the molecule, such as often occurs in organiccompounds; and, second, ionic changes. The intermolecular changescannot appropriately be discussed here, as they involve a somewhatdetailed knowledge of the classification and general behavior oforganic compounds; they will, therefore, be merely alluded to, andonly 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 itimparts to solutions is ascribed to the presence of the undissociatedbase. If an acid, such as HCl, is added to such a solution, the acidreacts with the indicator (neutralizes it) and a salt is formed, asindicated by the equation:

(M. O. )^{+}, OH^{-} + H^{+}, Cl^{-} --> (M. O. )^{+} Cl^{-} + (H_{2}O). 

This salt ionizes into (M. O. )^{+} (using this abbreviation for thepositive complex) and Cl^{-}; but simultaneously with this ionizationthere appears to be an internal rearrangement of the atoms whichresults in the production of a cation which may be designated as(M'. O'. )^{+}, and it is this which imparts a characteristic red colorto the solution. As these changes occur in the presence of even avery small excess of acid (that is, of H^{+} ions), it serves as thedesired index of their presence in the solution. If, now, an alkali, such as NaOH, is added to this reddened solution, the reverseseries of changes takes place. As soon as the free acid present isneutralized, the slightest excess of sodium hydroxide, acting asa strong base, sets free the weak, little-dissociated base of theindicator, and at the moment of its formation it reverts, because ofthe 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-dissociatedacid, which is colorless in neutral aqueous solution or in thepresence of free H^{+} ions. When an alkali is added to such asolution, even in slight excess, the anion of the salt which hasformed from the acid of the indicator undergoes a rearrangement of theatoms, and a new ion, (Ph')^{+}, is formed, which imparts a pink colorto 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, onthe other hand, occasions first the reversion to the colorless ion andthen 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 towardalkalies and phenolphthalein toward acids; the others occupyintermediate positions. That methyl orange should be most sensitivetoward alkalies is evident from the following considerations: Methylorange is a weak base and, therefore, but little dissociated. Itshould, then, be formed in the undissociated condition as soon as evena slight excess of OH^{-} ions is present in the solution, and thereshould be a prompt change from red to yellow as outlined above. On theother hand, it should be an unsatisfactory indicator for use with weakacids (acetic acid, for example) because the salts which it formswith such acids are, like all salts of that type, hydrolyzed to aconsiderable extent. This hydrolytic change is illustrated by theequation:

(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 plainthat hydrolysis is just the reverse of neutralization and must, accordingly, interfere with it. Salts of methyl orange with weak acidsare so far hydrolyzed that the end-point is uncertain, and methylorange cannot be used in the titration of such acids, while withthe very weak acids, such as carbonic acid or hydrogen sulphide(hydrosulphuric acid), the salts formed with methyl orange are, ineffect, completely hydrolyzed (i. E. , no neutralization occurs), andmethyl orange is accordingly scarcely affected by these acids. Thisexplains its usefulness, as referred to later, for the titration ofstrong acids, such as hydrochloric acid, even in the presence ofcarbonates or sulphides in solution. 

Phenolphthalein, on the other hand, should be, as it is, the best ofthe common indicators for use with weak acids. For, since it isitself a weak acid, it is very little dissociated, and its nearlyundissociated, colorless molecules are promptly formed as soon asthere 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 saltswhich it forms with weak alkalies are easily hydrolyzed, and as aconsequence 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 lesshydrolyzed and because the excess of OH^{-} ions from these bases alsodiminishes 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 isthe indicator most generally employed for the titration of organicacids. 

In general, it may be stated that when a strong acid, such ashydrochloric, sulphuric or nitric acid, is titrated against a strongbase, such as sodium hydroxide, potassium hydroxide, or bariumhydroxide, any of these indicators may be used, since very littlehydrolysis ensues. It has been noted above that the color change doesnot occur exactly at theoretical neutrality, from which it followsthat no two indicators will show exactly the same end-point when acidsand alkalis are brought together. It is plain, therefore, that thesame 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 volumeof an alkali solution (tenth-normal sodium hydroxide) required toproduce an alkaline end-point when run into 10 cc. Of tenth-normalsulphuric acid, diluted with 50 cc. Of water, using five drops of eachof 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 | YellowLacmoid | 10 | 10. 00 | Red | BlueLitmus | 10 | 10. 00 | Red | BlueRosalic acid | 10 | 10. 07 | Yellow | PinkPhenolphthalein| 10 | 10. 10 | Colorless | Pink====================================================================

It should also be stated that there are occasionally secondarychanges, other than those outlined above, which depend upon thetemperature and concentration of the solutions in which the indicatorsare used. These changes may influence the sensitiveness of anindicator. It is important, therefore, to take pains to useapproximately the same volume of solution when standardizing that islikely to be employed in analysis; and when it is necessary, as isoften the case, to titrate the solution at boiling temperature, thestandardization should take place under the same conditions. It isalso obvious that since some acid or alkali is required to react withthe indicator itself, the amount of indicator used should be uniformand not excessive. Usually a few drops of solution will suffice. 

The foregoing statements with respect to the behavior of indicatorspresent the subject in its simplest terms. Many substances other thanthose named may be employed, and they have been carefully studied todetermine the exact concentration of H^{+} ions at which the colorchange of each occurs. It is thus possible to select an indicatorfor a particular purpose with considerable accuracy. As data of thisnature do not belong in an introductory manual, reference is made tothe following papers or books in which a more extended treatment ofthe subject may be found:

Washburn, E. W. , Principles of Physical Chemistry (McGraw-Hill BookCo. ), (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 toVolumetric Analysis, !J. Am. Chem. Soc. !, 32 (1911), 815-861. 

Bjerrum, N. , General Discussion, !Z. Anal. Chem. !, 66 (1917), 13-28and 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 bydissolving 0. 05-0. 1 gram of the compound (also known as Orange III) ina few cubic centimeters of alcohol and diluting with water to 100 cc. A good grade of material should be secured. It can be successfullyused for the titration of hydrochloric, nitric, sulphuric, phosphoric, and sulphurous acids, and is particularly useful in the determinationof bases, such as sodium, potassium, barium, calcium, and ammoniumhydroxides, and even many of the weak organic bases. It can also beused for the determination, by titration with a standard solution ofa strong acid, of the salts of very weak acids, such as carbonates, sulphides, arsenites, borates, and silicates, because the weak acidswhich are liberated do not affect the indicator, and the reddening ofthe solution does not take place until an excess of the strong acidis added. It should be used in cold, not too dilute, solutions. Itssensitiveness is lessened in the presence of considerable quantitiesof the salts of the alkalies. 

A !phenolphthalein solution! is prepared by dissolving 1 gram of thepure compound in 100 cc. Of 95 per cent alcohol. This indicator isparticularly valuable in the determination of weak acids, especiallyorganic acids. It cannot be used with weak bases, even ammonia. Itis affected by carbonic acid, which must, therefore, be removed byboiling when other acids are to be measured. It can be used in hotsolutions. Some care is necessary to keep the volume of the solutionsto be titrated approximately uniform in standardization and inanalysis, and this volume should not in general exceed 125-150 cc. Forthe best results, since the compounds formed by the indicator undergochanges in very dilute solution which lessen its sensitiveness. 

The preparation of a !solution of litmus! which is suitable for useas an indicator involves the separation from the commercial litmus ofazolithmine, the true coloring principle. Soluble litmus tablets areoften obtainable, but the litmus as commonly supplied to the market ismixed with calcium carbonate or sulphate and compressed into lumps. Toprepare a solution, these are powdered and treated two or three timeswith alcohol, which dissolves out certain constituents which cause atroublesome intermediate color if not removed. The alcohol is decantedand drained off, after which the litmus is extracted with hot wateruntil exhausted. The solution is allowed to settle for some time, theclear liquid siphoned off, concentrated to one-third its volume andacetic acid added in slight excess. It is then concentrated to asirup, and a large excess of 95 per cent. Alcohol added to it. Thisprecipitates the blue coloring matter, which is filtered off, washedwith alcohol, and finally dissolved in a small volume of water anddiluted until about three drops of the solution added to 50 cc. Ofwater just produce a distinct color. This solution must be kept in anunstoppered bottle. It should be protected from dust by a loose plugof absorbent cotton. If kept in a closed bottle it soon undergoes areduction and loses its color, which, however, is often restored byexposure to the air. 

Litmus can be employed successfully with the strong acids and bases, and also with ammonium hydroxide, although the salts of the latterinfluence the indicator unfavorably if present in considerableconcentration. It may be employed with some of the stronger organicacids, 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 hydrochloricacid into a clean liter bottle, and dilute with distilled water to anapproximate volume of 1000 cc. Shake the solution vigorously for afull minute to insure uniformity. Be sure that the bottle is not toofull to permit of a thorough mixing, since lack of care at this pointwill be the cause of much wasted time (Note 1). 

Weigh out, upon a rough balance, 23 grams of sodium hydroxide (Note2). Dissolve the hydroxide in water in a beaker. Pour the solutioninto a liter bottle and dilute, as above, to approximately 1000 cc. This bottle should preferably have a rubber stopper, as the hydroxidesolution attacks the glass of the ground joint of a glass stopper, andmay cement the stopper to the bottle. Shake the solution as describedabove. 

[Note 1: The original solutions are prepared of a strength greaterthan 0. 5 N, as they are more readily diluted than strengthened iflater adjustment is desired. 

Too much care cannot be taken to insure perfect uniformity ofsolutions before standardization, and thoroughness in this respectwill, as stated, often avoid much waste of time. A solution oncethoroughly mixed remains uniform. ]

[Note 2: Commercial sodium hydroxide is usually impure and alwayscontains more or less carbonate; an allowance is therefore made forthis impurity by placing the weight taken at 23 grams per liter. Ifthe hydroxide is known to be pure, a lesser amount (say 21 grams) willsuffice. ]

COMPARISON OF ACID AND ALKALI SOLUTIONS

PROCEDURE. --Rinse a previously calibrated burette three times with thehydrochloric acid solution, using 10 cc. Each time, and allowing theliquid to run out through the tip to displace all water and airfrom that part of the burette. Then fill the burette with the acidsolution. Carry out the same procedure with a second burette, usingthe sodium hydroxide solution. 

The acid solution may be placed in a plain or in a glass-stopperedburette as may be more convenient, but the alkaline solution shouldnever be allowed to remain long in a glass-stoppered burette, as ittends to cement the stopper to the burette, rendering it useless. Itis preferable to use a plain burette for this solution. 

When the burettes are ready for use and all air bubbles displaced fromthe tip (see Note 2, page 17) note the exact position of the liquid ineach, and record the readings in the notebook. (Consult page 188. ) Runout from the burette into a beaker about 40 cc. Of the acid and addtwo drops of a solution of methyl orange; dilute the acid to about80 cc. And run out alkali solution from the other burette, stirringconstantly, until the pink has given place to a yellow. Wash down thesides of the beaker with a little distilled water if the solution hasspattered upon them, return the beaker to the acid burette, and addacid to restore the pink; continue these alternations until the pointis accurately fixed at which a single drop of either solutions servedto produce a distinct change of color. Select as the final end-pointthe appearance of the faintest pink tinge which can be recognized, orthe disappearance of this tinge, leaving a pure yellow; but alwaystitrate to the same point (Note 1). If the titration has occupied morethan the three minutes required for draining the sides of the burette, the final reading may be taken immediately and recorded in thenotebook. 

Refill the burettes and repeat the titration. From the records ofcalibration already obtained, correct the burette readings and makecorrections for temperature, if necessary. Obtain the ratio of thesodium hydroxide solution to that of hydrochloric acid by dividingthe number of cubic centimeters of acid used by the number of cubiccentimeters of alkali required for neutralization. The check resultsof the two titrations should not vary by more than two parts in onethousand (Note 2). If the variation in results is greater than this, refill the burettes and repeat the titration until satisfactory valuesare obtained. Use a new page in the notebook for each titration. Inaccurate values should not be erased or discarded. They should beretained and marked "correct" or "incorrect, " as indicated by thefinal outcome of the titrations. This custom should be rigidlyfollowed in all analytical work. 

[Note 1: The end-point should be chosen exactly at the point ofchange; any darker tint is unsatisfactory, since it is impossible tocarry shades of color in the memory and to duplicate them from day today. ]

[Note 2: While variation of two parts in one thousand in the valuesobtained by an inexperienced analyst is not excessive, the idea mustbe carefully avoided that this is a standard for accurate work to be!generally applied!. In many cases, after experience is gained, theallowable error is less than this proportion. In a few cases alarger variation is permissible, but these are rare and can onlybe recognized by an experienced analyst. It is essential that thebeginner should acquire at least the degree of accuracy indicated ifhe 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 acidsolutions 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 beststandard is sodium carbonate prepared from sodium bicarbonate byheating the latter at temperature between 270° and 300°C. Thebicarbonate is easily prepared in a pure state, and at thetemperatures named the decomposition takes place according to theequation

2HNaCO_{3} --> Na_{2}CO_{3} + H_{2}O + CO_{2}

and without loss of any carbon dioxide from the sodium carbonate, suchas may occur at higher temperatures. The process is carried out asdescribed below. 

PROCEDURE. --Place in a porcelain crucible about 6 grams (roughlyweighed) of the purest sodium bicarbonate obtainable. Rest thecrucible upon a triangle of iron or copper wire so placed within alarge crucible that there is an open air space of about three eighthsof an inch between them. The larger crucible may be of iron, nickel orporcelain, as may be most convenient. Insert the bulb of a thermometerreading to 350°C. In the bicarbonate, supporting it with a clamp sothat the bulb does not rest on the bottom of the crucible. Heatthe outside crucible, using a rather small flame, and raise thetemperature of the bicarbonate fairly rapidly to 270°C. Then regulatethe heat in such a way that the temperature rises !slowly! to 300°C. In the course of a half-hour. The bicarbonate should be frequentlystirred with a clean, dry, glass rod, and after stirring, should beheaped up around the bulb of the thermometer in such a way as to coverit. This will require attention during most of the heating, as thetemperature should not be permitted to rise above 310°C. For anylength of time. At the end of the half-hour remove the thermometer andtransfer the porcelain crucible, which now contains sodium carbonate, to a desiccator. When it is cold, transfer the carbonate to astoppered weighing tube or weighing-bottle. 

STANDARDIZATION

PROCEDURE. --Clean carefully the outside of a weighing-tube, orweighing-bottle, containing the pure sodium carbonate, taking careto handle it as little as possible after wiping. Weigh the tubeaccurately to 0. 0001 gram, and record the weight in the notebook. Holdthe tube over the top of a beaker (200-300 cc. ) and cautiously removethe stopper, making sure that no particles fall from it or from thetube elsewhere than in the beaker. Pour out from the tube a portionof the carbonate, replace the stopper and determine approximately howmuch has been removed. Continue this procedure until 1. 00 to 1. 10grams has been taken from the tube. Then weigh the tube accuratelyand record the weight under the first weight in the notebook. The difference in the two weights is the weight of the carbonatetransferred to the beaker. Proceed in the same way to transfer asecond portion of the carbonate from the tube to another beaker ofabout the same size as the first. The beakers should be labeled andplainly marked to correspond with the entries in the notebook. 

Pour over the carbonate in each beaker about 80 cc. Of water, stiruntil solution is complete, and add two drops of methyl orangesolution. Fill the burettes with the standard acid and alkalisolutions, noting the initial readings of the burettes and temperatureof the solutions. Run in acid from the burette, stirring and avoidingloss by effervescence, until the solution has become pink. Wash downthe sides of the beaker with a !little! water from a wash-bottle, andthen run in alkali from the other burette until the pink is replacedby yellow; then finish the titration as described on page 37. Note thereadings of the burettes after the proper interval, and record them inthe notebook. Repeat the procedure, using the second portion of sodiumcarbonate. Apply the necessary calibration corrections to the volumesof the solutions used, and correct for temperature if necessary. 

From the data obtained, calculate the volume of the hydrochloricacid solution which is equivalent to the volume of sodium hydroxidesolution used in this titration. Subtract this volume from the volumeof hydrochloric acid. The difference represents the volume of acidused to react with the sodium carbonate. Divide the weight of sodiumcarbonate by this volume in cubic centimeters, thus obtaining theweight of sodium carbonate equivalent to each cubic centimeter of theacid. 

From this weight it is possible to calculate the corresponding weightof HCl in each cubic centimeter of the acid, and in turn the relationof the acid to the normal. 

If, however, it is recalled that normal solutions are equivalent toeach other, it will be seen that the same result may be more readilyreached by dividing the weight in grams of sodium carbonate per cubiccentimeter just found by titration by the weight which would becontained 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 acidsolution to the normal is, therefore, calculated by dividing theweight of the carbonate to which each cubic centimeter of the acid isequivalent by 0. 0530. The standardization must be repeated until thevalues 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 sodiumhydroxide solution to a normal solution (Notes 1 and 2). 

[Note 1: In the foregoing procedure the acid solution is standardizedand the alkali solution referred to this standard by calculation. Itis equally possible, if preferred, to standardize the alkali solution. The standards in a common use for this purpose are purifiedoxalic 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 acidand the oxalates should be specially prepared to insure purity, the main difficulty lying in the preservation of the water ofcrystallization. 

It should be noted that the acid oxalate and the acid tartrate eachcontain one hydrogen atom replaceable by a base, while the tetroxalatecontains three such atoms and the oxalic acid two. Each of the twosalts first named behave, therefore, as monobasic acids, and thetetroxalate as a tribasic acid. ]

[Note 2: It is also possible to standardize a hydrochloric acidsolution by precipitating the chloride ions as silver chloride andweighing the precipitate, as prescribed under the analysis of sodiumchloride to be described later. Sulphuric acid solutions may bestandardized by precipitation of the sulphate ions as barium sulphateand weighing the ignited precipitate, but the results are not abovecriticism on account of the difficulty in obtaining large precipitatesof barium sulphate which are uncontaminated by inclosures or are notreduced on ignition. ]

DETERMINATION OF THE TOTAL ALKALINE STRENGTH OF SODA ASH

Soda ash is crude sodium carbonate. If made by the ammonia process itmay contain also sodium chloride, sulphate, and hydroxide; when madeby the Le Blanc process it may contain sodium sulphide, silicate, andaluminate, 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 sodiumcarbonate; i. E. , as though no other alkali were present. 

PROCEDURE. --In order to secure a sample which shall represent theaverage value of the ash, it is well to take at least 5 grams. As thisis too large a quantity for convenient titration, an aliquot portionof the solution is measured off, representing one fifth of the entirequantity. This is accomplished as follows: Weigh out on an analyticalbalance two samples of soda ash of about 5 grams each into beakersof about 500 cc. Capacity. (The weighings need be made to centigramsonly. ) Dissolve the ash in 75 cc. Of water, warming gently, and filteroff the insoluble residue; wash the filter by filling it at leastthree times with distilled water, and allowing it to drain, adding thewashings to the main filtrate. Cool the filtrate to approximately thestandard temperature of the laboratory, and transfer it to a 250 cc. Measuring flask, washing out the beaker thoroughly. Add distilledwater of laboratory temperature until the lowest point of the meniscusis level with the graduation on the neck of the flask and remove anydrops of water that may be on the neck above the graduation by meansof a strip of filter paper; make the solution thoroughly uniform bypouring it out into a dry beaker and back into the flask severaltimes. Measure off 50 cc. Of the solution in a measuring flask, orpipette, either of which before use should, unless they are dry on theinside, be rinsed out with at least two small portions of the soda ashsolution to displace any water. 

If a flask is used, fill it to the graduation with the soda ashsolution and remove any liquid from the neck above the graduation withfilter paper. Empty it into a beaker, and wash out the small flask, unless it is graduated for !delivery!, using small quantities ofwater, which are added to the liquid in the beaker. A second 50 cc. Portion from the main solution should be measured off into a secondbeaker. Dilute the solutions in each beaker to 100 cc. , add two dropsof a solution of methyl orange (Note 1) and titrate for the alkaliwith the standard hydrochloric acid solution, using the alkalisolution to complete the titration as already prescribed. 

From the volumes of acid and alkali employed, corrected for buretteerrors and temperature changes, and the data derived from thestandardization, calculate the percentage of alkali present, assumingit all to be present as sodium carbonate (Note 2). 

[Note 1: The hydrochloric acid sets free carbonic acid which isunstable and breaks down into water and carbon dioxide, most of whichescapes from the solution. Carbonic acid is a weak acid and, as such, does not yield a sufficient concentration of H^{+} ions to cause theindicator 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 sodaash may be determined by precipitating the carbonate by the additionof barium chloride, removing the barium carbonate by filtration, andtitrating the alkali in the filtrate. 

The caustic alkali may also be determined by first usingphenolphthalein as an indicator, which will show by its change frompink to colorless the point at which the caustic alkali has beenneutralized and the carbonate has been converted to bicarbonate, andthen adding methyl orange and completing the titration. The amount ofacid necessary to change the methyl orange to pink is a measure of onehalf of the carbonate present. The results of the double titrationfurnish the data necessary for the determination of the caustic alkaliand 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 grameach. Dissolve these in 50 cc. Of warm water. Add two drops ofphenolphthalein solution, and run in alkali from the burette until thesolution is pink; add acid from the other burette until the pink isjust destroyed, and then add 0. 3 cc. (not more) in excess. Heat thesolution to boiling for three minutes. If the pink returns during theboiling, discharge it with acid and again add 0. 3 cc. In excess andrepeat the boiling (Note 1). If the color does not then reappear, addalkali until it does, and a !drop or two! of acid in excess and boilagain for one minute (Note 2). If no color reappears during this time, complete the titration in the hot solution. The end-point should bethe faintest visible shade of color (or its disappearance), as thesame difficulty would exist here as with methyl orange if an attemptwere made to match shades of pink. 

From the corrected volume of alkali required to react with theoxalic 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 thestandard solution was made contains some sodium carbonate. This reactswith 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 ofthe indicator is discharged in cold solutions at the point at whichbicarbonate is formed. If, however, the solution is heated to boiling, the bicarbonate loses carbon dioxide and water, and reverts to sodiumcarbonate, 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, thecarbonate is ultimately all brought into reaction. 

The student should make sure that the difference in behavior of thetwo 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 procedureare strictly followed, no loss of acid need be feared, but the amountadded in excess should not be greater than 0. 3-0. 4 cc. ]

[Note 3: Attention has already been called to the fact that the colorchanges in the different indicators occur at varying concentrationsof H^{+} or OH^{-} ions. They do not indicate exact theoreticalneutrality, but a particular indicator always shows its color changeat a particular concentration of H^{+} or OH^{-} ions. The resultsof titration with a given indicator are, therefore, comparable. As amatter of fact, a small error is involved in the procedure as outlinedabove. The comparison of the acid and alkali solutions was made, usingmethyl orange as an indicator, while the titration of the oxalic acidis made with the use of phenolphthalein. For our present purposes thesmall error may be neglected but, if time permits, the student isrecommended to standardize the alkali solution against one of thesubstances named in Note 1, page 41, and also to ascertainthe comparative value of the acid and alkali solutions, usingphenolphthalein as indicator throughout, and conducting the titrationsas described above. This will insure complete accuracy. ]

II. OXIDATION PROCESSES

GENERAL DISCUSSION

In the oxidation processes of volumetric analysis standard solutionsof oxidizing agents and of reducing agents take the place of the acidand alkali solutions of the neutralization processes already studied. Just as an acid solution was the principal reagent in alkalimetry, andthe alkali solution used only to make certain of the end-point, thesolution of the oxidizing agent is the principal reagent for thetitration 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 arepotassium bichromate, potassium permangenate, potassium ferricyanide, iodine, ferric chloride, and sodium hypochlorite. 

The important reducing agents which are used in the form of standardsolutions are ferrous sulphate (or ferrous ammonium sulphate), oxalicacid, sodium thiosulphate, stannous chloride, arsenious acid, andpotassium cyanide. Other reducing agents, as sulphurous acid, sulphureted hydrogen, and zinc (nascent hydrogen), may take part inthe processes, but not as standard solutions. 

The most important combinations among the foregoing are: Potassiumbichromate and ferrous salts; potassium permanganate and ferroussalts; 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 andchromic compounds which are formed. 

The acid may be either hydrochloric or sulphuric, but the former isusually preferred, since it is by far the best solvent for iron andits compounds. The reaction in the presence of hydrochloric acid is asfollows:

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 uponthe equivalence of the reagents which they contain to 8 grams ofoxygen or 1 gram of hydrogen. A normal solution of an oxidizing agentshould, therefore, contain that amount per liter which is equivalentin oxidizing power to 8 grams of oxygen; a normal reducing solutionmust be equivalent in reducing power to 1 gram of hydrogen. In orderto determine what the amount per liter will be it is necessary to knowhow the reagents enter into reaction. The two solutions to be employedin the process under consideration are those of potassium bichromateand ferrous sulphate. The reaction between them, in the presence of anexcess 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 nowconcerned, be written in such a way as to show the oxides of theseelements in each, they would appear as follows: On the left-hand sideof the equation 6(FeO. SO_{3}) and K_{2}O. 2CrO_{3}; on the right-handside, 3(Fe_{2}O_{3}. 3SO_{3}) and Cr_{2}O_{3}. 3SO_{3}. A carefulinspection shows that there are three less oxygen atoms associatedwith chromium atoms on the right-hand side of the equation than on theleft-hand, but there are three more oxygen atoms associated with ironatoms on the right than on the left. In other words, a molecule ofpotassium bichromate has given up three atoms of oxygen for oxidationpurposes; 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, thenormal solution of potassium bichromate should contain per liter onesixth of 294. 2 grams or 49. 03 grams. 

A further inspection of the dissected compounds above shows that sixmolecules of FeO. SO_{3} were required to react with the three atoms ofoxygen from the bichromate. From the two equations

3H_{2} + 3O --> 3H_{2}O6(FeO. SO_{3}) + 3O --> 3(Fe_{2}O_{3}. 3SO_{3})

it is plain that one molecule of ferrous sulphate is equivalent to oneatom of hydrogen in reducing power; therefore one molecular weight ingrams of ferrous sulphate (151. 9) is equivalent to 1 gram ofhydrogen. Since the ferrous sulphate crystalline form has the formulaFeSO_{4}. 7H_{2}O, a normal reducing solution of this crystalline saltshould contain 277. 9 grams per liter. 

PREPARATION OF SOLUTIONS

!Approximate Strength 0. 1 N!

It is possible to purify commercial potassium bichromate byrecrystallization from hot water. It must then be dried and cautiouslyheated to fusion to expel the last traces of moisture, but notsufficiently high to expel any oxygen. The pure salt thus prepared, may be weighed out directly, dissolved, and the solution diluted in agraduated flask to a definite volume. In this case no standardizationis made, as the normal value can be calculated directly. It is, however, more generally customary to standardize a solution ofthe commercial salt by comparison with some substance of definitecomposition, as described below. 

PROCEDURE. --Pulverize about 5 grams of potassium bichromate of goodquality. Dissolve the bichromate in distilled water, transfer thesolution to a liter bottle, and dilute to approximately 1000 cc. Shakethoroughly until the solution is uniform. 

To prepare the solution of the reducing agent, pulverize about 28grams of ferrous sulphate (FeSO_{4}. 7H_{2}O) or about 40 grams offerrous ammonium sulphate (FeSO_{4}. (NH_{4})_{2}SO_{4}. 6H_{2}O) anddissolve in distilled water containing 5 cc. Of concentrated sulphuricacid. Transfer the solution to a liter bottle, add 5 cc. Concentratedsulphuric acid, make up to about 1000 cc. And shake vigorously toinsure uniformity. 

INDICATOR SOLUTION

No indicator is known which, like methyl orange, can be used withinthe solution, to show when the oxidation process is complete. Instead, an outside indicator solution is employed to which drops of thetitrated solution are transferred for testing. The reagent used ispotassium ferricyanide, which produces a blue precipitate (or color)with ferrous compounds as long as there are unoxidized ferrous ions inthe titrated solution. Drops of the indicator solution are placed upona glazed porcelain tile, or upon white cardboard which has been coatedwith paraffin to render it waterproof, and drops of the titratedsolution are transferred to the indicator on the end of a stirringrod. When the oxidation is nearly completed only very small amountsof the ferrous compounds remain unoxidized and the reaction with theindicator is no longer instantaneous. It is necessary to allow a brieftime to elapse before determining that no blue color is formed. Thirtyseconds is a sufficient interval, and should be adopted throughout theanalytical procedure. If left too long, the combined effect of lightand dust from the air will cause a reduction of the ferric compoundsalready formed and a resultant blue will appear which misleads theobserver with respect to the true end-point. 

The indicator solution must be highly diluted, otherwise its own colorinterferes with accurate observation. Prepare a fresh solution, asneeded each day, by dissolving a crystal of potassium ferricyanideabout the size of a pin's head in 25 cc. Of distilled water. The saltshould be carefully tested with ferric chloride for the presence offerrocyanides, which give a blue color with ferric salts. 

In case of need, the ferricyanide can be purified by adding to itssolution a little bromine water and recrystallizing the compound. 

COMPARISON OF OXIDIZING AND REDUCING SOLUTIONS

PROCEDURE. --Fill one burette with each of the solutions, observingthe general procedure with respect to cleaning and rinsing alreadyprescribed. The bichromate solution is preferably to be placed in aglass-stoppered burette. 

Run out from a burette into a beaker of about 300 cc. Capacity nearly40 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 bichromatesolution from another burette. Since both solutions are approximatelytenth-normal, 35 cc. Of the bichromate solution may be added withouttesting. Test at that point by removing a very small drop of theiron solution on the end of a stirring rod, mixing it with a drop ofindicator 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 indistilled water before returning it to the iron solution. As soon asthe blue appears to be less intense, add the bichromate solution insmall portions, finally a single drop at a time, until the point isreached at which no blue color appears after the lapse of thirtyseconds from the time of mixing solution and indicator. At the closeof the titration a large drop of the iron solution should be taken forthe test. To determine the end-point beyond any question, as soon asthe thirty seconds have elapsed remove another drop of the solutionof the same size as that last taken and mix it with the indicator, placing it beside the last previous test. If this last previous testshows a blue tint in comparison with the fresh mixture, the end-pointhas not been reached; if no difference can be noted the reaction iscomplete. Should the end-point be overstepped, a little more of theferrous solution may be added and the end-point definitely fixed. 

From the volumes of the solutions used, after applying corrections forburette readings, and, if need be, for the temperature of solutions, calculate the value of the ferrous solution in terms of the oxidizingsolution. 

[Note 1: The accuracy of the work may be much impaired by the removalof unnecessarily large quantities of solution for the tests. At thebeginning of the titration, while much ferrous iron is still present, the end of the stirring rod need only be moist with the solution; butat the close of the titration drops of considerable size may properlybe taken for the final tests. The stirring rod should be washed toprevent transfer of indicator to the main solution. This cautiousremoval of solution does not seriously affect the accuracy of thedetermination, as it will be noted that the volume of the titratedsolution is about 200 cc. And the portions removed are verysmall. Moreover, if the procedure is followed as prescribed, theconcentration of unoxidized iron decreases very rapidly as thetitration is carried out so that when the final tests are made, thoughlarge drops may be taken, the amount of ferrous iron is not sufficientto produce any appreciable error in results. 

If the end-point is determined as prescribed, it can be as accuratelyfixed as that of other methods; and if a ferrous solution is athand, the titration need consume hardly more time than that of thepermanganate process to be described later on. ]

STANDARDIZATION OF POTASSIUM BICHROMATE SOLUTIONS

!Selection of a Standard!

A substance which will serve satisfactorily as a standard foroxidizing solutions must possess certain specific properties: It mustbe of accurately known composition and definite in its behavior as areducing agent, and it must be permanent against oxidation in the air, at least for considerable periods. Such standards may take the form ofpure crystalline salts, such as ferrous ammonium sulphate, or may bein the form of iron wire or an iron ore of known iron content. It isnot necessary that the standard should be of 100 per cent purity, provided the content of the active reducing agent is known and nointerfering substances are present. 

The two substances most commonly used as standards for a bichromatesolution are ferrous ammonium sulphate and iron wire. A standard wireis to be purchased in the market which answers the purpose well, andits iron content may be determined for each lot purchased by a numberof gravimetric determinations. It may best be preserved in jarscontaining calcium chloride, but this must not be allowed to comeinto contact with the wire. It should, however, even then be examinedcarefully for rust before use. 

If pure ferrous ammonium sulphate is used as the standard, clearcrystals only should be selected. It is perhaps even better todetermine by gravimetric methods once for all the iron content of alarge commercial sample which has been ground and well mixed. Thissalt is permanent over long periods if kept in stoppered containers. 

STANDARDIZATION

PROCEDURE. --Weigh out two portions of iron wire of about 0. 24-0. 26gram each, examining the wire carefully for rust. It should be handledand wiped with filter paper (not touched by the fingers), shouldbe weighed on a watch-glass, and be bent in such a way as not tointerfere with the movement of the balance. 

Place 30 cc. Of hydrochloric acid (sp. Gr. 1. 12) in each of two 300cc. Erlenmeyer flasks, cover them with watch-glasses, and bring theacid just to boiling. Remove them from the flame and drop in theportions of wire, taking great care to avoid loss of liquid duringsolution. Boil for two or three minutes, keeping the flasks covered(Note 1), then wash the sides of the flasks and the watch-glass witha little water and add stannous chloride solution to the hot liquid!from a dropper! until the solution is colorless, but avoid more thana drop or two in excess (Note 2). Dilute with 150 cc. Of water andcool !completely!. When cold, add rapidly about 30 cc. Of mercuricchloride solution. Allow the solutions to stand about three minutesand then titrate without further delay (Note 3), add about 35 cc. Ofthe standard solution at once and finish the titration as prescribedabove, making use of the ferrous solution if the end-point should bepassed. 

From the corrected volumes of the bichromate solution required tooxidize the iron actually know to be present in the wire, calculatethe relation of the standard solution to the normal. 

Repeat the standardization until the results are concordant within atleast two parts in one thousand. 

[Note 1: The hydrochloric acid is added to the ferrous solutionto insure the presence of at least sufficient free acid for thetitration, as required by the equation on page 48. 

The solution of the wire in hot acid and the short boiling insure theremoval of compounds of hydrogen and carbon which are formed from thesmall amount of carbon in the iron. These might be acted upon by thebichromate if not expelled. ]

[Note 2: It is plain that all the iron must be reduced to the ferrouscondition before the titration begins, as some oxidation may haveoccurred from the oxygen of the air during solution. It is alsoevident that any excess of the agent used to reduce the iron must beremoved; otherwise it will react with the bichromate added later. 

The reagents available for the reduction of iron are stannouschloride, sulphurous acid, sulphureted hydrogen, and zinc; of thesestannous chloride acts most readily, the completion of the reactionis most easily noted, and the excess of the reagent is most readilyremoved. The latter object is accomplished by oxidation to stannicchloride by means of mercuric chloride added in excess, as themercuric salts have no effect upon ferrous iron or the bichromate. Thereactions 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 stannouschloride should not be present in great excess, otherwise a secondaryreaction takes place, resulting in the reduction of the mercurouschloride to metallic mercury:

SnCl_{2} + 2HgCl --> SnCl_{4} + 2Hg. 

The occurrence of this secondary reaction is indicated by thedarkening of the precipitate; and, since potassium bichromate oxidizesthis mercury slowly, solutions in which it has been precipitated areworthless as iron determinations. ]

[Note 3: The solution should be allowed to stand about three minutesafter the addition of mercuric chloride to permit the completedeposition of mercurous chloride. It should then be titrated withoutdelay to avoid possible reoxidation of the iron by the oxygen of theair. ]

DETERMINATION OF IRON IN LIMONITE

PROCEDURE. --Grind the mineral (Note 1) to a fine powder. Weigh outaccurately two portions of about 0. 5 gram (Note 2) into porcelaincrucibles; heat these crucibles to dull redness for ten minutes, allow them to cool, and place them, with their contents, in beakerscontaining 30 cc. Of dilute hydrochloric acid (sp. Gr. 1. 12). Heatat a temperature just below boiling until the undissolved residue iswhite 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 beremoved by filtration. If a dark residue remains, collect it on afilter, wash free from hydrochloric acid, and ignite the filter in aplatinum crucible (Note 3). Mix the ash with five times its weight ofsodium carbonate and heat to fusion; cool, and disintegrate the fusedmass with boiling water in the crucible. Unite this solution andprecipitate (if any) with the acid solution, taking care to avoid lossby effervescence. Wash out the crucible, heat the acid solutionto boiling, add stannous chloride solution until it is colorless, avoiding a large excess (Note 4); cool, and when !cold!, add 40 cc. Ofmercuric chloride solution, dilute to 200 cc. , and proceed with thetitration as already described. 

From the standardization data already obtained, and the known weightof the sample, calculate the percentage of iron (Fe) in the limonite. 

[Note 1: Limonite is selected as a representative of iron ores ingeneral. It is a native, hydrated oxide of iron. It frequently occursin or near peat beds and contains more or less organic matter which, if brought into solution, would be acted upon by the potassiumbichromate. This organic matter is destroyed by roasting. Since a hightemperature tends to lessen the solubility of ferric oxide, the heatshould not be raised above low redness. ]

[Note 2: It is sometimes advantageous to dissolve a large portion--say5 grams--and to take one tenth of it for titration. The sample willthen represent more closely the average value of the ore. ]

[Note 3: A platinum crucible may be used for the roasting of thelimonite and must be used for the fusion of the residue. When used, itmust not be allowed to remain in the acid solution of ferric chloridefor 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, thusintroducing an error. It should also be noted that copper and antimonyinterfere with the determination of iron by the bichromate process. ]

[Note 4: The quantity of stannous chloride required for the reductionof the iron in the limonite will be much larger than that added to thesolution of iron wire, in which the iron was mainly already in theferrous condition. It should, however, be added from a dropper toavoid an unnecessary excess. ]

DETERMINATION OF CHROMIUM IN CHROME IRON ORE

PROCEDURE. --Grind the chrome iron ore (Note 1) in an agate mortaruntil no grit is perceptible under the pestle. Weigh out two portionsof 0. 5 gram each into iron crucibles which have been scoured insideuntil bright (Note 2). Weigh out on a watch-glass (Note 3), using therough balances, 5 grams of dry sodium peroxide for each portion, andpour about three quarters of the peroxide upon the ore. Mix ore andflux by thorough stirring with a dry glass rod. Then cover the mixturewith the remainder of the peroxide. Place the crucible on a triangleand raise the temperature !slowly! to the melting point of the flux, using a low flame, and holding the lamp in the hand (Note 4). Maintainthe fusion for five minutes, and stir constantly with a stout ironwire, but do not raise the temperature above moderate redness (Notes 5and 6). 

Allow the crucible to cool until it can be comfortably handled (Note7) and then place it in a 300 cc. Beaker, and cover it with distilledwater (Note 8). The beaker must be carefully covered to avoid lossduring the disintegration of the fused mass. When the evolution ofgas ceases, rinse off and remove the crucible; then heat the solution!while still alkaline! to boiling for fifteen minutes. Allow theliquid to cool for a few minutes; then acidify with dilute sulphuricacid (1:5), adding 10 cc. In excess of the amount necessary todissolve the ferric hydroxide (Note 9). Dilute to 200 cc. , cool, addfrom a burette an excess of a standard ferrous solution, and titratefor the excess with a standard solution of potassium bichromate, usingthe outside indicator (Note 10). 

From the corrected volumes of the two standard solutions, and theirrelations to normal solutions, calculate the percentage of chromium inthe ore. 

[Note 1: Chrome iron ore is essentially a ferrous chromite, orcombination of FeO and Cr_{2}O_{3}. It must be reduced to a state offine subdivision to ensure a prompt reaction with the flux. ]

[Note 2: The scouring of the iron crucible is rendered much easier ifit is first heated to bright redness and plunged into cold water. Inthis process oily matter is burned off and adhering scale is caused tochip off when the hot crucible contracts rapidly in the cold water. ]

[Note 3: Sodium peroxide must be kept off of balance pans and shouldnot be weighed out on paper, as is the usual practice in the roughweighing of chemicals. If paper to which the peroxide is adhering isexposed to moist air it is likely to take fire as a result ofthe absorption of moisture, and consequent evolution of heat andliberation of oxygen. ]

[Note 4: The lamp should never be allowed to remain under thecrucible, as this will raise the temperature to a point at which thecrucible itself is rapidly attacked by the flux and burned through. ]

[Note 5: The sodium peroxide acts as both a flux and an oxidizingagent. The chromic oxide is dissolved by the flux and oxidized tochromic anhydride (CrO_{3}) which combines with the alkali to formsodium chromate. The iron is oxidized to ferric oxide. ]

[Note 6: The sodium peroxide cannot be used in porcelain, platinum, orsilver crucibles. It attacks iron and nickel as well; but cruciblesmade from these metals may be used if care is exercised to keep thetemperature as low as possible. Preference is here given to ironcrucibles, because the resulting ferric hydroxide is more readilybrought into solution than the nickelic oxide from a nickel crucible. The peroxide must be dry, and must be protected from any admixture ofdust, paper, or of organic matter of any kind, otherwise explosionsmay ensue. ]

[Note 7: When an iron crucible is employed it is desirable to allowthe fusion to become nearly cold before it is placed in water, otherwise scales of magnetic iron oxide may separate from thecrucible, 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 peroxideis decomposed with the evolution of oxygen. The subsequent boilinginsures the complete decomposition of the peroxide. Unless this iscomplete, hydrogen peroxide is formed when the solution is acidified, and this reacts with the bichromate, reducing it and introducing aserious error. ]

[Note 9: The addition of the sulphuric acid converts the sodiumchromate to bichromate, which behaves exactly like potassiumbichromate in acid solution. ]

[Note 10: If a standard solution of a ferrous salt is not at hand, aweight of iron wire somewhat in excess of the amount which would berequired if the chromite were pure FeO. Cr_{2}O_{3} may be weighed outand dissolved in sulphuric acid; after reduction of all the iron bystannous chloride and the addition of mercuric chloride, this solutionmay be poured into the chromate solution and the excess of irondetermined by titration with standard bichromate solution. ]

PERMANGANATE PROCESS FOR THE DETERMINATION OF IRON

Potassium permanganate oxidizes ferrous salts in cold, acid solutionpromptly and completely to the ferric condition, while in hot acidsolution it also enters into a definite reaction with oxalic acid, bywhich 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 ofpotassium permanganate depends; but besides iron and oxalic acid thepermanganate enters into reaction with antimony, tin, copper, mercury, and manganese (the latter only in neutral solution), by which thesemetals are changed from a lower to a higher state of oxidation; and italso 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 quantitativelywith the permanganate, as is the case with oxalic acid and theoxalates. 

Potassium permanganate is acted upon by hydrochloric acid; the actionis rapid in hot or concentrated solution (particularly in the presenceof iron salts, which appear to act as catalyzers, increasing thevelocity of the reaction), but slow in cold, dilute solutions. However, the greater solubility of iron compounds in hydrochloric acidmakes it desirable to use this acid as a solvent, and experiments madewith this end in view have shown that in cold, dilute hydrochloricacid solution, to which considerable quantities of manganous sulphateand an excess of phosphoric acid have been added, it is possible toobtain satisfactory results. 

It is also possible to replace the hydrochloric acid by evaporatingthe solutions with an excess of sulphuric acid until the latter fumes. This procedure is somewhat more time-consuming, but the end-point ofthe permanganate titration is more permanent. Both procedures aredescribed below. 

Potassium permanganate has an intense coloring power, and since thesolution resulting from the oxidation of the iron and the reduction ofthe permanganate is colorless, the latter becomes its own indicator. The slightest excess is indicated with great accuracy by the pinkcolor of the solution. 

PREPARATION OF A STANDARD SOLUTION

!Approximate Strength 0. 1 N!

A study of the reactions given above which represent the oxidation offerrous compounds by potassium permanganate, shows that there are 2molecules of KMnO_{4} and 10 molecules of FeSO_{4} on theleft-hand side, and 2 molecules of MnSO_{4} and 5 molecules ofFe_{2}(SO_{4})_{5} on the right-hand side. Considering only thesecompounds, and writing the formulas in such a way as to show theoxides 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 theferrous compound. Since 8 grams of oxygen is the basis of normaloxidizing solutions and 80 grams of oxygen are supplied by 316. 0 gramsof 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 thereagent may be carried out as follows:

PROCEDURE. --Dissolve about 3. 25 grams of potassium permanganatecrystals in approximately 1000 cc. Of distilled water in a largebeaker, or casserole. Heat slowly and when the crystals havedissolved, boil the solution for 10-15 minutes. Cover the solutionwith a watch-glass; allow it to stand until cool, or preferably overnight. Filter the solution through a layer of asbestos. Transfer thefiltrate to a liter bottle and mix thoroughly (Note 2). 

[Note 1: The reactions given on page 61 are those which take place inthe presence of an excess of acid. In neutral solutions the reductionof the permanganate is less complete, and, under these conditions, two gram-molecular weights of KMnO_{4} will furnish only 48 gramsof 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 forlong periods, and change more rapidly when first prepared than afterstanding some days. This change is probably caused by interactionwith the organic matter contained in all distilled water, except thatredistilled from an alkaline permanganate solution. The solutionsshould be protected from light and heat as far as possible, since bothinduce decomposition with a deposition of manganese dioxide, and ithas been shown that decomposition proceeds with considerable rapidity, with the evolution of oxygen, after the dioxide has begun to form. Ascommercial samples of the permanganate are likely to be contaminatedby the dioxide, it is advisable to boil and filter solutions throughasbestos before standardization, as prescribed above. Such solutionsare relatively stable. ]

COMPARISON OF PERMANGANATE AND FERROUS SOLUTIONS

PROCEDURE. --Fill a glass-stoppered burette with the permanganatesolution, observing the usual precautions, and fill a second burettewith the ferrous sulphate solution prepared for use with the potassiumbichromate. The permanganate solution cannot be used in burettes withrubber tips, as a reduction takes place upon contact with the rubber. The solution has so deep a color that the lower line of the meniscuscannot be detected; readings must therefore be made from the upperedge. 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 runin the permanganate solution to a slight permanent pink. Repeat, untilthe 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 admitof its direct weighing as a standard. On this account, and becauseof the uncertainties as to the permanence of its solutions, it isadvisable to standardize them against substances of known value. Thosein most common use are iron wire, ferrous ammonium sulphate, sodiumoxalate, oxalic acid, and some other derivatives of oxalic acid. With the exception of sodium oxalate, these all contain water ofcrystallization which may be lost on standing. They should, therefore, be freshly prepared, and with great care. At present, sodium oxalateis considered to be one of the most satisfactory standards. 

!Method A!

!Iron Standards!

The standardization processes employed when iron or its compounds areselected as standards differ from those applicable in connection withoxalate standards. The procedure which immediately follows is that inuse with iron standards. 

As in the case of the bichromate process, it is necessary to reducethe iron completely to the ferrous condition before titration. Thereducing agents available are zinc, sulphurous acid, or sulphuretedhydrogen. Stannous chloride may also be used when the titration ismade in the presence of hydrochloric acid. Since the excess of boththe gaseous reducing agents can only be expelled by boiling, withconsequent uncertainty regarding both the removal of the excess andthe reoxidation of the iron, zinc or stannous chlorides are the mostsatisfactory agents. For prompt and complete reduction it is essentialthat the iron solution should be brought into ultimate contact withthe zinc. This is brought about by the use of a modified Jonesreductor, as shown in Figure 1. This reductor is a standard apparatusand is used in other quantitative processes. 

[Illustration: Fig. 1]

The tube A has an inside diameter of 18 mm. And is 300 mm. Long; thesmall tube has an inside diameter of 6 mm. And extends 100 mm. Belowthe stopcock. At the base of the tube A are placed some pieces ofbroken 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 isused for Gooch filters, 1 mm. Thick. The tube is then filled with theamalgamated zinc (Note 1) to within 50 mm. Of the top, and on the zincis placed a plug of glass wool. If the top of the tube is not alreadyshaped like the mouth of a thistle-tube (B), a 60 mm. Funnel is fittedinto the tube with a rubber stopper and the reductor is connectedwith a suction bottle, F. The bottle D is a safety bottle toprevent contamination of the solution by water from the pump. Afterpreparation for use, or when left standing, the tube A should befilled with water, to prevent clogging of the zinc. 

[Note 1: The use of fine zinc in the reductor is not necessary andtends to clog the tube. Particles which will pass a 10-mesh sieve, butare 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 of25 cc. Of normal mercuric chloride solution, 25 cc. Of hydrochloricacid (sp. Gr. 1. 12) and 250 cc. Of water for two minutes. The solutionshould then be poured off and the zinc thoroughly washed. It is thenready for bottling and preservation under water. A small quantity ofglass wool is placed in the neck of the funnel to hold back foreignmaterial when the reductor is in use. ]

STANDARDIZATION

PROCEDURE. --Weigh out into Erlenmeyer flasks two portions of iron wireof about 0. 25 gram each. Dissolve these in hot dilute sulphuric acid(5 cc. Of concentrated acid and 100 cc. Of water), using a coveredflask to avoid loss by spattering. Boil the solution for two orthree minutes after the iron has dissolved to remove any volatilehydrocarbons. Meanwhile prepare the reductor for use as follows:Connect the vacuum bottle with the suction pump and pour into thefunnel at the top warm, dilute sulphuric acid, prepared by adding 5cc. Of concentrated sulphuric acid to 100 cc. Of distilled water. Seethat the stopcock (C) is open far enough to allow the acid to runthrough slowly. Continue to pour in acid until 200 cc. Have passedthrough, then close the stopcock !while a small quantity of liquidis still left in the funnel!. Discard the filtrate, and againpass through 100 cc. Of the warm, dilute acid. Test this with thepermanganate solution. A single drop should color it permanently; ifit does not, repeat the washing, until assured that the zinc is notcontaminated with appreciable quantities of reducing substances. Besure that no air enters the reductor (Note 1). 

Pour the iron solution while hot (but not boiling) through thereductor at a rate not exceeding 50 cc. Per minute (Notes 2 and 3). Wash out the beaker with dilute sulphuric acid, and follow the ironsolution without interruption with 175 cc. Of the warm acid andfinally with 75 cc. Of distilled water, leaving the funnel partiallyfilled. Remove the filter bottle and cool the solution quickly underthe water tap (Note 4), avoiding unnecessary exposure to the oxygen ofthe air. Add 10 cc. Of dilute sulphuric acid and titrate to a faintpink with the permanganate solution, adding it directly to thecontents of the vacuum flask. Should the end-point be overstepped, theferrous sulphate solution may be added. 

From the volume of the solution required to oxidize the iron inthe wire, calculate the relation to the normal of the permanganatesolution. The duplicate results should be concordant within two partsin 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 forsubsequent use after a very little washing; but a preliminary test isalways necessary to safeguard against error. 

If more than a small drop of permanganate solution is required tocolor 100 cc. Of the dilute acid after the reductor is well washed, anallowance must be made for the iron in the zinc. !Great care! must beused to prevent the access of air to the reductor after it has beenwashed 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 withthe zinc. The active agent may be considered to be !nascent! hydrogen, and it must be borne in mind that the visible bubbles are produced bymolecular hydrogen, which is without appreciable effect upon ferriciron. 

The rate at which the iron solution passes through the zinc should notexceed that prescribed, but the rate may be increased somewhat whenthe wash-water is added. It is well to allow the iron solution to runnearly, but not entirely, out of the funnel before the wash-wateris added. If it is necessary to interrupt the process, the completeemptying of the funnel can always be avoided by closing the stopcock. 

It is also possible to reduce the iron by treatment with zinc in aflask from which air is excluded. The zinc must be present in excessof the quantity necessary to reduce the iron and is finally completelydissolved. This method is, however, less convenient and more tediousthan the use of the reductor. ]

[Note 3: The dilute sulphuric acid for washing must be warmed readyfor use before the reduction of the iron begins, and it is of thefirst importance that the volume of acid and of wash-water shouldbe measured, and the volume used should always be the same in thestandardizations 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 themanganous sulphate formed during titration. If the solution turnsbrown, it is an evidence of insufficient acid, and more should beimmediately added. The results are likely to be less accurate in thiscase, however, as a consequence of secondary reactions between theferrous iron and the manganese dioxide thrown down. It is wiser todiscard such results and repeat the process. ]

[Note 5: The potassium permanganate may, of course, be diluted andbrought to an exactly 0. 1 N solution from the data here obtained. Thepercentage of iron in the iron wire must be taken into account in allcalculations. ]

!Method B!

!Oxalate Standards!

PROCEDURE. --Weigh out two portions of pure sodium oxalate of 0. 25-0. 3gram each into beakers of about 600 cc. Capacity. Add about 400 cc. Ofboiling water and 20 cc. Of manganous sulphate solution (Note 1). When the solution of the oxalate is complete, heat the liquid, ifnecessary, until near its boiling point (70-90°C. ) and run in thestandard permanganate solution drop by drop from a burette, stirringconstantly until an end-point is reached (Note 2). Make a blank testwith 20 cc. Of manganous sulphate solution and a volume of distilledwater equal to that of the titrated solution to determine the volumeof the permanganate solution required to produce a very slight pink. Deduct this volume from the amount of permanganate solution used inthe titration. 

From the data obtained, calculate the relation of the permanganatesolution 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 bydissolving 20 grams of MnSO_{4} in 200 cubic centimeters of water andadding 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 placequantitatively only in hot acid solutions. The temperatures must notfall below 70°C. ]

DETERMINATION OF IRON IN LIMONITE

!Method A!

The procedures, as here prescribed, are applicable to iron ores ingeneral, provided these ores contain no constituents which are reducedby zinc or stannous chloride and reoxidized by permanganates. Manyiron ores contain titanium, and this element among others doesinterfere with the determination of iron by the process described. If, however, the solutions of such ores are treated with sulphuretedhydrogen or sulphurous acid, instead of zinc or stannous chloride toreduce the iron, and the excess reducing agent removed by boiling, anaccurate determination of the iron can be made. 

PROCEDURE. --Grind the mineral to a fine powder. Weigh out two portionsof about 0. 5 gram each into small porcelain crucibles. Roast the oreat dull redness for ten minutes (Note 1), allow the crucibles to cool, and place them and their contents in casseroles containing 30 cc. Ofdilute hydrochloric acid (sp. Gr. 1. 12). 

Proceed with the solution of the ore, and the treatment of theresidue, if necessary, exactly as described for the bichromate processon page 56. When solution is complete, add 6 cc. Of concentratedsulphuric acid to each casserole, and evaporate on the steam bathuntil the solution is nearly colorless (Note 2). Cover the casserolesand heat over the flame of the burner, holding the casserole inthe hand and rotating it slowly to hasten evaporation and preventspattering, until the heavy white fumes of sulphuric anhydride arefreely 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 previouslywashed; proceed as described under standardization, taking painsto use the same volume and strength of acid and the same volume ofwash-water as there prescribed, and titrate with the permanganatesolution in the reductor flask, using the ferrous sulphate solution ifthe end-point should be overstepped. 

From the corrected volume of permanganate solution used, calculate thepercentage of iron (Fe) in the limonite. 

[Note 1: The preliminary roasting is usually necessary because, eventhough the sulphuric acid would subsequently char the carbonaceousmatter, certain nitrogenous bodies are not thereby rendered insolublein the acid, and would be oxidized by the permanganate. ]

[Note 2: The temperature of the steam bath is not sufficient tovolatilize sulphuric acid. Solutions may, therefore, be left toevaporate overnight without danger of evaporation to dryness. ]

[Note 3: The hydrochloric acid, both free and combined, is displacedby the less volatile sulphuric acid at its boiling point. Ferricsulphate separates at this point, since there is no water to holdit in solution and care is required to prevent bumping. The ferricsulphate usually has a silky appearance and is easily distinguishedfrom the flocculent silica which often remains undissolved. ]

!Zimmermann-Reinhardt Procedure!

!Method (B)!

PROCEDURE. --Grind the mineral to a fine powder. Weigh out two portionsof about 0. 5 gram each into small porcelain crucibles. Proceed withthe solution of the ore, treat the residue, if necessary, and reducethe iron by the addition of stannous chloride, followed by mercuricchloride, as described for the bichromate process on page 56. Dilutethe solution to about 400 cc. With cold water, add 10 cc. Of themanganous sulphate titrating solution (Note 1, page 68) and titratewith the standard potassium permanganate solution to a faint pink(Note 1). 

From the standardization data already obtained calculate thepercentage of iron (Fe) in the limonite. 

[Note 1: It has already been noted that hydrochloric acid reactsslowly in cold solutions with potassium permanganate. It is, however, possible to obtain a satisfactory, although somewhat fugitiveend-point in the presence of manganous sulphate and phosphoric acid. The explanation of the part played by these reagents is somewhatobscure as yet. It is possible that an intermediate manganic compoundis formed which reacts rapidly with the ferrous compounds--thus ineffect catalyzing the oxidizing process. 

While an excess of hydrochloric acid is necessary for the successfulreduction of the iron by stannous chloride, too large an amountshould be avoided in order to lessen the chance of reduction of thepermanganate by the acid during titration. ]

DETERMINATION OF THE OXIDIZING POWER OF PYROLUSITE

INDIRECT OXIDATION

Pyrolusite, when pure, consists of manganese dioxide. Its value as anoxidizing agent, and for the production of chlorine, depends upon thepercentage of MnO_{2} in the sample. This percentage is determinedby an indirect method, in which the manganese dioxide is reduced anddissolved by an excess of ferrous sulphate or oxalic acid in thepresence of sulphuric acid, and the unused excess determined bytitration with standard permanganate solution. 

PROCEDURE. --Grind the mineral in an agate mortar until no gritwhatever can be detected under the pestle (Note 1). Transfer it to astoppered weighing-tube, and weigh out two portions of about 0. 5 graminto beakers (400-500 cc. ) Read Note 2, and then calculate in eachcase the weight of oxalic acid (H_{2}C_{2}O_{4}. 2H_{2}O) required toreact 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! oxalicacid into the corresponding beakers, weighing the acid accurately andrecording the weight in the notebook. Pour into each beaker 25 cc. Ofwater and 50 cc. Of dilute sulphuric acid (1:5), cover and warm thebeaker and its contents gently until the evolution of carbon dioxideceases (Note 3). If a residue remains which is sufficiently colored toobscure the end-reaction of the permanganate, it must be removed byfiltration. 

Finally, dilute the solution to 200-300 cc. , heat the solution to atemperature just below boiling, add 15 cc. Of a manganese sulphatesolution and while hot, titrate for the excess of the oxalic acid withstandard permanganate solution (Notes 4 and 5). 

From the corrected volume of the solution required, calculate theamount of oxalic acid undecomposed by the pyrolusite; subtract thisfrom the total quantity of acid used, and calculate the weight ofmanganese dioxide which would react with the balance of the acid, andfrom this the percentage in the sample. 

[Note 1: The success of the analysis is largely dependent upon thefineness of the powdered mineral. If properly ground, solution shouldbe complete in fifteen minutes or less. ]

[Note 2: A moderate excess of oxalic acid above that required to reactwith the pyrolusite is necessary to promote solution; otherwise theresidual quantity of oxalic acid would be so small that the lastparticles of the mineral would scarcely dissolve. It is also desirablethat a sufficient excess of the acid should be present to react with aconsiderable volume of the permanganate solution during the titration, thus increasing the accuracy of the process. On the other hand, theexcess of oxalic acid should not be so large as to react with more ofthe permanganate solution than is contained in a 50 cc. Burette. Ifthe pyrolusite under examination is known to be of high grade, say 80per cent pure, or above the calculation of the oxalic acid needed maybe based upon an assumption that the mineral is all MnO_{2}. If thequality of the mineral is unknown, it is better to weigh out threeportions instead of two and to add to one of these the amount ofoxalic prescribed, assuming complete purity of the mineral. Then runin the permanganate solution from a pipette or burette to determineroughly the amount required. If the volume exceeds the contents of aburette, the amount of oxalic acid added to the other two portions isreduced accordingly. ]

[Note 3: Care should be taken that the sides of the beaker are notoverheated, as oxalic acid would be decomposed by heat alone ifcrystallization should occur on the sides of the vessel. Strongsulphuric acid also decomposes the oxalic acid. The dilute acidshould, therefore, be prepared before it is poured into the beaker. ]

[Note 4: Ferrous ammonium sulphate, ferrous sulphate, or iron wiremay be substituted for the oxalic acid. The reaction is then thefollowing:

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 withpotassium bichromate, if desired. Care is required to prevent theoxidation of the iron by the air, if ferrous salts are employed. ]

[Note 5: The oxidizing power of pyrolusite may be determined by othervolumetric processes, one of which is outlined in the followingreactions:

MnO_{2} + 4HCl --> MnCl_{2} + Cl_{2} + 2H_{2}OCl_{2} + 2KI --> I_{2} + 2KClI_{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 ofpotassium iodide. The liberated iodine is then determined by titrationwith sodium thiosulphate, as described on page 78. This is a directprocess, although it involves three steps. ]

IODIMETRY

The titration of iodine against sodium thiosulphate, with starch as anindicator, may perhaps be regarded as the most accurate of volumetricprocesses. The thiosulphate solution may be used in both acid andneutral solutions to measure free iodine and the latter may, in turn, serve as a measure of any substance capable of liberating iodine frompotassium iodide under suitable conditions for titration, as, forexample, in the process outlined in Note 5 on page 74. 

The fundamental reaction upon which iodometric processes are based isthe 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 inthe formation of the compound Na_{2}S_{4}O_{6}, called sodiumtetrathionate, is quantitatively exact, and differs in thatrespect from the action of chlorine or bromine, which oxidize thethiosulphate, but not quantitatively. 

NORMAL SOLUTIONS OF IODINE AND SODIUM THIOSULPHATE

If the formulas of sodium thiosulphate and sodium tetrathionate arewritten in a manner to show the atoms of oxygen associatedwith sulphur atoms in each, thus, 2(Na_{2}). S_{2}O_{2} andNa_{2}O. S_{4}O_{5}, it is plain that in the tetrathionate there arefive atoms of oxygen associated with sulphur, instead of the fourin the two molecules of the thiosulphate taken together. Although, therefore, the iodine contains no oxygen, the two atoms of iodinehave, in effect, brought about the addition of one oxygen atoms to thesulphur atoms. That is the same thing as saying that 253. 84 grams ofiodine (I_{2}) are equivalent to 16 grams of oxygen; hence, since 8grams of oxygen is the basis of normal solutions, 253. 84/2 or 126. 97grams of iodine should be contained in one liter of normal iodinesolution. By a similar course of reasoning the conclusion is reachedthat the normal solution of sodium thiosulphate should contain, per liter, its molecular weight in grams. As the thiosulphate incrystalline form has the formula Na_{2}S_{2}O_{3}. 5H_{2}O, this weightis 248. 12 grams. Tenth-normal or hundredth-normal solutions aregenerally used. 

PREPARATION OF STANDARD SOLUTIONS

!Approximate Strength, 0. 1 N!

PROCEDURE. --Weigh out on the rough balances 13 grams of commercialiodine. Place it in a mortar with 18 grams of potassium iodide andtriturate with small portions of water until all is dissolved. Dilutethe solution to 1000 cc. And transfer to a liter bottle and mixthoroughly (Note 1). [1]

[Footnote 1: It will be found more economical to have a considerablequantity of the solution prepared by a laboratory attendant, and tohave all unused solutions returned to the common stock. ]

Weigh out 25 grams of sodium thiosulphate, dissolve it in water whichhas been previously boiled and cooled, and dilute to 1000 cc. , alsowith boiled water. Transfer the solution to a liter bottle and mixthoroughly (Note 2). 

[Note 1: Iodine solutions react with water to form hydriodic acidunder the influence of the sunlight, and even at low room temperaturesthe iodine tends to volatilize from solution. They should, therefore, be protected from light and heat. Iodine solutions are not stable forlong periods under the best of conditions. They cannot be used inburettes with rubber tips, since they attack the rubber. ]

[Note 2: Sodium thiosulphate (Na_{2}S_{2}O_{3}. 5H_{2}O) israrely wholly pure as sold commercially, but may be purified byrecrystallization. The carbon dioxide absorbed from the air bydistilled water decomposes the salt, with the separation of sulphur. Boiled water which has been cooled out of contact with the air shouldbe 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 of0. 5 gram of this starch in 25 cc. Of boiling water is sufficient. Thesolution 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 about1 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 putit aside to settle. Decant the supernatant liquid through a filterand use the clear filtrate; 5 cc. Of this solution are needed for atitration. 

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 decolorizedby the thiosulphate (Note 1). 

[Note 1: The blue color which results when free iodine and starchare brought together is probably not due to the formation of a truechemical compound. It is regarded as a "solid solution" of iodine instarch. Although it is unstable, and easily destroyed by heat, itserves as an indicator for the presence of free iodine of remarkablesensitiveness, and makes the iodometric processes the mostsatisfactory of any in the field of volumetric analysis. ]

COMPARISON OF IODINE AND THIOSULPHATE SOLUTIONS

PROCEDURE. --Place the solutions in burettes (the iodine in aglass-stoppered burette), observing the usual precautions. Run out 40cc. Of the thiosulphate solution into a beaker, dilute with 150 cc. Ofwater, add 1 cc. To 2 cc. Of the soluble starch solution, and titratewith 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 temperaturechanges. 

STANDARDIZATION OF SOLUTIONS

Commercial iodine is usually not sufficiently pure to permit of itsuse as a standard for thiosulphate solutions or the direct preparationof a standard solution of iodine. It is likely to contain, besidemoisture, some iodine chloride, if chlorine was used to liberate theiodine when it was prepared. It may be purified by sublimation aftermixing it with a little potassium iodide, which reacts with the iodinechloride, forming potassium chloride and setting free the iodine. Thesublimed iodine is then dried by placing it in a closed container overconcentrated sulphuric acid. It may then be weighed in a stopperedweighing-tube and dissolved in a solution of potassium iodide in astoppered flask to prevent loss of iodine by volatilization. About 18grams of the iodide and twelve grams of iodine per liter are requiredfor an approximately tenth-normal solution. 

An iodine solution made from commercial iodine may also bestandardized against arsenious oxide (As_{4}O_{6}). This substancealso usually requires purification by sublimation before use. 

The substances usually employed for the standardization of athiosulphate solution are potassium bromate and metallic copper. Theformer is obtainable in pure condition or may be easily purified byre-crystallization. Copper wire of high grade is sufficiently pureto serve as a standard. Both potassium bromate and cupric salts insolution will liberate iodine from an iodide, which is then titratedwith 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 thiosulphatesolution are here described, and one for the direct standardization ofthe iodine solution. 

!Method A!

PROCEDURE. --Weigh out into 500 cc. Beakers two portions of about0. 150-0. 175 gram of potassium bromate. Dissolve each of these in 50cc. Of water, and add 10 cc. Of a potassium iodide solution containing3 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 ofwater), allow the solution to stand for three minutes, and dilute to150 cc. (Note 2). Run in thiosulphate solution from a burette untilthe color of the liberated iodine is nearly destroyed, and then add 1cc. Or 2 cc. Of starch solution, titrate to the disappearance of theiodo-starch blue, and finally add iodine solution until the coloris just restored. Make a blank test for the amount of thiosulphatesolution required to react with the iodine liberated by the iodatewhich is generally present in the potassium iodide solution, anddeduct this from the total volume used in the titration. 

From the data obtained, calculate the relation of the thiosulphatesolution to a normal solution, and subsequently calculate the similarvalue for the iodine solution. 

[Note 1:--Potassium iodide usually contains small amounts of potassiumiodate as impurity which, when the iodide is brought into an acidsolution, liberates iodine, just as does the potassium bromate used asa standard. It is necessary to determine the amount of thiosulphatewhich reacts with the iodine thus liberated by making a "blank test"with the iodide and acid alone. As the iodate is not always uniformlydistributed throughout the iodide, it is better to make up asufficient volume of a solution of the iodide for the purposes of thework in hand, and to make the blank test by using the same volume ofthe iodide solution as is added in the standardizing process. Theiodide solution should contain about 3 grams of the salt in 10 cc. ]

[Note 2: The color of the iodo-starch is somewhat less satisfactory inconcentrated solutions of the alkali salts, notably the iodides. Thedilution prescribed obviates this difficulty. ]

!Method B!

PROCEDURE. --Weigh out two portions of 0. 25-0. 27 gram of clean copperwire into 250 cc. Erlenmeyer flasks (Note 1). Add to each 5 cc. Ofconcentrated nitric acid (sp. Gr. 1. 42) and 25 cc. Of water, cover, and warm until solution is complete. Add 5 cc. Of bromine water andboil until the excess of bromine is expelled. Cool, and add strongammonia (sp. Gr. 0. 90) drop by drop until a deep blue color indicatesthe presence of an excess. Boil the solution until the deep blue isreplaced by a light bluish green, or a brown stain appears on thesides 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 potassiumiodide (Note 3) containing about 3 grams of the salt, and titratewith thiosulphate solution until the color of the liberated iodineis nearly destroyed. Then add 1-2 cc. Of freshly prepared starchsolution, and add thiosulphate solution, drop by drop, until the bluecolor 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, therequirements for its use as a conductor of electricity are such thatthe impurities constitute only a few hundredths of one per cent andare negligible for analytical purposes. ]

[Note 2: Ammonia neutralizes the free nitric acid. It should be addedin slight excess only, since the excess must be removed by boiling, which is tedious. If too much ammonia is present when acetic acid isadded, the resulting ammonium acetate is hydrolyzed, and the ammoniumhydroxide reacts with the iodine set free. ]

[Note 3: A considerable excess of potassium iodide is necessary forthe 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. 200gram each of pure arsenious oxide. Dissolve each of these in 10 cc. Ofsodium hydroxide solution, with stirring. Dilute the solutions to 150cc. And add dilute hydrochloric acid until the solutions contain a fewdrops in excess, and finally add to each a concentrated solution of5 grams of pure sodium bicarbonate (NaHCO_{3}) in water. Cover thebeakers before adding the bicarbonate, to avoid loss. Add the starchsolution and titrate with the iodine to the appearance of the blue ofthe iodo-starch, taking care not to pass the end-point by more than afew drops (Note 1). 

From the corrected volume of the iodine solution used to oxidize thearsenious oxide, calculate its relation to the normal. From theratio between the solutions, calculate the similar value for thethiosulphate solution. 

[Note 1: Arsenious oxide dissolves more readily in caustic alkali thanin a bicarbonate solution, but the presence of caustic alkali duringthe titration is not admissible. It is therefore destroyed by theaddition of acid, and the solution is then made neutral with thesolution of bicarbonate, part of which reacts with the acid, theexcess 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 alwaysfree from secondary reactions in the presence of even the weaklyalkaline bicarbonate, it is best to avoid the addition of anyconsiderable excess of iodine. Should the end-point be passed by a fewdrops, the thiosulphate may be used to correct it. ]

DETERMINATION OF COPPER IN ORES

Copper ores vary widely in composition from the nearly pure copperminerals, such as malachite and copper sulphide, to very low gradematerials which contain such impurities as silica, lead, iron, silver, sulphur, arsenic, and antimony. In nearly all varieties there will befound 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 copperfrom commonly occurring ores, and for the presence of their commonimpurities. For practice analyses it is advisable to select an ore ofa fair degree of purity. 

PROCEDURE. -- Weigh out two portions of about 0. 5 gram each of theore (which should be ground until no grit is detected) into 250 cc. Erlenmeyer flasks or small beakers. Add 10 cc. Of concentrated nitricacid (sp. Gr. 1. 42) and heat very gently until the ore is decomposedand the acid evaporated nearly to dryness (Note 1). Add 5 cc. Ofconcentrated hydrochloric acid (sp. Gr. 1. 2) and warm gently. Thenadd about 7 cc. Of concentrated sulphuric acid (sp. Gr. 1. 84) andevaporate over a free flame until the sulphuric acid fumes freely(Note 2). It has then displaced nitric and hydrochloric acid fromtheir compounds. 

Cool the flask or beaker, add 25 cc. Of water, heat the solutionto boiling, and boil for two minutes. Filter to remove insolublesulphates, silica and any silver that may have been precipitated assilver chloride, and receive the filtrate in a small beaker, washingthe precipitate and filter paper with warm water until the filtrateand washings amount to 75 cc. Bend a strip of aluminium foil (5 cm. X12 cm. ) into triangular form and place it on edge in the beaker. Coverthe beaker and boil the solution (being careful to avoid loss ofliquid by spattering) for ten minutes, but do not evaporate to smallvolume. 

Wash the cover glass and sides of the beaker. The copper should now bein the form of a precipitate at the bottom of the beaker or adheringloosely to the aluminium sheet. Remove the sheet, wash it carefullywith hydrogen sulphide water and place it in a small beaker. Decantthe solution through a filter, wash the precipitated copper twice bydecantation with hydrogen sulphide water, and finally transfer thecopper to the filter paper, where it is again washed thoroughly, beingcareful at all times to keep the precipitated copper covered with thewash water. Remove and discard the filtrate and place an Erlenmeyerflask under the funnel. Pour 15 cc. Of dilute nitric acid (sp. Gr. 1. 20) over the aluminium foil in the beaker, thus dissolving anyadhering copper. Wash the foil with hot water and remove it. Warm thisnitric acid solution and pour it slowly through the filter paper, thereby dissolving the copper on the paper, receiving the acidsolution in the Erlenmeyer flask. Before washing the paper, pour 5 cc. Of saturated bromine water (Note 3) through it and finally wash thepaper carefully with hot water and transfer any particles of copperwhich may be left on it to the Erlenmeyer flask. Boil to expel thebromine. Add concentrated ammonia drop by drop until the appearance ofa deep blue coloration indicates an excess. Boil until the deep blueis displaced by a light bluish green coloration, or until brown stainsform on the sides of the flask. Add 10 cc. Of strong acetic acid (Note4) and cool under the water tap. Add a solution containing about 3grams of potassium iodide, as in the standardization, and titrate withthiosulphate solution until the yellow of the liberated iodine isnearly discharged. Add 1-2 cc. Of freshly prepared starch solution andtitrate to the disappearance of the blue color. 

From the data obtained, calculate the percentage of copper (Cu) in theore. 

[Note 1: Nitric acid, because of its oxidizing power, is used as asolvent for the sulphide ores. As a strong acid it will also dissolvethe copper from carbonate ores. The hydrochloric acid is added todissolve oxides of iron and to precipitate silver and lead. Thesulphuric acid displaces the other acids, leaving a solutioncontaining sulphates only. It also, by its dehydrating action, renderssilica from silicates insoluble. ]

[Note 2: Unless proper precautions are taken to insure the correctconcentrations of acid the copper will not precipitate quantitativelyon the aluminium foil; hence care must be taken to follow directionscarefully at this point. Lead and silver have been almost completelyremoved as sulphate and chloride respectively, or they too wouldbe precipitated on the aluminium. Bismuth, though precipitated onaluminium, has no effect on the analysis. Arsenic and antimonyprecipitate on aluminium and would interfere with the titration ifallowed to remain in the lower state of oxidation. ]

[Note 3: Bromine is added to oxidize arsenious and antimoniouscompounds from the original sample, and to oxidize nitrous acid formedby the action of nitric acid on copper and copper sulphide. ]

[Note 4: This reaction can be carried out in the presence of sulphuricand hydrochloric acids as well as acetic acid, but in the presenceof these strong acids arsenic and antimonic acids may react with thehydriodic acid produced with the liberation of free iodine, therebyreversing the process and introducing an error. ]

DETERMINATION OF ANTIMONY IN STIBNITE

Stibnite is native antimony sulphide. Nearly pure samples of thismineral are easily obtainable and should be used for practice, sincemany impurities, notably iron, seriously interfere with the accuratedetermination of the antimony by iodometric methods. It is, moreover, essential that the directions with respect to amounts of reagentsemployed and concentration of solutions should be followed closely. 

PROCEDURE. --Grind the mineral with great care, and weigh out twoportions of 0. 35-0. 40 gram into small, dry beakers (100 cc. ). Cover the beakers and pour over the stibnite 5 cc. Of concentratedhydrochloric 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 ofpowdered tartaric acid (Note 2). Warm the solution on the water bathfor ten minutes longer, dilute the solution very cautiously by addingwater in portions of 5 cc. , stopping if the solution turns red. Itis possible that no coloration will appear, in which case cautiouslycontinue the dilution to 125 cc. If a red precipitate or colorationdoes appear, warm the solution until it is colorless, and again dilutecautiously to a total volume of 125 cc. And boil for a minute (Note3). 

If a white precipitate of the oxychloride separates during dilution(which should not occur if the directions are followed), it is best todiscard 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 3grams 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 containsan excess of the bicarbonate, and then add 1 cc. Or 2 cc. Of starchsolution and titrate with iodine solution to the appearance of theblue, avoiding excess (Notes 5 and 6). 

From the corrected volume of the iodine solution required to oxidizethe antimony, calculate the percentage of antimony (Sb) in thestibnite. 

[Note 1: Antimony chloride is volatile with steam from itsconcentrated solutions; hence these solutions must not be boiled untilthey have been diluted. ]

[Note 2: Antimony salts, such as the chloride, are readily hydrolyzed, and compounds such as SbOCl are formed which are often relativelyinsoluble; but in the presence of tartaric acid compounds with complexions are formed, and these are soluble. An excess of hydrochloric acidalso prevents precipitation of the oxychloride because the H^{+} ionsfrom the acid lessen the dissociation of the water and thus preventany considerable hydrolysis. ]

[Note 3: The action of hydrochloric acid upon the sulphide sets freesulphureted 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 itis not wholly driven out, a point is reached during dilution at whichthe antimony sulphide, being no longer held in solution by the acid, separates. If the dilution is immediately stopped and the solutionwarmed, this sulphide is again brought into solution and at the sametime more of the sulphureted hydrogen is expelled. This procedure mustbe continued until the sulphureted hydrogen is all removed, since itreacts with iodine. If no precipitation of the sulphide occurs, itis an indication that the sulphureted hydrogen was all expelled onsolution 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. Theammonia should not neutralize all of the acid. ]

[Note 5: The reaction which takes place during titration may beexpressed 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 ofthe iodo-starch is discharged after standing a few moments, the causemay be an insufficient quantity of sodium bicarbonate, leaving thesolution slightly acid, or a very slight precipitation of an antimonycompound which is slowly acted upon by the iodine when the latter ismomentarily present in excess. In either case it is better to discardthe analysis and to repeat the process, using greater care in theamounts of reagents employed. ]

CHLORIMETRY

The processes included under the term !chlorimetry! comprisethose employed to determine chlorine, hypochlorites, bromine, andhypobromites. The reagent employed is sodium arsenite in the presenceof sodium bicarbonate. The reaction in the case of the hypochloritesis

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; butcommercial 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 thesolution to 300 cc. , and make it faintly acid with dilute hydrochloricacid. 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 solutionof sodium bicarbonate, with gentle warming, but solution in sodiumhydroxide takes place much more rapidly, and the excess of thehydroxide is readily neutralized by hydrochloric acid, with subsequentaddition of the bicarbonate to maintain neutrality during thetitration. 

The indicator required for this process is made by dipping strips offilter paper in a starch solution prepared as described on page 76, to which 1 gram of potassium iodide has been added. These strips areallowed to drain and spread upon a watch-glass until dry. When touchedby a drop of the solution the paper turns blue until the hypochloritehas 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 aderivative of both hydrochloric and hypochlorous acids. Its formula isCaClOCl. Its use as a bleaching or disinfecting agent, or as a sourceof chlorine, depends upon the amount of hypochlorous acid which ityields when treated with a stronger acid. It is customary to expressthe value of bleaching powder in terms of "available chlorine, " bywhich is meant the chlorine present as hypochlorite, but not thechlorine present as chloride. 

PROCEDURE. --Weigh out from a stoppered test tube into a porcelainmortar about 3. 5 grams of bleaching powder (Note 1). Triturate thepowder in the mortar with successive portions of water until it iswell 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, orpipette, observing the precaution that the liquid removed shallcontain approximately its proportion of suspended matter. 

Empty the flask or pipette into a beaker and wash it out. Run in thearsenite solution from a burette until no further reaction takes placeon the starch-iodide paper when touched by a drop of the solution ofbleaching powder. Repeat the titration, using a second 25 cc. Portion. 

From the volume of solution required to react with the bleachingpowder, calculate the percentage of available chlorine in the latter, assuming the titration reaction to be that between chlorine andarsenious 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 bleachingpowder enters into the reaction. 

[Note 1: The powder must be triturated until it is fine, otherwise thelumps will inclose calcium hypochlorite, which will fail to react withthe arsenious acid. The clear supernatant liquid gives percentageswhich are below, and the sediment percentages which are above, theaverage. The liquid measured off should, therefore, carry with it itsproper proportion of the sediment, so far as that can be brought aboutby shaking the solution just before removal of the aliquot part fortitration. ]

[Note 2: Bleaching powder is easily acted upon by the carbonic acid inthe air, which liberates the weak hypochlorous acid. This, of course, results in a loss of available chlorine. The original material foranalysis should be kept in a closed container and protected form theair as far as possible. It is difficult to obtain analytical sampleswhich are accurately representative of a large quantity of thebleaching powder. The procedure, as outlined, will yield results whichare 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 oneof silver in nitric acid causes a deposition of silver thiocyanate asa white, curdy precipitate. If ferric nitrate is also present, theslightest excess of the thiocyanate over that required to combine withthe silver is indicated by the deep red which is characteristic of thethiocyanate 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 inthat it is but very little dissociated in aqueous solutions, and thecharacteristic color appears to be occasioned by the formation of theun-ionized ferric salt. 

The normal solution of potassium thiocyanate should contain an amountof 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 ammoniumthiocyanate is used, the amount is 76. 08 grams. To prepare thesolution for this determination, which should be approximately 0. 05N, dissolve about 5 grams of potassium thiocyanate, or 4 grams ofammonium thiocyanate, in a small amount of water; dilute this solutionto 1000 cc. In a liter bottle and mix as usual. 

Prepare 20 cc. Of a saturated solution of ferric alum and add 5 cc. Ofdilute nitric acid (sp. Gr. 1. 20). About 5 cc. Of this solution shouldbe used as an indicator. 

STANDARDIZATION

PROCEDURE. --Crush a small quantity of silver nitrate crystals in amortar (Note 1). Transfer them to a watch-glass and dry them for anhour at 110°C. , protecting them from dust or other organic matter(Note 2). Weigh out two portions of about 0. 5 gram each and dissolvethem in 50 cc. Of water. Add 10 cc. Of dilute nitric acid which hasbeen recently boiled to expel the lower oxides of nitrogen, if any, and then add 5 cc. Of the indicator solution. Run in the thiocyanatesolution from a burette, with constant stirring, allowing theprecipitate to settle occasionally to obtain an exact recognitionof the end-point, until a faint red tinge can be detected in thesolution. 

From the data obtained, calculate the relation of the thiocyanatesolution to the normal. 

[Note 1: The thiocyanate cannot be accurately weighed; its solutionsmust, therefore, be standardized against silver nitrate (or puresilver), either in the form of a standard solution or in small, weighed portions. ]

[Note 2: The crystals of silver nitrate sometimes inclose water whichis expelled on drying. If the nitrate has come into contact withorganic bodies it suffers a reduction and blackens during the heating. 

It is plain that a standard solution of silver nitrate (made byweighing out the crystals) is convenient or necessary if manytitrations of this nature are to be made. In the absence of such asolution the liability of passing the end-point is lessened by settingaside a small fraction of the silver solution, to be added near theclose of the titration. ]

DETERMINATION OF SILVER IN COIN

PROCEDURE. -- Weigh out two portions of the coin of about 0. 5 grameach. Dissolve them in 15 cc. Of dilute nitric acid (sp. Gr. 1. 2) andboil until all the nitrous compounds are expelled (Note 1). Cool thesolution, dilute to 50 cc. , and add 5 cc. Of the indicator solution, and titrate with the thiocyanate to the appearance of the faint redcoloration (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 acidsolutions and in the presence of copper, if the latter does not exceed70 per cent. Above that percentage it is necessary to add silver inknown quantity to the solution. The liquid must be cold at the time oftitration and entirely free from nitrous compounds, as these sometimescause a reddening of the indicator solution. All utensils, distilledwater, 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 asthose 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, theweighing of the sample; second, the solution of the sample; third, theseparation of some substance from solution containing, or bearing adefinite relation to, the constituent to be measured, under conditionswhich render this separation as complete as possible; and finally, the segregation of that substance, commonly by filtration, and thedetermination of its weight, or that of some stable product formedfrom it on ignition. For example, the gravimetric determination ofaluminium is accomplished by solution of the sample, by precipitationin the form of hydroxide, collection of the hydroxide upon a filter, complete removal by washing of all foreign soluble matter, and theburning of the filter and ignition of the precipitate to aluminiumoxide, in which condition it is weighed. 

Among the operations which are common to nearly all gravimetricanalyses are precipitation, washing of precipitates, ignition ofprecipitates, and the use of desiccators. In order to avoid burdensomerepetitions in the descriptions of the various gravimetric procedureswhich follow, certain general instructions are introduced at thispoint. These instructions must, therefore, be considered to be as mucha part of all subsequent procedures as the description of apparatus, reagents, or manipulations. 

The analytical balance, the fundamentally important instrument ingravimetric analysis, has already been described on pages 11 to 15. 

PRECIPITATION

For successful quantitative precipitations those substances areselected which are least soluble under conditions which can be easilyestablished, and which separate from solution in such a state thatthey can be filtered readily and washed free from admixed material. In general, the substances selected are the same as those alreadyfamiliar to the student of Qualitative Analysis. 

When possible, substances are selected which separate in crystallineform, since such substances are less likely to clog the pores offilter paper and can be most quickly washed. In order to increase thesize of the crystals, which further promotes filtration and washing, it is often desirable to allow a precipitate to remain for some timein contact with the solution from which it has separated. The solutionis often kept warm during this period of "digestion. " The smallcrystals gradually disappear and the larger crystals increase in size, probably as the result of the force known as surface tension, whichtends to reduce the surface of a given mass of material to a minimum, combined with a very slightly greater solubility of small crystals ascompared with the larger ones. 

Amorphous substances, such as ferric hydroxide, aluminium hydroxide, or silicic acid, separate in a gelatinous form and are relativelydifficult to filter and wash. Substances of this class also exhibita tendency to form, with pure water, what are known as colloidalsolutions. To prevent this as far as possible, they are washed withsolutions of volatile salts, as will be described in some of thefollowing procedures. 

In all precipitations the reagent should be added slowly, withconstant stirring, and should be hot when circumstances permit. The slow addition is less likely to occasion contamination of theprecipitate by the inclosure of other substances which may be in thesolution, or of the reagent itself. 

FUNNELS AND FILTERS

Filtration in analytical processes is most commonly effected throughpaper filters. In special cases these may be advantageously replacedby an asbestos filter in a perforated porcelain or platinum crucible, commonly known, from its originator, as a "Gooch filter. " Theoperation and use of a filter of this type is described on page 103. Porous crucibles of a material known as alundum may also be employedto advantage in special cases. 

The glass funnels selected for use with paper filters should have anangle as near 60° as possible, and a narrow stem about six inches inlength. The filters employed should be washed filters, i. E. , thosewhich have been treated with hydrochloric and hydrofluoric acids, andwhich on incineration leave a very small and definitely known weightof ash, generally about . 00003 gram. Such filters are readilyobtainable on the market. 

The filter should be carefully folded to fit the funnel according toeither of the two well-established methods described in the Appendix. It should always be placed so that the upper edge of the paperis about one fourth inch below the top of the funnel. Under nocircumstances 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 continuouscolumn of liquid which, by its hydrostatic pressure, produces a gentlesuction, thus materially promoting the rapidity of filtration. Unlessthe filter allows free passage of water under these conditions, it islikely to give much trouble when a precipitate is placed upon it. 

The use of a suction pump to promote filtration is rarely altogetheradvantageous in quantitative analysis, if paper filters are employed. The tendency of the filter to break, unless the point of the filterpaper is supported by a perforated porcelain cone or a small "hardenedfilter" of parchment, and the tendency of the precipitates to passthrough the pores of the filter, more than compensate for the possiblegain in time. On the other hand, filtration by suction may be usefulin the case of precipitates which do not require ignition beforeweighing, or in the case of precipitates which are to be discardedwithout weighing. This is best accomplished with the aid of thespecial apparatus called a Gooch filter referred to above. 

FILTRATION AND WASHING OF PRECIPITATES

Solutions should be filtered while hot, as far as possible, sincethe passage of a liquid through the pores of a filter is retarded byfriction, and this, for water at 100°C. , is less than one sixth of theresistance at 0°C. 

When the filtrate is received in a beaker, the stem of the funnelshould touch the side of the receiving vessel to avoid loss byspattering. Neglect of this precaution is a frequent source of error. 

The vessels which contain the initial filtrate should !always! bereplaced by clean ones, properly labeled, before the washing of aprecipitate begins. In many instances a finely divided precipitatewhich shows no tendency to pass through the filter at first, while thesolution is relatively dense, appears at once in the washings. Undersuch conditions the advantages accruing from the removal of the firstfiltrate are obvious, both as regards the diminished volume requiringrefiltration, and also the smaller number of washings subsequentlyrequired. 

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. Thesupernatant, clear wash-water is then decanted through the filter, so far as practicable without disturbing the precipitate, and a newportion of wash-water is added. This procedure can be employed tospecial advantage with gelatinous precipitates, which fill up thepores of the filter paper. As the medium from which the precipitateis to settle becomes less dense it subsides less readily, and itultimately becomes necessary to transfer it to the filter and completethe washing there. 

A precipitate should never completely fill a filter. The wash-watershould 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-watereach time, and allowing it to drain completely after each addition;but that when a precipitate is to be washed with the !least possiblevolume! of liquid the latter should be applied in repeated !small!quantities. 

Gelatinous precipitates should not be allowed to dry before completeremoval of foreign matter is effected. They are likely to shrink andcrack, and subsequent additions of wash-water pass through thesechannels only. 

All filtrates and wash-waters without exception must be properlytested. !This lies at the foundation of accurate work!, and thestudent should clearly understand that it is only by the invariableapplication of this rule that assurance of ultimate reliability canbe secured. Every original filtrate must be tested to prove completeprecipitation of the compound to be separated, and the wash-watersmust also be tested to assure complete removal of foreign material. Intesting the latter, the amount first taken should be but a fewdrops if the filtrate contains material which is to be subsequentlydetermined. When, however, the washing of the filter and precipitateis nearly completed the amount should be increased, and for the finaltest not less than 3 cc. Should be used. 

It is impossible to trust to one's judgment with regard to the washingof precipitates; the washings from !each precipitate! of a seriessimultaneously treated must be tested, since the rate of washing willoften differ materially under apparently similar conditions, !Noexception can ever be made to this rule!. 

The habit of placing a clean common filter paper under the receivingbeaker during filtration is one to be commended. On this paper arecord of the number of washings can very well be made as the portionsof wash-water are added. 

It is an excellent practice, when possible, to retain filtrates andprecipitates until the completion of an analysis, in order that, incase of question, they may be examined to discover sources of error. 

For the complete removal of precipitates from containing vessels, itis often necessary to rub the sides of these vessels to loosen theadhering particles. This can best be done by slipping over the end ofa 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 withsilica tubes, to support the crucible or other utensils. The cover ofthe desiccator should be made air-tight by the use of a thin coatingof vaseline. 

Pumice moistened with concentrated sulphuric acid may be used in placeof the calcium chloride, and is essential in special cases; but formost purposes the calcium chloride, if renewed occasionally and notallowed to cake together, is practically efficient and does not slopabout when the desiccator is moved. 

Desiccators should never remain uncovered for any length of time. Thedehydrating agents rapidly lose their efficiency on exposure to theair. 

CRUCIBLES

It is often necessary in quantitative analysis to employ fluxes tobring into solution substances which are not dissolved by acids. Thefluxes in most common use are sodium carbonate and sodium or potassiumacid sulphate. In gravimetric analysis it is usually necessary toignite the separated substance after filtration and washing, in orderto remove moisture, or to convert it through physical or chemicalchanges into some definite and stable form for weighing. Cruciblesto be used in fusion processes must be made of materials which willwithstand the action of the fluxes employed, and crucibles to be usedfor ignitions must be made of material which will not undergo anypermanent change during the ignition, since the initial weight of thecrucible must be deducted from the final weight of the crucible andproduct to obtain the weight of the ignited substance. The threematerials which satisfy these conditions, in general, are platinum, porcelain, and silica. 

Platinum crucibles have the advantage that they can be employed athigh temperatures, but, on the other hand, these crucibles can neverbe used when there is a possibility of the reduction to the metallicstate of metals like lead, copper, silver, or gold, which would alloywith and ruin the crucible. When platinum crucibles are used withcompounds of arsenic or phosphorus, special precautions are necessaryto prevent damage. This statement applies to both fusions andignitions. 

Fusions with sodium carbonate can be made only in platinum, sinceporcelain or silica crucibles are attacked by this reagent. Acidsulphate fusions, which require comparatively low temperatures, cansometimes be made in platinum, although platinum is slightly attackedby the flux. Porcelain or silica crucibles may be used with acidfluxes. 

Silica crucibles are less likely to crack on heating than porcelaincrucibles on account of their smaller coefficient of expansion. Ignition of substances not requiring too high a temperature may bemade 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 ispracticable, and this is the custom in private, research or commerciallaboratories. Platinum has, however, become so valuable that it isliable to theft unless constantly under the protection of the user. Asconstant protection is often difficult in instructional laboratories, it is advisable, in order to avoid serious monetary losses, to useporcelain or silica crucibles whenever these will give satisfactoryservice. When platinum utensils are used the danger of theft shouldalways be kept in mind. 

PREPARATION OF CRUCIBLES FOR USE

All crucibles, of whatever material, must always be cleaned, ignitedand allowed to cool in a desiccator before weighing, since all bodiesexposed to the air condense on their surfaces a layer of moisturewhich increases their weight. The amount and weight of this moisturevaries with the humidity of the atmosphere, and the latter may changefrom hour to hour. The air in the desiccator (see above) is kept ata constant and low humidity by the drying agent which it contains. Bodies which remain in a desiccator for a sufficient time (usually20-30 minutes) retain, therefore, on their surfaces a constant weightof moisture which is the same day after day, thus insuring constantconditions. 

Hot objects, such as ignited crucibles, should be allowed to cool inthe air until, when held near the skin, but little heat is noticeable. If this precaution is not taken, the air within the desiccator isstrongly heated and expands before the desiccator is covered. As thetemperature falls, the air contracts, causing a reduction of airpressure within the covered vessel. When the cover is removed (whichis often rendered difficult) the inrush of air from the outside maysweep light particles out of a crucible, thus ruining an entireanalysis. 

Constant heating of platinum causes a slight crystallization of thesurface which, if not removed, penetrates into the crucible. Gentlepolishing of the surface destroys the crystalline structure andprevents further damage. If sea sand is used for this purpose, greatcare is necessary to keep it from the desk, since beakers are easilyscratched by it, and subsequently crack on heating. 

Platinum crucibles stained in use may often be cleaned by the fusionin them of potassium or sodium acid sulphate, or by heating withammonium chloride. If the former is used, care should be taken notto heat so strongly as to expel all of the sulphuric acid, since thenormal sulphates sometimes expand so rapidly on cooling as to splitthe crucible. The fused material should be poured out, while hot, onto a !dry! tile or iron surface. 

IGNITION OF PRECIPITATES

Most precipitates may, if proper precautions are taken, be ignitedwithout previous drying. If, however, such precipitates can be driedwithout loss of time to the analyst (as, for example, over night), itis well to submit them to this process. It should, nevertheless, beremembered that a partially dried precipitate often requires more careduring ignition than a thoroughly moist one. 

The details of the ignition of precipitates vary so much with thecharacter of the precipitate, its moisture content, and temperature towhich it is to be heated, that these details will be given under thevarious 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 sodiumchloride, handling it as little as possible with the moist fingers, and weigh it accurately to 0. 0001 gram, recording the weight at oncein the notebook (see Appendix). Hold the tube over the top of a beaker(200-300 cc. ), and cautiously remove the stopper, noting carefullythat no particles fall from it, or from the tube, elsewhere than intothe beaker. Pour out a small portion of the chloride, replace thestopper, and determine by approximate weighing how much has beenremoved. Continue this procedure until 0. 25-0. 30 gram has been takenfrom the tube, then weigh accurately and record the weight beneath thefirst in the notebook. The difference of the two weights representsthe weight of the chloride taken for analysis. Again weigh a secondportion of 0. 25-0. 30 gram into a second beaker of the same size as thefirst. The beakers should be plainly marked to correspond with theentries in the notebook. Dissolve each portion of the chloride in 150cc. Of distilled water and add about ten drops of dilute nitric acid(sp. Gr. 1. 20) (Note 2). Calculate the volume of silver nitratesolution required to effect complete precipitation in each case, and add slowly about 5 cc. In excess of that amount, with constantstirring. Heat the solutions cautiously to boiling, stirringoccasionally, and continue the heating and stirring until theprecipitates settle promptly, leaving a nearly clear supernatantliquid (Note 3). This heating should not take place in direct sunlight(Note 4). The beaker should be covered with a watch-glass, and bothboiling and stirring so regulated as to preclude any possibility ofloss of material. Add to the clear liquid one or two drops of silvernitrate solution, to make sure that an excess of the reagent ispresent. If a precipitate, or cloudiness, appears as the drops fallinto the solution, heat again, and stir until the whole precipitatehas 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 porcelaincrucible (Gooch crucible), making sure that when the crucible isgently 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 therubber tubing is stretched over the top of the funnel and then drawnup 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 thefilter bottle with the suction pump. Suspend some finely dividedasbestos, 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, andthen pour some of the mixture cautiously into the crucible until aneven felt of asbestos, not over 1/32 inch in thickness, is formed. Agentle suction must be applied while preparing this felt. Wash thefelt thoroughly by passing through it distilled water until all fineor loose particles are removed, increasing the suction at the lastuntil no more water can be drawn out of it; place on top of the feltthe small, perforated porcelain disc and hold it in place by pouring avery thin layer of asbestos over it, washing the whole carefully;then place the crucible in a small beaker, and place both in a dryingcloset at 100-110°C. For thirty to forty minutes. Cool the cruciblein a desiccator, and weigh. Heat again for twenty to thirty minutes, cool, and again weigh, repeating this until the weight is constantwithin 0. 0003 gram. The filter is then ready for use. 

Place the crucible in the funnel, and apply a gentle suction, !afterwhich! the solution to be filtered may be poured in without disturbingthe asbestos felt. When pouring liquid onto a Gooch filter hold thestirring-rod at first well down in the crucible, so that the liquiddoes not fall with any force upon the asbestos, and afterward keep thecrucible 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 theprecipitate twice by decantation with warm water; then transfer itto the filter with the aid of a stirring-rod with a rubber tip and astream from the wash-bottle. 

Examine the first portions of the filtrate which pass through thefilter with great care for asbestos fibers, which are most likely tobe lost at this point. Refilter the liquid if any fibers are visible. Finally, wash the precipitate thoroughly with warm water until freefrom soluble silver salts. To test the washings, disconnect thesuction at the flask and remove the funnel or filter tube from thesuction flask. Hold the end of the tube over the mouth of a small testtube and add from a wash-bottle 2-3 cc. Of water. Allow the water todrip through into the test tube and add a drop of dilute hydrochloricacid. 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 constantwithin 0. 0003 gram, as described for the preparation of the filter. Deduct the weight of the dry crucible from the final weight, and fromthe weight of silver chloride thus obtained calculate the percentageof chlorine in the sample of sodium chloride. 

[Note 1: The washed asbestos for this type of filter is prepared bydigesting in concentrated hydrochloric acid, long-fibered asbestoswhich has been cut in pieces of about 0. 5 cm. In length. Afterdigestion, the asbestos is filtered off on a filter plate and washedwith hot, distilled water until free from chlorides. A small portionof the asbestos is shaken with water, forming a thin suspension, whichis bottled and kept for use. ]

[Note 2: The nitric acid is added before precipitation to lessen thetendency of the silver chloride to carry down with it other substanceswhich might be precipitated from a neutral solution. A large excess ofthe acid would exert a slight solvent action upon the chloride. ]

[Note 3: The solution should not be boiled after the addition of thenitric acid before the presence of an excess of silver nitrate isassured, since a slight interaction between the nitric acid and thesodium chloride is possible, by which a loss of chlorine, either assuch or as hydrochloric acid, might ensue. The presence of an excessof the precipitant can usually be recognized at the time of itsaddition, by the increased readiness with which the precipitatecoagulates and settles. ]

[Note 4: The precipitate should not be exposed to strong sunlight, since under those conditions a reduction of the silver chloride ensueswhich is accompanied by a loss of chlorine. The superficial alterationwhich the chloride undergoes in diffused daylight is not sufficientto materially affect the accuracy of the determination. It should benoted, however, that a slight error does result from the effect oflight upon the silver chloride precipitate and in cases in which thegreatest obtainable accuracy is required, the procedure describedunder "Method B" should be followed, in which this slight reduction ofthe silver chloride is corrected by subsequent treatment with nitricand hydrochloric acids. ]

[Note 5: The asbestos used in the Gooch filter should be of the finestquality and capable of division into minute fibrous particles. Acoarse felt is not satisfactory. ]

[Note 6: The precipitate must be washed with warm water until it isabsolutely free from silver and sodium nitrates. It may be assumedthat the sodium salt is completely removed when the wash-water showsno evidence of silver. It must be borne in mind that silver chlorideis somewhat soluble in hydrochloric acid, and only a single dropshould be added. The washing should be continued until no cloudinesswhatever can be detected in 3 cc. Of the washings. 

Silver chloride is but slightly soluble in water. The solubilityvaries with its physical condition within small limits, and isabout 0. 0018 gram per liter at 18°C. For the curdy variety usuallyprecipitated. The chloride is also somewhat soluble in solutions ofmany chlorides, in solutions of silver nitrate, and in concentratednitric acid. 

As a matter of economy, the filtrate, which contains whatever silvernitrate was added in excess, may be set aside. The silver can beprecipitated as chloride and later converted into silver nitrate. ]

[Note 7: The use of the Gooch filter commends itself strongly when aconsiderable number of halogen determinations are to be made, sincesuccessive portions of the silver halides may be filtered on the samefilter, without the removal of the preceding portions, until thecrucible is about two thirds filled. If the felt is properly prepared, filtration and washing are rapidly accomplished on this filter, andthis, combined with the possibility of collecting several precipitateson the same filter, is a strong argument in favor of its use with anybut gelatinous precipitates. ]

!Method B. With the Use of a Paper Filter!

PROCEDURE. --Weigh out two portions of sodium chloride of about0. 25-0. 3 gram each and proceed with the precipitation of the silverchloride as described under Method A above. When the chloride is readyfor filtration prepare two 9 cm. Washed paper filters (see Appendix). Pour the liquid above the precipitates through the filters, wash twiceby decantation and transfer the precipitates to the filters, finallywashing them until free from silver solution as described. The funnelshould then be covered with a moistened filter paper by stretching itover the top and edges, to which it will adhere on drying. It shouldbe properly labeled with the student's name and desk number, and thenplaced in a drying closet, at a temperature of about 100-110°C. , untilcompletely dry. 

The perfectly dry filter is then opened over a circular piece ofclean, smooth, glazed paper about six inches in diameter, placed upona larger piece about twelve inches in diameter. The precipitate isremoved from the filter as completely as possible by rubbing the sidesgently together, or by scraping them cautiously with a feather whichhas been cut close to the quill and is slightly stiff (Note 1). Ineither case, care must be taken not to rub off any considerablequantity of the paper, nor to lose silver chloride in the form ofdust. Cover the precipitate on the glazed paper with a watch-glass toprevent loss of fine particles and to protect it from dust from theair. Fold the filter paper carefully, roll it into a small cone, andwind loosely around !the top! a piece of small platinum wire (Note 2). Hold the filter by the wire over a small porcelain crucible (which hasbeen cleaned, ignited, cooled in a desiccator, and weighed), igniteit, and allow the ash to fall into the crucible. Place the crucibleupon a clean clay triangle, on its side, and ignite, with a lowflame well at its base, until all the carbon of the filter has beenconsumed. Allow the crucible to cool, add two drops of concentratednitric acid and one drop of concentrated hydrochloric acid, and heat!very cautiously!, to avoid spattering, until the acids have beenexpelled; then transfer the main portion of the precipitate from theglazed paper to the cooled crucible, placing the latter on the largerpiece of glazed paper and brushing the precipitate from thesmaller piece into it, sweeping off all particles belonging to thedetermination. 

Moisten the precipitate with two drops of concentrated nitric acid andone drop of concentrated hydrochloric acid, and again heat with greatcaution until the acids are expelled and the precipitate is white, when the temperature is slowly raised until the silver chloride justbegins to fuse at the edges (Note 3). The crucible is then cooled in adesiccator and weighed, after which the heating (without the additionof acids) is repeated, and it is again weighed. This must be continueduntil the weight is constant within 0. 0003 gram in two consecutiveweighings. Deduct the weight of the crucible, and calculate thepercentage of chlorine in the sample of sodium chloride taken foranalysis. 

[Note 1: The separation of the silver chloride from the filter isessential, since the burning carbon of the paper would reduce aconsiderable quantity of the precipitate to metallic silver, and itscomplete reconversion to the chloride within the crucible, by means ofacids, would be accompanied by some difficulty. The small amount ofsilver reduced from the chloride adhering to the filter paper afterseparating the bulk of the precipitate, and igniting the paperas prescribed, can be dissolved in nitric acid, and completelyreconverted to chloride by hydrochloric acid. The subsequent additionof the two acids to the main portion of the precipitate restores thechlorine to any chloride which may have been partially reduced by thesunlight. The excess of the acids is volatilized by heating. ]

[Note 2: The platinum wire is wrapped around the top of the filterduring its incineration to avoid contact with any reduced silver fromthe reduction of the precipitate. If the wire were placed nearer theapex, 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 hightemperatures. The temperature of fusion is not always sufficientto destroy filter shreds; hence these should not be allowed tocontaminate 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 thesample (Note 1) of about 1 gram each and dissolve these in 50 cc. Ofwater, to which 1 cc. Of dilute hydrochloric acid (sp. Gr. 1. 12) hasbeen added (Note 2). Heat the solution to boiling, and while at theboiling point add concentrated nitric acid (sp. Gr. 1. 42), !drop bydrop! (noting the volume used), until the brown coloration, whichappears after the addition of a part of the nitric acid, gives placeto a yellow or red (Note 3). Avoid a large excess of nitric acid, butbe sure that the action is complete. Pour this solution cautiouslyinto about 200 cc. Of water, containing a slight excess of ammonia. Calculate for this purpose the amount of aqueous ammonia required toneutralize the hydrochloric and nitric acids added (see Appendix fordata), and also to precipitate the iron as ferric hydroxide from theweight of the ferrous ammonium sulphate taken for analysis, assumingit to be pure (Note 4). The volume thus calculated will be in excessof that actually required for precipitation, since the acids are inpart consumed in the oxidation process, or are volatilized. Heat thesolution to boiling, and allow the precipitated ferric hydroxide tosettle. Decant the clear liquid through a washed filter (9 cm. ), keeping as much of the precipitate in the beaker as possible. Washtwice 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 aspossible and noting the volume used. Collect the solution in thebeaker in which precipitation took place. Add 1 cc. Of nitric acid(sp. Gr. 1. 42), boil for a few moments, and again pour into acalculated excess of ammonia. 

Wash the precipitate twice by decantation, and finally transfer it tothe original filter. Wash continuously with hot water until finally3 cc. Of the washings, acidified with nitric acid (Note 5), showno evidences of the presence of chlorides when tested with silvernitrate. The filtrate and washings are combined with those from thefirst precipitation and treated for the determination of sulphur, asprescribed 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 ofwater of crystallization; and also those which are red, indicatingthe presence of ferric iron. If, on the other hand, the value of anaverage sample of material is desired, it is preferable to grind thewhole together, mix thoroughly, and take a sample from the mixture foranalysis. ]

[Note 2: When aqueous solutions of ferrous compounds are heated in theair, oxidation of the Fe^{++} ions to Fe^{+++} ions readily occurs inthe absence of free acid. The H^{+} and OH^{-} ions from water areinvolved in the oxidation process and the result is, in effect, theformation of some ferric hydroxide which tends to separate. Moreover, at the boiling temperature, the ferric sulphate produced by theoxidation hydrolyzes in part with the formation of a basic ferricsulphate, which also tends to separate from solution. The addition ofthe hydrochloric acid prevents the formation of ferric hydroxide, andso far reduces the ionization of the water that the hydrolysis of theferric sulphate is also prevented, and no precipitation occurs onheating. ]

[Note 3: The nitric acid, after attaining a moderate strength, oxidizes the Fe^{++} ions to Fe^{+++} ions with the formation of anintermediate nitroso-compound similar in character to that formed inthe "ring-test" for nitrates. The nitric oxide is driven out by heat, and the solution then shows by its color the presence of ferriccompounds. A drop of the oxidized solution should be tested ona watch-glass with potassium ferricyanide, to insure a completeoxidation. This oxidation of the iron is necessary, since Fe^{++} ionsare not completely precipitated by ammonia. 

The ionic changes which are involved in this oxidation are perhapsmost 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 caseeither the nitric or the hydrochloric acid. The full equation on whichthis 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 theoxidation. ]

[Note 4: The ferric hydroxide precipitate tends to carry down somesulphuric acid in the form of basic ferric sulphate. This tendency islessened if the solution of the iron is added to an excess of OH^{-}ions from the ammonium hydroxide, since under these conditionsimmediate and complete precipitation of the ferric hydroxide ensues. A gradual neutralization with ammonia would result in the localformation of a neutral solution within the liquid, and subsequentdeposition of a basic sulphate as a consequence of a local deficiencyof OH^{-} ions from the NH_{4}OH and a partial hydrolysis of theferric salt. Even with this precaution the entire absence of sulphatesfrom the first iron precipitate is not assured. It is, therefore, redissolved and again thrown down by ammonia. The organic matter ofthe filter paper may occasion a partial reduction of the iron duringsolution, with consequent possibility of incomplete subsequentprecipitation with ammonia. The nitric acid is added to reoxidize thisiron. 

To avoid errors arising from the solvent action of ammoniacalliquids upon glass, the iron precipitate should be filtered withoutunnecessary delay. ]

[Note 5: The washings from the ferric hydroxide are acidified withnitric acid, before testing with silver nitrate, to destroy theammonia 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 dryduring the washing. ]

!Ignition of the Iron Precipitate!

Heat a platinum or porcelain crucible, cool it in a desiccator andweigh, repeating until a constant weight is obtained. 

Fold the top of the filter paper over the moist precipitate of ferrichydroxide and transfer it cautiously to the crucible. Wipe the insideof the funnel with a small fragment of washed filter paper, ifnecessary, and place the paper in the crucible. 

Incline the crucible on its side, on a triangle supported on aring-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 flameand protecting it from drafts of air by means of a chimney. The heatfrom the burner is thus reflected into the crucible and driesthe precipitate without danger of loss as the result of a suddengeneration of steam within the mass of ferric hydroxide. As the dryingprogresses the burner may be gradually moved toward the base of thecrucible and the flame increased until the paper of the filter beginsto char and finally to smoke, as the volatile matter is expelled. Thisis known as "smoking off" a filter, and the temperature should not beraised sufficiently high during this process to cause the paper toignite, as the air currents produced by the flame of the blazing papermay carry away particles of the precipitate. 

When the paper is fully charred, move the burner to the base of thecrucible and raise the temperature to the full heat of the burner forfifteen minutes, with the crucible still inclined on its side, butwithout the cover (Note 1). Finally set the crucible upright in thetriangle, cover it, and heat at the full temperature of a blast lampor other high temperature burner. Cool and weigh in the usual manner(Note 2). Repeat the strong heating until the weight is constantwithin 0. 0003 gram. 

From the weight of ferric oxide (Fe_{2}O_{3}) calculate the percentageof iron (Fe) in the sample (Note 3). 

[Note 1: These directions for the ignition of the precipitate must beclosely followed. A ready access of atmospheric oxygen is of specialimportance to insure the reoxidation to ferric oxide of any iron whichmay be reduced to magnetic oxide (Fe_{3}O_{4}) during the combustionof the filter. The final heating over the blast lamp is essentialfor the complete expulsion of the last traces of water from thehydroxide. ]

[Note 2: Ignited ferric oxide is somewhat hygroscopic. On this accountthe weighings must be promptly completed after removal from thedesiccator. In all weighings after the first it is well to place theweights upon the balance-pan before removing the crucible from thedesiccator. It is then only necessary to move the rider to obtain theweight. ]

[Note 3: The gravimetric determination of aluminium or chromium iscomparable with that of iron just described, with the additionalprecaution that the solution must be boiled until it contains but avery slight excess of ammonia, since the hydroxides of aluminium andchromium are more soluble than ferric hydroxide. 

The most important properties of these hydroxides, from a quantitativestandpoint, other than those mentioned, are the following: All areprecipitable by the hydroxides of sodium and potassium, but alwaysinclose some of the precipitant, and should be reprecipitated withammonium hydroxide before ignition to oxides. Chromium and aluminiumhydroxides dissolve in an excess of the caustic alkalies and formanions, probably of the formula AlO_2^{-} and CrO_{2}^{-}. Chromiumhydroxide is reprecipitated from this solution on boiling. When firstprecipitated the hydroxides are all readily soluble in acids, butaluminium hydroxide dissolves with considerable difficulty afterstanding or boiling for some time. The precipitation of the hydroxidesis promoted by the presence of ammonium chloride, but is partiallyor 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 hydroxideabout 0. 6 gram of anhydrous sodium carbonate; cover the beaker, andthen add dilute hydrochloric acid (sp. Gr. 1. 12) in moderate excessand evaporate to dryness on the water bath. Add 10 cc. Of concentratedhydrochloric acid (sp. Gr. 1. 20) to the residue, and again evaporateto dryness on the bath. Dissolve the residue in water, filter if notclear, transfer to a 700 cc. Beaker, dilute to about 400 cc. , andcautiously add hydrochloric acid until the solution shows a distinctlyacid reaction (Note 1). Heat the solution to boiling, and add !veryslowly! and with constant stirring, 20 cc. In excess of the calculatedamount of a hot barium chloride solution, containing about 20 gramsBaCl_{2}. 2H_{2}O per liter (Notes 2 and 3). Continue the boiling forabout two minutes, allow the precipitate to settle, and decant theliquid at the end of half an hour (Note 4). Replace the beakercontaining the original filtrate by a clean beaker, wash theprecipitated sulphate by decantation with hot water, and subsequentlyupon the filter until it is freed from chlorides, testing the washingsas described in the determination of iron. The filter is thentransferred to a platinum or porcelain crucible and ignited, asdescribed above, until the weight is constant (Note 5). From theweight of barium sulphate (BaSO_{4}) obtained, calculate thepercentage 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 ofdissociation of sulphuric acid in the presence of the H^{+} ions ofthe hydrochloric acid, and possibly because of the formation of acomplex anion made up of barium and chlorine; hence only the smallestexcess should be added over the amount required to acidify thesolution. ]

[Note 2: The ionic changes involved in the precipitation of bariumsulphate are very simple:

Ba^{++} + SO_{4}^{--} --> [BaSO_{4}]

This case affords one of the best illustrations of the effect of anexcess of a precipitant in decreasing the solubility of a precipitate. If the conditions are considered which exist at the moment when justenough of the Ba^{++} ions have been added to correspond to theSO_{4}^{--} ions in the solution, it will be seen that nearly all ofthe barium sulphate has been precipitated, and that the small amountwhich then remains in the solution which is in contact with theprecipitate must represent a saturated solution for the existingtemperature, and that this solution is comparable with a solution ofsugar to which more sugar has been added than will dissolve. Itshould be borne in mind that the quantity of barium sulphate inthis !saturated solution is a constant quantity! for the existingconditions. The dissolved barium sulphate, like any electrolyte, isdissociated, 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 constantvalue (which is very small), it may be eliminated, when the expressionbecomes !Conc'n Ba^{++} x Conc'n SO_{4}^{--} = Const. !, which isthe "solubility product" of BaSO_{4}. If, now, an excess of theprecipitant, a soluble barium salt, is added in the form of arelatively concentrated solution (the slight change of volume of a fewcubic centimeters may be disregarded for the present discussion)the concentration of the Ba^{++} ions is much increased, and as aconsequence the !Conc'n SO_{4}! must decrease in proportion if thevalue of the expression is to remain constant, which is a requisitecondition if the law of mass action upon which our argument dependsholds true. In other words, SO_{4}^{--} ions must combine with someof the added Ba^{++} ions to form [BaSO_{4}]; but it will be recalledthat the solution is already saturated with BaSO_{4}, and this freshlyformed quantity must, therefore, separate and add itself to theprecipitate. This is exactly what is desired in order to insuremore complete precipitation and greater accuracy, and leads to theconclusion that the larger the excess of the precipitant added themore successful the analysis; but a practical limit is placed uponthe quantity of the precipitant which may be properly added by otherconditions, 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 solutionfrom which it separates, even when these other substances arerelatively soluble, and including the barium chloride used as theprecipitant. This is also notably true in the case of nitrates andchlorates of the alkalies, and of ferric compounds; and, since in thisanalysis ammonium nitrate has resulted from the neutralization of theexcess of the nitric acid added to oxidize the iron, it is essentialthat this should be destroyed by repeated evaporation with arelatively large quantity of hydrochloric acid. During evaporation amutual decomposition of the two acids takes place, and the nitric acidis finally decomposed and expelled by the excess of hydrochloric acid. 

Iron is usually found in the precipitate of barium sulphate whenthrown down from hot solutions in the presence of ferric salts. This, according to Kuster and Thiel (!Zeit. Anorg. Chem. !, 22, 424), is dueto the formation of a complex ion (Fe(SO_{4})_{2}) which precipitateswith the Ba^{++} ion, while Richards (!Zeit. Anorg. Chem. !, 23, 383)ascribes it to hydrolytic action, which causes the formation of abasic ferric complex which is occluded in the barium precipitate. Whatever the character of the compound may be, it has been shown thatit loses sulphuric anhydride upon ignition, causing low results, eventhough the precipitate contains iron. 

The contamination of the barium sulphate by iron is much less in thepresence of ferrous than ferric salts. If, therefore, the sulphuralone were to be determined in the ferrous ammonium sulphate, theprecipitation by barium might be made directly from an aqueoussolution of the salt, which had been made slightly acid withhydrochloric acid. ]

[Note 4: The precipitation of the barium sulphate is probably completeat the end of a half-hour, and the solution may safely be filtered atthe expiration of that time if it is desired to hasten the analysis. 

As already noted, many precipitates of the general character of thissulphate tend to grow more coarsely granular if digested for some timewith the liquid from which they have separated. It is therefore wellto allow the precipitate to stand in a warm place for several hours, if practicable, to promote ease of filtration. The filtrate andwashings should always be carefully examined for minute quantities ofthe sulphate which may pass through the pores of the filter. This isbest accomplished by imparting to the filtrate a gentle rotary motion, when the sulphate, if present, will collect at the center of thebottom of the beaker. ]

[Note 5: A reduction of barium sulphate to the sulphide may veryreadily be caused by the reducing action of the burning carbon of thefilter, and much care should be taken to prevent any considerablereduction from this cause. Subsequent ignition, with ready accessof air, reconverts the sulphide to sulphate unless a considerablereduction has occurred. In the latter case it is expedient to add oneor two drops of sulphuric acid and to heat cautiously until the excessof acid is expelled. ]

[Note 6: Barium sulphate requires about 400, 000 parts of water forits 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 about0. 5 gram of the sulphate. Mix each in the crucible with five to sixtimes its weight of anhydrous sodium carbonate. This can best be doneby placing the crucible on a piece of glazed paper and stirring themixture with a clean, dry stirring-rod, which may finally be wiped offwith a small fragment of filter paper, the latter being placed in thecrucible. Cover the crucible and heat until a quiet, liquid fusionensues. Remove the burner, and tip the crucible until the fused massflows nearly to its mouth. Hold it in that position until the mass hassolidified. When cold, the material may usually be detached in a lumpby tapping the crucible or gently pressing it near its upper edge. Ifit still adheres, a cubic centimeter or so of water may be placed inthe cold crucible and cautiously brought to boiling, when the cakewill 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 furtherwill dissolve, decant the solution from the residue of bariumcarbonate (Note 1). Pour over the residue 20 cc. Of a solution ofsodium carbonate and 10 cc. Of water and heat to gentle boiling forabout three minutes (Note 2). Filter off the carbonate and wash itwith hot water, testing the slightly acidified washings for sulphateand preserving any precipitates which appear in these tests. Acidifythe filtrate with hydrochloric acid until just acid, bring to boiling, and slowly add hot barium chloride solution, as in the precedingdetermination. Add also any tests from the washings in whichprecipitates have appeared. Filter, wash, ignite, and weigh. 

From the weight of barium sulphate, calculate the percentage ofsulphur (S) in the sample. 

[Note 1: This alkaline fusion is much employed to disintegratesubstances ordinarily insoluble in acids into two components, oneof which is water soluble and the other acid soluble. The reactioninvolved is:

BaSO_{4} + Na_{2}CO_{3}, --> BaCO_{3}, + Na_{2}SO_{4}. 

As the sodium sulphate is soluble in water, and the barium carbonateinsoluble, a separation between them is possible and the sulphur canbe determined in the water-soluble portion. 

It should be noted that this method can be applied to the purificationof a precipitate of barium sulphate if contaminated by most of thesubstances mentioned in Note 3 on page 114. The impurities pass intothe water solution together with the sodium sulphate, but, beingpresent in such minute amounts, do not again precipitate with thebarium sulphate. ]

[Note 2: The barium carbonate is boiled with sodium carbonate solutionbefore filtration because the reaction above is reversible; and it isonly by keeping the sodium carbonate present in excess until nearlyall of the sodium sulphate solution has been removed by filtrationthat the reversion of some of the barium carbonate to barium sulphateis prevented. This is an application of the principle of mass action, in which the concentration of the reagent (the carbonate ion) iskept as high as practicable and that of the sulphate ion as low aspossible, in order to force the reaction in the desired direction (seeAppendix). ]

DETERMINATION OF PHOSPHORIC ANHYDRIDE IN APATITE

The mineral apatite is composed of calcium phosphate, associated withcalcium chloride, or fluoride. Specimens are easily obtainable whichare nearly pure and leave on treatment with acid only a slightsiliceous residue. 

For the purpose of gravimetric determination, phosphoric acid isusually precipitated from ammoniacal solutions in the form ofmagnesium ammonium phosphate which, on ignition, is converted intomagnesium pyrophosphate. Since the calcium phosphate of the apatiteis also insoluble in ammoniacal solutions, this procedure cannot beapplied directly. The separation of the phosphoric acid from thecalcium must first be accomplished by precipitation in the form ofammonium phosphomolybdate in nitric acid solution, using ammoniummolybdate as the precipitant. The "yellow precipitate, " as it is oftencalled, is not always of a definite composition, and therefore notsuitable for direct weighing, but may be dissolved in ammonia, and thephosphoric acid thrown out as magnesium ammonium phosphate from thesolution. 

Of the substances likely to occur in apatite, silicic acid aloneinterferes with the precipitation of the phosphoric acid in nitricacid solution. 

PRECIPITATION OF AMMONIUM PHOSPHOMOLYBDATE

PROCEDURE. --Grind the mineral in an agate mortar until no grit isperceptible. Transfer the substance to a weighing-tube, and weigh outtwo portions, not exceeding 0. 20 gram each (Note 1) into two beakersof about 200 cc. Capacity. Pour over them 20 cc. Of dilute nitric acid(sp. Gr. 1. 2) and warm gently until solvent action has apparentlyceased. Evaporate the solution cautiously to dryness, heat the residuefor about an hour at 100-110°C. , and treat it again with nitric acidas described above; separate the residue of silica by filtration ona small filter (7 cm. ) and wash with warm water, using as little aspossible (Note 2). Receive the filtrate in a beaker (200-500 cc. ). Test the washings with ammonia for calcium phosphate, but add all suchtests in which a precipitate appears to the original nitrate (Note 3). The filtrate and washings must be kept as small as possible and shouldnot exceed 100 cc. In volume. Add aqueous ammonia (sp. Gr. 0. 96) untilthe precipitate of calcium phosphate first produced just fails toredissolve, and then add a few drops of nitric acid until this isagain brought into solution (Note 4). Warm the solution until itcannot be comfortably held in the hand (about 60°C. ) and, afterremoval of the burner, add 75 cc. Of ammonium molybdate solution whichhas been !gently! warmed, but which must be perfectly clear. Allowthe mixture to stand at a temperature of about 50 or 60°C. For twelvehours (Notes 5 and 6). Filter off the yellow precipitate on a 9 cm. Filter, and wash by decantation with a solution of ammonium nitratemade acid with nitric acid. [1] Allow the precipitate to remain in thebeaker as far as possible. Test the washings for calcium with ammoniaand ammonium oxalate (Note 3). 

[Footnote 1: This solution is prepared as follows: Mix 100 cc. Ofammonia 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 fora 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 slightlysoluble under the conditions of operation. Consequently theunavoidable errors of analysis are greater in this determination thanin those which have preceded it, and some divergence may be expectedin duplicate analyses. It is obvious that the larger the amount ofsubstance taken for analysis the less will be the relative loss orgain due to unavoidable experimental errors; but, in this instance, acheck is placed upon the amount of material which may be taken both bythe bulk of the resulting precipitate of ammonium phosphomolybdateand by the excessive amount of ammonium molybdate required to effectcomplete separation of the phosphoric acid, since a liberal excessabove the theoretical quantity is demanded. Molybdic acid is one ofthe more expensive reagents. ]

[Note 2: Soluble silicic acid would, if present, partially separatewith the phosphomolybdate, although not in combination withmolybdenum. Its previous removal by dehydration is thereforenecessary. ]

[Note 3: When washing the siliceous residue the filtrate may be testedfor calcium by adding ammonia, since that reagent neutralizes theacid which holds the calcium phosphate in solution and causesprecipitation; but after the removal of the phosphoric acid incombination with the molybdenum, the addition of an oxalate isrequired to show the presence of calcium. ]

[Note 4: An excess of nitric acid exerts a slight solventaction, while ammonium nitrate lessens the solubility; hence theneutralization of the former by ammonia. ]

[Note 5: The precipitation of the phosphomolybdate takes place morepromptly in warm than in cold solutions, but the temperature shouldnot exceed 60°C. During precipitation; a higher temperature tends toseparate molybdic acid from the solution. This acid is nearly white, and its deposition in the filtrate on long standing should not bemistaken for a second precipitation of the yellow precipitate. Theaddition of 75 cc. Of ammonium molybdate solution insures the presenceof a liberal excess of the reagent, but the filtrate should be testedas in all quantitative procedures. 

The precipitation is probably complete in many cases in less thantwelve hours; but it is better, when practicable, to allow thesolution to stand for this length of time. Vigorous shaking orstirring promotes the separation of the precipitate. ]

[Note 6: The composition of the "yellow precipitate" undoubtedlyvaries 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 theconditions prescribed in the procedure. Whatever other variations mayoccur in its composition, the ratio of 12 MoO_{3}:1 P seems tohold, and this fact is utilized in volumetric processes for thedetermination of phosphorus, in which the molybdenum is reduced toa lower oxide and reoxidized by a standard solution of potassiumpermanganate. In principle, the procedure is comparable with thatdescribed for the determination of iron by permanganate. ]

PRECIPITATION OF MAGNESIUM AMMONIUM PHOSPHATE

PROCEDURE. --Dissolve the precipitate of phosphomolybdate upon thefilter by pouring through it dilute aqueous ammonia (one volume ofdilute ammonia (sp. Gr. 0. 96) and three volumes of water, whichshould be carefully measured), and receive the solution in the beakercontaining the bulk of the precipitate. The total volume of nitrateand washings should not much exceed 100 cc. Acidify the solution withdilute hydrochloric acid, and heat it nearly to boiling. Calculate thevolume of magnesium ammonium chloride solution ("magnesia mixture")required to precipitate the phosphoric acid, assuming 40 per centP_{2}O_{5} in the apatite. Measure out about 5 cc. In excess of thisamount, and pour it into the acid solution. Then add slowly diluteammonium hydroxide (1 volume of strong ammonia (sp. Gr. 0. 90) and 9volumes of water), stirring constantly until a precipitate forms. Thenadd a volume of filtered, concentrated ammonia (sp. Gr. 0. 90) equal toone third of the volume of liquid in the beaker (Note 1). Allow thewhole to cool. The precipitated magnesium ammonium phosphate shouldthen be definitely crystalline in appearance (Note 2). (If it isdesired to hasten the precipitation, the solution may be cooled, firstin cold and then in ice-water, and stirred !constantly! for half anhour, when precipitation will usually be complete. )

Decant the clear liquid through a filter, and transfer the precipitateto the filter, using as wash-water a mixture of one volume ofconcentrated ammonia and three volumes of water. It is not necessaryto clean the beaker completely or to wash the precipitate thoroughlyat this point, as it is necessary to purify it by reprecipitation. 

[Note 1: Magnesium ammonium phosphate is not a wholly insolublesubstance, even under the most favorable analytical conditions. Itis least soluble in a liquid containing one fourth of its volume ofconcentrated aqueous ammonia (sp. Gr. 0. 90) and this proportion shouldbe carefully maintained as prescribed in the procedure. On account ofthis slight solubility the volume of solutions should be kept as smallas possible and the amount of wash-water limited to that absolutelyrequired. 

A large excess of the magnesium solution tends both to throw outmagnesium hydroxide (shown by a persistently flocculent precipitate)and to cause the phosphate to carry down molybdic acid. The tendencyof the magnesium precipitate to carry down molybdic acid is alsoincreased if the solution is too concentrated. The volume should notbe less than 90 cc. , nor more than 125 cc. , at the time of the firstprecipitation with the magnesia mixture. ]

[Note 2: The magnesium ammonium phosphate should be perfectlycrystalline, and will be so if the directions are followed. The slowaddition of the reagent is essential, and the stirring not less so. Stirring promotes the separation of the precipitate and the formationof larger crystals, and may therefore be substituted for digestion inthe cold. The stirring-rod must not be allowed to scratch the glass, as the crystals adhere to such scratches and are removed withdifficulty. ]

REPRECIPITATION AND IGNITION OF MAGNESIUM AMMONIUM PHOSPHATE

A single precipitation of the magnesium compound in the presence ofmolybdenum compounds rarely yields a pure product. The molybdenum canbe removed by solution of the precipitate in acid and precipitationof the molybdenum by sulphureted hydrogen, after which the magnesiumprecipitate may be again thrown down. It is usually more satisfactoryto dissolve the magnesium precipitate and reprecipitate the phosphateas magnesium ammonium phosphate as described below. 

PROCEDURE. --Dissolve the precipitate from the filter in a littledilute hydrochloric acid (sp. Gr. 1. 12), allowing the acid solution torun into the beaker in which the original precipitation was made (Note1). Wash the filter with water until the wash-water shows no test forchlorides, but avoid an unnecessary amount of wash-water. Add tothe solution 2 cc. (not more) of magnesia mixture, and then diluteammonium hydroxide solution (sp. Gr. 0. 96), drop by drop, withconstant stirring, until the liquid smells distinctly of ammonia. Stirfor 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 thesolution to stand for some hours, and then filter off the magnesiumammonium phosphate, which should be distinctly crystalline incharacter. Wash the precipitate with dilute ammonia water, asprescribed above, until, finally, 3 cc. Of the washings, afteracidifying with nitric acid, show no evidence of chlorides. Test bothfiltrates for complete precipitation by adding a few cubic centimetersof magnesia mixture and allowing them to stand for some time. 

Transfer the moist precipitate to a weighed porcelain or platinumcrucible and ignite, using great care to raise the temperature slowlywhile drying the filter in the crucible, and to insure the readyaccess of oxygen during the combustion of the filter paper, thusguarding against a possible reduction of the phosphate, which wouldresult in disastrous consequences both to the crucible, if ofplatinum, and the analysis. Do not raise the temperature abovemoderate redness until the precipitate is white. (Keep this precautionwell in mind. ) Ignite finally at the highest temperature of theTirrill burner, and repeat the heating until the weight is constant. If the ignited precipitate is persistently discolored by particles ofunburned carbon, moisten the mass with a drop or two of concentratednitric acid and heat cautiously, finally igniting strongly. Theacid will dissolve magnesium pyrophosphate from the surface of theparticles of carbon, which will then burn away. Nitric acid also aidsas an oxidizing agent in supplying oxygen for the combustion of thecarbon. 

From the weight of magnesium pyrophosphate (Mg_{2}P_{2}O_{7})obtained, calculate the phosphoric anhydride (P_{2}O_{5}) in thesample of apatite. 

[Note 1: The ionic change involved in the precipitation of themagnesium compound is

PO_{4}^{---} + NH_{4}^{+} + Mg^{++} --> [MgNH_{4}PO_{4}]. 

The magnesium ammonium phosphate is readily dissolved by acids, eventhose which are no stronger than acetic acid. This is accounted forby the fact that two of the ions into which phosphoric acid maydissociate, the HPO_{4}^{--} or H_{2}PO_{4}^{-} ions, exhibit thecharacteristics of very weak acids, in that they show almost notendency to dissociate further into H^{+} and PO_{4}^{--} ions. Consequently the ionic changes which occur when the magnesium ammoniumphosphate is brought into contact with an acid may be typified by thereaction:

H^{+} + Mg^{++} + NH_{4}^{+} + PO_{4}^{---} --> Mg^{++} + NH_{4}^{+} +HPO_{4}^{--};

that is, the PO_{4}^{--} ions and the H^{+} ions lose their identityin the formation of the new ion, HPO_{4}^{--}, and this continuesuntil the magnesium ammonium phosphate is entirely dissolved. ]

[Note 2: During ignition the magnesium ammonium phosphate losesammonia 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 withgreat care during the ignition of this precipitate. The danger herelies in a possible reduction of the phosphate by the carbon of thefilter paper, or by the ammonia evolved, which may act as a reducingagent. 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 marblethrough the dolomitic limestones, containing varying amounts ofmagnesium, to the impure varieties, which contain also ferrous andmanganous 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 ofsulphides or sulphates, phosphates, and titanates, and the alkali oreven the heavy metals. No attempt is made in the following proceduresto provide a complete quantitative scheme which would take intoaccount all of these constituents. Such a scheme for a completeanalysis of a limestone may be found in Bulletin No. 700 of the UnitedStates Geological Survey. It is assumed that, for these practicedeterminations, a limestone is selected which contains only the morecommon constituents first enumerated above. 

DETERMINATION OF MOISTURE

The determination of the amount of moisture in minerals or ores isoften of great importance. Ores which have been exposed to the weatherduring shipment may have absorbed enough moisture to appreciablyaffect the results of analysis. Since it is essential that the sellerand buyer should make their analyses upon comparable material, it iscustomary for each analyst to determine the moisture in the sampleexamined, and then to calculate the percentages of the variousconstituents with reference to a sample dried in the air, or at atemperature a little above 100°C. , which, unless the ore has undergonechemical change because of the wetting, should be the same before andafter shipment. 

PROCEDURE. --Spread 25 grams of the powdered sample on a weighedwatch-glass; weigh to the nearest 10 milligrams only and heat at105°C. ; weigh at intervals of an hour, after cooling in a desiccator, until the loss of weight after an hour's heating does not exceed10 milligrams. It should be noted that a variation in weight of 10milligrams in a total weight of 25 grams is no greater relatively thana 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 thepowder 25 cc. Of water, and then add 50 cc. Of dilute hydrochloricacid (sp. Gr. 1. 12) in small portions, warming gently, until nothingfurther appears to dissolve (Note 2). Evaporate to dryness on thewater bath. Pour over the residue a mixture of 5 cc. Of water and5 cc. Of concentrated hydrochloric acid (sp. Gr. 1. 2) and againevaporate to dryness, and finally heat for at least an hour ata temperature of 110°C. Pour over this residue 50 cc. Of dilutehydrochloric acid (one volume acid (sp. Gr. 1. 12) to five volumeswater), and boil for about five minutes; then filter and wash twicewith the dilute hydrochloric acid, and then with hot water untilfree from chlorides. Transfer the filter and contents to a porcelaincrucible, dry carefully over a low flame, and ignite to constantweight. The residue represents the insoluble matter and the silicafrom any soluble silicates (Note 3). 

Calculate the combined percentage of these in the limestone. 

[Note 1: The relatively large weight (5 grams) taken for analysisinsures greater accuracy in the determination of the ingredients whichare present in small proportions, and is also more likely to be arepresentative sample of the material analyzed. ]

[Note 2: It is plain that the amount of the insoluble residue and alsoits character will often depend upon the strength of acid used forsolution of the limestone. It cannot, therefore, be regarded asrepresenting any well-defined constituent, and its determination isessentially empirical. ]

[Note 3: It is probable that some of the silicates present are whollyor partly decomposed by the acid, and the soluble silicic acid mustbe 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 waterbath for a short time tends to leave the silica in the form of apowder, which promotes subsequent filtration. The siliceous residueis washed first with dilute acid to prevent hydrolytic changes, whichwould result in the formation of appreciable quantities of insolublebasic iron or aluminium salts on the filter when washing with hotwater. 

If it is desired to determine the percentage of silica separately, theignited residue should be mixed in a platinum crucible with about sixtimes its weight of anhydrous sodium carbonate, and the proceduregiven on page 151 should be followed. The filtrate from the silica isthen 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 ammoniumhydroxide until the solution just smells distinctly of ammonia, but donot 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 brominewater, and heat again for a few minutes. Finally add 10 cc. Ofammonium chloride solution and keep the solution warm until it barelysmells of ammonia; then filter promptly (Note 2). Wash the filtertwice with hot water, then (after replacing the receiving beaker) pourthrough it 25 cc. Of hot, dilute hydrochloric acid (one volume diluteHCl [sp. Gr. 1. 12] to five volumes water). A brown residue insolublein the acid may be allowed to remain on the filter. Wash the filterfive times with hot water, add to the filtrate ammonium hydroxideand bromine water as described above, and repeat the precipitation. Collect the precipitate on the filter already used, wash it free fromchlorides with hot water, and ignite and weigh as described for ferrichydroxide on page 110. The residue after ignition consists of ferricoxide, alumina, and mangano-manganic oxide (Mn_{3}O_{4}), if manganeseis 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 solutionsserves to oxidize any ferrous hydroxide to ferric hydroxide and toprecipitate manganese as MnO(OH)_{2}. The solution must contain notmore than a bare excess of hydroxyl ions (ammonium hydroxide) when itis filtered, on account of the tendency of the aluminium hydroxide toredissolve. 

The solution should not be strongly ammoniacal when the bromine isadded, as strong ammonia reacts with the bromine, with the evolutionof nitrogen. ]

[Note 2: The precipitate produced by ammonium hydroxide and bromineshould be filtered off promptly, since the alkaline solution absorbscarbon dioxide from the air, with consequent partial precipitationof the calcium as carbonate. This is possible even under the mostfavorable conditions, and for this reason the iron precipitate isredissolved and again precipitated to free it from calcium. When theprecipitate is small, this reprecipitation may be omitted. ]

[Note 3: In the absence of significant amounts of manganese the ironand aluminium may be separately determined by fusion of the mixedignited precipitate, after weighing, with about ten times its weightof 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 theweight of ferric oxide corresponding to the amount of iron found. 

If a separate determination of the iron, aluminium, and manganeseis desired, the mixed precipitate may be dissolved in acid beforeignition, and the separation effected by special methods (see, forexample, Fay, !Quantitative Analyses!, First Edition, pp. 15-19 and23-27). ]

DETERMINATION OF CALCIUM

PROCEDURE. --To the combined filtrates from the double precipitation ofthe 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, andfill the flask with distilled water until the lowest point of themeniscus is exactly level with the mark on the neck of the flask. Carefully remove any drops of water which are on the inside of theneck of the flask above the graduation by means of a strip of filterpaper, make the solution uniform by pouring it out into a dry beakerand back into the flask several times. Measure off one fifth of thissolution as follows (Note 1): Pour into a 100 cc. Graduated flaskabout 10 cc. Of the solution, shake the liquid thoroughly over theinner surface of the small flask, and pour it out. Repeat the sameoperation. Fill the 100 cc. Flask until the lowest point of themeniscus is exactly level with the mark on its neck, remove any dropsof solution from the upper part of the neck with filter paper, andpour the solution into a beaker (400-500 cc. ). Wash out the flask withsmall quantities of water until it is clean, adding these to the 100cc. Of solution. When the duplicate portion of 100 cc. Is measured outfrom the solution, remember that the flask must be rinsed out twicewith that solution, as prescribed above, before the measurement ismade. (A 100 cc. Pipette may be used to measure out the aliquotportions, if preferred. )

Dilute each of the measured portions to 250 cc. With distilled water, heat the whole to boiling, and add ammonium oxalate solution slowlyin moderate excess, stirring well. Boil for two minutes; allow theprecipitated calcium oxalate to settle for a half-hour, and decantthrough a filter. Test the filtrate for complete precipitation byadding a few cubic centimeters of the precipitant, allowing it tostand for fifteen minutes. If no precipitate forms, make the solutionslightly acid with hydrochloric acid (Note 2); see that it is properlylabeled and reserve it to be combined with the filtrate from thesecond calcium oxalate precipitation (Notes 3 and 4). 

Redissolve the calcium oxalate in the beaker with warm hydrochloricacid, pouring the acid through the filter. Wash the filter five timeswith water, and finally pour through it aqueous ammonia. Dilute thesolution to 250 cc. , bring to boiling, and add 1 cc. Ammonium oxalatesolution (Note 5) and ammonia in slight excess; boil for two minutes, and set aside for a half-hour. Filter off the calcium oxalate upon thefilter first used, and wash free from chlorides. The filtrate shouldbe made barely acid with hydrochloric acid and combined with thefiltrate from the first precipitation. Begin at once the evaporationof the solutions for the determination of magnesium as describedbelow. 

The precipitate of calcium oxalate may be converted into calcium oxideby ignition without previous drying. After burning the filter, it maybe ignited for three quarters of an hour in a platinum crucible atthe highest heat of the Bunsen or Tirrill burner, and finally for tenminutes at the blast lamp (Note 6). Repeat the heating over the blastlamp until the weight is constant. As the calcium oxide absorbsmoisture from the air, it must (after cooling) be weighed as rapidlyas possible. 

The precipitate may, if preferred, be placed in a weighted porcelaincrucible. After burning off the filter and heating for ten minutes thecalcium precipitate may be converted into calcium sulphate by placing2 cc. Of dilute sulphuric acid in the crucible (cold), heating thecovered crucible very cautiously over a low flame to drive off theexcess of acid, and finally at redness to constant weight (Note 7). 

From the weight of the oxide or sulphate, calculate the percentage ofthe calcium (Ca) in the limestone, remembering that only one fifth ofthe 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 beproperly treated. The solution is, therefore, diluted to a definitevolume (500 cc. ), and exactly one fifth (100 cc. ) is measured off in agraduated flask or by means of a pipette. ]

[Note 2: The filtrate from the calcium oxalate should be made slightlyacid immediately after filtration, in order to avoid the solventaction of the alkaline liquid upon the glass. ]

[Note 3: The accurate quantitative separation of calcium and magnesiumas oxalates requires considerable care. The calcium precipitateusually carries down with it some magnesium, and this can bestbe removed by redissolving the precipitate after filtration, andreprecipitation in the presence of only the small amount of magnesiumwhich was included in the first precipitate. When, however, theproportion of magnesium is not very large, the second precipitation ofthe calcium can usually be avoided by precipitating it from a ratherdilute solution (800 cc. Or so) and in the presence of a considerableexcess of the precipitant, that is, rather more than enough to convertboth the magnesium and calcium into oxalates. ]

[Note 4: The ionic changes involved in the precipitation of calciumas oxalate are exceedingly simple, and the principles discussed inconnection with the barium sulphate precipitation on page 113 alsoapply here. The reaction is

C_{2}O_{4}^{--} + Ca^{++} --> [CaC_{2}O_{4}]. 

Calcium oxalate is nearly insoluble in water, and only very slightlysoluble in acetic acid, but is readily dissolved by the strong mineralacids. This behavior with acids is explained by the fact that oxalicacid is a stronger acid than acetic acid; when, therefore, the oxalateis brought into contact with the latter there is almost no tendency todiminish the concentration of C_{2}O_{4}^{--} ions by the formation ofan acid less dissociated than the acetic acid itself, and practicallyno solvent action ensues. When a strong mineral acid is present, however, the ionization of the oxalic acid is much reduced by the highconcentration of the H^{+} ions from the strong acid, the formationof the undissociated acid lessens the concentration of theC_{2}O_{4}^{--} ions in solution, more of the oxalate passes intosolution to re-establish equilibrium, and this process repeats itselfuntil all is dissolved. 

The oxalate is immediately reprecipitated from such a solution on theaddition of OH^{-} ions, which, by uniting with the H^{+} ions of theacids (both the mineral acid and the oxalic acid) to form water, leavethe Ca^{++} and C_{2}O_{4}^{--} ions in the solution to recombine toform [CaC_{2}O_{4}], which is precipitated in the absence of theH^{+} ions. It is well at this point to add a small excess ofC_{2}O_{4}^{--} ions in the form of ammonium oxalate to decrease thesolubility of the precipitate. 

The oxalate precipitate consists mainly of CaC_{2}O_{4}. H_{2}O whenthrown down. ]

[Note 5: The small quantity of ammonium oxalate solution is addedbefore the second precipitation of the calcium oxalate to insurethe presence of a slight excess of the reagent, which promotes theseparation of the calcium compound. ]

[Note 6: On ignition the calcium oxalate loses carbon dioxide andcarbon 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 maybe brought about in a platinum crucible at the highest temperature ofa Tirrill burner, but it is well to ignite larger quantities than thisover the blast lamp until the weight is constant. ]

[Note 7: The heat required to burn the filter, and that subsequentlyapplied as described, will convert most of the calcium oxalate tocalcium 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 tosulphate is to be preferred, as a complete conversion to oxide isdifficult to accomplish. ]

[Note 8: The determination of the calcium may be completedvolumetrically by washing the calcium oxalate precipitate fromthe filter into dilute sulphuric acid, warming, and titratingthe liberated oxalic acid with a standard solution of potassiumpermanganate as described on page 72. When a considerable number ofanalyses are to be made, this procedure will save much of the timeotherwise required for ignition and weighing. ]

DETERMINATION OF MAGNESIUM

PROCEDURE. --Evaporate the acidified filtrates from the calciumprecipitates until the salts begin to crystallize, but do !not!evaporate to dryness (Note 1). Dilute the solution cautiously untilthe salts are brought into solution, adding a little acid if thesolution has evaporated to very small volume. The solution should becarefully examined at this point and must be filtered if a precipitatehas appeared. Heat the clear solution to boiling; remove the burnerand add 25 cc. Of a solution of disodium phosphate. Then add slowlydilute ammonia (1 volume strong ammonia (sp. Gr. 0. 90) and 9 volumeswater) as long as a precipitate continues to form. Finally, add avolume of concentrated ammonia (sp. Gr. 0. 90) equal to one third ofthe volume of the solution, and allow the whole to stand for abouttwelve hours. 

Decant the solution through a filter, wash it with dilute ammoniawater, proceeding as prescribed for the determination of phosphoricanhydride on page 122, including; the reprecipitation (Note 2), except that 3 cc. Of disodium phosphate solution are added before thereprecipitation of the magnesium ammonium phosphate instead ofthe magnesia mixture there prescribed. From the weight of thepyrophosphate, calculate the percentage of magnesium oxide (MgO) inthe sample of limestone. Remember that the pyrophosphate finallyobtained is from one fifth of the original sample. 

[Note 1: The precipitation of the magnesium should be made in as smallvolume as possible, and the ratio of ammonia to the total volume ofsolution should be carefully provided for, on account of the relativesolubility of the magnesium ammonium phosphate. This matter hasbeen fully discussed in connection with the phosphoric anhydridedetermination. ]

[Note 2: The first magnesium ammonium phosphate precipitate is rarelywholly crystalline, as it should be, and is not always of the propercomposition when precipitated in the presence of such large amounts ofammonium salts. The difficulty can best be remedied by filtering theprecipitate and (without washing it) redissolving in a small quantityof hydrochloric acid, from which it may be again thrown down byammonia after adding a little disodium phosphate solution. If theflocculent character was occasioned by the presence of magnesiumhydroxide, the second precipitation, in a smaller volume containingfewer salts, will often result more favorably. 

The removal of iron or alumina from a contaminated precipitate isa matter involving a long procedure, and a redetermination of themagnesium from a new sample, with additional precautions, is usuallyto be preferred. ]

DETERMINATION OF CARBON DIOXIDE

!Absorption Apparatus!

[Illustration: Fig. 3]

The apparatus required for the determination of the carbon dioxideshould be arranged as shown in the cut (Fig. 3). The flask (A) isan ordinary wash bottle, which should be nearly filled with dilutehydrochloric 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 itslower end bent upward and drawn out to small bore, so that the carbondioxide evolved from the limestone cannot bubble back into (b). Theevolution flask should preferably be a wide-mouthed Soxhlet extractionflask of about 150 cc. Capacity because of the ease with which tubesand stoppers may be fitted into the neck of a flask of this type. Theflask should be fitted with a two-hole rubber stopper. The condenser(C) may consist of a tube with two or three large bulbs blown init, for use as an air-cooled condenser, or it may be a smallwater-jacketed condenser. The latter is to be preferred if a number ofdeterminations 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 concentratedsulphuric acid (sp. Gr. 1. 84). The short rubber tubing (d) connectsthe first U-tube to a second U-tube (E) which is filled with smalldust-free lumps of dry calcium chloride, with a small, loose plug ofcotton at the top of each arm. Both tubes should be closed by corkstoppers, the tops of which are cut off level with, or preferablyforced a little below, the top of the U-tube, and then neatly sealedwith sealing wax. 

The carbon dioxide may be absorbed in a tube containing soda lime(F) or in a Geissler bulb (F') containing a concentrated solutionof potassium hydroxide (Note 2). The tube (F) is a glass-stopperedside-arm U-tube in which the side toward the evolution flask and onehalf of the other side are filled with small, dust-free lumps of sodalime of good quality (Note 3). Since soda lime contains considerablemoisture, the other half of the right side of the tube is filled withsmall lumps of dry, dust-free calcium chloride to retain the moisturefrom the soda lime. Loose plugs of cotton are placed at the top ofeach arm and between the soda lime and the calcium chloride. 

The Geissler bulb (F'), if used, should be filled with potassiumhydroxide solution (1 part of solid potassium hydroxide dissolved intwo parts of water) until each small bulb is about two thirds full(Note 4). A small tube containing calcium chloride is connected withthe Geissler bulb proper by a ground joint and should be wired to thebulb for safety. This is designed to retain any moisture from thehydroxide solution. A piece of clean, fine copper wire is so attachedto the bulb that it can be hung from the hook above a balance pan, orother 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. Thiswill prevent the absorption of water vapor by (F) or (F') and servesas an aid in regulating the flow of air through the apparatus. (H) isan aspirator bottle of about four liters capacity, filled with water;(k) is a safety tube and a means of refilling (H); (h) is a screwclamp, and (K) a U-tube filled with soda lime. 

[Note 1: The air current, which is subsequently drawn through theapparatus, to sweep all of the carbon dioxide into the absorptionapparatus, is likely to carry with it some hydrochloric acid fromthe evolution flask. This acid is retained by the silver sulphatesolution. The addition of concentrated sulphuric acid to this solutionreduces its vapor pressure so far that very little water is carried onby the air current, and this slight amount is absorbed by the calciumchloride in (E). As the calcium chloride frequently contains a smallamount of a basic material which would absorb carbon dioxide, it isnecessary to pass carbon dioxide through (E) for a short time and thendrive all the gas out with a dry air current for thirty minutes beforeuse. ]

[Note 2: Soda-lime absorption tubes are to be preferred if asatisfactory quality of soda lime is available and the number ofdeterminations to be made successively is small. The potash bulbs willusually permit of a larger number of successive determinations withoutrefilling, but they require greater care in handling and in theanalytical procedure. ]

[Note 3: Soda lime is a mixture of sodium and calcium hydroxides. Bothcombine with carbon dioxide to form carbonates, with the evolutionof water. Considerable heat is generated by the reaction, and thetemperature of the tube during absorption serves as a rough index ofthe progress of the reaction through the mass of soda lime. 

It is essential that soda lime of good quality for analytical purposesshould be used. The tube should not contain dust, as this is likely tobe swept away. ]

[Note 4: The solution of the hydroxide for use in the Geissler bulbmust be highly concentrated to insure complete absorption of thecarbon dioxide and also to reduce the vapor pressure of the solution, thus lessening the danger of loss of water with the air which passesthrough the bulbs. The small quantity of moisture which is thencarried out of the bulbs is held by the calcium chloride in theprolong tube. The best form of absorption bulb is that to which theprolong tube is attached by a ground glass joint. 

After the potassium hydroxide is approximately half consumed in thefirst bulb of the absorption apparatus, potassium bicarbonate isformed, and as it is much less soluble than the carbonate, it oftenprecipitates. Its formation is a warning that the absorbing power ofthe 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 thebalance case, and wiping it carefully with a lintless cloth, takingcare to handle it as little as possible after wiping (Note 1). Connectthe absorption apparatus with (e) and (f). If a soda-lime tube isused, 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 rubbertube from the flask (A), making sure that the tubes (a) and (b) do notcontain any hydrochloric acid, close the pinchcocks (a) and (k) andopen (h). No bubbles should pass through (D) or (G) after a fewseconds. When assured that the fittings are tight, close (h) and open(a) cautiously to admit air to restore atmospheric pressure. Thisprecaution is essential, as a sudden inrush of air will project liquidfrom (D) or (F'). Reconnect the rubber tube with the flask (A). Open the pinchcocks (a) and (k) and blow over about 10 cc. Of thehydrochloric acid from (A) into (B). When the action of the acidslackens, blow over (slowly) another 10 cc. 

The rate of gas evolution should not exceed for more than a fewseconds that at which about two bubbles per second pass through (G)(Note 2). Repeat the addition of acid in small portions until theaction upon the limestone seems to be at an end, taking care to close(a) after each addition of acid (Note 3). Disconnect (A) and connectthe 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 waythat about two bubbles per second pass through (G). Place a smallflame under (B) and !slowly! raise the contents to boiling and boilfor three minutes. Then remove the burner from under (B) and continueto draw air through the apparatus for 20-30 minutes, or until (H)is emptied (Note 4). Remove the absorption apparatus, closing thestopcocks on (F) or stoppering the open ends of (F'), leave theapparatus in the balance case for at least thirty minutes, wipe itcarefully and weigh, after opening the stopcocks (or removing plugs). The increase in weight is due to absorption of CO_{2}, from which itspercentage in the sample may be calculated. 

After cleaning (B) and refilling (H), the apparatus is ready for theduplicate analysis. 

[Note 1: The absorption tubes or bulbs have large surfaces on whichmoisture may collect. By allowing them to remain in the balance casefor some time before weighing, the amount of moisture absorbed on thesurface is as nearly constant as practicable during two weighings, anda uniform temperature is also assured. The stopcocks of the U-tubeshould be opened, or the plugs used to close the openings of theGeissler bulb should be removed before weighing in order that the aircontents 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 completelyretained by the absorbents. ]

[Note 3: The essential ionic changes involved in this procedure arethe following: It is assumed that the limestone, which is typified bycalcium carbonate, is very slightly soluble in water, and the ionsresulting are Ca^{++} and CO_{3}^{--}. In the presence of H^{+} ionsof the mineral acid, the CO_{3}^{--} ions form [H_{2}CO_{3}]. Thisis not only a weak acid which, by its formation, diminishes theconcentration of the CO_{3}^{--} ions, thus causing more of thecarbonate to dissolve to re-establish equilibrium, but it is also anunstable compound and breaks down into carbon dioxide and water. ]

[Note 4: Carbon dioxide is dissolved by cold water, but the gas isexpelled by boiling, and, together with that which is distributedthrough the apparatus, is swept out into the absorption bulb by thecurrent of air. This air is purified by drawing it through the tube(K) containing soda lime, which removes any carbon dioxide which maybe 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 toanother through solutions of electrolytes, the individual ions presentin these solutions tend to move toward the electrode of oppositeelectrical charge to that which each ion bears, and to be dischargedby that electrode. Whether or not such discharge actually occurs inthe case of any particular ion depends upon the potential (voltage) ofthe current which is passing through the solution, since for each ionthere is, under definite conditions, a minimum potential below whichthe discharge of the ion cannot be effected. By taking advantageof differences in discharge-potentials, it is possible to effectseparations of a number of the metallic ions by electrolysis, and atthe same time to deposit the metals in forms which admit of directweighing. In this way the slower procedures of precipitation andfiltration may frequently be avoided. The following paragraphs presenta brief statement of the fundamental principles and conditionsunderlying electro-analysis. 

The total energy of an electric current as it passes through asolution 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 measuredin amperes, and corresponds to the volume of water passing across-section of a stream in a given time interval; and third, theresistance of the conducting medium, which is measured in ohms. Therelation between these three factors is expressed by Ohm's law, namely, that !I = E/R!, when I is current strength, E potential, and Rresistance. It is plain that, for a constant resistance, thestrength of the current and its potential are mutually and directlyinterdependent. 

As already stated, the applied electrical potential determines whetheror not deposition of a metal upon an electrode actually occurs. Thecurrent strength determines the rate of deposition and the physicalcharacteristics of the deposit. The resistance of the solution isgenerally so small as to fall out of practical consideration. 

Approximate deposition-potentials have been determined for a numberof the metallic elements, and also for hydrogen and some of theacid-forming radicals. The values given below are those requiredfor deposition from normal solutions at ordinary temperatureswith reference to a hydrogen electrode. They must be regarded asapproximate, since several disturbing factors and some secondaryreactions render difficult their exact application under theconditions 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 anormal solution of copper sulphate a minimum potential equal to thealgebraic sum of the deposition-potentials of copper ions and sulphateions must be applied, that is, +1. 56 volts. The deposition of zincfrom a solution of zinc sulphate would require +2. 67 volts, but, sincethe deposition of hydrogen from sulphuric acid solution requires only+1. 90 volts, the quantitative deposition of zinc by electrolysis froma sulphuric acid solution of a zinc salt is not practicable. On theother hand, silver, if present in a solution of copper sulphate, woulddeposit with the copper. 

The foregoing examples suffice to illustrate the application of theprinciple of deposition potentials, but it must further be notedthat the values stated apply to normal solutions of the compounds inquestion, that is, to solutions of considerable concentrations. As theconcentration of the ions diminishes, and hence fewer ions approachthe electrodes, somewhat higher voltages are required to attract anddischarge them. From this it follows that the concentrations should bekept as high as possible to effect complete deposition in the leastpracticable time, or else the potentials applied must be progressivelyincreased as deposition proceeds. In practice, the desired result isobtained by starting with small volumes of solution, using as large anelectrode surface as possible, and by stirring the solution to bringthe ions in contact with the electrodes. This is, in general, a moreconvenient procedure than that of increasing the potential of thecurrent during electrolysis, although that method is also used. 

As already stated, those ions in a solution of electrolytes will firstbe discharged which have the lowest deposition potentials, and solong as these ions are present around the electrode in considerableconcentration they, almost alone, are discharged, but, as theirconcentration diminishes, other ions whose deposition potentials arehigher but still within that of the current applied, will also beginto separate. For example, from a nitric acid solution of coppernitrate, the copper ions will first be discharged at the cathode, butas they diminish in concentration hydrogen ions from the acid (orwater) will be also discharged. Since the hydrogen thus liberated is areducing agent, the nitric acid in the solution is slowly reduced toammonia, and it may happen that if the current is passed through for along time, such a solution will become alkaline. Oxygen is liberatedat the anode, but, since there is no oxidizable substance presentaround that electrode, it escapes as oxygen gas. It should be notedthat, 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 themetals are preferable to those of the chlorides, since liberatedchlorine attacks the electrodes. In some cases, as for example, thatof silver, solution of salts forming complex ions, like that ofthe double cyanide of silver and potassium, yield better metallicdeposits. 

Most metals are deposited as such upon the cathode; a few, notablylead and manganese, separate in the form of dioxides upon the anode. It is evidently important that the deposited material should be sofirmly adherent that it can be washed, dried, and weighed withoutloss in handling. To secure these conditions it is essential that thecurrent density (that is, the amount of current per unit of area ofthe electrodes) shall not be too high. In prescribing analyticalconditions it is customary to state the current strength in "normaldensities" 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 spongyor loosely adherent and falls off on subsequent treatment. This placesa practical limit to the current density to be employed, for a givenelectrode surface. The cause of the unsatisfactory character ofthe deposit is apparently sometimes to be found in the coincidentliberation of considerable hydrogen and sometimes in the failure ofthe rapidly deposited material to form a continuous adherent surface. The effect of rotating electrodes upon the character of the deposit isreferred to below. 

The negative ions of an electrolyte are attracted to the anode and aredischarged on contact with it. Anions such as the chloride ion yieldchlorine atoms, from which gaseous chlorine molecules are formedand escape. The radicals which compose such ions as NO_{3}^{-} orSO_{4}^{--} are not capable of independent existence after discharge, and break down into oxygen and N_{2}O_{5} and SO_{3} respectively. Theoxygen escapes and the anhydrides, reacting with water, re-form nitricand sulphuric acids. 

The law of Faraday expresses the relation between current strength andthe quantities of the decomposition products which, under constantconditions, appear at the electrodes, namely, that a given quantityof electricity, acting for a given time, causes the separation ofchemically equivalent quantities of the various elements or radicals. For example, since 107. 94 grams of silver is equivalent to 1. 008 gramsof hydrogen, and that in turn to 8 grams of oxygen, or 31. 78 grams ofcopper, the quantity of electricity which will cause the deposit of107. 94 grams of silver in a given time will also separate the weightsjust indicated of the other substances. Experiments show that acurrent of one ampere passing for one second, i. E. , a coulomb ofelectricity, causes the deposition of 0. 001118 gram of silver from anormal solution of a silver salt. The number of coulombs required todeposit 107. 94 grams is 107. 94/0. 001118 or 96, 550 and the same numberof 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 firstappear that Faraday's law could thus be used as a basis for thecalculation of the time required for the deposition of a givenquantity of an electrolyte from solution, it must be remembered thatthe law expresses what occurs when the concentration of the ions inthe solution is kept constant, as, for example, when the anode ina silver salt solution is a plate of metallic silver. Under theconditions of electro-analysis the concentration of the ions isconstantly diminishing as deposition proceeds and the time actuallyrequired for complete deposition of a given weight of material bya current of constant strength is, therefore, greater than thatcalculated on the basis of the law as stated above. 

The electrodes employed in electro-analysis are almost exclusivelyof platinum, since that metal alone satisfactorily resists chemicalaction of the electrolytes, and can be dried and weighed withoutchange in composition. The platinum electrodes may be used in theform of dishes, foil or gauze. The last, on account of the ease ofcirculation of the electrolyte, its relatively large surface inproportion to its weight and the readiness with which it can be washedand dried, is generally preferred. 

Many devices have been described by the use of which the electrodeupon which deposition occurs can be mechanically rotated. This has aneffect parallel to that of greatly increasing the electrode surfaceand also provides a most efficient means of stirring the solution. With such an apparatus the amperage may be increased to 5 or even 10amperes and depositions completed with great rapidity and accuracy. Itis desirable, whenever practicable, to provide a rotating or stirringdevice, since, for example, the time consumed in the deposition of theamount of copper usually found in analysis may be reduced from the20 to 24 hours required with stationary electrodes, and unstirredsolutions, 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. Dissolvethe 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 thebeaker and the bottom of the watch-glass with water and dilute thesolution to about 50 cc. Carefully heat to boiling and boil for aminute 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 waterand finally with 95 per cent alcohol (Note 3). The alcohol may beignited and burned off. The electrodes are then cooled in a desiccatorand weighed. Connect the electrodes with the binding posts (or otherdevice for connection with the electric circuit) in such a way thatthe copper will be deposited upon the electrode with the largersurface, which is made the cathode. The beaker containing the solutionshould then be raised into place from below the electrodes until thelatter reach nearly to the bottom of the beaker. The support for thebeaker must be so arranged that it can be easily raised or lowered. 

If the electrolytic apparatus is provided with a mechanism for therotation of the electrode or stirring of the electrolyte, proceed asfollows: Arrange the resistance in the circuit to provide a directcurrent of about one ampere. Pass this current through the solutionto be electrolyzed, and start the rotating mechanism. Keep the beakercovered as completely as possible, using a split watch-glass (or otherdevice) to avoid loss by spattering. When the solution is colorless, which is usually the case after about 35 minutes, rinse off the coverglass, wash down the sides of the beaker, add about 0. 30 gram of ureaand continue the electrolysis for another five minutes (Notes 4 and5). 

If stationary electrodes are employed, the current strength should beabout 0. 1 ampere, which may, after 12 to 15 hours, be increased to 0. 2ampere. The time required for complete deposition is usually from 20to 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 addedin this case. 

When the deposition of the copper appears to be complete, stop therotating mechanism and slowly lower the beaker with the left hand, directing at the same time a stream of water from a wash bottle onboth electrodes. Remove the beaker, shut off the current, and, ifnecessary, complete the washing of the electrodes (Note 6). Rinse theelectrodes cautiously with alcohol and heat them in a hot closet untilthe alcohol has just evaporated, but no longer, since the copper islikely to oxidize at the higher temperature. (The alcohol may beremoved by ignition if care is taken to keep the electrodes in motionin the air so that the copper deposit is not too strongly heated atany one point. )

Test the solution in the beaker for copper as follows, rememberingthat it is to be used for subsequent determinations of iron and zinc:Remove about 5 cc. And add a slight excess of ammonia. Compare themixture with some distilled water, holding both above a white surface. The solution should not show any tinge of blue. If the presence ofcopper is indicated, add the test portion to the main solution, evaporate the whole to a volume of about 100 cc. , and againelectrolyze with clean electrodes (Note 7). 

After cooling the electrodes in a desiccator, weigh them and from theweight of copper on the cathode and of lead dioxide (PbO_{2}) on theanode, calculate the percentage of copper (Cu) and of lead (Pb) in thebrass. 

[Note 1: It is obvious that the brass taken for analysis should beuntarnished, which can be easily assured, when wire is used, byscouring with emery. If chips or borings are used, they should be wellmixed, and the sample for analysis taken from different parts of themixture. ]

[Note 2: If a white residue remains upon treatment of the alloy withnitric acid, it indicates the presence of tin. The material is not, therefore, a true brass. This may be treated as follows: Evaporate thesolution to dryness, moisten the residue with 5 cc. Of dilute nitricacid (sp. Gr. 1. 2) and add 50 cc. Of hot water. Filter off themeta-stannic acid, wash, ignite in porcelain and weigh as SnO_{2}. This oxide is never wholly free from copper and must be purified foran 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 beforeusing, and those portions upon which the metal will deposit should notbe touched with the fingers after cleaning. ]

[Note 4: Of the ions in solution, the H^{+}, Cu^{++}, Zn^{++}, andFe^{+++} ions tend to move toward the cathode. The NO_{3}^{-} ions andthe lead, probably in the form of PbO_{2}^{--} ions, move toward theanode. At the cathode the Cu^{++} ions are discharged and plate out asmetallic copper. This alone occurs while the solution is relativelyconcentrated. Later on, H^{+} ions are also discharged. In thepresence of considerable quantities of H^{+} ions, as in this acidsolution, no Zn^{++} or Fe^{+++} ions are discharged because of theirgreater deposition potentials. At the anode the lead is deposited asPbO_{2} and oxygen is evolved. 

For the reasons stated on page 141 care must be taken that thesolution does not become alkaline if the electrolysis is longcontinued. ]

[Note 5: Urea reacts with nitrous acid, which may be formed in thesolution as a result of the reducing action of the liberated hydrogen. Its removal promotes the complete precipitation of the copper. Thereaction 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 fromthe nitric acid solution before the circuit is broken to preventre-solution of the copper. 

If several solutions are connected in the same circuit it is obviousthat some device must be used to close the circuit as soon as thebeaker is removed. ]

[Note 7: The electrodes upon which the copper has been depositedmay be cleaned by immersion in warm nitric acid. To remove the leaddioxide, 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. 1per cent) which, unless removed, is precipitated as phosphate andweighed with the zinc. 

PROCEDURE. --To the solution from the precipitation of copper andlead by electrolysis, add dilute ammonia (sp. Gr. 0. 96) until theprecipitate of zinc hydroxide which first forms re-dissolves, leavingonly a slight red precipitate of ferric hydroxide. Filter off theiron precipitate, using a washed filter, and wash five times with hotwater. Test a portion of the last washing with a dilute solution ofammonium sulphide to assure complete removal of the zinc. 

The precipitate may then be ignited and weighed as ferric oxide, asdescribed on page 110. 

Calculate the percentage of iron (Fe) in the brass. 

DETERMINATION OF ZINC

PROCEDURE. --Acidify the filtrate from the iron determination withdilute nitric acid. Concentrate it to 150 cc. Add to the cold solutiondilute ammonia (sp. Gr. 0. 96) cautiously until it barely smells ofammonia; 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 theblue of the litmus just changes to red. It is important that thispoint should not be overstepped. Heat the solution nearly to boilingand pour into it slowly a filtered solution of di-ammonium hydrogenphosphate[1] containing a weight of the phosphate about equalto twelve times that of the zinc to be precipitated. (For thiscalculation the approximate percentage of zinc is that found bysubtracting the sum of the percentages of the copper, lead and ironfrom 100 per cent. ) Keep the solution just below boiling for fifteenminutes, stirring frequently (Note 2). If at the end of this time theamorphous precipitate has become crystalline, allow the solution tocool for about four hours, although a longer time does no harm (Note3), and filter upon an asbestos filter in a porcelain Gooch crucible. The filter is prepared as described on page 103, and should be driedto constant weight at 105°C. 

[Footnote 1: The ammonium phosphate which is commonly obtainablecontains some mono-ammonium salt, and this is not satisfactory as aprecipitant. It is advisable, therefore, to weigh out the amount ofthe salt required, dissolve it in a small volume of water, add a dropof phenolphthalein solution, and finally add dilute ammonium hydroxidesolution cautiously until the solution just becomes pink, but do notadd an excess. ]

Wash the precipitate until free from sulphates with a warm 1 per centsolution of the di-ammonium phosphate, and then five times with 50 percent alcohol (Note 4). Dry the crucible and precipitate for an hour at105°C. , and finally to constant weight (Note 5). The filtrate shouldbe made alkaline with ammonia and tested for zinc with a few drops ofammonium 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 inammonia. It is, therefore, necessary to precipitate the zinc in anearly neutral solution, which is more accurately obtained by addinga drop of a litmus solution to the liquid than by the use of litmuspaper. ]

[Note 2: The precipitate which first forms is amorphous, and may havea variable composition. On standing it becomes crystalline and thenhas the composition ZnNH_{4}PO_{4}. The precipitate then settlesrapidly and is apt to occasion "bumping" if the solution is heated toboiling. Stirring promotes the crystallization. ]

[Note 3: In a carefully neutralized solution containing a considerableexcess of the precipitant, and also ammonium salts, the separationof the zinc is complete after standing four hours. The ionic changesconnected with the precipitation of the zinc as zinc ammoniumphosphate are similar to those described for magnesium ammoniumphosphate, except that the zinc precipitate is soluble in an excess ofammonium hydroxide, probably as a result of the formation of complexions of the general character Zn(NH_{3})_{4}^{++}. ]

[Note 4: The precipitate is washed first with a dilute solution of thephosphate to prevent a slight decomposition of the precipitate (as aresult of hydrolysis) if hot water alone is used. The alcohol is addedto the final wash-water to promote the subsequent drying. ]

[Note 5: The precipitate may be ignited and weighed asZn_{2}P_{2}O_{7}, by cautiously heating the porcelain Gooch cruciblewithin a nickel or iron crucible, used as a radiator. The heatingmust be very slow at first, as the escaping ammonia may reduce theprecipitate 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 cubiccentimeters of dilute nitric acid, and the zinc reprecipitated asphosphate, filtered off, dried, and weighed, and the weight added tothat of the main precipitate. ]

[Note 7: It has been found that some samples of asbestos are actedupon by the phosphate solution and lose weight. An error from thissource may be avoided by determining the weight of the crucibleand filter after weighing the precipitate. For this purpose theprecipitate may be dissolved in dilute nitric acid, the asbestoswashed thoroughly, and the crucible reweighed. ]

[Note 8. The details of this method of precipitation of zinc are fullydiscussed 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 andsome of the cements, a comparatively few can be completely decomposedby treatment with acids, but by far the larger number require fusionwith an alkaline flux to effect decomposition and solutionfor analysis. The procedure given below applies to silicatesundecomposable by acids, of which the mineral feldspar is taken as atypical example. Modifications of the procedure, which are applicableto 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) untilno grittiness is to be detected, or, better, until it will entirelypass through a sieve made of fine silk bolting cloth. The sieve may bemade by placing a piece of the bolting cloth over the top of a smallbeaker in which the ground mineral is placed, holding the cloth inplace by means of a rubber band below the lip of the beaker. Byinverting the beaker over clean paper and gently tapping it, the fineparticles pass through the sieve, leaving the coarser particles withinthe beaker. These must be returned to the mortar and ground, and theprocess of sifting and grinding repeated until the entire samplepasses through the sieve. 

[Note 1: If the sample of feldspar for analysis is in the massive orcrystalline form, it should be crushed in an iron mortar until thepieces are about half the size of a pea, and then transferred to asteel mortar, in which they are reduced to a coarse powder. A woodenmallet 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 continueduntil the whole of the portion of the mineral originally taken hasbeen ground so that it will pass the bolting cloth, otherwise thesifted portion does not represent an average sample, the softeringredients, if foreign matter is present, being first reduced topowder. For this reason it is best to start with not more than thequantity of the feldspar needed for analysis. The mineral must bethoroughly mixed after the grinding. ]

FUSION AND SOLUTION

PROCEDURE. --Weigh into platinum crucibles two portions of the groundfeldspar of about 0. 8 gram each. Weigh on rough balances two portionsof anhydrous sodium carbonate, each amounting to about six times theweight of the feldspar taken for analysis (Note 1). Pour about threefourths of the sodium carbonate into the crucible, place the latter ona piece of clean, glazed paper, and thoroughly mix the substance andthe flux by carefully stirring for several minutes with a dry glassrod, the end of which has been recently heated and rounded in a flameand slowly cooled. The rod may be wiped off with a small fragment offilter paper, which may be placed in the crucible. Place the remainingfourth of the carbonate on the top of the mixture. Cover the crucible, heat it to dull redness for five minutes, and then gradually increasethe heat to the full capacity of a Bunsen or Tirrill burner fortwenty minutes, or until a quiet, liquid fusion is obtained (Note 2). Finally, heat the sides and cover strongly until any material whichmay have collected upon them is also brought to fusion. 

Allow the crucible to cool, and remove the fused mass as directed onpage 116. Disintegrate the mass by placing it in a previously preparedmixture 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 andlid by means of a little hydrochloric acid, adding this acid to themain solution (Notes 4 and 5). 

[Note 1: Quartz, and minerals containing very high percentages ofsilica, may require eight or ten parts by weight of the flux to insurea satisfactory decomposition. ]

[Note 2: During the fusion the feldspar, which, when pure, is asilicate of aluminium and either sodium or potassium, but usuallycontains some iron, calcium, and magnesium, is decomposed by thealkaline flux. The sodium of the latter combines with the silicic acidof the silicate, with the evolution of carbon dioxide, while about twothirds of the aluminium forms sodium aluminate and the remainderis converted into basic carbonate, or the oxide. The calcium andmagnesium, if present, are changed to carbonates or oxides. 

The heat is applied gently to prevent a too violent reaction whenfusion first takes place. ]

[Note 3: The solution of a silicate by a strong acid is the result ofthe combination of the H^{+} ions of the acid and the silicate ionsof the silicate to form a slightly ionized silicic acid. As aconsequence, the concentration of the silicate ions in the solution isreduced nearly to zero, and more silicate dissolves to re-establishthe disturbed equilibrium. This process repeats itself until all ofthe silicate is brought into solution. 

Whether the resulting solution of the silicate contains ortho-silicicacid (H_{4}SiO_{4}) or whether it is a colloidal solution of someother less hydrated acid, such as meta-silicic acid (H_{2}SiO_{3}), is a matter that is still debatable. It is certain, however, that thegelatinous material which readily separates from such solutions is ofthe nature of a hydrogel, that is, a colloid which is insoluble inwater. This substance when heated to 100°C. , or higher, is completelydehydrated, leaving only the anhydride, SiO_{2}. The changes may berepresented 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 bythe escaping carbon dioxide during the fusion. The crucible musttherefore be kept covered as much as possible and the lid carefullycleaned. ]

[Note 5: A gritty residue remaining after the disintegration ofthe fused mass by acid indicates that the substance has been butimperfectly decomposed. Such a residue should be filtered, washed, dried, ignited, and again fused with the alkaline flux; or, if thequantity of material at hand will permit, it is better to reject theanalysis, and to use increased care in grinding the mineral and inmixing it with the flux. ]

DEHYDRATION AND FILTRATION

PROCEDURE. --Evaporate the solution of the fusion to dryness, stirringfrequently until the residue is a dry powder. Moisten the residue with5 cc. Of strong hydrochloric acid (sp. Gr. 1. 20) and evaporate againto dryness. Heat the residue for at least one hour at a temperatureof 110°C. (Note 1). Again moisten the residue with concentratedhydrochloric acid, warm gently, making sure that the acid comes intocontact with the whole of the residue, dilute to about 200 cc. Andbring to boiling. Filter off the silica without much delay (Note 2), and wash five times with warm dilute hydrochloric acid (one partdilute acid (1. 12 sp. Gr. ) to three parts of water). Allow the filterto drain for a few moments, then place a clean beaker below the funneland wash with water until free from chlorides, discarding thesewashings. Evaporate the original filtrate to dryness, dehydrate at110°C. For one hour (Note 3), and proceed as before, using a secondfilter to collect the silica after the second dehydration. Wash thisfilter with warm, dilute hydrochloric acid (Note 4), and finally withhot water until free from chlorides. 

[Note 1: The silicic acid must be freed from its combination witha base (sodium, in this instance) before it can be dehydrated. The excess of hydrochloric acid accomplishes this liberation. Bydisintegrating the fused mass with a considerable volume of diluteacid the silicic acid is at first held in solution to a large extent. Immediate treatment of the fused mass with strong acid is likelyto cause a semi-gelatinous silicic acid to separate at once and toinclose alkali salts or alumina. 

A flocculent residue will often remain after the decomposition of thefused mass is effected. This is usually partially dehydrated silicicacid and does not require further treatment at this point. Theprogress of the dehydration is indicated by the behavior of thesolution, which as evaporation proceeds usually gelatinizes. On thisaccount it is necessary to allow the solution to evaporate on a steambath, or to stir it vigorously, to avoid loss by spattering. ]

[Note 2: To obtain an approximately pure silica, the residue afterevaporation must be thoroughly extracted by warming with hydrochloricacid, and the solution freely diluted to prevent, as far as possible, the inclosure of the residual salts in the particles of silica. Thefiltration should take place without delay, as the dehydrated silicaslowly dissolves in hydrochloric acid on standing. ]

[Note 3: It has been shown by Hillebrand that silicic acid cannot becompletely dehydrated by a single evaporation and heating, nor byseveral such treatments, unless an intermediate filtration of thesilica occurs. If, however, the silica is removed and the filtratesare again evaporated and the residue heated, the amount of silicaremaining in solution is usually negligible, although severalevaporations and filtrations are required with some silicates toinsure absolute accuracy. 

It is probable that temperatures above 100°C. Are not absolutelynecessary to dehydrate the silica; but it is recommended, as tendingto leave the silica in a better condition for filtration than whenthe lower temperature of the water bath is used. This, and many otherpoints in the analysis of silicates, are fully discussed by Dr. Hillebrand in the admirable monograph on "The Analysis of Silicate andCarbonate Rocks, " Bulletin No. 700 of the United States GeologicalSurvey. 

The double evaporation and filtration spoken of above are essentialbecause of the relatively large amount of alkali salts (sodiumchloride) present after evaporation. For the highest accuracy in thedetermination of silica, or of iron and alumina, it is also necessaryto examine for silica the precipitate produced in the filtrate byammonium hydroxide by fusing it with acid potassium sulphate andsolution of the fused mass in water. The insoluble silica is filtered, washed, and weighed, and the weight added to the weight of silicapreviously obtained. ]

[Note 4: Aluminium and iron are likely to be thrown down as basicsalts from hot, very dilute solutions of their chlorides, as a resultof hydrolysis. If the silica were washed only with hot water, thesolution of these chlorides remaining in the filter after the passageof the original filtrate would gradually become so dilute as to throwdown basic salts within the pores of the filter, which would remainwith the silica. To avoid this, an acid wash-water is used until thealuminium and iron are practically removed. The acid is then removedby water. ]

IGNITION AND TESTING OF SILICA

PROCEDURE. --Transfer the two washed filters belonging to eachdetermination to a platinum crucible, which need not be previouslyweighed, and burn off the filter (Note 1). Ignite for thirty minutesover the blast lamp with the cover on the crucible, and then forperiods of ten minutes, until the weight is constant. 

When a constant weight has been obtained, pour into the crucible about3 cc. Of water, and then 3 cc. Of hydrofluoric acid. !This must bedone in a hood with a good draft and great care must be taken not tocome into contact with the acid or to inhale its fumes (Note 2!). 

If the precipitate has dissolved in this quantity of acid, add twodrops of concentrated sulphuric acid, and heat very slowly (alwaysunder the hood) until all the liquid has evaporated, finally ignitingto redness. Cool in a desiccator, and weigh the crucible and residue. Deduct this weight from the previous weight of crucible and impuresilica, and from the difference calculate the percentage of silica inthe sample (Note 3). 

[Note 1: The silica undergoes no change during the ignition beyond theremoval of all traces of water; but Hillebrand (!loc. Cit. !) has shownthat the silica holds moisture so tenaciously that prolonged ignitionover the blast lamp is necessary to remove it entirely. This finelydivided, ignited silica tends to absorb moisture, and should beweighed quickly. ]

[Note 2: Notwithstanding all precautions, the ignited precipitate ofsilica is rarely wholly pure. It is tested by volatilisation of thesilica as silicon fluoride after solution in hydrofluoric acid, and, if the analysis has been properly conducted, the residue, aftertreatment 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 originalprecipitate in the form of oxides. The addition of the sulphuricacid displaces the hydrofluoric acid, and it may be assumed that theresulting sulphates (usually of iron or aluminium) are converted tooxides by the final ignition. 

It is obvious that unless the sulphuric and hydrofluoric acids usedare known to leave no residue on evaporation, a quantity equal to thatemployed in the analysis must be evaporated and a correction appliedfor any residue found. ]

[Note 4: If the silicate to be analyzed is shown by a previousqualitative examination to be completely decomposable, it may bedirectly treated with hydrochloric acid, the solution evaporated todryness, and the silica dehydrated and further treated as described inthe case of the feldspar after fusion. 

A silicate which gelatinizes on treatment with acids should be mixedfirst with a little water, and the strong acid added in small portionswith stirring, otherwise the gelatinous silicic acid inclosesparticles of the original silicate and prevents decomposition. Thewater, by separating the particles and slightly lessening the rapidityof action, prevents this difficulty. This procedure is one whichapplies in general to the solution of fine mineral powders in acids. 

If a small residue remains undecomposed by the treatment of thesilicate with acid, this may be filtered, washed, ignited and fusedwith sodium carbonate and a solution of the fused mass added to theoriginal acid solution. This double procedure has an advantage, inthat it avoids adding so large a quantity of sodium salts as isrequired for disintegration of the whole of the silicate by the fusionmethod. ]

PART IV

STOICHIOMETRY

The problems with which the analytical chemist has to deal are not, asa 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 thesmall number of types into which these problems may be grouped, eachproblem is approached as an individual puzzle, unrelated to othersalready solved or explained. This attitude of mind should be carefullyavoided. 

It is obvious that ability to make the calculations necessary forthe interpretation of analytical data is no less important than themanipulative skill required to obtain them, and that a moderate timespent in the careful study of the solutions of the typical problemswhich follow may save much later embarrassment. 

1. It is often necessary to calculate what is known as a "chemicalfactor, " or its equivalent logarithmic value called a "log factor, "for the conversion of the weight of a given chemical substance into anequivalent weight of another substance. This is, in reality, a verysimple problem in proportion, making use of the atomic or molecularweights of the substances in question which are chemically equivalentto each other. One of the simplest cases of this sort is thefollowing: What is the factor for the conversion of a given weight ofbarium 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 inthe molecule and the atomic weight of S is 32. 1. The chemical factoris, therefore, 32. 1/233. 5, or 0. 1375 and the weight of S correspondingto a given weight of BaSO_{4} is found by multiplying the weight ofBaSO_{4} by this factor. If the problem takes the form, "What isthe factor for the conversion of a given weight of ferric oxide(Fe_{2}O_{3}) into ferrous oxide (FeO), or of a given weight ofmangano-manganic oxide (Mn_{3}O_{4}) into manganese (Mn)?" theprinciple involved is the same, but it must then be noted that, in thefirst instance, each molecule of Fe_{2}O_{3} will be equivalent to twomolecules of FeO, and in the second instance that each molecule ofMn_{3}O_{4} is equivalent to three atoms of Mn. The respective factorsthen 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 typeof problem are extremely simple. It is only necessary to observecarefully the chemical equivalents. It is plainly incorrect to expressthe ratio of ferrous to ferric oxide as (FeO/Fe_{2}O_{3}), since eachmolecule of the ferric oxide will yield two molecules of the ferrousoxide. Mistakes of this sort are easily made and constitute one of themost 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 thefactor 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})?" Todetermine the chemical equivalents in such an instance it is necessaryto 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 molecularquantities 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 thefactor required for the conversion of a given weight of potassiumpermanganate (KMnO_{4}) into an equivalent weight of potassiumbichromate (K_{2}Cr_{2}O_{7}), each acting as an oxidizing agentagainst 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 with10FeSO_{4}, while K_{2}Cr_{2}O_{7} reacts with 6FeSO_{4}. These arenot equivalent, but if the first equation is multiplied by 3 and thesecond by 5 the number of molecules of FeSO_{4} is then the same inboth, 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. Theseare obviously chemically equivalent and the desired factor isexpressed 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 solutionsaccording to the principles just explained, when several successivereactions are involved. Such problems may be solved by a series ofproportions, but it is usually possible to eliminate the commonfactors and solve but a single one. For example, the amount of MnO_{2}in a sample of the mineral pyrolusite may be determined by dissolvingthe mineral in hydrochloric acid, absorbing the evolved chlorine in asolution of potassium iodide, and measuring the liberated iodineby titration with a standard solution of sodium thiosulphate. Thereactions involved are:

MnO_{2} + 4HCl --> MnCl_{2} + 2H_{2}O + Cl_{2}Cl_{2} + 2KI --> I_{2} + 2KClI_{2} + 2Na_{2}S_{2}O_{3} --> 2NaI + Na_{2}S_{4}O_{6}

Assuming that the weight of thiosulphate corresponding to thevolume of sodium thiosulphate solution used is known, what is thecorresponding weight of manganese dioxide? From the reactions givenabove, 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 remains2Na_{2}S_{2}O_{3}:MnO_{2} = 316. 4:86. 9, and the factor for theconversion of thiosulphate into an equivalent of manganese dioxide is86. 9/316. 4. 

4. To calculate the volume of a reagent required for a specificoperation, it is necessary to know the exact reaction which is to bebrought about, and, as with the calculation of factors, to keep inmind the molecular relations between the reagent and the substancereacted upon. For example, to estimate the weight of barium chloridenecessary 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 thetwo atoms of sulphur will form upon oxidation a molecule of sulphuricacid or a sulphate, which, in turn, will require a molecule of thebarium chloride for precipitation. To determine the quantity of thebarium chloride required, it is necessary to include in its molecularweight the water of crystallization, since this is inseparable fromthe chloride when it is weighed. This applies equally to other similarinstances. 

If the strength of an acid is expressed in percentage by weight, dueregard must be paid to its specific gravity. For example, hydrochloricacid (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 forthe separation of the constituents of a mixture of substances bymaking what is called an "indirect analysis. " For example, in theanalysis of silicate rocks, the sodium and potassium present may beobtained in the form of their chlorides and weighed together. If theweight of such a mixture is known, and also the percentage of chlorinepresent, it is possible to calculate the amount of each chloride inthe mixture. Let it be assumed that the weight of the mixed chloridesis 0. 15 gram, and that it contains 53 per cent of chlorine. 

The simplest solution of such a problem is reached through algebraicmethods. The weight of chlorine is evidently 0. 15 x 0. 53, or 0. 0795gram. Let x represent the weight of sodium chloride present and ythat of potassium chloride. The molecular weight of NaCl is 58. 5 andthat 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. 0625gram. The weight of KCl is found by subtracting this from 0. 15. 

The above is one of the most common types of indirect analyses. Othersare more complex but they can be reduced to algebraic expressions andsolved by their aid. It should, however, be noted that the resultsobtained by these indirect methods cannot be depended upon for highaccuracy, since slight errors in the determination of the commonconstituent, as chlorine in the above mixture, will cause considerablevariations in the values found for the components. They should not beemployed when direct methods are applicable, if accuracy is essential. 

PROBLEMS

(The reactions necessary for the solution of these problems are eitherstated with the problem or may be found in the earlier text. In thecalculations from which the answers are derived, the atomic weightsgiven on page 195 have been employed, using, however, only the firstdecimal but increasing this by 1 when the second decimal is 5 orabove. Thus, 39. 1 has been taken as the atomic weight of potassium, 32. 1 for sulphur, etc. This has been done merely to secure uniformityof treatment, and the student should remember that it is always wellto take into account the degree of accuracy desired in a particularinstance in determining the number of decimal places to retain. Four-place logarithms were employed in the calculations. Where fourfigures 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 issolved. )

VOLUMETRIC ANALYSIS

1. How many grams of pure potassium hydroxide are required for exactly1 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 sodiumbicarbonate 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 aphosphate; (b) as a sodium salt?

!Answers!: (a) 47. 33; (b) 71. 0. 

5. A sample of aqueous hydrochloric acid has a specific gravityof 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. 20containing 39. 80 per cent HCl by weight) are required to furnish 36. 45grams of the gaseous compound?

!Answer!: 76. 33 cc. 

7. A given solution contains 0. 1063 equivalents of hydrochloric acidin 976 cc. What is its normal value?

!Answer!: 0. 1089 N. 

8. In standardizing a hydrochloric acid solution it is found that47. 26 cc. Of hydrochloric acid are exactly equivalent to 1. 216 gramsof pure sodium carbonate, using methyl orange as an indicator. What isthe normal value of the hydrochloric acid?

!Answer!: 0. 4855 N. 

9. Convert 42. 75 cc. Of 0. 5162 normal hydrochloric acid to theequivalent volume of normal hydrochloric acid. 

!Answer!: 22. 07 cc. 

10. A solution containing 25. 27 cc. Of 0. 1065 normal hydrochloric acidis added to one containing 92. 21 cc. Of 0. 5431 normal sulphuric acidand 50 cc. Of exactly normal potassium hydroxide added from a pipette. Is the solution acid or alkaline? How many cubic centimeters of0. 1 normal acid or alkali must be added to exactly neutralize thesolution?

!Answer!: 27. 6 cc. Alkali (solution is acid). 

11. By experiment the normal value of a sulphuric acid solution isfound to be 0. 5172. Of this acid 39. 65 cc. Are exactly equivalent to21. 74 cc. Of a standard alkali solution. What is the normal value ofthe alkali?

!Answer!: 0. 9432 N. 

12. A solution of sulphuric acid is standardized against a sample ofcalcium carbonate which has been previously accurately analyzed andfound to contain 92. 44% CaCO_{3} and no other basic material. Thesample 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 sodiumhydroxide. 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 portionwhen an excess of silver nitrate is added? (b) how many grams of pureanhydrous sodium carbonate (Na_{2}CO_{3}) will be neutralized by aportion of it? (c) how many grams of silver will there be in thesilver chloride formed when an excess of silver nitrate is added to aportion? (d) how many grams of iron will be dissolved to FeCl_{2} by aportion of it? (e) how many grams of magnesium chloride will be formedand how many grams of carbon dioxide liberated when an excess ofmagnesium 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 solutiondiluted to exactly 1 liter, and 40 cc. Are neutralized with 20 cc. Of a potassium carbonate solution, what is the normal value of thecarbonate solution?

!Answer!: 0. 7084 N. 

15. How many cubic centimeters of 0. 3 normal sulphuric acid will berequired to neutralize (a) 30 cc. Of 0. 5 normal potassium hydroxide;(b) to neutralize 30 cc. Of 0. 5 normal barium hydroxide; (c) toneutralize 20 cc. Of a solution containing 10. 02 grams of potassiumbicarbonate per 100 cc. ; (d) to give a precipitate of barium sulphateweighing 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 1cc. Is equivalent to 0. 1027 gram of pure sodium carbonate to make itexactly 1. 250 normal. 700 cc. Of the solution are available. To whatvolume must it be diluted?

!Answer!: 1084 cc. 

17. Given the following data: 1 cc. Of NaOH = 1. 117 cc. HCl. The HClis 0. 4876 N. How much water must be added to 100 cc. Of the alkali tomake it exactly 0. 5 N. ?

!Answer!: 9. 0 cc. 

18. What is the normal value of a sulphuric acid solution which has aspecific 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), afterignition in platinum to convert it to the double carbonate, istitrated with sulphuric acid, using methyl orange as an indicator. From the following data calculate the percentage purity of the sample:

Wt. Sample = 0. 9500 gramH_{2}SO_{4} used = 43. 65 cc. NaOH used = 1. 72 cc. 1 cc. H_{2}SO_{4} = 1. 064 cc. NaOHNormal value NaOH = 0. 1321 N. 

!Answer!: 87. 72 cc. 

20. One gram of a mixture of 50% sodium carbonate and 50% potassiumcarbonate is dissolved in water, and 17. 36 cc. Of 1. 075 N acid isadded. Is the resulting solution acid or alkaline? How many cubiccentimeters of 1. 075 N acid or alkali will have to be added to makethe solution exactly neutral?

!Answers!: Acid; 1. 86 cc. Alkali. 

21. In preparing an alkaline solution for use in volumetric work, ananalyst, because of shortage of chemicals, mixed exactly 46. 32 gramsof pure KOH and 27. 64 grams of pure NaOH, and after dissolving inwater, diluted the solution to exactly one liter. How many cubiccentimeters of 1. 022 N hydrochloric acid are necessary to neutralize50 cc. Of the basic solution?

!Answer!: 74. 18 cc. 

22. One gram of crude ammonium salt is treated with strong potassiumhydroxide solution. The ammonia liberated is distilled and collectedin 50 cc. Of 0. 5 N acid and the excess titrated with 1. 55 cc. Of 0. 5 Nsodium hydroxide. Calculate the percentage of NH_{3} in the sample. 

!Answer!: 41. 17%. 

23. In titrating solutions of alkali carbonates in the presence ofphenolphthalein, the color change takes place when the carbonate hasbeen converted to bicarbonate. In the presence of methyl orange, thecolor change takes place only when the carbonate has been completelyneutralized. From the following data, calculate the percentages ofNa_{2}CO_{3} and NaOH in an impure mixture. Weight of sample, 1. 500grams; HCl (0. 5 N) required for phenolphthalein end-point, 28. 85 cc. ;HCl (0. 5 N) required to complete the titration after adding methylorange, 23. 85 cc. 

!Answers!: 6. 67% NaOH; 84. 28% Na_{2}CO_{3}. 

24. A sample of sodium carbonate containing sodium hydroxide weighs1. 179 grams. It is titrated with 0. 30 N hydrochloric acid, usingphenolphthalein in cold solution as an indicator and becomes colorlessafter the addition of 48. 16 cc. Methyl orange is added and 24. 08 cc. Are needed for complete neutralization. What is the percentage of NaOHand 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 phenolphthaleinend-point, 26. 40 cc. ; after adding an excess of acid and boiling outthe carbon dioxide, the total volume of 0. 25 N hydrochloric acidrequired 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 thenormality of the acid in order that the number of cubic centimeters ofacid 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 thatthe number of cubic centimeters of 0. 5 N acid used may be equal to onethird the percentage of K_{2}CO_{3}?

!Answer!: 1. 152 grams. 

28. What weight of cream of tartar must have been taken for analysisin order to have obtained 97. 60% KHC_{4}H_{4}O_{6} in an analysisinvolving 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 gramCaCO_{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 ofpotassium carbonate from the following data: Weight of sample = 1. 00gram; HCl sol. Used = 55. 90 cc. ; NaOH sol. Used = 0. 42 cc. ; 1 cc. NaOHsol. = 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 ofH_{2}C_{2}O_{4}. 2H_{2}O = 0. 2460 gram; NaOH solution used = 41. 03cc. ; HCl solution used = 0. 63; 1 cc. NaOH solution = 1. 190 cc. HClsolution. (Analysis); Weight of sample 0. 1200 gram; HCl used = 36. 38cc. ; NaOH used = 6. 20 cc. 

!Answer!: 97. 97%. 

31. It is desired to dilute a solution of hydrochloric acid to exactly0. 05 N. The following data are given: 44. 97 cc. Of the hydrochloricacid are equivalent to 43. 76 cc. Of the NaOH solution. The NaOHis standardized against a pure potassium tetroxalate(KHC_{2}O_{4}. H_{2}C_{2}O_{4}. 2H_{2}O) weighing 0. 2162 gram andrequires 49. 14 cc. How many cc. Of water must be added to 1000 cc. Ofthe aqueous hydrochloric acid?

!Answer!: 11 cc. 

32. How many cubic centimeters of 3 N phosphoric acid must be added to300 cc. Of 0. 4 N phosphoric acid in order that the resulting solutionmay be 0. 6 N?

!Answer!: 25 cc. 

33. To oxidize the iron in 1 gram ofFeSO_{4}(NH_{4})_{2}SO_{4}. 6H_{2}O (mol. Wgt. 392) requires 3 cc. Ofa given solution of HNO_{3}. What is the normality of the nitricacid 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 thesame pressure is liberated from a 1 gram sample of dolomite, by addingan excess of hydrochloric acid, as can be liberated by the addition of35 cc. Of 0. 5 N hydrochloric acid to an excess of any pure or impurecarbonate. 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 thechloride in the calcium chloride formed by the action of 100 cc. Of0. 1072 N hydrochloric acid on an excess of calcium carbonate, and howmany 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 onanalysis to contain 7. 62% water, 2. 38% K_{2}CO_{3}. And 90% KOH. Whatweight of residue will be obtained if one gram of this sample is addedto 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 statementthat they contained either pure NaOH; pure Na_{2}CO_{3}; pureNaHCO_{3}, or mixtures of these substances. From the following dataidentify them:

Sample I. On adding phenolphthalein to a solution of the substance, itgave no color to the solution. 

Sample II. On titrating with standard acid, it required 15. 26 cc. Fora 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 thepink of phenolphthalein disappeared, and on the addition of methylorange 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 requiredexactly 30 cc. Of the same acid for neutralization, using methylorange. 

!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 followingdata 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 1cc. 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 tothe air for some time, is dissolved in water and diluted to exactly500 cc. One hundred cubic centimeters of the solution, when titratedwith 0. 1062 N hydrochloric acid, using methyl orange as an indicator, requires 38. 60 cc. For complete neutralization. Barium chloride inexcess is added to a second portion of 100 cc. Of the solution, whichis diluted to exactly 250 cc. , allowed to stand and filtered. Twohundred cubic centimeters of this filtrate require 29. 62 cc. Of 0. 1062N hydrochloric acid for neutralization, using phenolphthalein as anindicator. 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 beenexposed to the air) was titrated against a standard acid using methylorange as an indicator, and was found to be exactly 0. 1 N. Thissolution was used in the analysis of a material sold at 2 cents perpound per cent of an acid constituent A, and always mixed so thatit was supposed to contain 15% of A, on the basis of the analyst'sreport. Owing to the carelessness of the analyst's assistant, thesodium hydroxide solution was used with phenolphthalein as anindicator in cold solution in making the analyses. The concernmanufacturing this material sells 600 tons per year, and when themistake was discovered it was estimated that at the end of a yearthe error in the use of indicators would either cost them or theircustomers $6000. Who would lose and why? Assuming the impure NaOH usedoriginally in making the titrating solution consisted of NaOH andNa_{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 ironwire, 99. 85% pure, 42. 42 cc. Of the solution were added. The weight ofthe wire used was 0. 22 gram. 3. 27 cc. Of a ferrous sulphate solutionhaving a normal value as a reducing agent of 0. 1011 were addedto complete the titration. Calculate the normal value of theK_{2}Cr_{2}O_{7}. 

!Answer!: 0. 1006 N. 

42. What weight of iron ore containing 56. 2% Fe should be taken tostandardize an approximately 0. 1 N oxidizing solution, if not morethan 47 cc. Are to be used?

!Answer!: 0. 4667 gram. 

43. One tenth gram of iron wire, 99. 78% pure, is dissolved inhydrochloric acid and the iron oxidized completely with bromine water. How many grams of stannous chloride are there in a liter of solutionif it requires 9. 47 cc. To just reduce the iron in the above? Whatis the normal value of the stannous chloride solution as a reducingagent?

!Answer!: 17. 92 grams; 0. 1888 N. 

44. One gram of an oxide of iron is fused with potassium acid sulphateand the fusion dissolved in acid. The iron is reduced with stannouschloride, mercuric chloride is added, and the iron titrated with anormal K_{2}Cr_{2}O_{7} solution. 12. 94 cc. Were used. What is theformula 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 reductionwith stannous chloride could be oxidized by bichromate to a statecorresponding to an XO_{4}^{-} anion, compute the oxide, or valence, corresponding to the reduced state from the following data: 0. 3266gram of the pure element, after being dissolved, was reduced withstannous chloride and oxidized by 40 cc. Of K_{2}Cr_{2}O_{7}, of whichone 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 thefollowing 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} andAl_{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 limefrom the above the carbon dioxide is reduced to 3. 00%. How many cubiccentimeters of normal KMnO_{4} will be required to determine thecalcium oxide volumetrically in a 1 gram sample of the lime?

!Answer!: 20. 08 cc. 

49. If 100 cc. Of potassium bichromate solution (10 gramK_{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) aremixed, and the resulting solution titrated with 0. 2121 N KMnO_{4}, howmany cubic centimeters of the KMnO_{4} solution will be required tooxidize the iron?

!Answer!: 5. 70 cc. 

50. If a 0. 5000 gram sample of limonite containing 59. 50 per centFe_{2}O_{3} requires 40 cc. Of KMnO_{4} to oxidize the iron, whatis 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. 9000gram of oxalic acid. The excess oxalic acid requires 23. 95 cc. Ofpermanganate (1 cc. = 0. 03038 gram FeSO_{4}. 7H_{2}O). What is thepercentage of MnO_{2}, in the sample?

!Answer!: 84. 47%. 

52. A solution contains 50 grams ofKHC_{2}O_{4}. H_{2}C_{2}O_{4}. 2H_{2}O per liter. What is the normalvalue 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}, asample weighing 0. 5000 gram is taken and the iron is reduced withsulphurous acid. On account of failure to boil out all the excessSO_{2}, 38. 60 cubic centimeters of 0. 1 N KMnO_{4} were required totitrate the solution. What was the error, percentage error, and whatweight 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 asan 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, whenstandardized gravimetrically as silver chloride, yields a precipitateweighing 0. 5465 gram. 24. 35 cc. Of the hydrochloric acid are exactlyequivalent to 30. 17 cc. Of KHC_{2}O_{4}. H_{2}C_{2}O_{4}. 2H_{2}Osolution. How much water must be added to a liter of the oxalatesolution 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 toexactly 1000 cc. The normal value of the oxalate solution when used asan acid is 0. 1315. Calculate the ratio of tetroxalate to oxalate usedin making up the solution and the normal value of the solution as areducing 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 formerto be 0. 07500 N as an alkali and the latter exactly 0. 1 N as anoxidizing agent. By coincidence, exactly 47. 26 cc. Were used in eachstandardization. Find the ratio of the oxalate used in theNaOH standardization to the oxalate used in the permanganatestandardization. 

!Answer!: 1:1. 

58. A sample of apatite weighing 0. 60 gram is analyzed for itsphosphoric anhydride content. If the phosphate is precipitated as(NH_{4})_{3}PO_{4}. 12MoO_{3}, and the precipitate (after solution andreduction of the MoO_{3} to Mo_{24}O_{37}), requires 100 cc. Of normalKMnO_{4} to oxidize it back to MoO_{3}, what is the percentage ofP_{2}O_{5}?

!Answer!: 33. 81%. 

59. In the analysis of a sample of steel weighing 1. 881 grams thephosphorus was precipitated with ammonium molybdate and the yellowprecipitate was dissolved, reduced and titrated with KMnO_{4}. If thesample 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. 007188gram 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. Equivalentto 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 totitrate the iodine liberated from an excess of potassium iodidesolution 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 thesodium 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 thiosulphatesolution is allowed to stand. One per cent by weight of thethiosulphate is decomposed by the carbonic acid present in thesolution. To what volume must the solution be diluted to make itexactly 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 foranalysis. A student titrates it with a solution of iodine of which 1cc. Is equivalent to 0. 004950 gram of As_{2}O_{3}. Due to an error onhis part in standardization, the student's analysis shows the sampleto contain 70. 32% Sb. Calculate the true normal value of the iodinesolution, 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 anexcess of hydrochloric acid, the liberated chlorine is passed intopotassium iodide and the liberated iodine is titrated with sodiumthiosulphate solution (49. 66 grams of pure Na_{2}S_{2}O_{3}. 5H_{2}Oper liter). If 38. 72 cc. Are required, what volume of 0. 25 normalpermanganate solution will be required in an indirect determinationin which a similar sample is reduced with 0. 9012 gramH_{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 sulphuras hydrogen sulphide, precipitating cadmium sulphide by passing theliberated hydrogen sulphide through ammoniacal cadmium chloridesolution, and decomposing the CdS with acid in the presence of ameasured amount of standard iodine, the following data are obtained:Sample, 5. 027 grams; cc. Na_{2}S_{2}O_{3} sol. = 12. 68; cc. Iodinesol. = 15. 59; 1 cc. Iodine sol. = 1. 086 cc. Na_{2}S_{2}O_{3} sol. ; 1cc. Na_{2}S_{2}O_{3}= 0. 005044 gram Cu. Calculate the percentage ofsulphur. (H_{2}S + I_{2} = 2HI + S. )

!Answer!: 0. 107%. 

66. Given the following data, calculate the percentage of iron ina sample of crude ferric chloride weighing 1. 000 gram. The iodineliberated by the reaction 2FeCl_{3}+ 2HI = 2HCl + 2FeCl_{2} + I_{2} isreduced by the addition of 50 cc. Of sodium thiosulphate solution andthe excess thiosulphate is titrated with standard iodine and requires7. 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. Arsenitesolution = 0. 005160 gram As_{2}O_{3}. 

!Answer!: 23. 77%. 

67. Sulphide sulphur was determined in a sample of reduced bariumsulphate by the evolution method, in which the sulphur was evolved ashydrogen sulphide and was passed into CdCl_{2} solution, the acidifiedprecipitate being titrated with iodine and thiosulphate. Sample, 5. 076grams; 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 oxidizethe same amount of oxalic acid as 37. 12 cc. Of a permanganatesolution, 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% pureMnO_{2}. The iodine liberated from KI by the manganese dioxide issufficient to react with 46. 24 cc. Of Na_{2}S_{2}O_{3} sol. What isthe normal value of the thiosulphate?

!Answer!:: 0. 1105 N. 

70. In the volumetric analysis of silver coin (90% Ag), using a0. 5000 gram sample, what is the least normal value that a potassiumthiocyanate solution may have and not require more than 50 cc. Ofsolution in the analysis?

!Answer!: 0. 08339 N. 

71. A mixture of pure lithium chloride and barium bromide weighing0. 6 gram is treated with 45. 15 cubic centimeters of 0. 2017 N silvernitrate, and the excess titrated with 25 cc. Of 0. 1 N KSCN solution, using ferric alum as an indicator. Calculate the percentage of brominein the sample. 

!Answer!: 40. 11%. 

72. A mixture of the chlorides of sodium and potassium from 0. 5000gram of a feldspar weighs 0. 1500 gram, and after solution in waterrequires 22. 71 cc. Of 0. 1012 N silver nitrate for the precipitation ofthe chloride ions. What are the percentages of Na_{2}O and K_{2}O inthe 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 excessof 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) MgOin Mg_{2}P_{2}O_{7}; (c) P_{2}O_{5} in Mg_{2}P_{2}O_{7}; (d) Fe inFe_{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 ofMnO_{2}?

!Answer!: 0. 8774 gram. 

77. If a sample of silver coin weighing 0. 2500 gram gives aprecipitate of AgCl weighing 0. 2991 gram, what weight of AgI couldhave been obtained from the same weight of sample, and what is thepercentage of silver in the coin?

!Answers!: 0. 4898 gr. ; 90. 05%. 

78. How many cubic centimeters of hydrochloric acid (sp. Gr. 1. 13containing 25. 75% HCl by weight) are required to exactly neutralize25 cc. Of ammonium hydroxide (sp. Gr. . 90 containing 28. 33% NH_{3} byweight)?

!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 precipitatethe 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 ignitedprecipitate?

!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 ofFeSO_{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) isdissolved 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 neutralizethe 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 added13. 00 cc. Of aqueous ammonia (sp. Gr. 0. 90 containing 28. 4% NH_{3} byweight). How many cubic centimeters of sulphuric acid (sp. Gr. 1. 18containing 24. 7% H_{2}SO_{4} by weight) must be added to the mixturein order to bring the aluminium into solution?

!Answer!: 34. 8 cc. 

83. How many cubic centimeters of sulphurous acid (sp. Gr. 1. 04containing 75 grams SO_{2} per liter) are required to reduce theiron 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 bichromatecontaining 26. 30 grams of K_{2}Cr_{2}O_{7} per liter must be takenin order to yield 0. 6033 gram of Cr_{2}O_{3} after reduction andprecipitation 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. 946containing 13. 88% NH_{3} by weight) are required to precipitatethe iron as Fe(OH)_{3} from a sample of pureFeSO_{4}. (NH_{4})_{2}SO_{4}. 6H_{2}O, which requires 0. 34 cc. Of nitricacid (sp. Gr. 1. 350 containing 55. 79% HNO_{3} by weight) for oxidationof the iron? (See problem No. 80 for reaction. )

!Answer!: 4. 74 cc. 

86. In the analysis of an iron ore by solution, oxidation andprecipitation of the iron as Fe(OH)_{3}, what weight of sample must betaken for analysis so that each one hundredth of a gram of the ignitedprecipitate of Fe_{2}O_{3} shall represent one tenth of one per centof iron?

!Answer!: 6. 99 grams. 

87. What weight in grams of impure ferrous ammonium sulphate shouldbe taken for analysis so that the number of centigrams of BaSO_{4}obtained will represent five times the percentage of sulphur in thesample?

!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}, thepercentage of Fe_{2}O_{4} in the sample may be found by multiplyingthe 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} toarsenic acid, it is precipitated with "magnesia mixture" (MgCl_{2} +2NH_{4}Cl). If exactly 12. 6 cc. Of the mixture are required, how manygrams 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 ofthe sample gives 0. 4013 gram of Mg_{2}P_{2}O_{7}, how many cubiccentimeters of ammonium oxalate solution (containing 40 grams of(NH_{4})_{2}C_{2}O_{4}. H_{2}O per liter) would be required toprecipitate 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 purecalcium carbonate, when treated with an excess of hydrochloric acid, yields 0. 3117 gram of carbon dioxide, calculate the percentage ofmagnesium oxide and of calcium oxide in the sample. 

!Answers!: 13. 22% MgO; 40. 54% CaO. 92. The calcium in a sample ofdolomite weighing 0. 9380 gram is precipitated as calcium oxalate andignited to calcium oxide. What volume of gas, measured over waterat 20°C. And 765 mm. Pressure, is given off during ignition, if theresulting oxide weighs 0. 2606 gram? (G. M. V. = 22. 4 liters; V. P. Waterat 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 thelimestone, assuming complete conversion to oxide. What weight ofMg_{2}P_{2}O_{7} could be obtained from a 3-gram sample of thelimestone?

!Answers!: 1. 565 tons; 0. 2044 gram. 

94. A sample of dolomite is analyzed for calcium by precipitatingas the oxalate and igniting the precipitate. The ignited product isassumed to be CaO and the analyst reports 29. 50% Ca in the sample. Owing to insufficient ignition, the product actually contained 8% ofits weight of CaCO_{3}. What is the correct percentage of calcium inthe sample, and what is the percentage error?

!Answers!: 28. 46%; 3. 65% error. 

95. What weight of impure calcite (CaCO_{3}) should be taken foranalysis so that the volume in cubic centimeters of CO_{2} obtained bytreating with acid, measured dry at 18°C. And 763 mm. , shall equal thepercentage of CaO in the sample?

!Answer!: 0. 2359 gram. 

96. How many cubic centimeters of HNO_{3} (sp. Gr. 1. 13 containing21. 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 thenitric acid to NO by each constituent? What fraction of this volume ofacid is used for oxidation?

!Answers!: 55. 06 cc. ; 25%. 

97. What weight of metallic copper will be deposited from a cupricsalt solution by a current of 1. 5 amperes during a period of 45minutes, assuming 100% current efficiency? (1 Faraday = 96, 500coulombs. )

!Answer!: 1. 335 grams. 

98. In the electrolysis of a 0. 8000 gram sample of brass, there isobtained 0. 0030 gram of PbO_{2}, and a deposit of metallic copperexactly equal in weight to the ignited precipitate of Zn_{2}P_{2}O_{7}subsequently obtained from the solution. What is the percentagecomposition 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) weighing0. 9400 gram is dissolved in nitric acid. The lead is determined byweighing as PbSO_{4}, the copper by electrolysis and the zinc byprecipitation with (NH_{4})_{2}HPO_{4} in a neutral solution. 

(a) Calculate the cubic centimeters of nitric acid (sp. Gr. 1. 42containing 69. 90% HNO_{3} by weight) required to just dissolve thebrass, assuming reduction to NO. 

!Answer!: 2. 48 cc. 

(b) Calculate the cubic centimeters of sulphuric acid (sp. Gr. 1. 84containing 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 weightafter the copper has been deposited. 

!Answer!: 11. 6116 grams. 

(e) Calculate the grams of (NH_{4})_{2}HPO_{4} required to precipitatethe 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 ismade in weighing a 2. 5 gram portion by which 0. 001 gram too much isweighed out, what percentage error in the zinc determination wouldresult? What volume of a solution of sodium hydrogen phosphate, containing 90 grams of Na_{2}HPO_{4}. 12H_{2}O per liter, would berequired to precipitate the zinc as ZnNH_{4}PO_{4} and what weight ofprecipitate 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 itsonly impurity, weighs 0. 5000 gram and loses 0. 1000 gram on ignition. What volume of disodium phosphate solution (containing 90 gramsNa_{2}HPO_{4}. 12H_{2}O per liter) will be required to precipitate themagnesium as magnesium ammonium phosphate?

!Answer!: 9. 07 cc. 

102. 2. 62 cubic centimeters of nitric acid (sp. Gr. 1. 42 containing69. 80% HNO_{2} by weight) are required to just dissolve a sampleof 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 everycase, calculate the weight of the brass and the cubic centimeters ofacid used as acid. 

!Answer!: 0. 992 gram; 1. 97 cc. 

103. One gram of a mixture of silver chloride and silver bromide isfound to contain 0. 6635 gram of silver. What is the percentage ofbromine?

!Answer!: 21. 30%. 

104. A precipitate of silver chloride and silver bromide weighs 0. 8132gram. On heating in a current of chlorine, the silver bromide isconverted to silver chloride, and the mixture loses 0. 1450 gramin weight. Calculate the percentage of chlorine in the originalprecipitate. 

!Answer!: 6. 13%. 

105. A sample of feldspar weighing 1. 000 gram is fused and the silicadetermined. The weight of silica is 0. 6460 gram. This is fused with 4grams of sodium carbonate. How many grams of the carbonate actuallycombined with the silica in fusion, and what was the loss in weightdue 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 gramsis transformed into mixed sulphates, weighing 5. 023 grams. Calculatethe 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 fewof its applications are intended for reference in connection with theexplanations which are given in the Notes accompanying the variousprocedures. The reader who desires a more extended discussion of thefundamental theory and its uses is referred to such books as Talbotand Blanchard's !Electrolytic Dissociation Theory! (MacmillanCompany), or Alexander Smith's !Introduction to General InorganicChemistry! (Century Company). 

The !electrolytic dissociation theory!, as propounded by Arrhenius in1887, 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 orgroups of atoms, as, for example, H^{+} and Br^{-} from hydrobromicacid, Na^{+} and OH^{-} from sodium hydroxide, 2NH_{4}^{+} andSO_{4}^{--} from ammonium sulphate. The unit charge is that which isdissociated with a hydrogen ion. Those upon other ions vary in signand number according to the chemical character and valence of theatoms or radicals of which the ions are composed. In any solution theaggregate of the positive charges upon the positive ions (!cations!)must always balance the aggregate negative charges upon the negativeions (!anions!). 

It is assumed that the Na^{+} ion, for example, differs from thesodium atom in behavior because of the very considerable electricalcharge which it carries and which, as just stated, must, in anelectrically neutral solution, be balanced by a corresponding negativecharge on some other ion. When an electric current is passed through asolution of an electrolyte the ions move with and convey the current, and when the cations come into contact with the negatively chargedcathode they lose their charges, and the resulting electricallyneutral atoms (or radicals) are liberated as such, or else enter atonce into chemical reaction with the components of the solution. 

Two ions of identically the same composition but with differentelectrical charges may exhibit widely different properties. Forexample, the ion MnO_{4}^{-} from permanganates yields a purple-redsolution and differs in its chemical behavior from the ionMnO_{4}^{--} from manganates, the solutions of which are green. 

The chemical changes upon which the procedures of analytical chemistrydepend are almost exclusively those in which the reacting substancesare electrolytes, and analytical chemistry is, therefore, essentiallythe chemistry of the ions. The percentage dissociation of the sameelectrolyte tends to increase with increasing dilution of itssolution, although not in direct proportion. The percentagedissociation of different electrolytes in solutions of equivalentconcentrations (such, for example, as normal solutions) varies widely, as is indicated in the following tables, in which approximate figuresare 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 electricalconductivity 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 theconcentration of the H^{+} ions and is measured by the percentagedissociation in solutions of equivalent concentration. The commonmineral acids are largely dissociated and therefore give a relativelyhigh concentration of H^{+} ions, and are commonly known as "strongacids. " The organic acids, on the other hand, belong generally to thegroup of "weak acids. "

!Bases! yield OH^{-} ions in water solution, and the comparativestrength of the bases is measured by their relative dissociation insolutions of equivalent concentration. Ammonium hydroxide is a weakbase, as shown in the table above, while the hydroxides of sodium andpotassium exhibit strongly basic properties. 

Ionic reactions are all, to a greater or less degree, !reversiblereactions!. A typical example of an easily reversible reaction is thatrepresenting the changes in ionization which an electrolyte such asacetic 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 wasstated above, the ionization increases with dilution, the reactionthen proceeding from left to right, while concentration of thesolution occasions a partial reassociation of the ions, and thereaction proceeds from right to left. To understand the principleunderlying these changes it is necessary to consider first theconditions which prevail when a solution of acetic acid, which hasbeen stirred until it is of uniform concentration throughout, has cometo a constant temperature. A careful study of such solutions has shownthat there is a definite state of equilibrium between the constituentsof the solution; that is, there is a definite relation between theundissociated acetic acid and its ions, which is characteristic forthe prevailing conditions. It is not, however, assumed that this is acondition of static equilibrium, but rather that there is continualdissociation and association, as represented by the opposingreactions, the apparent condition of rest resulting from the fact thatthe amount of change in one direction during a given time is exactlyequal to that in the opposite direction. A quantitative study ofthe amount of undissociated acid, and of H^{+} ions andC_{2}H_{3}O_{2}^{-} ions actually to be found in a large number ofsolutions of acetic acid of varying dilution (assuming them to be ina condition of equilibrium at a common temperature), has shown thatthere is always a definite relation between these three quantitieswhich 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 betweenthe product of the concentrations of the ions and the concentration ofthe undissociated acid when conditions of equilibrium prevail. 

It has been found, further, that a similar statement may be maderegarding all reversible reactions, which may be expressed in generalterms thus: The rate of chemical change is proportional to the productof 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, thenthe product of the concentrations of the substances entering intothe reaction stands in a constant ratio to the product of theconcentrations of the resulting substances, as given in the expressionabove 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 acidapplies to a wide range of dilutions, provided the temperature remainsconstant. If the temperature changes the value of the constant changessomewhat, but is again uniform for different dilutions at thattemperature. The following data are given for temperatures of about18°C. [1]

========================================================================== | | | | MOLAL | FRACTION | MOLAL CONCENTRA- | MOLAL CONCENTRA- | VALUE OFCONCENTRATION | 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 agram-molecule per liter of the undissociated acid, and to fractions ofthe corresponding quantities of H^{+} and C_{2}H_{3}O_{2}^{-} ionsper liter which would result from the complete dissociation of agram-molecule of acetic acid. The values calculated for the constantare subject to some variation on account of experimental errors indetermining the percentage ionized in each case, but the approximateagreement between the values found for molal and centimolal (onehundredfold dilution) is significant. 

The figures given also illustrate the general principle, that the!relative! ionization of an electrolyte increases with the dilution ofits solution. If we consider what happens during the (usually) briefperiod of dilution of the solution from molal to 0. 1 molal, forexample, it will be seen that on the addition of water the conditionsof concentration which led to equality in the rate of change, andhence to equilibrium in the molal solution, cease to exist; and sincethe dissociating tendency increases with dilution, as just stated, it is true at the first instant after the addition of water that theconcentration of the undissociated acid is too great to bepermanent under the new conditions of dilution, and the reaction, HC_{2}H_{3}O_{2} H^{+} + C_{2}H_{3}O_{2}^{-}, will proceed fromleft to right with great rapidity until the respective concentrationsadjust themselves to the new conditions. 

That which is true of this reaction is also true of all reversiblereactions, namely, that any change of conditions which occasionsan increase or a decrease in concentration of one or more of thecomponents causes the reaction to proceed in one direction or theother until a new state of equilibrium is established. This principleis constantly applied throughout the discussion of the applicationsof the ionic theory in analytical chemistry, and it should be clearlyunderstood that whenever an existing state of equilibrium is disturbedas a result of changes of dilution or temperature, or as a consequenceof chemical changes which bring into action any of the constituents ofthe solution, thus altering their concentrations, there is always atendency to re-establish this equilibrium in accordance with the law. Thus, if a base is added to the solution of acetic acid the H^{+} ionsthen unite with the OH^{-} ions from the base to form undissociatedwater. The concentration of the H^{+} ions is thus diminished, andmore 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 withchloride ions, or barium ions react with sulphate ions. In the formercase the dissociation reaction of the silver nitrate is AgNO_{3} Ag^{+} + NO_{3}^{-}, and as soon as the Ag^{+} ions unite with theCl^{-} ions the concentration of the former is diminished, more of theAgNO_{3} dissociates, and this process goes on until the Ag^{+} ionsare practically all removed from the solution, if the Cl^{-} ions arepresent in sufficient quantity. 

For the sake of accuracy it should be stated that the mass law cannotbe rigidly applied to solutions of those electrolytes which arelargely dissociated. While the explanation of the deviation fromquantitative exactness in these cases is not known, the law is stillof marked service in developing analytical methods along more logicallines than was formerly practicable. It has not seemed wise to qualifyeach statement made in the Notes to indicate this lack of quantitativeexactness. The student should recognize its existence, however, andwill realize its significance better as his knowledge of physicalchemistry increases. 

If we apply the mass law to the case of a substance of smallsolubility, such as the compounds usually precipitated in quantitativeanalysis, we derive what is known as the !solubility product!, asfollows: Taking silver chloride as an example, and remembering that itis not absolutely insoluble in water, the equilibrium expression forits 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 thesolid precipitate must be saturated for the existing temperature, andthe quantity of undissociated AgCl in the solution is definite andconstant for that temperature. Since it is a constant, it may beeliminated, and the expression becomes !Conc'n Ag^{+} x Conc'nCl^{-} = Constant!, and this is known as the solubility product. Noprecipitation of a specific substance will occur until the product ofthe concentrations of its ions in a solution exceeds the solubilityproduct for that substance; whenever that product is exceededprecipitation must follow. 

It will readily be seen that if a substance which yields an ion incommon with the precipitated compound is added to such a solution ashas just been described, the concentration of that ion isincreased, and as a result the concentration of the other ion mustproportionately decrease, which can only occur through the formationof some of the undissociated compound which must separate from thealready saturated solution. This explains why the addition of anexcess of the precipitant is often advantageous in quantitativeprocedures. Such a case is discussed at length in Note 2 on page 113. 

Similarly, the ionization of a specific substance in solution tends todiminish on the addition of another substance with a common ion, as, for instance, the addition of hydrochloric acid to a solutionof hydrogen sulphide. Hydrogen sulphide is a weak acid, and theconcentration of the hydrogen ions in its aqueous solutions is verysmall. 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 amarked increase in the concentration of the H^{+} ions, such as wouldresult from the addition of even a small amount of the highly ionizedhydrochloric acid, displaces the point of equilibrium and some of theS^{--} ions unite with H^{+} ions to form undissociated H_{2}S. Thisis of much importance in studying the reactions in which hydrogensulphide is employed, as in qualitative analysis. By a parallel courseof reasoning it will be seen that the addition of a salt of a weakacid 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 wherechemical action is involved, it should be remembered that no substanceis completely insoluble in water, and that those products of achemical 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 formundissociated water. The concentration of the OH^{-} ions is thusdiminished, more Mg(OH)_{2} dissociates, the solution is no longersaturated with the undissociated compound, and more of the soliddissolves. This process repeats itself with great rapidity until, ifsufficient acid is present, the solid passes completely into solution. 

Exactly the same sort of process takes place if calcium oxalate, forexample, is dissolved in hydrochloric acid. The C_{2}O_{4}^{--} ionsunite with the H^{+} ions to form undissociated oxalic acid, the acidbeing less dissociated than normally in the presence of the H^{+} ionsfrom the hydrochloric acid (see statements regarding hydrogen sulphideabove). As the undissociated oxalic acid forms, the concentration ofthe C_{2}O_{4}^{--} ions lessens and more CaC_{2}O_{4} dissolves, as described for the Mg(OH)_{2} above. Numerous instances of theapplications 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. Ofdissociated water in 10^{7} liters), yet under some conditions theymay give rise to important consequences. The term !hydrolysis! isapplied 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, whendissolved in water gives an alkaline solution because some of theH^{+} ions from the water unite with CN^{-} ions to form (HCN), whichis a very weak acid, and is but very slightly dissociated. Potassiumhydroxide, which might form from the OH^{-} ions, is so largelydissociated that the OH^{-} ions remain as such in the solution. Theunion of the H^{+} ions with the CN^{-} ions to form the undissociatedHCN diminishes the concentration of the H^{+} ions, and more waterdissociates (H_{2}O H^{+} + OH^{-}) to restore the equilibrium. It is clear, however, that there must be a gradual accumulation ofOH^{-} ions in the solution as a result of these changes, causing thesolution to exhibit an alkaline reaction, and also that ultimately thefurther dissociation of the water will be checked by the presence ofthese ions, just as the dissociation of the H_{2}S was lessened by theaddition of HCl. 

An exactly opposite result follows the solution of such a salt asAl_{2}(SO_{4})_{3} in water. In this case the acid is strong and thebase weak, and the OH^{-} ions form the little dissociated Al(OH)_{3}, while the H^{+} ions remain as such in the solution, sulphuric acidbeing extensively dissociated. The solution exhibits an acid reaction. 

Such hydrolytic processes as the above are of great importance inanalytical chemistry, especially in the understanding of the action ofindicators in volumetric analysis. (See page 32. )

The impelling force which causes an element to pass from the atomicto the ionic condition is termed !electrolytic solution pressure!, orionization tension. This force may be measured in terms of electricalpotential, and the table below shows the relative values for a numberof elements. 

In general, an element with a greater solution pressure tends to causethe deposition of an element of less solution pressure when placed ina solution of its salt, as, for instance, when a strip of zinc oriron is placed in a solution of a copper salt, with the resultingprecipitation of metallic copper. 

Hydrogen is included in the table, and its position should be notedwith reference to the other common elements. For a more extendeddiscussion 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. 13Calcium Ca^{++} | | Hydrogen H^{+} | -0. 28Magnesium Mg^{++} | | Bismuth Bi^{+++}|Aluminum A1^{+++} | +1. 00 | Antimony | -0. 75Manganese Mn^{++} | | Arsenic |Zinc Zn^{++} | +0. 49 | Copper Cu^{++} | -0. 61Cadmium Cd^{++} | +0. 14 | Mercury Hg^{+} | -1. 03Iron Fe^{++} | +0. 063 | Silver Ag^{+} | -1. 05Cobalt 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 alongthe radius at right angles to the original fold, a cone is formed onopening, the angle of which is 60°. Funnels for analytical use aresupposed to have the same angle, but are rarely accurate. It ispossible, however, with care, to fit a filter thus folded into afunnel in such a way as to prevent air from passing down between thepaper and the funnel to break the column of liquid in the stem, which aids greatly, by its gentle suction, in promoting the rate offiltration. 

Such a filter has, however, the disadvantage that there are threethicknesses of paper back of half of its filtering surface, as aconsequence of which one half of a precipitate washes or drains moreslowly. Much time may be saved in the aggregate by learning to fold afilter in such a way as to improve its effective filtering surface. The directions which follow, though apparently complicated on firstreading, are easily applied and easily remembered. Use a 6-inch filterfor practice. Place four dots on the filter, two each on diameterswhich are at right angles to each other. Then proceed as follows:(1) Fold the filter evenly across one of the diameters, creasing itcarefully; (2) open the paper, turn it over, rotate it 90° to theright, bring the edges together and crease along the other diameter;(3) open, and rotate 45° to the right, bring edges together, andcrease evenly; (4) open, and rotate 90° to the right, and creaseevenly; (5) open, turn the filter over, rotate 22-(1/2)° to the right, and crease evenly; (6) open, rotate 45° to the right and creaseevenly; (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 theright hand, fold the second crease to the left over on it, and dothe same with each of the other dots. Place it, thus folded, in thefunnel, moisten it, and fit to the side of the funnel. The filter willthen have four short segments where there are three thicknessesand four where there is one thickness, but the latter are evenlydistributed around its circumference, thus greatly aiding the passageof liquids through the paper and hastening both filtration and washingof 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 | CorrectedFinal Reading HCl | 48. 17 48. 08 | 43. 20 43. 14Initial Reading HCl | 0. 12 . 12 | . 17 . 17 | ----- ----- | ----- ----- | 47. 96 | 42. 97 | | | Corrected | CorrectedFinal Reading HCl | 46. 36 46. 29 | 40. 51 40. 37Initial 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. 77Initial Reading HCl | . 00 . 00 | . 04 . 04 | ----- ----- | ----- ----- | 39. 83 | 49. 73 | |Final Reading NaOH | . 26 . 26 | . 67 . 67Initial Reading NaOH | . 12 . 12 | . 36 . 36 | --- --- | --- --- | . 14 | . 31 | | | . 14 | . 31Corrected cc. HCl | 39. 83 - ----- = 39. 68 | 49. 73 - ----- = 49. 40 | . 9286 | . 9286 | |log sample | 0. 0025 | 0. 0983colog cc | 8. 4014 - 10 | 8. 3063 - 10colog 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 usein 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. 25Barium chloride, BaCl_{2}. 2H_{2}O 25 0. 2N 0. 1Magnesium ammonium chloride (of MgCl_{2}) 71 1. 5N 0. 75Mercuric chloride, HgCl_{2} 45 0. 33N 0. 66Potassium hydroxide, KOH (sp. Gr. 1. 27) 480Potassium thiocyanate, KSCN 5 0. 05N 0. 55Silver nitrate, AgNO_{3} 21 0. 125N 0. 125Sodium hydroxide, NaOH 100 2. 5N 2. 5Sodium carbonate. Na_{2}CO_{3} 159 3N 1. 5Sodium 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 addinga slight excess of acid from time to time. A strip of metallic tin iskept in the bottle. 

A solution of ammonium molybdate is best prepared as follows: Stir100 grams of molybdic acid (MoO_{3}) into 400 cc. Of cold, distilledwater. Add 80 cc. Of concentrated ammonium hydroxide (sp. Gr. 0. 90). Filter, and pour the filtrate slowly, with constant stirring, into amixture of 400 cc. Concentrated nitric acid (sp. Gr. 1. 42) and 600cc. 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 ammoniaare based upon percentages given in the Standard Tables of theManufacturing Chemists' Association of the United States [!J. S. C. I. !, 24 (1905), 787-790]. All gravities are taken at 15. 5°C. And comparedwith water at the same temperature. 

Aqueous ammonia (sp. Gr. 0. 96) contains 9. 91 per cent NH_{3} byweight, 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} byweight, and corresponds to a 5. 6 N and 5. 6 molal solution. 

Hydrochloric acid (sp. Gr. 1. 12) contains 23. 81 per cent HCl byweight, and corresponds to a 7. 3 N and 7. 3 molal solution. 

Hydrochloric acid (sp. Gr. 1. 20) contains 39. 80 per cent HCl byweight, 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} byweight, 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} byweight, and corresponds to a 5. 9 N or 2. 95 molal solution. 

The term !normal! (N), as used above, has the same significance asin volumetric analyses. The molal solution is assumed to contain onemolecular 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!, thirdedition. 

CORRECTIONS FOR CHANGE OF TEMPERATURE OF STANDARD SOLUTIONS

The values below are average values computed from data relating to aconsiderable number of solutions. They are sufficiently accurate foruse in chemical analyses, except in the comparatively few caseswhere the highest attainable accuracy is demanded in chemicalinvestigations. The expansion coefficients should then be carefullydetermined for the solutions employed. For a compilation of theexisting 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 thetemperature of the solution varies from the standard temperatureselected for the laboratory. The total correction thus found issubtracted from the observed burette reading if the temperature ishigher than the standard, or added, if it is lower. Corrections arenot usually necessary for variations of temperature of 2°C. Or less. 

 INTERNATIONAL ATOMIC WEIGHTS

========================================================== | | | | 1920 | | 1920_________________|_________|___________________|__________ | | |Aluminium Al | 27. 1 | Molybdenum Mo | 96. 0Antimony Sb | 120. 2 | Neodymium Nd | 144. 3Argon A | 39. 9 | Neon Ne | 20. 2Arsenic As | 74. 96 | Nickel Ni | 58. 68Barium Ba | 137. 37 | Nitrogen N | 14. 008Bismuth Bi | 208. 0 | Osmium Os | 190. 9Boron B | 11. 0 | Oxygen O | 16. 00Bromine Br | 79. 92 | Palladium Pd | 106. 7Cadmium Cd | 112. 40 | Phosphorus P | 31. 04Caesium Cs | 132. 81 | Platinum Pt | 195. 2Calcium Ca | 40. 07 | Potassium K | 39. 10Carbon C | 12. 005 | Praseodymium Pr | 140. 9Cerium Ce | 140. 25 | Radium Ra | 226. 0Chlorine Cl | 35. 46 | Rhodium Rh | 102. 9Chromium Cr | 52. 0 | Rubidium Rb | 85. 45Cobalt Co | 58. 97 | Ruthenium Ru | 101. 7Columbium Cb | 93. 1 | Samarium Sm | 150. 4Copper Cu | 63. 57 | Scandium Sc | 44. 1Dysprosium Dy | 162. 5 | Selenium Se | 79. 2Erbium Er | 167. 7 | Silicon Si | 28. 3Europium Eu | 152. 0 | Silver Ag | 107. 88Fluorine Fl | 19. 0 | Sodium Na | 23. 00Gadolinium Gd | 157. 3 | Strontium Sr | 87. 63Gallium Ga | 69. 9 | Sulphur S | 32. 06Germanium Ge | 72. 5 | Tantalum Ta | 181. 5Glucinum Gl | 9. 1 | Tellurium Te | 127. 5Gold Au | 197. 2 | Terbium Tb | 159. 2Helium He | 4. 00 | Thallium Tl | 204. 0Hydrogen H | 1. 008 | Thorium Th | 232. 4Indium In | 114. 8 | Thulium Tm | 168. 5Iodine I | 126. 92 | Tin Sn | 118. 7Iridium Ir | 193. 1 | Titanium Ti | 48. 1Iron Fe | 55. 84 | Tungsten W | 184. 0Krypton Kr | 82. 92 | Uranium U | 238. 2Lanthanum La | 139. 0 | Vanadium V | 51. 0Lead Pb | 207. 2 | Xenon Xe | 130. 2Lithium Li | 6. 94 | Ytterbium Yb | 173. 5Lutecium Lu | 175. 0 | Yttrium Y | 88. 7Magnesium Mg | 24. 32 | Zinc Zn | 65. 37Manganese Mn | 54. 93 | Zirconium Zr | 90. 6Mercury Hg | 200. 6 | |==========================================================

INDEX

AcidimetryAcid solutions, normal standardAcids, definition ofAcids, weak, action of other acids on action of salts onAccuracy demandedAlkalimetryAlkali solutions, normal standardAlumina, determination of in stibniteAmmonium nitrate, acidAnalytical chemistry, subdivisions ofAntimony, determination of, in stibniteApatite, analysis ofAsbestos filtersAtomic weights, table of

Balances, essential features of use and care ofBarium sulphate, determination of sulphur inBases, definition ofBichromate process for ironBleaching powder, analysis ofBrass, analysis ofBurette, description of calibration of cleaning of reading of

Calcium, determination of, in limestoneCalibration, definition of of burettes of flasksCarbon dioxide, determination of, in limestoneChlorimetryChlorine, gravimetric determination ofChrome iron ore, analysis ofCoin, determination of silver inColloidal solution of precipitatesColorimetric analyses, definition ofCopper, determination of, in brass determination of in copper oresCrucibles, use ofCrystalline precipitates

Densities of waterDeposition potentialsDesiccatorsDirect methodsDissociation, degree of

Economy of timeElectrolytic dissociation, theory ofElectrolytic separations, principles ofEnd-point, definition ofEquilibrium, chemicalEvaporation of liquids

Faraday's lawFeldspar, analysis ofFerrous ammonium sulphate, analysis ofFilters, folding of how fittedFiltrates, testing ofFiltrationFlasks, graduation ofFunnelsFusions, removal of from crucibles

General directions for gravimetric analysis volumetric analysisGooch filterGravimetric analysis, definition of

Hydrochloric acid, standardization ofHydrolysis

Ignition of precipitatesIndicators, definition of for acidimetry preparation ofIndirect methodsInsoluble matter, determination of in limestoneIntegrityIodimetryIons, definition ofIron, gravimetric determination of volumetric determination of

Jones reductor

Lead, determination of in brassLimestone, analysis ofLimonite, determination of iron inLiquids, evaporation of transfer ofLitmusLogarithms

Magnesium, determination ofMass action, law ofMeasuring instrumentsMethyl orangeMoisture, determination of in limestone

Neutralization methodsNormal solutions, acid and alkali oxidizing agents reducing agentsNotebooks, sample pages of

Oxalic acid, determination of strength ofOxidation processesOxidizing power of pyrolusite

Permanganate process for ironPhenolphthaleinPhosphoric anhydride, determination ofPipette, calibration of description ofPlatinum crucibles, care ofPrecipitates, colloidal crystalline ignition of separation from filter washing ofPrecipitationPrecipitation methods (volumetric)ProblemsPyrolusite, oxidizing power of

Quantitative Analyses, subdivisions of

Reagents, strength ofReducing solution, normalReductor, JonesReversible reactions

Silica, determination of, in limestone determination of, in silicates purification ofSilicic acid, dehydration ofSilver, determination of in coinSoda ash, alkaline strength ofSodium chloride, determination of chlorine inSolubility productSolution pressureSolutions, normal standardStandardization, definition ofStandard solutions, acidimetry and alkalimetry chlorimetry iodimetry oxidizing and reducing agents thiocyanateStarch solutionsStibnite, determination of antimony inStirring rodsStoichiometryStrength of reagentsSuction, use ofSulphur, determination of in ferrous ammonium sulphate in barium sulphate

Temperature, corrections forTesting of washingsTheory of electrolytic dissociationThiocyanate process for silverTitration, definition ofTransfer of liquids

Volumetric analysis, definition of general directions

Wash-bottlesWashed filtersWashing of precipitatesWashings, testing ofWater, ionization of densities ofWeights, care of

Zimmermann-Reinhardt method for ironZinc, determination of, in brass