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Title: An Introductory Course of Quantitative Chemical Analysis

Author: Henry Paul Talbot

Release date: June 30, 2004 [eBook #12787]
Most recently updated: October 28, 2024

Language: English

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*** START OF THE PROJECT GUTENBERG EBOOK AN
INTRODUCTORY COURSE OF QUANTITATIVE CHEMICAL
ANALYSIS ***

[Transcriber's notes: In the chemical equations, superscripts are

indicated with a ^ and subscripts are indicated with a . The affected item
is enclosed in curly brackets {}. Examples are H^{+} for hydrogen ion and
H{2}O for water. Since the underscore is already being used in this project,
italics are designated by an exclamation point before and after the italicized
word or phrase.]

AN INTRODUCTORY COURSE
OF

QUANTITATIVE

CHEMICAL ANALYSIS

WITH

EXPLANATORY NOTES

BY

HENRY P. TALBOT

PROFESSOR OF INORGANIC CHEMISTRY AT THE MASSACHUSETTS INSTITUTE OF
TECHNOLOGY

SIXTH EDITION, COMPLETELY REWRITTEN

PREFACE

This Introductory Course of Quantitative Analysis has been prepared to
meet the needs of students who are just entering upon the subject, after a
course of qualitative analysis. It is primarily intended to enable the student
to work successfully and intelligently without the necessity for a larger
measure of personal assistance and supervision than can reasonably be
given to each member of a large class. To this end the directions are given
in such detail that there is very little opportunity for the student to go astray;
but the manual is not, the author believes, on this account less adapted for
use with small classes, where the instructor, by greater personal influence,
can stimulate independent thought on the part of the pupil.

The method of presentation of the subject is that suggested by Professor
A.A. Noyes' excellent manual of Qualitative Analysis. For each analysis the
procedure is given in considerable detail, and this is accompanied by
explanatory notes, which are believed to be sufficiently expanded to enable
the student to understand fully the underlying reason for each step
prescribed. The use of the book should, nevertheless, be supplemented by
classroom instruction, mainly of the character of recitations, and the student
should be taught to consult larger works. The general directions are
intended to emphasize those matters upon which the beginner in
quantitative analysis must bestow special care, and to offer helpful
suggestions. The student can hardly be expected to appreciate the force of

all the statements contained in these directions, or, indeed, to retain them all
in the memory after a single reading; but the instructor, by frequent
reference to special paragraphs, as suitable occasion presents itself, can
soon render them familiar to the student.

The analyses selected for practice are those comprised in the first course
of quantitative analysis at the Massachusetts Institute of Technology, and
have been chosen, after an experience of years, as affording the best
preparation for more advanced work, and as satisfactory types of
gravimetric and volumetric methods. From the latter point of view, they
also seem to furnish the best insight into quantitative analysis for those
students who can devote but a limited time to the subject, and who may
never extend their study beyond the field covered by this manual. The
author has had opportunity to test the efficiency of the course for use with
such students, and has found the results satisfactory.

In place of the usual custom of selecting simple salts as material for
preliminary practice, it has been found advantageous to substitute, in most
instances, approximately pure samples of appropriate minerals or industrial
products. The difficulties are not greatly enhanced, while the student gains
in practical experience.

The analytical procedures described in the following pages have been
selected chiefly with reference to their usefulness in teaching the subject,
and with the purpose of affording as wide a variety of processes as is
practicable within an introductory course of this character. The scope of the
manual precludes any extended attempt to indicate alternative procedures,
except through general references to larger works on analytical chemistry.
The author is indebted to the standard works for many suggestions for
which it is impracticable to make specific acknowledgment; no
considerable credit is claimed by him for originality of procedure.

For many years, as a matter of convenience, the classes for which this
text was originally prepared were divided, one part beginning with
gravimetric processes and the other with volumetric analyses. After a
careful review of the experience thus gained the conclusion has been
reached that volumetric analysis offers the better approach to the subject.
Accordingly the arrangement of the present (the sixth) edition of this
manual has been changed to introduce volumetric procedures first. Teachers
who are familiar with earlier editions will, however, find that the order of
presentation of the material under the various divisions is nearly the same
as that previously followed, and those who may still prefer to begin the
course of instruction with gravimetric processes will, it is believed, be able
to follow that order without difficulty.

Procedures for the determination of sulphur in insoluble sulphates, for the
determination of copper in copper ores by iodometric methods, for the
determination of iron by permanganate in hydrochloric acid solutions, and
for the standardization of potassium permanganate solutions using sodium
oxalate as a standard, and of thiosulphate solutions using copper as a
standard, have been added. The determination of silica in silicates
decomposable by acids, as a separate procedure, has been omitted.

The explanatory notes have been rearranged to bring them into closer
association with the procedures to which they relate. The number of
problems has been considerably increased.

The author wishes to renew his expressions of appreciation of the kindly
reception accorded the earlier editions of this manual. He has received
helpful suggestions from so many of his colleagues within the Institute, and
friends elsewhere, that his sense of obligation must be expressed to them
collectively. He is under special obligations to Professor L.F. Hamilton for
assistance in the preparation of the present edition.

HENRY P. TALBOT

!Massachusetts Institute of Technology, September, 1921!.

CONTENTS

PART I. INTRODUCTION

SUBDIVISIONS OF ANALYTICAL CHEMISTRY

GENERAL DIRECTIONS
Accuracy and Economy of Time; Notebooks; Reagents; Wash-bottles;
Transfer of Liquids

PART II. VOLUMETRIC ANALYSIS

GENERAL DISCUSSION
Subdivisions; The Analytical Balance; Weights; Burettes;
Calibration of Measuring Devices
GENERAL DIRECTIONS
Standard and Normal Solutions

!I. Neutralization Methods!

ALKALIMETRY AND ACIDIMETRY
Preparation and Standardization of Solutions; Indicators
STANDARDIZATION OF HYDROCHLORIC ACID
DETERMINATION OF TOTAL ALKALINE STRENGTH OF SODA ASH
DETERMINATION OF ACID STRENGTH OF OXALIC ACID

!II. Oxidation Processes!

GENERAL DISCUSSION BICHROMATE PROCESS FOR THE DETERMINATION OF
IRON DETERMINATION OF IRON IN LIMONITE BY THE BICHROMATE PROCESS
DETERMINATION OF CHROMIUM IN CHROME IRON ORE PERMANGANATE
PROCESS FOR THE DETERMINATION OF IRON DETERMINATION OF IRON IN
LIMONITE BY THE PERMANGANATE PROCESS DETERMINATION OF IRON IN
LIMONITE BY THE ZIMMERMANN-REINHARDT PROCESS DETERMINATION OF
THE OXIDIZING POWER OF PYROLUSITE IODIMETRY DETERMINATION OF
COPPER IN ORES DETERMINATION OF ANTIMONY IN STIBNITE CHLORIMETRY
DETERMINATION OF AVAILABLE CHLORINE IN BLEACHING POWDER

!III. Precipitation Methods!

DETERMINATION OF SILVER BY THE THIOCYANATE PROCESS

PART III. GRAVIMETRIC ANALYSIS

GENERAL DIRECTIONS
Precipitation; Funnels and Filters; Filtration and Washing of
Precipitates; Desiccators; Crucibles and their Preparation
for Use; Ignition of Precipitates
DETERMINATION OF CHLORINE IN SODIUM CHLORIDE
DETERMINATION OF IRON AND OF SULPHUR IN FERROUS
AMMONIUM SULPHATE
DETERMINATION OF SULPHUR IN BARIUM SULPHATE
DETERMINATION OF PHOSPHORIC ANHYDRIDE IN APATITE
ANALYSIS OF LIMESTONE
Determination of Moisture; Insoluble Matter and Silica; Ferric
Oxide and Alumina; Calcium; Magnesium; Carbon Dioxide
ANALYSIS OF BRASS
Electrolytic Separations; Determination of Lead, Copper, Iron
and Zinc.
DETERMINATION OF SILICA IN SILICATES

PART IV. STOICHIOMETRY

SOLUTIONS OF TYPICAL PROBLEMS PROBLEMS

APPENDIX

ELECTROLYTIC DISSOCIATION THEORY FOLDING OF A FILTER PAPER SAMPLE
NOTEBOOK PAGES STRENGTH OF REAGENTS DENSITIES AND VOLUMES OF
WATER CORRECTIONS FOR CHANGE OF TEMPERATURE OF STANDARD
SOLUTIONS ATOMIC WEIGHTS LOGARITHM TABLES

QUANTITATIVE CHEMICAL ANALYSIS

PART I

INTRODUCTION

SUBDIVISIONS OF ANALYTICAL CHEMISTRY

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

It is assumed that the operations of qualitative analysis are familiar to the
student, who will find that the reactions made use of in quantitative
processes are frequently the same as those employed in qualitative analyses
with respect to both precipitation and systematic separation from interfering
substances; but it should be noted that the conditions must now be regulated
with greater care, and in such a manner as to insure the most complete
separation possible. For example, in the qualitative detection of sulphates
by precipitation as barium sulphate from acid solution it is not necessary, in
most instances, to take into account the solubility of the sulphate in
hydrochloric acid, while in the quantitative determination of sulphates by

this reaction this solubility becomes an important consideration. The
operations of qualitative analysis are, therefore, the more accurate the
nearer they are made to conform to quantitative conditions.

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

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

solution involved in the reaction. This is a !volumetric! process, and is
equivalent to weighing without the use of a balance.

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

In !colorimetric! analyses the substance to be determined is converted
into some compound which imparts to its solutions a distinct color, the
intensity of which must vary in direct proportion to the amount of the
compound in the solution. Such solutions are compared with respect to
depth of color with standard solutions containing known amounts of the
colored compound, or of other similar color-producing substance which has
been found acceptable as a color standard. Colorimetric methods are, in
general, restricted to the determinations of very small quantities, since only
in dilute solutions are accurate comparisons of color possible.

GENERAL DIRECTIONS

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

ACCURACY AND ECONOMY OF TIME

The fundamental conception of quantitative analysis implies a necessity for
all possible care in guarding against loss of material or the introduction of
foreign matter. The laboratory desk, and all apparatus, should be
scrupulously neat and clean at all times. A sponge should always be ready at
hand, and desk and filter-stands should be kept dry and in good order.
Funnels should never be allowed to drip upon the base of the stand.
Glassware should always be wiped with a clean, lintless towel just before
use. All filters and solutions should be covered to protect them from dust,
just as far as is practicable, and every drop of solution or particle of
precipitate must be regarded as invaluable for the success of the analysis.

An economical use of laboratory hours is best secured by acquiring a
thorough knowledge of the character of the work to be done before
undertaking it, and then by so arranging the work that no time shall be
wasted during the evaporation of liquids and like time-consuming
operations. To this end the student should read thoughtfully not only the
!entire! procedure, but the explanatory notes as well, before any step is
taken in the analysis. The explanatory notes furnish, in general, the reasons
for particular steps or precautions, but they also occasionally contain details
of manipulation not incorporated, for various reasons, in the procedure.
These notes follow the procedures at frequent intervals, and the exact points
to which they apply are indicated by references. The student should realize
that a !failure to study the notes will inevitably lead to mistakes, loss of
time, and an inadequate understanding of the subject!.

All analyses should be made in duplicate, and in general a close
agreement of results should be expected. It should, however, be
remembered that a close concordance of results in "check analyses" is not
conclusive evidence of the accuracy of those results, although the
probability of their accuracy is, of course, considerably enhanced. The
satisfaction in obtaining "check results" in such analyses must never be
allowed to interfere with the critical examination of the procedure
employed, nor must they ever be regarded as in any measure a substitute for
absolute truth and accuracy.

In this connection it must also be emphasized that only the operator
himself can know the whole history of an analysis, and only he can know
whether his work is worthy of full confidence. No work should be
continued for a moment after such confidence is lost, but should be
resolutely discarded as soon as a cause for distrust is fully established. The
student should, however, determine to put forth his best efforts in each
analysis; it is well not to be too ready to condone failures and to "begin

again," as much time is lost in these fruitless attempts. Nothing less than
!absolute integrity! is or can be demanded of a quantitative analyst, and any
disregard of this principle, however slight, is as fatal to success as lack of
chemical knowledge or inaptitude in manipulation can possibly be.

NOTEBOOKS

Notebooks should contain, beside the record of observations, descriptive
notes. All records of weights should be placed upon the right-hand page,
while that on the left is reserved for the notes, calculations of factors, or the
amount of reagents required.

The neat and systematic arrangement of the records of analyses is of the
first importance, and is an evidence of careful work and an excellent
credential. Of two notebooks in which the results may be, in fact, of equal
value as legal evidence, that one which is neatly arranged will carry with it
greater weight.

All records should be dated, and all observations should be recorded at
once in the notebook. The making of records upon loose paper is a practice
to be deprecated, as is also that of copying original entries into a second
notebook. The student should accustom himself to orderly entries at the
time of observation. Several sample pages of systematic records are to be
found in the Appendix. These are based upon experience; but other
arrangements, if clear and orderly, may prove equally serviceable. The
student is advised to follow the sample pages until he is in a position to plan
out a system of his own.

REAGENTS

The habit of carefully testing reagents, including distilled water, cannot be
too early acquired or too constantly practiced; for, in spite of all reasonable
precautionary measures, inferior chemicals will occasionally find their way
into the stock room, or errors will be made in filling reagent bottles. The
student should remember that while there may be others who share the
responsibility for the purity of materials in the laboratory of an institution,
the responsibility will later be one which he must individually assume.

The stoppers of reagent bottles should never be laid upon the desk, unless
upon a clean watch-glass or paper. The neck and mouth of all such bottles
should be kept scrupulously clean, and care taken that no confusion of
stoppers occurs.

WASH-BOTTLES

Wash-bottles for distilled water should be made from flasks of about 750
cc. capacity and be provided with gracefully bent tubes, which should not
be too long. The jet should be connected with the tube entering the wash-
bottle by a short piece of rubber tubing in such a way as to be flexible, and
should deliver a stream about one millimeter in diameter. The neck of the
flask may be wound with cord, or covered with wash-leather, for greater
comfort when hot water is used. It is well to provide several small wash-
bottles for liquids other than distilled water, which should invariably be
clearly labeled.

TRANSFER OF LIQUIDS

Liquids should never be transferred from one vessel to another, nor to a
filter, without the aid of a stirring rod held firmly against the side or lip of
the vessel. When the vessel is provided with a lip it is not usually necessary

to use other means to prevent the loss of liquid by running down the side;
whenever loss seems imminent a !very thin! layer of vaseline, applied with
the finger to the edge of the vessel, will prevent it. The stirring rod down
which the liquid runs should never be drawn upward in such a way as to
allow the solution to collect on the under side of the rim or lip of a vessel.

The number of transfers of liquids from one vessel to another during an
analysis should be as small as possible to avoid the risk of slight losses.
Each vessel must, of course, be completely washed to insure the transfer of
all material; but it should be remembered that this can be accomplished
better by the use of successive small portions of wash-water (perhaps 5-10
cc.), if each wash-water is allowed to drain away for a few seconds, than by
the addition of large amounts which unnecessarily increase the volume of
the solutions, causing loss of time in subsequent filtrations or evaporations.

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

EVAPORATION OF LIQUIDS

The greatest care must be taken to prevent loss of solutions during
processes of evaporation, either from too violent ebullition, from
evaporation to dryness and spattering, or from the evolution of gas during
the heating. In general, evaporation upon the steam bath is to be preferred to
other methods on account of the impossibility of loss by spattering. If the
steam baths are well protected from dust, solutions should be left without
covers during evaporation; but solutions which are boiled upon the hot
plate, or from which gases are escaping, should invariably be covered. In
any case a watch-glass may be supported above the vessel by means of a

glass triangle, or other similar device, and the danger of loss of material or
contamination by dust thus be avoided. It is obvious that evaporation is
promoted by the use of vessels which admit of the exposure of a broad
surface to the air.

Liquids which contain suspended matter (precipitates) should always be
cautiously heated, since the presence of the solid matter is frequently the
occasion of violent "bumping," with consequent risk to apparatus and
analysis.

PART II

VOLUMETRIC ANALYSIS

The processes of volumetric analysis are, in general, simpler than those
of gravimetric analysis and accordingly serve best as an introduction to the
practice of quantitative analysis. For their execution there are required, first,
an accurate balance with which to weigh the material for analysis; second,
graduated instruments in which to measure the volume of the solutions
employed; third, standard solutions, that is, solutions the value of which is
accurately known; and fourth, indicators, which will furnish accurate
evidence of the point at which the desired reaction is completed. The nature
of the indicators employed will be explained in connection with the
different analyses.

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

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

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

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

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

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

MEASURING INSTRUMENTS

THE ANALYTICAL BALANCE

For a complete discussion of the physical principles underlying the
construction and use of balances, and the various methods of weighing, the
student is referred to larger manuals of Quantitative Analysis, such as those
of Fresenius, or Treadwell-Hall, and particularly to the admirable discussion
of this topic in Morse's !Exercises in Quantitative Chemistry!.

The statements and rules of procedure which follow are sufficient for the
intelligent use of an analytical balance in connection with processes
prescribed in this introductory manual. It is, however, imperative that the

student should make himself familiar with these essential features of the
balance, and its use. He should fully realize that the analytical balance is a
delicate instrument which will render excellent service under careful
treatment, but such treatment is an essential condition if its accuracy is to be
depended upon. He should also understand that no set of rules, however
complete, can do away with the necessity for a sense of personal
responsibility, since by carelessness he can render inaccurate not only his
own analyses, but those of all other students using the same balance.

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

1. The balance case is mounted on three brass legs, which should
preferably rest in glass cups, backed with rubber to prevent slipping. The
front legs are adjustable as to height and are used to level the balance case;
the rear leg is of permanent length.

2. The front of the case may be raised to give access to the balance. In
some makes doors are provided also at the ends of the balance case.

3. The balance beam is mounted upon an upright in the center of the case
on the top of which is an inlaid agate plate. To the center of the beam there
is attached a steel or agate knife-edge on which the beam oscillates when it
rests on the agate plate.

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

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

6. At the base of the post, usually in the rear, is a spirit-level.

7. Within the upright is a mechanism, controlled by a knob at the front of
the balance case, which is so arranged as to raise the entire beam slightly
above the level at which the knife-edges are in contact with the agate plates.
When the balance is not in use the beam must be supported by this device
since, otherwise, the constant jarring to which a balance is inevitably
subjected, will soon dull the knife-edges, and lessen the sensitiveness of the
balance.

8. A small weight, or bob, is attached to the pointer (or sometimes to the
beam) by which the center of gravity of the beam and its attachments may
be regulated. The center of gravity must lie very slightly below the level of
the agate plates to secure the desired sensitiveness of the balance. This is
provided for when the balance is set up and very rarely requires alteration.
The student should never attempt to change this adjustment.

9. Below the balance pans are two pan-arrests operated by a button from
the front of the case. These arrests exert a very slight upward pressure upon
the pans and minimize the displacement of the beam when objects or
weights are being placed upon the pans.

10. A movable rod, operated from one end of the balance case, extends
over the balance beam and carries a small wire weight, called a rider. By
means of this rod the rider can be placed upon any desired division of the
scale on the balance beam. Each numbered division on the beam
corresponds to one milligram, and the use of the rider obviates the placing
of very small fractional weights on the balance pan.

If a new rider is purchased, or an old one replaced, care must be taken
that its weight corresponds to the graduations on the beam of the balance on
which it is to be used. The weight of the rider in milligrams must be equal
to the number of large divisions (5, 6, 10, or 12) between the central knife-
edge and the knife-edge at the end of the beam. It should be noted that on
some balances the last division bears no number. Each new rider should be
tested against a 5 or 10-milligram weight.

In some of the most recent forms of the balance a chain device replaces
the smaller weights and the use of the rider as just described.

Before using a balance, it is always best to test its adjustment. This is
absolutely necessary if the balance is used by several workers; it is always a
wise precaution under any conditions. For this purpose, brush off the
balance pans with a soft camel's hair brush. Then note (1) whether the
balance is level; (2) that the mechanism for raising and lowering the beams
works smoothly; (3) that the pan-arrests touch the pans when the beam is
lowered; and (4) that the needle swings equal distances on either side of the
zero-point when set in motion without any load on the pans. If the latter
condition is not fulfilled, the balance should be adjusted by means of the
adjusting screw at the end of the beam unless the variation is not more than
one division on the scale; it is often better to make a proper allowance for
this small zero error than to disturb the balance by an attempt at correction.
Unless a student thoroughly understands the construction of a balance he
should never attempt to make adjustments, but should apply to the
instructor in charge.

The object to be weighed should be placed on the left-hand balance pan
and the weights upon the right-hand pan. Every substance which could
attack the metal of the balance pan should be weighed upon a watch-glass,

and all objects must be dry and cold. A warm body gives rise to air currents
which vitiate the accuracy of the weighing.

The weights should be applied in the order in which they occur in the
weight-box (not at haphazard), beginning with the largest weight which is
apparently required. After a weight has been placed upon the pan the beam
should be lowered upon its knife-edges, and, if necessary, the pan-arrests
depressed. The movement of the pointer will then indicate whether the
weight applied is too great or too small. When the weight has been
ascertained, by the successive addition of small weights, to the nearest 5 or
10 milligrams, the weighing is completed by the use of the rider. The
correct weight is that which causes the pointer to swing an equal number of
divisions to the right and left of the zero-point, when the pointer traverses
not less than five divisions on either side.

The balance case should always be closed during the final weighing,
while the rider is being used, to protect the pans from the effect of air
currents.

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

The beam should never be set in motion by lowering it forcibly upon the
knife-edges, nor by touching the pans, but rather by lifting the rider (unless
the balance be provided with some of the newer devices for the purpose),
and the swing should be arrested only when the needle approaches zero on
the scale, otherwise the knife-edges become dull. For the same reason the
beam should never be left upon its knife-edges, nor should weights be

removed from or placed on the pans without supporting the beam, except in
the case of the small fractional weights.

When the process of weighing has been completed, the weight should be
recorded in the notebook by first noting the vacant spaces in the weight-
box, and then checking the weight by again noting the weights as they are
removed from the pan. This practice will often detect and avoid errors. It is
obvious that the weights should always be returned to their proper places in
the box, and be handled only with pincers.

It should be borne in mind that if the mechanism of a balance is deranged
or if any substance is spilled upon the pans or in the balance case, the
damage should be reported at once. In many instances serious harm can be
averted by prompt action when delay might ruin the balance.

Samples for analysis are commonly weighed in small tubes with cork
stoppers. Since the stoppers are likely to change in weight from the varying
amounts of moisture absorbed from the atmosphere, it is necessary to
confirm the recorded weight of a tube which has been unused for some time
before weighing out a new portion of substance from it.

WEIGHTS

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

The sets of weights purchased from reputable dealers are usually
sufficiently accurate for analytical work. It is not necessary that such a set
should be strictly exact in comparison with the absolute standard of weight,

provided they are relatively correct among themselves, and provided the
same set of weights is used in all weighings made during a given analysis.
The analyst should assure himself that the weights in a set previously
unfamiliar to him are relatively correct by a few simple tests. For example,
he should make sure that in his set two weights of the same denomination
(i.e., two 10-gram weights, or the two 100-milligram weights) are actually
equal and interchangeable, or that the 500-milligram weight is equal to the
sum of the 200, 100, 100, 50, 20, 20 and 10-milligram weights combined,
and so on. If discrepancies of more than a few tenths of a milligram
(depending upon the total weight involved) are found, the weights should
be returned for correction. The rider should also be compared with a 5 or
10-milligram weight.

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

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

Weights are inevitably subject to corrosion if not properly protected at all
times, and are liable to damage unless handled with great care. It is obvious
that anything which alters the weight of a single piece in an analytical set
will introduce an error in every weighing made in which that piece is used.
This source of error is often extremely obscure and difficult to detect. The
only safeguard against such errors is to be found in scrupulous care in
handling and protection on the part of the analyst, and an equal insistence
that if several analysts use the same set of weights, each shall realize his
responsibility for the work of others as well as his own.

BURETTES

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

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

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

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

PREPARATION OF A BURETTE FOR USE

The inner surface of a burette must be thoroughly cleaned in order that the
liquid as drawn out may drain away completely, without leaving drops upon
the sides. This is best accomplished by treating the inside of the burette with
a warm solution of chromic acid in concentrated sulphuric acid, applied as
follows: If the burette is of the "plain" type, first remove the rubber tip and
force the lower end of the burette into a medium-sized cork stopper. Nearly
fill the burette with the chromic acid solution, close the upper end with a
cork stopper and tip the burette backward and forward in such a way as to
bring the solution into contact with the entire inner surface. Remove the
stopper and pour the solution into a stock bottle to be kept for further use,
and rinse out the burette with water several times. Unless the water then
runs freely from the burette without leaving drops adhering to the sides, the
process must be repeated (Note 1).

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

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

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

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

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

READING OF A BURETTE

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

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

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

The eye soon becomes accustomed to estimating the tenths of the
divisions. If the paper is held as directed, two divisions below the meniscus,
one whole division is visible to correct the judgment. It is not well to
attempt to bring the meniscus exactly to a division mark on the burette.
Such readings are usually less accurate than those in which the tenths of a
division are estimated.]

CALIBRATION OF GLASS MEASURING DEVICES

If accuracy of results is to be attained, the correctness of all measuring
instruments must be tested. None of the apparatus offered for sale can be
implicitly relied upon except those more expensive instruments which are
accompanied by a certificate from the !National Bureau of Standards! at
Washington, or other equally authentic source.

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

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

There is, unfortunately, no uniform standard of volume which has been
adopted for general use in all laboratories. It has been variously proposed to
consider the volume of 1000 grams of water at 4°, 15.5°, 16°, 17.5°, and
even 20°C., as a liter for practical purposes, and to consider the cubic
centimeter to be one one-thousandth of that volume. The true liter is the
volume of 1000 grams of water at 4°C.; but this is obviously a lower
temperature than that commonly found in laboratories, and involves the
constant use of corrections if taken as a laboratory standard. Many
laboratories use 15.5°C. (60° F.) as the working standard. It is plain that any
temperature which is deemed most convenient might be chosen for a
particular laboratory, but it cannot be too strongly emphasized that all
measuring instruments, including burettes, pipettes, and flasks, should be
calibrated at that temperature in order that the contents of each burette,
pipette, etc., shall be comparable with that of every other instrument, thus
permitting general interchange and substitution. For example, it is obvious

that if it is desired to remove exactly 50 cc. from a solution which has been
diluted to 500 cc. in a graduated flask, the 50 cc. flask or pipette used to
remove the fractional portion must give a correct reading at the same
temperature as the 500 cc. flask. Similarly, a burette used for the titration of
the 50 cc. of solution removed should be calibrated under the same
conditions as the measuring flasks or pipettes employed with it.

The student should also keep constantly in mind the fact that all
volumetric operations, to be exact, should be carried out as nearly at a
constant temperature as is practicable. The spot selected for such work
should therefore be subject to a minimum of temperature variations, and
should have as nearly the average temperature of the laboratory as is
possible. In all work, whether of calibration, standardization, or analysis,
the temperature of the liquids employed must be taken into account, and if
the temperature of these liquids varies more than 3° or 4° from the standard
temperature chosen for the laboratory, corrections must be applied for errors
due to expansion or contraction, since volumes of a liquid measured at
different times are comparable only under like conditions as to temperature.
Data to be used for this purpose are given in the Appendix. Neglect of this
correction is frequently an avoidable source of error and annoyance in
otherwise excellent work. The temperature of all solutions at the time of
standardization should be recorded to facilitate the application of
temperature corrections, if such are necessary at any later time.

CALIBRATION OF THE BURETTES

Two burettes, one at least of which should have a glass stopper, are required
throughout the volumetric work. Both burettes should be calibrated by the
student to whom they are assigned.

PROCEDURE.—Weigh a 50 cc., flat-bottomed flask (preferably a light-
weight flask), which must be dry on the outside, to the nearest centigram.
Record the weight in the notebook. (See Appendix for suggestions as to
records.) Place the flask under the burette and draw out into it about 10 cc.
of water, removing any drop on the tip by touching it against the inside of
the neck of the flask. Do not attempt to stop exactly at the 10 cc. mark, but
do not vary more than 0.1 cc. from it. Note the time, and at the expiration of
three minutes (or longer) read the burette accurately, and record the reading
in the notebook (Note 1). Meanwhile weigh the flask and water to
centigrams and record its weight (Note 2). Draw off the liquid from 10 cc.
to about 20 cc. into the same flask without emptying it; weigh, and at the
expiration of three minutes take the reading, and so on throughout the
length of the burette. When it is completed, refill the burette and check the
first calibration.

The differences in readings represent the apparent volumes, the
differences in weights the true volumes. For example, if an apparent volume
of 10.05 cc. is found to weigh 10.03 grams, it may be assumed with
sufficient accuracy that the error in that 10 cc. amounts to -0.02 cc., or
-0.002 for each cubic centimeter (Note 3).

In the calculation of corrections the temperature of the water must be
taken into account, if this varies more than 4°C. from the laboratory
standard temperature, consulting the table of densities of water in the
Appendix.

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

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

[Note 2: A comparatively rough balance, capable of weighing to
centigrams, is sufficiently accurate for use in calibrations, for a moment's
reflection will show that it would be useless to weigh the water with an
accuracy greater than that of the readings taken on the burette. The latter
cannot exceed 0.01 cc. in accuracy, which corresponds to 0.01 gram.

The student should clearly understand that !all other weighings!, except
those for calibration, should be made accurately to 0.0001 gram, unless
special directions are given to the contrary.

Corrections for temperature variations of less than 4°C. are negligible, as
they amount to less than 0.01 gram for each 10 grams of water withdrawn.]

[Note 3: Should the error discovered in any interval of 10 cc. on the
burette exceed 0.10 cc., it is advisable to weigh small portions (even 1 cc.)
to locate the position of the variation of bore in the tube rather than to
distribute the correction uniformly over the corresponding 10 cc. The latter
is the usual course for small corrections, and it is convenient to calculate the
correction corresponding to each cubic centimeter and to record it in the
form of a table or calibration card, or to plot a curve representing the
values.

Burettes may also be calibrated by drawing off the liquid in successive
portions through a 5 cc. pipette which has been accurately calibrated, as a
substitute for weighing. If many burettes are to be tested, this is a more
rapid method.]

PIPETTES

A !pipette! may consist of a narrow tube, in the middle of which is blown a
bulb of a capacity a little less than that which it is desired to measure by the
pipette; or it may be a miniature burette, without the stopcock or rubber tip
at the lower extremity. In either case, the flow of liquid is regulated by the
pressure of the finger on the top, which governs the admission of the air.

Pipettes are usually already graduated when purchased, but they require
calibration for accurate work.

CALIBRATION OF PIPETTES

PROCEDURE.—Clean the pipette. Draw distilled water into it by sucking
at the upper end until the water is well above the graduation mark. Quickly
place the forefinger over the top of the tube, thus preventing the entrance of
air and holding the water in the pipette. Cautiously admit a little air by
releasing the pressure of the finger, and allow the level of the water to fall
until the lowest point of the meniscus is level with the graduation. Hold the
water at that point by pressure of the finger and then allow the water to run
out from the pipette into a small tared, or weighed, beaker or flask. After a
definite time interval, usually two to three minutes, touch the end of the
pipette against the side of the beaker or flask to remove any liquid adhering
to it (Note 1). The increase in weight of the flask in grams represents the
volume of the water in cubic centimeters delivered by the pipette. Calculate
the necessary correction.

[Note 1: A definite interval must be allowed for draining, and a definite
practice adopted with respect to the removal of the liquid which collects at
the end of the tube, if the pipette is designed to deliver a specific volume
when emptied. This liquid may be removed at the end of a definite interval

either by touching the side of the vessel or by gently blowing out the last
drops. Either practice, when adopted, must be uniformly adhered to.]

FLASKS

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

GRADUATION OF FLASKS

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

PROCEDURE.—To calibrate a flask for !contents!, proceed as follows:
Clean the flask, using a chromic acid solution, and dry it carefully outside
and inside. Tare it accurately; pour water into the flask until the weight of
the latter counterbalances weights on the opposite pan which equal in grams
the number of cubic centimeters of water which the flask is to contain.
Remove any excess of water with the aid of filter paper (Note 1). Take the

flask from the balance, stopper it, place it in a bath at the desired
temperature, usually 15.5° or 17.5°C., and after an hour mark on the neck
with a diamond the location of the lowest point of the meniscus (Note 2).
The mark may be etched upon the flask by hydrofluoric acid, or by the use
of an etching ink now commonly sold on the market.

To graduate a flask which is designed to !deliver! a specified volume,
proceed as follows: Clean the flask as usual and wipe all moisture from the
outside. Fill it with distilled water. Pour out the water and allow the water to
drain from the flask for three minutes. Counterbalance the flask with
weights to the nearest centigram. Add weights corresponding in grams to
the volume desired, and add distilled water to counterbalance these weights.
An excess of water, or water adhering to the neck of the flask, may be
removed by means of a strip of clean filter paper. Stopper the flask, place it
in a bath at 15.5°C. or 17.5°C. and, after an hour, mark the location of the
lowest point of the meniscus, as described above.

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

[Note 2: Other methods are employed which involve the use of calibrated
apparatus from which the desired volume of water may be run into the dry
flask and the position of the meniscus marked directly upon it. For a
description of a procedure which is most convenient when many flasks are
to be calibrated, the student is referred to the !Am. Chem J.!, 16, 479.]

GENERAL DIRECTIONS FOR VOLUMETRIC
ANALYSES

It cannot be too strongly emphasized that for the success of analyses
uniformity of practice must prevail throughout all volumetric work with
respect to those factors which can influence the accuracy of the
measurement of liquids. For example, whatever conditions are imposed
during the calibration of a burette, pipette, or flask (notably the time
allowed for draining), must also prevail whenever the flask or burette is
used.

The student should also be constantly watchful to insure parallel
conditions during both standardization and analyst with respect to the final
volume of liquid in which a titration takes place. The value of a standard
solution is only accurate under the conditions which prevailed when it was
standardized. It is plain that the standard solutions must be scrupulously
protected from concentration or dilution, after their value has been
established. Accordingly, great care must be taken to thoroughly rinse out
all burettes, flasks, etc., with the solutions which they are to contain, in
order to remove all traces of water or other liquid which could act as a
diluent. It is best to wash out a burette at least three times with small
portions of a solution, allowing each to run out through the tip before
assuming that the burette is in a condition to be filled and used. It is, of

course, possible to dry measuring instruments in a hot closet, but this is
tedious and unnecessary.

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

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

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

STANDARD SOLUTIONS

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

NORMAL SOLUTIONS

Standard solutions may be made of a purely empirical strength dictated
solely by convenience of manipulation, or the concentration may be chosen
with reference to a system which is applicable to all solutions, and based
upon chemical equivalents. Such solutions are called !Normal Solutions!
and contain such an amount of the reacting substance per liter as is
equivalent in its chemical action to one gram of hydrogen, or eight grams of
oxygen. Solutions containing one half, one tenth, or one one-hundredth of
this quantity per liter are called, respectively, half-normal, tenth-normal, or
hundredth-normal solutions.

Since normal solutions of various reagents are all referred to a common
standard, they have an advantage not possessed by empirical solutions,
namely, that they are exactly equivalent to each other. Thus, a liter of a
normal solution of an acid will exactly neutralize a liter of a normal alkali
solution, and a liter of a normal oxidizing solution will exactly react with a
liter of a normal reducing solution, and so on.

Beside the advantage of uniformity, the use of normal solutions
simplifies the calculations of the results of analyses. This is particularly true
if, in connection with the normal solution, the weight of substance for
analysis is chosen with reference to the atomic or molecular weight of the
constituent to be determined. (See problem 26.)

The preparation of an !exactly! normal, half-normal, or tenth-normal
solution requires considerable time and care. It is usually carried out only
when a large number of analyses are to be made, or when the analyst has
some other specific purpose in view. It is, however, a comparatively easy
matter to prepare standard solutions which differ but slightly from the
normal or half-normal solution, and these have the advantage of practical

equality; that is, two approximately half-normal solutions are more
convenient to work with than two which are widely different in strength. It
is, however, true that some of the advantage which pertains to the use of
normal solutions as regards simplicity of calculations is lost when using
these approximate solutions.

The application of these general statements will be made clear in
connection with the use of normal solutions in the various types of
volumetric processes which follow.

I. NEUTRALIZATION METHODS

ALKALIMETRY AND ACIDIMETRY

GENERAL DISCUSSION

!Standard Acid Solutions! may be prepared from either hydrochloric,
sulphuric, or oxalic acid. Hydrochloric acid has the advantage of forming
soluble compounds with the alkaline earths, but its solutions cannot be
boiled without danger of loss of strength; sulphuric acid solutions may be
boiled without loss, but the acid forms insoluble sulphates with three of the
alkaline earths; oxalic acid can be accurately weighed for the preparation of
solutions, and its solutions may be boiled without loss, but it forms
insoluble oxalates with three of the alkaline earths and cannot be used with
certain of the indicators.

!Standard Alkali Solutions! may be prepared from sodium or potassium
hydroxide, sodium carbonate, barium hydroxide, or ammonia. Of sodium
and potassium hydroxide, it may be said that they can be used with all
indicators, and their solutions may be boiled, but they absorb carbon
dioxide readily and attack the glass of bottles, thereby losing strength;
sodium carbonate may be weighed directly if its purity is assured, but the
presence of carbonic acid from the carbonate is a disadvantage with many
indicators; barium hydroxide solutions may be prepared which are entirely
free from carbon dioxide, and such solutions immediately show by
precipitation any contamination from absorption, but the hydroxide is not
freely soluble in water; ammonia does not absorb carbon dioxide as readily
as the caustic alkalies, but its solutions cannot be boiled nor can they be

used with all indicators. The choice of a solution must depend upon the
nature of the work in hand.

A !normal acid solution! should contain in one liter that quantity of the
reagent which represents 1 gram of hydrogen replaceable by a base. For
example, the normal solution of hydrochloric acid (HCl) should contain
36.46 grams of gaseous hydrogen chloride, since that amount furnishes the
requisite 1 gram of replaceable hydrogen. On the other hand, the normal
solution of sulphuric acid (H_{2}SO_{4}) should contain only 49.03
grams, i.e., one half of its molecular weight in grams.

A !normal alkali solution! should contain sufficient alkali in a liter to
replace 1 gram of hydrogen in an acid. This quantity is represented by the
molecular weight in grams (40.01) of sodium hydroxide (NaOH), while a
sodium carbonate solution (Na_{2}CO_{3}) should contain but one half
the molecular weight in grams (i.e., 53.0 grams) in a liter of normal
solution.

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

INDICATORS

It has already been pointed out that the purpose of an indicator is to mark
(usually by a change of color) the point at which just enough of the titrating
solution has been added to complete the chemical change which it is
intended to bring about. In the neutralization processes which are employed
in the measurement of alkalies (!alkalimetry!) or acids (!acidimetry!) the
end-point of the reaction should, in principle, be that of complete neutrality.

Expressed in terms of ionic reactions, it should be the point at which the
H^{+} ions from an acid[Note 1] unite with a corresponding number of
OH^{-} ions from a base to form water molecules, as in the equation

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

It is not usually possible to realize this condition of exact neutrality, but it
is possible to approach it with sufficient exactness for analytical purposes,
since substances are known which, in solution, undergo a sharp change of
color as soon as even a minute excess of H^{+} or OH^{-} ions are present.
Some, as will be seen, react sharply in the presence of H^{+} ions, and
others with OH^{-} ions. These substances employed as indicators are
usually organic compounds of complex structure and are closely allied to
the dyestuffs in character.

[Note 1: A knowledge on the part of the student of the ionic theory as
applied to aqueous solutions of electrolytes is assumed. A brief outline of
the more important applications of the theory is given in the Appendix.]

BEHAVIOR OF ORGANIC INDICATORS

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

In the following discussion of the principles underlying the behavior of
the indicators as a class, methyl orange and phenolphthalein will be taken as
types. It has just been pointed out that indicators are bodies of complicated
structure. In the case of the two indicators named, the changes which they
undergo have been carefully studied by Stieglitz (!J. Am. Chem. Soc.!, 25,
1112) and others, and it appears that the changes involved are of two sorts:
First, a rearrangement of the atoms within the molecule, such as often

occurs in organic compounds; and, second, ionic changes. The
intermolecular changes cannot appropriately be discussed here, as they
involve a somewhat detailed knowledge of the classification and general
behavior of organic compounds; they will, therefore, be merely alluded to,
and only the ionic changes followed.

Methyl orange is a representative of the group of indicators which, in
aqueous solutions, behave as weak bases. The yellow color which it imparts
to solutions is ascribed to the presence of the undissociated base. If an acid,
such as HCl, is added to such a solution, the acid reacts with the indicator
(neutralizes it) and a salt is formed, as indicated by the equation:

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

This salt ionizes into (M.o.)^{+} (using this abbreviation for the positive
complex) and Cl^{-}; but simultaneously with this ionization there appears
to be an internal rearrangement of the atoms which results in the production
of a cation which may be designated as (M'.o'.)^{+}, and it is this which
imparts a characteristic red color to the solution. As these changes occur in
the presence of even a very small excess of acid (that is, of H^{+} ions), it
serves as the desired index of their presence in the solution. If, now, an
alkali, such as NaOH, is added to this reddened solution, the reverse series
of changes takes place. As soon as the free acid present is neutralized, the
slightest excess of sodium hydroxide, acting as a strong base, sets free the
weak, little-dissociated base of the indicator, and at the moment of its
formation it reverts, because of the rearrangement of the atoms, to the
yellow form:

OH^{-} + (M'.o'.)^{+} —> [M'.o'.OH] —> [M.o.OH].

Phenolphthalein, on the other hand, is a very weak, little-dissociated acid,
which is colorless in neutral aqueous solution or in the presence of free
H^{+} ions. When an alkali is added to such a solution, even in slight
excess, the anion of the salt which has formed from the acid of the indicator
undergoes a rearrangement of the atoms, and a new ion, (Ph')^{+}, is
formed, which imparts a pink color to the solution:

H^{+}, (Ph)^{-} + Na^{+}, OH^{-} —> (H_{2}O) + Na^{+}, (Ph)^{-}
—> Na^{+}, (Ph')^{-}

The addition of the slightest excess of an acid to this solution, on the
other hand, occasions first the reversion to the colorless ion and then the
setting free of the undissociated acid of the indicator:

H^{+}, (Ph')^{-} —> H^{+}, (Ph)^{-} —> (HPh).

Of the common indicators methyl orange is the most sensitive toward
alkalies and phenolphthalein toward acids; the others occupy intermediate
positions. That methyl orange should be most sensitive toward alkalies is
evident from the following considerations: Methyl orange is a weak base
and, therefore, but little dissociated. It should, then, be formed in the
undissociated condition as soon as even a slight excess of OH^{-} ions is
present in the solution, and there should be a prompt change from red to
yellow as outlined above. On the other hand, it should be an unsatisfactory
indicator for use with weak acids (acetic acid, for example) because the
salts which it forms with such acids are, like all salts of that type,
hydrolyzed to a considerable extent. This hydrolytic change is illustrated by
the equation:

(M.o.)^{+} C_{2}H_{3}O_{2}^{-} + H^{+}, OH^{-} —> [M.o.OH] +
H^{+},
C_{2}H_{3}O_{2}^{-}.

Comparison of this equation with that on page 30 will make it plain that
hydrolysis is just the reverse of neutralization and must, accordingly,
interfere with it. Salts of methyl orange with weak acids are so far
hydrolyzed that the end-point is uncertain, and methyl orange cannot be
used in the titration of such acids, while with the very weak acids, such as
carbonic acid or hydrogen sulphide (hydrosulphuric acid), the salts formed
with methyl orange are, in effect, completely hydrolyzed (i.e., no
neutralization occurs), and methyl orange is accordingly scarcely affected
by these acids. This explains its usefulness, as referred to later, for the
titration of strong acids, such as hydrochloric acid, even in the presence of
carbonates or sulphides in solution.

Phenolphthalein, on the other hand, should be, as it is, the best of the
common indicators for use with weak acids. For, since it is itself a weak
acid, it is very little dissociated, and its nearly undissociated, colorless
molecules are promptly formed as soon as there is any free acid (that is, free
H^{+} ions) in the solution. This indicator cannot, however, be successfully
used with weak bases, even ammonium hydroxide; for, since it is weak
acid, the salts which it forms with weak alkalies are easily hydrolyzed, and
as a consequence of this hydrolysis the change of color is not sharp. This
indicator can, however, be successfully used with strong bases, because the
salts which it forms with such bases are much less hydrolyzed and because
the excess of OH^{-} ions from these bases also diminishes the hydrolytic
action of water.

This indicator is affected by even so weak an acid as carbonic acid,
which must be removed by boiling the solution before titration. It is the
indicator most generally employed for the titration of organic acids.

In general, it may be stated that when a strong acid, such as hydrochloric,
sulphuric or nitric acid, is titrated against a strong base, such as sodium

hydroxide, potassium hydroxide, or barium hydroxide, any of these
indicators may be used, since very little hydrolysis ensues. It has been noted
above that the color change does not occur exactly at theoretical neutrality,
from which it follows that no two indicators will show exactly the same
end-point when acids and alkalis are brought together. It is plain, therefore,
that the same indicator must be employed for both standardization and
analysis, and that, if this is done, accurate results are obtainable.

The following table (Note 1) illustrates the variations in the volume of an
alkali solution (tenth-normal sodium hydroxide) required to produce an
alkaline end-point when run into 10 cc. of tenth-normal sulphuric acid,
diluted with 50 cc. of water, using five drops of each of the different
indicator solutions.

====================================================
================ | | | | INDICATOR | N/10 | N/10 |COLOR IN
ACID|COLOR IN ALKA- | H_{2}SO_{4}| NaOH |SOLUTION |LINE
SOLUTION
_______________|____________|__________|_____________|_________
_____ | cc. | cc. | cc. | Methyl orange | 10 | 9.90 | Red | Yellow Lacmoid | 10 |
10.00 | Red | Blue Litmus | 10 | 10.00 | Red | Blue Rosalic acid | 10 | 10.07 |
Yellow | Pink Phenolphthalein| 10 | 10.10 | Colorless | Pink
======================================================
==============

It should also be stated that there are occasionally secondary changes,
other than those outlined above, which depend upon the temperature and
concentration of the solutions in which the indicators are used. These
changes may influence the sensitiveness of an indicator. It is important,
therefore, to take pains to use approximately the same volume of solution
when standardizing that is likely to be employed in analysis; and when it is

necessary, as is often the case, to titrate the solution at boiling temperature,
the standardization should take place under the same conditions. It is also
obvious that since some acid or alkali is required to react with the indicator
itself, the amount of indicator used should be uniform and not excessive.
Usually a few drops of solution will suffice.

The foregoing statements with respect to the behavior of indicators
present the subject in its simplest terms. Many substances other than those
named may be employed, and they have been carefully studied to determine
the exact concentration of H^{+} ions at which the color change of each
occurs. It is thus possible to select an indicator for a particular purpose with
considerable accuracy. As data of this nature do not belong in an
introductory manual, reference is made to the following papers or books in
which a more extended treatment of the subject may be found:

Washburn, E.W., Principles of Physical Chemistry (McGraw-Hill Book
Co.), (Second Edition, 1921), pp. 380-387.

Prideaux, E.B.R., The Theory and Use of Indicators (Constable & Co.,
Ltd.), (1917).

Salm, E., A Study of Indicators, !Z. physik. Chem.!, 57 (1906), 471-501.

Stieglitz, J., Theories of Indicators, !J. Am. Chem. Soc.!, 25 (1903),
1112-1127.

Noyes, A.A., Quantitative Applications of the Theory of Indicators to
Volumetric Analysis, !J. Am. Chem. Soc.!, 32 (1911), 815-861.

Bjerrum, N., General Discussion, !Z. Anal. Chem.!, 66 (1917), 13-28 and
81-95.

Ostwald, W., Colloid Chemistry of Indicators, !Z. Chem. Ind.
Kolloide!, 10 (1912), 132-146.

[Note 1: Glaser, !Indikatoren der Acidimetrie und Alkalimetrie!.
Wiesbaden, 1901.]

PREPARATION OF INDICATOR SOLUTIONS

A !methyl orange solution! for use as an indicator is commonly made by
dissolving 0.05-0.1 gram of the compound (also known as Orange III) in a
few cubic centimeters of alcohol and diluting with water to 100 cc. A good
grade of material should be secured. It can be successfully used for the
titration of hydrochloric, nitric, sulphuric, phosphoric, and sulphurous acids,
and is particularly useful in the determination of bases, such as sodium,
potassium, barium, calcium, and ammonium hydroxides, and even many of
the weak organic bases. It can also be used for the determination, by
titration with a standard solution of a strong acid, of the salts of very weak
acids, such as carbonates, sulphides, arsenites, borates, and silicates,
because the weak acids which are liberated do not affect the indicator, and
the reddening of the solution does not take place until an excess of the
strong acid is added. It should be used in cold, not too dilute, solutions. Its
sensitiveness is lessened in the presence of considerable quantities of the
salts of the alkalies.

A !phenolphthalein solution! is prepared by dissolving 1 gram of the pure
compound in 100 cc. of 95 per cent alcohol. This indicator is particularly
valuable in the determination of weak acids, especially organic acids. It
cannot be used with weak bases, even ammonia. It is affected by carbonic
acid, which must, therefore, be removed by boiling when other acids are to
be measured. It can be used in hot solutions. Some care is necessary to keep
the volume of the solutions to be titrated approximately uniform in

standardization and in analysis, and this volume should not in general
exceed 125-150 cc. for the best results, since the compounds formed by the
indicator undergo changes in very dilute solution which lessen its
sensitiveness.

The preparation of a !solution of litmus! which is suitable for use as an
indicator involves the separation from the commercial litmus of
azolithmine, the true coloring principle. Soluble litmus tablets are often
obtainable, but the litmus as commonly supplied to the market is mixed
with calcium carbonate or sulphate and compressed into lumps. To prepare
a solution, these are powdered and treated two or three times with alcohol,
which dissolves out certain constituents which cause a troublesome
intermediate color if not removed. The alcohol is decanted and drained off,
after which the litmus is extracted with hot water until exhausted. The
solution is allowed to settle for some time, the clear liquid siphoned off,
concentrated to one-third its volume and acetic acid added in slight excess.
It is then concentrated to a sirup, and a large excess of 95 per cent. alcohol
added to it. This precipitates the blue coloring matter, which is filtered off,
washed with alcohol, and finally dissolved in a small volume of water and
diluted until about three drops of the solution added to 50 cc. of water just
produce a distinct color. This solution must be kept in an unstoppered
bottle. It should be protected from dust by a loose plug of absorbent cotton.
If kept in a closed bottle it soon undergoes a reduction and loses its color,
which, however, is often restored by exposure to the air.

Litmus can be employed successfully with the strong acids and bases,
and also with ammonium hydroxide, although the salts of the latter
influence the indicator unfavorably if present in considerable concentration.
It may be employed with some of the stronger organic acids, but the use of
phenolphthalein is to be preferred.

PREPARATION OF STANDARD SOLUTIONS

!Hydrochloric Acid and Sodium Hydroxide. Approximate Strength!, 0.5 N

PROCEDURE.—Measure out 40 cc. of concentrated, pure hydrochloric
acid into a clean liter bottle, and dilute with distilled water to an
approximate volume of 1000 cc. Shake the solution vigorously for a full
minute to insure uniformity. Be sure that the bottle is not too full to permit
of a thorough mixing, since lack of care at this point will be the cause of
much wasted time (Note 1).

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

[Note 1: The original solutions are prepared of a strength greater than 0.5
N, as they are more readily diluted than strengthened if later adjustment is
desired.

Too much care cannot be taken to insure perfect uniformity of solutions
before standardization, and thoroughness in this respect will, as stated, often
avoid much waste of time. A solution once thoroughly mixed remains
uniform.]

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

COMPARISON OF ACID AND ALKALI SOLUTIONS

PROCEDURE.—Rinse a previously calibrated burette three times with the
hydrochloric acid solution, using 10 cc. each time, and allowing the liquid
to run out through the tip to displace all water and air from that part of the
burette. Then fill the burette with the acid solution. Carry out the same
procedure with a second burette, using the sodium hydroxide solution.

The acid solution may be placed in a plain or in a glass-stoppered burette
as may be more convenient, but the alkaline solution should never be
allowed to remain long in a glass-stoppered burette, as it tends to cement
the stopper to the burette, rendering it useless. It is preferable to use a plain
burette for this solution.

When the burettes are ready for use and all air bubbles displaced from the
tip (see Note 2, page 17) note the exact position of the liquid in each, and
record the readings in the notebook. (Consult page 188.) Run out from the
burette into a beaker about 40 cc. of the acid and add two drops of a
solution of methyl orange; dilute the acid to about 80 cc. and run out alkali
solution from the other burette, stirring constantly, until the pink has given
place to a yellow. Wash down the sides of the beaker with a little distilled
water if the solution has spattered upon them, return the beaker to the acid
burette, and add acid to restore the pink; continue these alternations until
the point is accurately fixed at which a single drop of either solutions served
to produce a distinct change of color. Select as the final end-point the
appearance of the faintest pink tinge which can be recognized, or the
disappearance of this tinge, leaving a pure yellow; but always titrate to the
same point (Note 1). If the titration has occupied more than the three
minutes required for draining the sides of the burette, the final reading may
be taken immediately and recorded in the notebook.

Refill the burettes and repeat the titration. From the records of calibration
already obtained, correct the burette readings and make corrections for
temperature, if necessary. Obtain the ratio of the sodium hydroxide solution
to that of hydrochloric acid by dividing the number of cubic centimeters of
acid used by the number of cubic centimeters of alkali required for
neutralization. The check results of the two titrations should not vary by
more than two parts in one thousand (Note 2). If the variation in results is
greater than this, refill the burettes and repeat the titration until satisfactory
values are obtained. Use a new page in the notebook for each titration.
Inaccurate values should not be erased or discarded. They should be
retained and marked "correct" or "incorrect," as indicated by the final
outcome of the titrations. This custom should be rigidly followed in all
analytical work.

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

[Note 2: While variation of two parts in one thousand in the values
obtained by an inexperienced analyst is not excessive, the idea must be
carefully avoided that this is a standard for accurate work to be !generally
applied!. In many cases, after experience is gained, the allowable error is
less than this proportion. In a few cases a larger variation is permissible, but
these are rare and can only be recognized by an experienced analyst. It is
essential that the beginner should acquire at least the degree of accuracy
indicated if he is to become a successful analyst.]

STANDARDIZATION OF HYDROCHLORIC
ACID

SELECTION AND PREPARATION OF STANDARD

The selection of the best substance to be used as a standard for acid
solutions has been the subject of much controversy. The work of Lunge
(!Ztschr. angew. Chem.! (1904), 8, 231), Ferguson (!J. Soc. Chem. Ind.!
(1905), 24, 784), and others, seems to indicate that the best standard is
sodium carbonate prepared from sodium bicarbonate by heating the latter at
temperature between 270° and 300°C. The bicarbonate is easily prepared in
a pure state, and at the temperatures named the decomposition takes place
according to the equation

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

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

PROCEDURE.—Place in a porcelain crucible about 6 grams (roughly
weighed) of the purest sodium bicarbonate obtainable. Rest the crucible
upon a triangle of iron or copper wire so placed within a large crucible that
there is an open air space of about three eighths of an inch between them.
The larger crucible may be of iron, nickel or porcelain, as may be most

convenient. Insert the bulb of a thermometer reading to 350°C. in the
bicarbonate, supporting it with a clamp so that the bulb does not rest on the
bottom of the crucible. Heat the outside crucible, using a rather small flame,
and raise the temperature of the bicarbonate fairly rapidly to 270°C. Then
regulate the heat in such a way that the temperature rises !slowly! to 300°C.
in the course of a half-hour. The bicarbonate should be frequently stirred
with a clean, dry, glass rod, and after stirring, should be heaped up around
the bulb of the thermometer in such a way as to cover it. This will require
attention during most of the heating, as the temperature should not be
permitted to rise above 310°C. for any length of time. At the end of the half-
hour remove the thermometer and transfer the porcelain crucible, which
now contains sodium carbonate, to a desiccator. When it is cold, transfer the
carbonate to a stoppered weighing tube or weighing-bottle.

STANDARDIZATION

PROCEDURE.—Clean carefully the outside of a weighing-tube, or
weighing-bottle, containing the pure sodium carbonate, taking care to
handle it as little as possible after wiping. Weigh the tube accurately to
0.0001 gram, and record the weight in the notebook. Hold the tube over the
top of a beaker (200-300 cc.) and cautiously remove the stopper, making
sure that no particles fall from it or from the tube elsewhere than in the
beaker. Pour out from the tube a portion of the carbonate, replace the
stopper and determine approximately how much has been removed.
Continue this procedure until 1.00 to 1.10 grams has been taken from the
tube. Then weigh the tube accurately and record the weight under the first
weight in the notebook. The difference in the two weights is the weight of
the carbonate transferred to the beaker. Proceed in the same way to transfer
a second portion of the carbonate from the tube to another beaker of about

the same size as the first. The beakers should be labeled and plainly marked
to correspond with the entries in the notebook.

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

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

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

If, however, it is recalled that normal solutions are equivalent to each
other, it will be seen that the same result may be more readily reached by
dividing the weight in grams of sodium carbonate per cubic centimeter just
found by titration by the weight which would be contained in the same

volume of a normal solution of sodium carbonate. A normal solution of
sodium carbonate contains 53.0 grams per liter, or 0.0530 gram per cc. (see
page 29). The relation of the acid solution to the normal is, therefore,
calculated by dividing the weight of the carbonate to which each cubic
centimeter of the acid is equivalent by 0.0530. The standardization must be
repeated until the values obtained agree within, at most, two parts in one
thousand.

When the standard of the acid solution has been determined, calculate,
from the known ratio of the two solutions, the relation of the sodium
hydroxide solution to a normal solution (Notes 1 and 2).

[Note 1: In the foregoing procedure the acid solution is standardized and
the alkali solution referred to this standard by calculation. It is equally
possible, if preferred, to standardize the alkali solution. The standards in a
common use for this purpose are purified oxalic acid
(H_{2}C_{2}O_{4}.2H_{2}O), potassium acid oxalate
(KHC_{2}O_{4}.H_{2}O or KHC_{2}O_{4}), potassium tetroxalate
(KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O), or potassium acid tartrate
(KHC_{4}O_{6}), with the use of a suitable indicator. The oxalic acid and
the oxalates should be specially prepared to insure purity, the main
difficulty lying in the preservation of the water of crystallization.

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

[Note 2: It is also possible to standardize a hydrochloric acid solution by
precipitating the chloride ions as silver chloride and weighing the
precipitate, as prescribed under the analysis of sodium chloride to be
described later. Sulphuric acid solutions may be standardized by

precipitation of the sulphate ions as barium sulphate and weighing the
ignited precipitate, but the results are not above criticism on account of the
difficulty in obtaining large precipitates of barium sulphate which are
uncontaminated by inclosures or are not reduced on ignition.]

DETERMINATION OF THE TOTAL
ALKALINE STRENGTH OF SODA ASH

Soda ash is crude sodium carbonate. If made by the ammonia process it
may contain also sodium chloride, sulphate, and hydroxide; when made by
the Le Blanc process it may contain sodium sulphide, silicate, and
aluminate, and other impurities. Some of these, notably the hydroxide,
combine with acids and contribute to the total alkaline strength, but it is
customary to calculate this strength in terms of sodium carbonate; i.e., as
though no other alkali were present.

PROCEDURE.—In order to secure a sample which shall represent the
average value of the ash, it is well to take at least 5 grams. As this is too
large a quantity for convenient titration, an aliquot portion of the solution is
measured off, representing one fifth of the entire quantity. This is
accomplished as follows: Weigh out on an analytical balance two samples
of soda ash of about 5 grams each into beakers of about 500 cc. capacity.
(The weighings need be made to centigrams only.) Dissolve the ash in 75
cc. of water, warming gently, and filter off the insoluble residue; wash the
filter by filling it at least three times with distilled water, and allowing it to
drain, adding the washings to the main filtrate. Cool the filtrate to
approximately the standard temperature of the laboratory, and transfer it to a
250 cc. measuring flask, washing out the beaker thoroughly. Add distilled
water of laboratory temperature until the lowest point of the meniscus is

level with the graduation on the neck of the flask and remove any drops of
water that may be on the neck above the graduation by means of a strip of
filter paper; make the solution thoroughly uniform by pouring it out into a
dry beaker and back into the flask several times. Measure off 50 cc. of the
solution in a measuring flask, or pipette, either of which before use should,
unless they are dry on the inside, be rinsed out with at least two small
portions of the soda ash solution to displace any water.

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

From the volumes of acid and alkali employed, corrected for burette
errors and temperature changes, and the data derived from the
standardization, calculate the percentage of alkali present, assuming it all to
be present as sodium carbonate (Note 2).

[Note 1: The hydrochloric acid sets free carbonic acid which is unstable
and breaks down into water and carbon dioxide, most of which escapes
from the solution. Carbonic acid is a weak acid and, as such, does not yield
a sufficient concentration of H^{+} ions to cause the indicator to change to
a pink (see page 32).

The chemical changes involved may be summarized as follows:

2H^{+}, 2Cl^{-} + 2Na^{+}, CO_{3}^{—} —> 2Na^{+}, 2Cl^{-} +
[H_{2}CO_{3}] —> H_{2}O + CO_{2}]

[Note 2: A determination of the alkali present as hydroxide in soda ash
may be determined by precipitating the carbonate by the addition of barium
chloride, removing the barium carbonate by filtration, and titrating the
alkali in the filtrate.

The caustic alkali may also be determined by first using phenolphthalein
as an indicator, which will show by its change from pink to colorless the
point at which the caustic alkali has been neutralized and the carbonate has
been converted to bicarbonate, and then adding methyl orange and
completing the titration. The amount of acid necessary to change the methyl
orange to pink is a measure of one half of the carbonate present. The results
of the double titration furnish the data necessary for the determination of
the caustic alkali and of the carbonate in the sample.]

DETERMINATION OF THE ACID STRENGTH
OF OXALIC ACID

PROCEDURE.—Weigh out two portions of the acid of about 1 gram each.
Dissolve these in 50 cc. of warm water. Add two drops of phenolphthalein
solution, and run in alkali from the burette until the solution is pink; add
acid from the other burette until the pink is just destroyed, and then add 0.3
cc. (not more) in excess. Heat the solution to boiling for three minutes. If
the pink returns during the boiling, discharge it with acid and again add 0.3
cc. in excess and repeat the boiling (Note 1). If the color does not then
reappear, add alkali until it does, and a !drop or two! of acid in excess and
boil again for one minute (Note 2). If no color reappears during this time,
complete the titration in the hot solution. The end-point should be the
faintest visible shade of color (or its disappearance), as the same difficulty
would exist here as with methyl orange if an attempt were made to match
shades of pink.

From the corrected volume of alkali required to react with the oxalic
acid, calculate the percentage of the crystallized acid
(H_{2}C_{2}O_{4}.2H_{2}O) in the sample (Note 3).

[Note 1: All commercial caustic soda such as that from which the
standard solution was made contains some sodium carbonate. This reacts

with the oxalic acid, setting free carbonic acid, which, in turn, forms
sodium bicarbonate with the remaining carbonate:

H_{2}CO_{3} + Na_{2}CO_{3} —> 2HNaCO_{3}.

This compound does not hydrolyze sufficiently to furnish enough
OH^{-} ions to cause phenolphthalein to remain pink; hence, the color of
the indicator is discharged in cold solutions at the point at which
bicarbonate is formed. If, however, the solution is heated to boiling, the
bicarbonate loses carbon dioxide and water, and reverts to sodium
carbonate, which causes the indicator to become again pink:

2HNaCO_{3} —> H_{2}O + CO_{2} + Na_{2}CO_{3}.

By adding successive portions of hydrochloric acid and boiling, the
carbonate is ultimately all brought into reaction.

The student should make sure that the difference in behavior of the two
indicators, methyl orange and phenolphthalein, is understood.]

[Note 2: Hydrochloric acid is volatilized from aqueous solutions, except
such as are very dilute. If the directions in the procedure are strictly
followed, no loss of acid need be feared, but the amount added in excess
should not be greater than 0.3-0.4 cc.]

[Note 3: Attention has already been called to the fact that the color
changes in the different indicators occur at varying concentrations of H^{+}
or OH^{-} ions. They do not indicate exact theoretical neutrality, but a
particular indicator always shows its color change at a particular
concentration of H^{+} or OH^{-} ions. The results of titration with a
given indicator are, therefore, comparable. As a matter of fact, a small error
is involved in the procedure as outlined above. The comparison of the acid

and alkali solutions was made, using methyl orange as an indicator, while
the titration of the oxalic acid is made with the use of phenolphthalein. For
our present purposes the small error may be neglected but, if time permits,
the student is recommended to standardize the alkali solution against one of
the substances named in Note 1, page 41, and also to ascertain the
comparative value of the acid and alkali solutions, using phenolphthalein as
indicator throughout, and conducting the titrations as described above. This
will insure complete accuracy.]

II. OXIDATION PROCESSES

GENERAL DISCUSSION

In the oxidation processes of volumetric analysis standard solutions of
oxidizing agents and of reducing agents take the place of the acid and alkali
solutions of the neutralization processes already studied. Just as an acid
solution was the principal reagent in alkalimetry, and the alkali solution
used only to make certain of the end-point, the solution of the oxidizing
agent is the principal reagent for the titration of substances exerting a
reducing action. It is, in general, true that oxidizable substances are
determined by !direct! titration, while oxidizing substances are determined
by !indirect! titration.

The important oxidizing agents employed in volumetric solutions are
potassium bichromate, potassium permangenate, potassium ferricyanide,
iodine, ferric chloride, and sodium hypochlorite.

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

The most important combinations among the foregoing are: Potassium
bichromate and ferrous salts; potassium permanganate and ferrous salts;
potassium permanganate and oxalic acid, or its derivatives; iodine and
sodium thiosulphate; hypochlorites and arsenious acid.

BICHROMATE PROCESS FOR THE
DETERMINATION OF IRON

Ferrous salts may be promptly and completely oxidized to ferric salts, even
in cold solution, by the addition of potassium bichromate, provided
sufficient acid is present to hold in solution the ferric and chromic
compounds which are formed.

The acid may be either hydrochloric or sulphuric, but the former is
usually preferred, since it is by far the best solvent for iron and its
compounds. The reaction in the presence of hydrochloric acid is as follows:

6FeCl_{2} + K_{2}Cr_{2}O_{7} + 14HCl —> 6FeCl_{3} + 2CrCl_{3}
+ 2KCl + 7H_{2}O.

NORMAL SOLUTIONS OF OXIDIZING OR REDUCING AGENTS

It will be recalled that the system of normal solutions is based upon the
equivalence of the reagents which they contain to 8 grams of oxygen or 1
gram of hydrogen. A normal solution of an oxidizing agent should,
therefore, contain that amount per liter which is equivalent in oxidizing
power to 8 grams of oxygen; a normal reducing solution must be equivalent
in reducing power to 1 gram of hydrogen. In order to determine what the
amount per liter will be it is necessary to know how the reagents enter into

reaction. The two solutions to be employed in the process under
consideration are those of potassium bichromate and ferrous sulphate. The
reaction between them, in the presence of an excess of sulphuric acid, may
be expressed as follows:

6FeSO_{4} + K_{2}Cr_{2}O_{7} + 7H_{2}SO_{4} —> 3Fe_{2}
(SO_{4}){3} + K{2}SO_{4} + Cr_{2}(SO_{4}){3} + 7H{2}O.

If the compounds of iron and chromium, with which alone we are now
concerned, be written in such a way as to show the oxides of these elements
in each, they would appear as follows: On the left-hand side of the equation
6(FeO.SO_{3}) and K_{2}O.2CrO_{3}; on the right-hand side,
3(Fe_{2}O_{3}.3SO_{3}) and Cr_{2}O_{3}.3SO_{3}. A careful
inspection shows that there are three less oxygen atoms associated with
chromium atoms on the right-hand side of the equation than on the left-
hand, but there are three more oxygen atoms associated with iron atoms on
the right than on the left. In other words, a molecule of potassium
bichromate has given up three atoms of oxygen for oxidation purposes; i.e.,
a molecular weight in grams of the bichromate (294.2) will furnish 3 X 16
or 48 grams of oxygen for oxidation purposes. As this 48 grams is six times
8 grams, the basis of the system, the normal solution of potassium
bichromate should contain per liter one sixth of 294.2 grams or 49.03
grams.

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

3H_{2} + 3O —> 3H_{2}O 6(FeO.SO_{3}) + 3O —>
3(Fe_{2}O_{3}.3SO_{3})

it is plain that one molecule of ferrous sulphate is equivalent to one atom
of hydrogen in reducing power; therefore one molecular weight in grams of
ferrous sulphate (151.9) is equivalent to 1 gram of hydrogen. Since the
ferrous sulphate crystalline form has the formula FeSO_{4}.7H_{2}O, a
normal reducing solution of this crystalline salt should contain 277.9 grams
per liter.

PREPARATION OF SOLUTIONS

!Approximate Strength 0.1 N!

It is possible to purify commercial potassium bichromate by
recrystallization from hot water. It must then be dried and cautiously heated
to fusion to expel the last traces of moisture, but not sufficiently high to
expel any oxygen. The pure salt thus prepared, may be weighed out directly,
dissolved, and the solution diluted in a graduated flask to a definite volume.
In this case no standardization is made, as the normal value can be
calculated directly. It is, however, more generally customary to standardize
a solution of the commercial salt by comparison with some substance of
definite composition, as described below.

PROCEDURE.—Pulverize about 5 grams of potassium bichromate of
good quality. Dissolve the bichromate in distilled water, transfer the
solution to a liter bottle, and dilute to approximately 1000 cc. Shake
thoroughly until the solution is uniform.

To prepare the solution of the reducing agent, pulverize about 28 grams
of ferrous sulphate (FeSO_{4}.7H_{2}O) or about 40 grams of ferrous
ammonium sulphate (FeSO_{4}.(NH_{4}){2}SO{4}.6H_{2}O) and
dissolve in distilled water containing 5 cc. of concentrated sulphuric acid.

Transfer the solution to a liter bottle, add 5 cc. concentrated sulphuric acid,
make up to about 1000 cc. and shake vigorously to insure uniformity.

INDICATOR SOLUTION

No indicator is known which, like methyl orange, can be used within the
solution, to show when the oxidation process is complete. Instead, an
outside indicator solution is employed to which drops of the titrated
solution are transferred for testing. The reagent used is potassium
ferricyanide, which produces a blue precipitate (or color) with ferrous
compounds as long as there are unoxidized ferrous ions in the titrated
solution. Drops of the indicator solution are placed upon a glazed porcelain
tile, or upon white cardboard which has been coated with paraffin to render
it waterproof, and drops of the titrated solution are transferred to the
indicator on the end of a stirring rod. When the oxidation is nearly
completed only very small amounts of the ferrous compounds remain
unoxidized and the reaction with the indicator is no longer instantaneous. It
is necessary to allow a brief time to elapse before determining that no blue
color is formed. Thirty seconds is a sufficient interval, and should be
adopted throughout the analytical procedure. If left too long, the combined
effect of light and dust from the air will cause a reduction of the ferric
compounds already formed and a resultant blue will appear which misleads
the observer with respect to the true end-point.

The indicator solution must be highly diluted, otherwise its own color
interferes with accurate observation. Prepare a fresh solution, as needed
each day, by dissolving a crystal of potassium ferricyanide about the size of
a pin's head in 25 cc. of distilled water. The salt should be carefully tested
with ferric chloride for the presence of ferrocyanides, which give a blue
color with ferric salts.

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

COMPARISON OF OXIDIZING AND REDUCING SOLUTIONS

PROCEDURE.—Fill one burette with each of the solutions, observing the
general procedure with respect to cleaning and rinsing already prescribed.
The bichromate solution is preferably to be placed in a glass-stoppered
burette.

Run out from a burette into a beaker of about 300 cc. capacity nearly 40
cc. of the ferrous solution, add 15 cc. of dilute hydrochloric acid (sp. gr.
1.12) and 150 cc. of water and run in the bichromate solution from another
burette. Since both solutions are approximately tenth-normal, 35 cc. of the
bichromate solution may be added without testing. Test at that point by
removing a very small drop of the iron solution on the end of a stirring rod,
mixing it with a drop of indicator on the tile (Note 1). If a blue precipitate
appears at once, 0.5 cc. of the bichromate solution may be added before
testing again. The stirring rod which has touched the indicator should be
dipped in distilled water before returning it to the iron solution. As soon as
the blue appears to be less intense, add the bichromate solution in small
portions, finally a single drop at a time, until the point is reached at which
no blue color appears after the lapse of thirty seconds from the time of
mixing solution and indicator. At the close of the titration a large drop of the
iron solution should be taken for the test. To determine the end-point
beyond any question, as soon as the thirty seconds have elapsed remove
another drop of the solution of the same size as that last taken and mix it
with the indicator, placing it beside the last previous test. If this last
previous test shows a blue tint in comparison with the fresh mixture, the
end-point has not been reached; if no difference can be noted the reaction is

complete. Should the end-point be overstepped, a little more of the ferrous
solution may be added and the end-point definitely fixed.

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

[Note 1: The accuracy of the work may be much impaired by the removal
of unnecessarily large quantities of solution for the tests. At the beginning
of the titration, while much ferrous iron is still present, the end of the
stirring rod need only be moist with the solution; but at the close of the
titration drops of considerable size may properly be taken for the final tests.
The stirring rod should be washed to prevent transfer of indicator to the
main solution. This cautious removal of solution does not seriously affect
the accuracy of the determination, as it will be noted that the volume of the
titrated solution is about 200 cc. and the portions removed are very small.
Moreover, if the procedure is followed as prescribed, the concentration of
unoxidized iron decreases very rapidly as the titration is carried out so that
when the final tests are made, though large drops may be taken, the amount
of ferrous iron is not sufficient to produce any appreciable error in results.

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

STANDARDIZATION OF POTASSIUM BICHROMATE
SOLUTIONS

!Selection of a Standard!

A substance which will serve satisfactorily as a standard for oxidizing
solutions must possess certain specific properties: It must be of accurately
known composition and definite in its behavior as a reducing agent, and it
must be permanent against oxidation in the air, at least for considerable
periods. Such standards may take the form of pure crystalline salts, such as
ferrous ammonium sulphate, or may be in the form of iron wire or an iron
ore of known iron content. It is not necessary that the standard should be of
100 per cent purity, provided the content of the active reducing agent is
known and no interfering substances are present.

The two substances most commonly used as standards for a bichromate
solution are ferrous ammonium sulphate and iron wire. A standard wire is to
be purchased in the market which answers the purpose well, and its iron
content may be determined for each lot purchased by a number of
gravimetric determinations. It may best be preserved in jars containing
calcium chloride, but this must not be allowed to come into contact with the
wire. It should, however, even then be examined carefully for rust before
use.

If pure ferrous ammonium sulphate is used as the standard, clear crystals
only should be selected. It is perhaps even better to determine by
gravimetric methods once for all the iron content of a large commercial
sample which has been ground and well mixed. This salt is permanent over
long periods if kept in stoppered containers.

STANDARDIZATION

PROCEDURE.—Weigh out two portions of iron wire of about 0.24-0.26
gram each, examining the wire carefully for rust. It should be handled and
wiped with filter paper (not touched by the fingers), should be weighed on a

watch-glass, and be bent in such a way as not to interfere with the
movement of the balance.

Place 30 cc. of hydrochloric acid (sp. gr. 1.12) in each of two 300 cc.
Erlenmeyer flasks, cover them with watch-glasses, and bring the acid just to
boiling. Remove them from the flame and drop in the portions of wire,
taking great care to avoid loss of liquid during solution. Boil for two or
three minutes, keeping the flasks covered (Note 1), then wash the sides of
the flasks and the watch-glass with a little water and add stannous chloride
solution to the hot liquid !from a dropper! until the solution is colorless, but
avoid more than a drop or two in excess (Note 2). Dilute with 150 cc. of
water and cool !completely!. When cold, add rapidly about 30 cc. of
mercuric chloride solution. Allow the solutions to stand about three minutes
and then titrate without further delay (Note 3), add about 35 cc. of the
standard solution at once and finish the titration as prescribed above,
making use of the ferrous solution if the end-point should be passed.

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

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

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

The solution of the wire in hot acid and the short boiling insure the
removal of compounds of hydrogen and carbon which are formed from the

small amount of carbon in the iron. These might be acted upon by the
bichromate if not expelled.]

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

The reagents available for the reduction of iron are stannous chloride,
sulphurous acid, sulphureted hydrogen, and zinc; of these stannous chloride
acts most readily, the completion of the reaction is most easily noted, and
the excess of the reagent is most readily removed. The latter object is
accomplished by oxidation to stannic chloride by means of mercuric
chloride added in excess, as the mercuric salts have no effect upon ferrous
iron or the bichromate. The reactions involved are:

2FeCl_{3} + SnCl_{2} —> 2FeCl_{2} + SnCl_{4} SnCl_{2} +
2HgCl_{2} —> SnCl_{4} + 2HgCl

The mercurous chloride is precipitated.

It is essential that the solution should be cold and that the stannous
chloride should not be present in great excess, otherwise a secondary
reaction takes place, resulting in the reduction of the mercurous chloride to
metallic mercury:

SnCl_{2} + 2HgCl —> SnCl_{4} + 2Hg.

The occurrence of this secondary reaction is indicated by the darkening
of the precipitate; and, since potassium bichromate oxidizes this mercury

slowly, solutions in which it has been precipitated are worthless as iron
determinations.]

[Note 3: The solution should be allowed to stand about three minutes
after the addition of mercuric chloride to permit the complete deposition of
mercurous chloride. It should then be titrated without delay to avoid
possible reoxidation of the iron by the oxygen of the air.]

DETERMINATION OF IRON IN LIMONITE

PROCEDURE.—Grind the mineral (Note 1) to a fine powder. Weigh out
accurately two portions of about 0.5 gram (Note 2) into porcelain crucibles;
heat these crucibles to dull redness for ten minutes, allow them to cool, and
place them, with their contents, in beakers containing 30 cc. of dilute
hydrochloric acid (sp. gr. 1.12). Heat at a temperature just below boiling
until the undissolved residue is white or until solvent action has ceased. If
the residue is white, or known to be free from iron, it may be neglected and
need not be removed by filtration. If a dark residue remains, collect it on a
filter, wash free from hydrochloric acid, and ignite the filter in a platinum
crucible (Note 3). Mix the ash with five times its weight of sodium
carbonate and heat to fusion; cool, and disintegrate the fused mass with
boiling water in the crucible. Unite this solution and precipitate (if any)
with the acid solution, taking care to avoid loss by effervescence. Wash out
the crucible, heat the acid solution to boiling, add stannous chloride solution
until it is colorless, avoiding a large excess (Note 4); cool, and when !cold!,
add 40 cc. of mercuric chloride solution, dilute to 200 cc., and proceed with
the titration as already described.

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

[Note 1: Limonite is selected as a representative of iron ores in general. It
is a native, hydrated oxide of iron. It frequently occurs in or near peat beds

and contains more or less organic matter which, if brought into solution,
would be acted upon by the potassium bichromate. This organic matter is
destroyed by roasting. Since a high temperature tends to lessen the
solubility of ferric oxide, the heat should not be raised above low redness.]

[Note 2: It is sometimes advantageous to dissolve a large portion—say 5
grams—and to take one tenth of it for titration. The sample will then
represent more closely the average value of the ore.]

[Note 3: A platinum crucible may be used for the roasting of the limonite
and must be used for the fusion of the residue. When used, it must not be
allowed to remain in the acid solution of ferric chloride for any length of
time, since the platinum is attacked and dissolved, and the platinic chloride
is later reduced by the stannous chloride, and in the reduced condition
reacts with the bichromate, thus introducing an error. It should also be noted
that copper and antimony interfere with the determination of iron by the
bichromate process.]

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

DETERMINATION OF CHROMIUM IN
CHROME IRON ORE

PROCEDURE.—Grind the chrome iron ore (Note 1) in an agate mortar
until no grit is perceptible under the pestle. Weigh out two portions of 0.5
gram each into iron crucibles which have been scoured inside until bright
(Note 2). Weigh out on a watch-glass (Note 3), using the rough balances, 5
grams of dry sodium peroxide for each portion, and pour about three
quarters of the peroxide upon the ore. Mix ore and flux by thorough stirring
with a dry glass rod. Then cover the mixture with the remainder of the
peroxide. Place the crucible on a triangle and raise the temperature !slowly!
to the melting point of the flux, using a low flame, and holding the lamp in
the hand (Note 4). Maintain the fusion for five minutes, and stir constantly
with a stout iron wire, but do not raise the temperature above moderate
redness (Notes 5 and 6).

Allow the crucible to cool until it can be comfortably handled (Note 7)
and then place it in a 300 cc. beaker, and cover it with distilled water (Note
8). The beaker must be carefully covered to avoid loss during the
disintegration of the fused mass. When the evolution of gas ceases, rinse off
and remove the crucible; then heat the solution !while still alkaline! to
boiling for fifteen minutes. Allow the liquid to cool for a few minutes; then
acidify with dilute sulphuric acid (1:5), adding 10 cc. in excess of the
amount necessary to dissolve the ferric hydroxide (Note 9). Dilute to 200

cc., cool, add from a burette an excess of a standard ferrous solution, and
titrate for the excess with a standard solution of potassium bichromate,
using the outside indicator (Note 10).

From the corrected volumes of the two standard solutions, and their
relations to normal solutions, calculate the percentage of chromium in the
ore.

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

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

[Note 3: Sodium peroxide must be kept off of balance pans and should
not be weighed out on paper, as is the usual practice in the rough weighing
of chemicals. If paper to which the peroxide is adhering is exposed to moist
air it is likely to take fire as a result of the absorption of moisture, and
consequent evolution of heat and liberation of oxygen.]

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

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

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

[Note 7: When an iron crucible is employed it is desirable to allow the
fusion to become nearly cold before it is placed in water, otherwise scales
of magnetic iron oxide may separate from the crucible, which by slowly
dissolving in acid form ferrous sulphate, which reduces the chromate.]

[Note 8: Upon treatment with water the chromate passes into solution,
the ferric hydroxide remains undissolved, and the excess of peroxide is
decomposed with the evolution of oxygen. The subsequent boiling insures
the complete decomposition of the peroxide. Unless this is complete,
hydrogen peroxide is formed when the solution is acidified, and this reacts
with the bichromate, reducing it and introducing a serious error.]

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

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

PERMANGANATE PROCESS FOR THE
DETERMINATION OF IRON

Potassium permanganate oxidizes ferrous salts in cold, acid solution
promptly and completely to the ferric condition, while in hot acid solution it
also enters into a definite reaction with oxalic acid, by which the latter is
oxidized to carbon dioxide and water.

The reactions involved are these:

10FeSO_{4} + 2KMnO_{4} + 8H_{2}S_{4} —> 5Fe_{2}(SO_{4}){3}
+ K{2}SO_{4} + 2MnSO_{4} + 8H_{2}O

5C_{2}H_{2}O_{4}(2H_{2}O) + 2KMnO_{4} +3H_{2}SO_{4} —>
K_{2}SO_{4} + 2MnSO_{4} + 10CO_{2} + 1 H_{2}O.

These are the fundamental reactions upon which the extensive use of
potassium permanganate depends; but besides iron and oxalic acid the
permanganate enters into reaction with antimony, tin, copper, mercury, and
manganese (the latter only in neutral solution), by which these metals are
changed from a lower to a higher state of oxidation; and it also reacts with
sulphurous acid, sulphureted hydrogen, nitrous acid, ferrocyanides, and
most soluble organic bodies. It should be noted, however, that very few of
these organic compounds react quantitatively with the permanganate, as is
the case with oxalic acid and the oxalates.

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

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

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

PREPARATION OF A STANDARD SOLUTION

!Approximate Strength 0.1 N!

A study of the reactions given above which represent the oxidation of
ferrous compounds by potassium permanganate, shows that there are 2
molecules of KMnO_{4} and 10 molecules of FeSO_{4} on the left-hand
side, and 2 molecules of MnSO_{4} and 5 molecules of Fe_{2}
(SO_{4})_{5} on the right-hand side. Considering only these compounds,

and writing the formulas in such a way as to show the oxides of the
elements in each, the equation becomes:

K_{2}O.Mn_{2}O_{7} + 10(FeO.SO_{3}) —> K_{2}O.SO_{3} +
2(MnO.SO_{3}) + 5(Fe_{2}O_{3}.3SO_{3}).

From this it appears that two molecules of KMnO_{4} (or 316.0 grams)
have given up five atoms (or 80 grams) of oxygen to oxidize the ferrous
compound. Since 8 grams of oxygen is the basis of normal oxidizing
solutions and 80 grams of oxygen are supplied by 316.0 grams of
KMnO_{4}, the normal solution of the permanganate should contain, per
liter, 316.0/10 grams, or 31.60 grams (Note 1).

The preparation of an approximately tenth-normal solution of the reagent
may be carried out as follows:

PROCEDURE.—Dissolve about 3.25 grams of potassium permanganate
crystals in approximately 1000 cc. of distilled water in a large beaker, or
casserole. Heat slowly and when the crystals have dissolved, boil the
solution for 10-15 minutes. Cover the solution with a watch-glass; allow it
to stand until cool, or preferably over night. Filter the solution through a
layer of asbestos. Transfer the filtrate to a liter bottle and mix thoroughly
(Note 2).

[Note 1: The reactions given on page 61 are those which take place in the
presence of an excess of acid. In neutral solutions the reduction of the
permanganate is less complete, and, under these conditions, two gram-
molecular weights of KMnO_{4} will furnish only 48 grams of oxygen. A
normal solution for use under these conditions should, therefore, contain
316.0/6 grams, or 52.66 grams.]

[Note 2: Potassium permanganate solutions are not usually stable for
long periods, and change more rapidly when first prepared than after
standing some days. This change is probably caused by interaction with the
organic matter contained in all distilled water, except that redistilled from
an alkaline permanganate solution. The solutions should be protected from
light and heat as far as possible, since both induce decomposition with a
deposition of manganese dioxide, and it has been shown that decomposition
proceeds with considerable rapidity, with the evolution of oxygen, after the
dioxide has begun to form. As commercial samples of the permanganate are
likely to be contaminated by the dioxide, it is advisable to boil and filter
solutions through asbestos before standardization, as prescribed above.
Such solutions are relatively stable.]

COMPARISON OF PERMANGANATE AND FERROUS
SOLUTIONS

PROCEDURE.—Fill a glass-stoppered burette with the permanganate
solution, observing the usual precautions, and fill a second burette with the
ferrous sulphate solution prepared for use with the potassium bichromate.
The permanganate solution cannot be used in burettes with rubber tips, as a
reduction takes place upon contact with the rubber. The solution has so deep
a color that the lower line of the meniscus cannot be detected; readings
must therefore be made from the upper edge. Run out into a beaker about
40 cc. of the ferrous solution, dilute to about 100 cc., add 10 cc. of dilute
sulphuric acid, and run in the permanganate solution to a slight permanent
pink. Repeat, until the ratio of the two solutions is satisfactorily established.

STANDARDIZATION OF A POTASSIUM PERMANGANATE
SOLUTION

!Selection of a Standard!

Commercial potassium permanganate is rarely sufficiently pure to admit
of its direct weighing as a standard. On this account, and because of the
uncertainties as to the permanence of its solutions, it is advisable to
standardize them against substances of known value. Those in most
common use are iron wire, ferrous ammonium sulphate, sodium oxalate,
oxalic acid, and some other derivatives of oxalic acid. With the exception of
sodium oxalate, these all contain water of crystallization which may be lost
on standing. They should, therefore, be freshly prepared, and with great
care. At present, sodium oxalate is considered to be one of the most
satisfactory standards.

!Method A!

!Iron Standards!

The standardization processes employed when iron or its compounds are
selected as standards differ from those applicable in connection with oxalate
standards. The procedure which immediately follows is that in use with iron
standards.

As in the case of the bichromate process, it is necessary to reduce the iron
completely to the ferrous condition before titration. The reducing agents
available are zinc, sulphurous acid, or sulphureted hydrogen. Stannous
chloride may also be used when the titration is made in the presence of
hydrochloric acid. Since the excess of both the gaseous reducing agents can
only be expelled by boiling, with consequent uncertainty regarding both the
removal of the excess and the reoxidation of the iron, zinc or stannous
chlorides are the most satisfactory agents. For prompt and complete

reduction it is essential that the iron solution should be brought into
ultimate contact with the zinc. This is brought about by the use of a
modified Jones reductor, as shown in Figure 1. This reductor is a standard
apparatus and is used in other quantitative processes.

[Illustration: Fig. 1]

The tube A has an inside diameter of 18 mm. and is 300 mm. long; the
small tube has an inside diameter of 6 mm. and extends 100 mm. below the
stopcock. At the base of the tube A are placed some pieces of broken glass
or porcelain, covered by a plug of glass wool about 8 mm. thick, and upon
this is placed a thin layer of asbestos, such as is used for Gooch filters, 1
mm. thick. The tube is then filled with the amalgamated zinc (Note 1) to
within 50 mm. of the top, and on the zinc is placed a plug of glass wool. If
the top of the tube is not already shaped like the mouth of a thistle-tube (B),
a 60 mm. funnel is fitted into the tube with a rubber stopper and the
reductor is connected with a suction bottle, F. The bottle D is a safety bottle
to prevent contamination of the solution by water from the pump. After
preparation for use, or when left standing, the tube A should be filled with
water, to prevent clogging of the zinc.

[Note 1: The use of fine zinc in the reductor is not necessary and tends to
clog the tube. Particles which will pass a 10-mesh sieve, but are retained by
one of 20 meshes to the inch, are most satisfactory. The zinc can be
amalgamated by stirring or shaking it in a mixture of 25 cc. of normal
mercuric chloride solution, 25 cc. of hydrochloric acid (sp. gr. 1.12) and
250 cc. of water for two minutes. The solution should then be poured off
and the zinc thoroughly washed. It is then ready for bottling and
preservation under water. A small quantity of glass wool is placed in the
neck of the funnel to hold back foreign material when the reductor is in
use.]

STANDARDIZATION

PROCEDURE.—Weigh out into Erlenmeyer flasks two portions of iron
wire of about 0.25 gram each. Dissolve these in hot dilute sulphuric acid (5
cc. of concentrated acid and 100 cc. of water), using a covered flask to
avoid loss by spattering. Boil the solution for two or three minutes after the
iron has dissolved to remove any volatile hydrocarbons. Meanwhile prepare
the reductor for use as follows: Connect the vacuum bottle with the suction
pump and pour into the funnel at the top warm, dilute sulphuric acid,
prepared by adding 5 cc. of concentrated sulphuric acid to 100 cc. of
distilled water. See that the stopcock (C) is open far enough to allow the
acid to run through slowly. Continue to pour in acid until 200 cc. have
passed through, then close the stopcock !while a small quantity of liquid is
still left in the funnel!. Discard the filtrate, and again pass through 100 cc.
of the warm, dilute acid. Test this with the permanganate solution. A single
drop should color it permanently; if it does not, repeat the washing, until
assured that the zinc is not contaminated with appreciable quantities of
reducing substances. Be sure that no air enters the reductor (Note 1).

Pour the iron solution while hot (but not boiling) through the reductor at
a rate not exceeding 50 cc. per minute (Notes 2 and 3). Wash out the beaker
with dilute sulphuric acid, and follow the iron solution without interruption
with 175 cc. of the warm acid and finally with 75 cc. of distilled water,
leaving the funnel partially filled. Remove the filter bottle and cool the
solution quickly under the water tap (Note 4), avoiding unnecessary
exposure to the oxygen of the air. Add 10 cc. of dilute sulphuric acid and
titrate to a faint pink with the permanganate solution, adding it directly to
the contents of the vacuum flask. Should the end-point be overstepped, the
ferrous sulphate solution may be added.

From the volume of the solution required to oxidize the iron in the wire,
calculate the relation to the normal of the permanganate solution. The
duplicate results should be concordant within two parts in one thousand.

[Note 1: The funnel of the reductor must never be allowed to empty. If it
is left partially filled with water the reductor is ready for subsequent use
after a very little washing; but a preliminary test is always necessary to
safeguard against error.

If more than a small drop of permanganate solution is required to color
100 cc. of the dilute acid after the reductor is well washed, an allowance
must be made for the iron in the zinc. !Great care! must be used to prevent
the access of air to the reductor after it has been washed out ready for use. If
air enters, hydrogen peroxide forms, which reacts with the permanganate,
and the results are worthless.]

[Note 2: The iron is reduced to the ferrous condition by contact with the
zinc. The active agent may be considered to be !nascent! hydrogen, and it
must be borne in mind that the visible bubbles are produced by molecular
hydrogen, which is without appreciable effect upon ferric iron.

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

It is also possible to reduce the iron by treatment with zinc in a flask from
which air is excluded. The zinc must be present in excess of the quantity
necessary to reduce the iron and is finally completely dissolved. This

method is, however, less convenient and more tedious than the use of the
reductor.]

[Note 3: The dilute sulphuric acid for washing must be warmed ready for
use before the reduction of the iron begins, and it is of the first importance
that the volume of acid and of wash-water should be measured, and the
volume used should always be the same in the standardizations and all
subsequent analyses.]

[Note 4: The end-point is more permanent in cold than hot solutions,
possibly because of a slight action of the permanganate upon the
manganous sulphate formed during titration. If the solution turns brown, it
is an evidence of insufficient acid, and more should be immediately added.
The results are likely to be less accurate in this case, however, as a
consequence of secondary reactions between the ferrous iron and the
manganese dioxide thrown down. It is wiser to discard such results and
repeat the process.]

[Note 5: The potassium permanganate may, of course, be diluted and
brought to an exactly 0.1 N solution from the data here obtained. The
percentage of iron in the iron wire must be taken into account in all
calculations.]

!Method B!

!Oxalate Standards!

PROCEDURE.—Weigh out two portions of pure sodium oxalate of 0.25-
0.3 gram each into beakers of about 600 cc. capacity. Add about 400 cc. of
boiling water and 20 cc. of manganous sulphate solution (Note 1). When the
solution of the oxalate is complete, heat the liquid, if necessary, until near

its boiling point (70-90°C.) and run in the standard permanganate solution
drop by drop from a burette, stirring constantly until an end-point is reached
(Note 2). Make a blank test with 20 cc. of manganous sulphate solution and
a volume of distilled water equal to that of the titrated solution to determine
the volume of the permanganate solution required to produce a very slight
pink. Deduct this volume from the amount of permanganate solution used
in the titration.

From the data obtained, calculate the relation of the permanganate
solution to the normal. The reaction involved is:

5Na_{2}C_{2}O_{4} + 2KMnO_{4} + 8H_{2}SO_{4} —>
5Na_{2}SO_{4} + K_{2}SO_{4} + 2MnSO_{4} + 10CO_{2} + 8H_{2}O

[Note 1: The manganous sulphate titrating solution is made by dissolving
20 grams of MnSO_{4} in 200 cubic centimeters of water and adding 40
cc. of concentrated sulphuric acid (sp. gr. 1.84) and 40 cc. or phosphoric
acid (85%).]

[Note 2: The reaction between oxalates and permanganates takes place
quantitatively only in hot acid solutions. The temperatures must not fall
below 70°C.]

DETERMINATION OF IRON IN LIMONITE

!Method A!

The procedures, as here prescribed, are applicable to iron ores in general,
provided these ores contain no constituents which are reduced by zinc or
stannous chloride and reoxidized by permanganates. Many iron ores contain
titanium, and this element among others does interfere with the
determination of iron by the process described. If, however, the solutions of
such ores are treated with sulphureted hydrogen or sulphurous acid, instead
of zinc or stannous chloride to reduce the iron, and the excess reducing
agent removed by boiling, an accurate determination of the iron can be
made.

PROCEDURE.—Grind the mineral to a fine powder. Weigh out two
portions of about 0.5 gram each into small porcelain crucibles. Roast the
ore at dull redness for ten minutes (Note 1), allow the crucibles to cool, and
place them and their contents in casseroles containing 30 cc. of dilute
hydrochloric acid (sp. gr. 1.12).

Proceed with the solution of the ore, and the treatment of the residue, if
necessary, exactly as described for the bichromate process on page 56.
When solution is complete, add 6 cc. of concentrated sulphuric acid to each
casserole, and evaporate on the steam bath until the solution is nearly
colorless (Note 2). Cover the casseroles and heat over the flame of the

burner, holding the casserole in the hand and rotating it slowly to hasten
evaporation and prevent spattering, until the heavy white fumes of sulphuric
anhydride are freely evolved (Note 3). Cool the casseroles, add 100 cc. of
water (measured), and boil gently until the ferric sulphate is dissolved; pour
the warm solution through the reductor which has been previously washed;
proceed as described under standardization, taking pains to use the same
volume and strength of acid and the same volume of wash-water as there
prescribed, and titrate with the permanganate solution in the reductor flask,
using the ferrous sulphate solution if the end-point should be overstepped.

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

[Note 1: The preliminary roasting is usually necessary because, even
though the sulphuric acid would subsequently char the carbonaceous matter,
certain nitrogenous bodies are not thereby rendered insoluble in the acid,
and would be oxidized by the permanganate.]

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

[Note 3: The hydrochloric acid, both free and combined, is displaced by
the less volatile sulphuric acid at its boiling point. Ferric sulphate separates
at this point, since there is no water to hold it in solution and care is
required to prevent bumping. The ferric sulphate usually has a silky
appearance and is easily distinguished from the flocculent silica which
often remains undissolved.]

!Zimmermann-Reinhardt Procedure!

!Method (B)!

PROCEDURE.—Grind the mineral to a fine powder. Weigh out two
portions of about 0.5 gram each into small porcelain crucibles. Proceed with
the solution of the ore, treat the residue, if necessary, and reduce the iron by
the addition of stannous chloride, followed by mercuric chloride, as
described for the bichromate process on page 56. Dilute the solution to
about 400 cc. with cold water, add 10 cc. of the manganous sulphate
titrating solution (Note 1, page 68) and titrate with the standard potassium
permanganate solution to a faint pink (Note 1).

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

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

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

DETERMINATION OF THE OXIDIZING
POWER OF PYROLUSITE

INDIRECT OXIDATION

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

PROCEDURE.—Grind the mineral in an agate mortar until no grit
whatever can be detected under the pestle (Note 1). Transfer it to a
stoppered weighing-tube, and weigh out two portions of about 0.5 gram into
beakers (400-500 cc.) Read Note 2, and then calculate in each case the
weight of oxalic acid (H_{2}C_{2}O_{4}.2H_{2}O) required to react with
the weights of pyrolusite taken. The reaction involved is

MnO_{2} + H_{2}C_{2}O_{4}(2H_{2}O) + H_{2}SO_{4} —>
MnSO_{4} + 2CO_{2} + 4H_{2}O.

Weigh out about 0.2 gram in excess of this quantity of !pure! oxalic acid
into the corresponding beakers, weighing the acid accurately and recording

the weight in the notebook. Pour into each beaker 25 cc. of water and 50 cc.
of dilute sulphuric acid (1:5), cover and warm the beaker and its contents
gently until the evolution of carbon dioxide ceases (Note 3). If a residue
remains which is sufficiently colored to obscure the end-reaction of the
permanganate, it must be removed by filtration.

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

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

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

[Note 2: A moderate excess of oxalic acid above that required to react
with the pyrolusite is necessary to promote solution; otherwise the residual
quantity of oxalic acid would be so small that the last particles of the
mineral would scarcely dissolve. It is also desirable that a sufficient excess
of the acid should be present to react with a considerable volume of the
permanganate solution during the titration, thus increasing the accuracy of
the process. On the other hand, the excess of oxalic acid should not be so
large as to react with more of the permanganate solution than is contained
in a 50 cc. burette. If the pyrolusite under examination is known to be of
high grade, say 80 per cent pure, or above the calculation of the oxalic acid
needed may be based upon an assumption that the mineral is all MnO_{2}.

If the quality of the mineral is unknown, it is better to weigh out three
portions instead of two and to add to one of these the amount of oxalic
prescribed, assuming complete purity of the mineral. Then run in the
permanganate solution from a pipette or burette to determine roughly the
amount required. If the volume exceeds the contents of a burette, the
amount of oxalic acid added to the other two portions is reduced
accordingly.]

[Note 3: Care should be taken that the sides of the beaker are not
overheated, as oxalic acid would be decomposed by heat alone if
crystallization should occur on the sides of the vessel. Strong sulphuric acid
also decomposes the oxalic acid. The dilute acid should, therefore, be
prepared before it is poured into the beaker.]

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

2 FeSO_{4} + MnO_{2} + 2H_{2}SO_{4} —> Fe_{2}(SO_{4}){3} +
2H{2}O

The excess of ferrous iron may also be determined by titration with
potassium bichromate, if desired. Care is required to prevent the oxidation
of the iron by the air, if ferrous salts are employed.]

[Note 5: The oxidizing power of pyrolusite may be determined by other
volumetric processes, one of which is outlined in the following reactions:

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

The chlorine generated by the pyrolusite is passed into a solution of
potassium iodide. The liberated iodine is then determined by titration with
sodium thiosulphate, as described on page 78. This is a direct process,
although it involves three steps.]

IODIMETRY

The titration of iodine against sodium thiosulphate, with starch as an
indicator, may perhaps be regarded as the most accurate of volumetric
processes. The thiosulphate solution may be used in both acid and neutral
solutions to measure free iodine and the latter may, in turn, serve as a
measure of any substance capable of liberating iodine from potassium
iodide under suitable conditions for titration, as, for example, in the process
outlined in Note 5 on page 74.

The fundamental reaction upon which iodometric processes are based is
the following:

I_{2} + 2 Na_{2}S_{2}O_{3} —> 2 NaI + Na_{2}S_{4}O_{6}.

This reaction between iodine and sodium thiosulphate, resulting in the
formation of the compound Na_{2}S_{4}O_{6}, called sodium
tetrathionate, is quantitatively exact, and differs in that respect from the
action of chlorine or bromine, which oxidize the thiosulphate, but not
quantitatively.

NORMAL SOLUTIONS OF IODINE AND SODIUM THIOSULPHATE

If the formulas of sodium thiosulphate and sodium tetrathionate are
written in a manner to show the atoms of oxygen associated with sulphur
atoms in each, thus, 2(Na_{2}).S_{2}O_{2} and Na_{2}O.S_{4}O_{5}, it

is plain that in the tetrathionate there are five atoms of oxygen associated
with sulphur, instead of the four in the two molecules of the thiosulphate
taken together. Although, therefore, the iodine contains no oxygen, the two
atoms of iodine have, in effect, brought about the addition of one oxygen
atoms to the sulphur atoms. That is the same thing as saying that 253.84
grams of iodine (I_{2}) are equivalent to 16 grams of oxygen; hence, since
8 grams of oxygen is the basis of normal solutions, 253.84/2 or 126.97
grams of iodine should be contained in one liter of normal iodine solution.
By a similar course of reasoning the conclusion is reached that the normal
solution of sodium thiosulphate should contain, per liter, its molecular
weight in grams. As the thiosulphate in crystalline form has the formula
Na_{2}S_{2}O_{3}.5H_{2}O, this weight is 248.12 grams. Tenth-normal
or hundredth-normal solutions are generally used.

PREPARATION OF STANDARD SOLUTIONS

!Approximate Strength, 0.1 N!

PROCEDURE.—Weigh out on the rough balances 13 grams of
commercial iodine. Place it in a mortar with 18 grams of potassium iodide
and triturate with small portions of water until all is dissolved. Dilute the
solution to 1000 cc. and transfer to a liter bottle and mix thoroughly (Note
1).[1]

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

Weigh out 25 grams of sodium thiosulphate, dissolve it in water which
has been previously boiled and cooled, and dilute to 1000 cc., also with

boiled water. Transfer the solution to a liter bottle and mix thoroughly (Note
2).

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

[Note 2: Sodium thiosulphate (Na_{2}S_{2}O_{3}.5H_{2}O) is rarely
wholly pure as sold commercially, but may be purified by recrystallization.
The carbon dioxide absorbed from the air by distilled water decomposes the
salt, with the separation of sulphur. Boiled water which has been cooled out
of contact with the air should be used in preparing solutions.]

INDICATOR SOLUTION

The starch solution for use as an indicator must be freshly prepared. A
soluble starch is obtainable which serves well, and a solution of 0.5 gram of
this starch in 25 cc. of boiling water is sufficient. The solution should be
filtered while hot and is ready for use when cold.

If soluble starch is not at hand, potato starch may be used. Mix about 1
gram with 5 cc. of cold water to a smooth paste, pour 150 cc. of !boiling!
water over it, warm for a moment on the hot plate, and put it aside to settle.
Decant the supernatant liquid through a filter and use the clear filtrate; 5 cc.
of this solution are needed for a titration.

The solution of potato starch is less stable than the soluble starch. The
solid particles of the starch, if not removed by filtration, become so colored

by the iodine that they are not readily decolorized by the thiosulphate (Note
1).

[Note 1: The blue color which results when free iodine and starch are
brought together is probably not due to the formation of a true chemical
compound. It is regarded as a "solid solution" of iodine in starch. Although
it is unstable, and easily destroyed by heat, it serves as an indicator for the
presence of free iodine of remarkable sensitiveness, and makes the
iodometric processes the most satisfactory of any in the field of volumetric
analysis.]

COMPARISON OF IODINE AND THIOSULPHATE SOLUTIONS

PROCEDURE.—Place the solutions in burettes (the iodine in a glass-
stoppered burette), observing the usual precautions. Run out 40 cc. of the
thiosulphate solution into a beaker, dilute with 150 cc. of water, add 1 cc. to
2 cc. of the soluble starch solution, and titrate with the iodine to the
appearance of the blue of the iodo-starch. Repeat until the ratio of the two
solutions is established, remembering all necessary corrections for burettes
and for temperature changes.

STANDARDIZATION OF SOLUTIONS

Commercial iodine is usually not sufficiently pure to permit of its use as a
standard for thiosulphate solutions or the direct preparation of a standard
solution of iodine. It is likely to contain, beside moisture, some iodine
chloride, if chlorine was used to liberate the iodine when it was prepared. It
may be purified by sublimation after mixing it with a little potassium
iodide, which reacts with the iodine chloride, forming potassium chloride
and setting free the iodine. The sublimed iodine is then dried by placing it

in a closed container over concentrated sulphuric acid. It may then be
weighed in a stoppered weighing-tube and dissolved in a solution of
potassium iodide in a stoppered flask to prevent loss of iodine by
volatilization. About 18 grams of the iodide and twelve grams of iodine per
liter are required for an approximately tenth-normal solution.

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

The substances usually employed for the standardization of a
thiosulphate solution are potassium bromate and metallic copper. The
former is obtainable in pure condition or may be easily purified by re-
crystallization. Copper wire of high grade is sufficiently pure to serve as a
standard. Both potassium bromate and cupric salts in solution will liberate
iodine from an iodide, which is then titrated with the thiosulphate solution.

The reactions involved are the following:

(a) KBrO_{3} + 6KI + 3H_{2}SO_{4} —> KBr + 3I_{2} +
3K_{2}SO_{4} + 3H_{2}O,

(b) 3Cu + 8HNO_{3} —> 3Cu(NO_{3}){2} + 2NO + 4H{2}O,
2Cu(NO_{3}){2} + 4KI —> 2CuI + 4KNO{3} + I_{2}.

Two methods for the direct standardization of the sodium thiosulphate
solution are here described, and one for the direct standardization of the
iodine solution.

!Method A!

PROCEDURE.—Weigh out into 500 cc. beakers two portions of about
0.150-0.175 gram of potassium bromate. Dissolve each of these in 50 cc. of
water, and add 10 cc. of a potassium iodide solution containing 3 grams of
the salt in that volume (Note 1). Add to the mixture 10 cc. of dilute
sulphuric acid (1 volume of sulphuric acid with 5 volumes of water), allow
the solution to stand for three minutes, and dilute to 150 cc. (Note 2). Run
in thiosulphate solution from a burette until the color of the liberated iodine
is nearly destroyed, and then add 1 cc. or 2 cc. of starch solution, titrate to
the disappearance of the iodo-starch blue, and finally add iodine solution
until the color is just restored. Make a blank test for the amount of
thiosulphate solution required to react with the iodine liberated by the
iodate which is generally present in the potassium iodide solution, and
deduct this from the total volume used in the titration.

From the data obtained, calculate the relation of the thiosulphate solution
to a normal solution, and subsequently calculate the similar value for the
iodine solution.

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

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

!Method B!

PROCEDURE.—Weigh out two portions of 0.25-0.27 gram of clean
copper wire into 250 cc. Erlenmeyer flasks (Note 1). Add to each 5 cc. of
concentrated nitric acid (sp. gr. 1.42) and 25 cc. of water, cover, and warm
until solution is complete. Add 5 cc. of bromine water and boil until the
excess of bromine is expelled. Cool, and add strong ammonia (sp. gr. 0.90)
drop by drop until a deep blue color indicates the presence of an excess.
Boil the solution until the deep blue is replaced by a light bluish green, or a
brown stain appears on the sides of the flask (Note 2). Add 10 cc. of strong
acetic acid (sp. gr. 1.04), cool under the water tap, and add a solution of
potassium iodide (Note 3) containing about 3 grams of the salt, and titrate
with thiosulphate solution until the color of the liberated iodine is nearly
destroyed. Then add 1-2 cc. of freshly prepared starch solution, and add
thiosulphate solution, drop by drop, until the blue color is discharged.

From the data obtained, including the "blank test" of the iodide, calculate
the relation of the thiosulphate solution to the normal.

[Note 1: While copper wire of commerce is not absolutely pure, the
requirements for its use as a conductor of electricity are such that the
impurities constitute only a few hundredths of one per cent and are
negligible for analytical purposes.]

[Note 2: Ammonia neutralizes the free nitric acid. It should be added in
slight excess only, since the excess must be removed by boiling, which is
tedious. If too much ammonia is present when acetic acid is added, the

resulting ammonium acetate is hydrolyzed, and the ammonium hydroxide
reacts with the iodine set free.]

[Note 3: A considerable excess of potassium iodide is necessary for the
prompt liberation of iodine. While a large excess will do no harm, the cost
of this reagent is so great that waste should be avoided.]

!Method C!

PROCEDURE.—Weigh out into 500 cc. beakers two portions of 0.175-
0.200 gram each of pure arsenious oxide. Dissolve each of these in 10 cc. of
sodium hydroxide solution, with stirring. Dilute the solutions to 150 cc. and
add dilute hydrochloric acid until the solutions contain a few drops in
excess, and finally add to each a concentrated solution of 5 grams of pure
sodium bicarbonate (NaHCO_{3}) in water. Cover the beakers before
adding the bicarbonate, to avoid loss. Add the starch solution and titrate
with the iodine to the appearance of the blue of the iodo-starch, taking care
not to pass the end-point by more than a few drops (Note 1).

From the corrected volume of the iodine solution used to oxidize the
arsenious oxide, calculate its relation to the normal. From the ratio between
the solutions, calculate the similar value for the thiosulphate solution.

[Note 1: Arsenious oxide dissolves more readily in caustic alkali than in a
bicarbonate solution, but the presence of caustic alkali during the titration is
not admissible. It is therefore destroyed by the addition of acid, and the
solution is then made neutral with the solution of bicarbonate, part of which
reacts with the acid, the excess remaining in solution.

The reaction during titration is the following:

Na_{3}AsO_{3} + I_{2} + 2NaHCO_{3} —> Na_{3}AsO_{4} + 2NaI
+ 2CO_{2}
+ H_{2}O

As the reaction between sodium thiosulphate and iodine is not always
free from secondary reactions in the presence of even the weakly alkaline
bicarbonate, it is best to avoid the addition of any considerable excess of
iodine. Should the end-point be passed by a few drops, the thiosulphate may
be used to correct it.]

DETERMINATION OF COPPER IN ORES

Copper ores vary widely in composition from the nearly pure copper
minerals, such as malachite and copper sulphide, to very low grade
materials which contain such impurities as silica, lead, iron, silver, sulphur,
arsenic, and antimony. In nearly all varieties there will be found a siliceous
residue insoluble in acids. The method here given, which is a modification
of that described by A.H. Low (!J. Am. Chem. Soc.! (1902), 24, 1082),
provides for the extraction of the copper from commonly occurring ores,
and for the presence of their common impurities. For practice analyses it is
advisable to select an ore of a fair degree of purity.

PROCEDURE.— Weigh out two portions of about 0.5 gram each of the
ore (which should be ground until no grit is detected) into 250 cc.
Erlenmeyer flasks or small beakers. Add 10 cc. of concentrated nitric acid
(sp. gr. 1.42) and heat very gently until the ore is decomposed and the acid
evaporated nearly to dryness (Note 1). Add 5 cc. of concentrated
hydrochloric acid (sp. gr. 1.2) and warm gently. Then add about 7 cc. of
concentrated sulphuric acid (sp. gr. 1.84) and evaporate over a free flame
until the sulphuric acid fumes freely (Note 2). It has then displaced nitric
and hydrochloric acid from their compounds.

Cool the flask or beaker, add 25 cc. of water, heat the solution to boiling,
and boil for two minutes. Filter to remove insoluble sulphates, silica and
any silver that may have been precipitated as silver chloride, and receive the

filtrate in a small beaker, washing the precipitate and filter paper with warm
water until the filtrate and washings amount to 75 cc. Bend a strip of
aluminium foil (5 cm. x 12 cm.) into triangular form and place it on edge in
the beaker. Cover the beaker and boil the solution (being careful to avoid
loss of liquid by spattering) for ten minutes, but do not evaporate to small
volume.

Wash the cover glass and sides of the beaker. The copper should now be
in the form of a precipitate at the bottom of the beaker or adhering loosely
to the aluminium sheet. Remove the sheet, wash it carefully with hydrogen
sulphide water and place it in a small beaker. Decant the solution through a
filter, wash the precipitated copper twice by decantation with hydrogen
sulphide water, and finally transfer the copper to the filter paper, where it is
again washed thoroughly, being careful at all times to keep the precipitated
copper covered with the wash water. Remove and discard the filtrate and
place an Erlenmeyer flask under the funnel. Pour 15 cc. of dilute nitric acid
(sp. gr. 1.20) over the aluminium foil in the beaker, thus dissolving any
adhering copper. Wash the foil with hot water and remove it. Warm this
nitric acid solution and pour it slowly through the filter paper, thereby
dissolving the copper on the paper, receiving the acid solution in the
Erlenmeyer flask. Before washing the paper, pour 5 cc. of saturated
bromine water (Note 3) through it and finally wash the paper carefully with
hot water and transfer any particles of copper which may be left on it to the
Erlenmeyer flask. Boil to expel the bromine. Add concentrated ammonia
drop by drop until the appearance of a deep blue coloration indicates an
excess. Boil until the deep blue is displaced by a light bluish green
coloration, or until brown stains form on the sides of the flask. Add 10 cc.
of strong acetic acid (Note 4) and cool under the water tap. Add a solution
containing about 3 grams of potassium iodide, as in the standardization, and
titrate with thiosulphate solution until the yellow of the liberated iodine is

nearly discharged. Add 1-2 cc. of freshly prepared starch solution and titrate
to the disappearance of the blue color.

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

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

[Note 2: Unless proper precautions are taken to insure the correct
concentrations of acid the copper will not precipitate quantitatively on the
aluminium foil; hence care must be taken to follow directions carefully at
this point. Lead and silver have been almost completely removed as
sulphate and chloride respectively, or they too would be precipitated on the
aluminium. Bismuth, though precipitated on aluminium, has no effect on
the analysis. Arsenic and antimony precipitate on aluminium and would
interfere with the titration if allowed to remain in the lower state of
oxidation.]

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

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

DETERMINATION OF ANTIMONY IN
STIBNITE

Stibnite is native antimony sulphide. Nearly pure samples of this mineral
are easily obtainable and should be used for practice, since many impurities,
notably iron, seriously interfere with the accurate determination of the
antimony by iodometric methods. It is, moreover, essential that the
directions with respect to amounts of reagents employed and concentration
of solutions should be followed closely.

PROCEDURE.—Grind the mineral with great care, and weigh out two
portions of 0.35-0.40 gram into small, dry beakers (100 cc.). Cover the
beakers and pour over the stibnite 5 cc. of concentrated hydrochloric acid
(sp. gr. 1.20) and warm gently on the water bath (Note 1). When the residue
is white, add to each beaker 2 grams of powdered tartaric acid (Note 2).
Warm the solution on the water bath for ten minutes longer, dilute the
solution very cautiously by adding water in portions of 5 cc., stopping if the
solution turns red. It is possible that no coloration will appear, in which case
cautiously continue the dilution to 125 cc. If a red precipitate or coloration
does appear, warm the solution until it is colorless, and again dilute
cautiously to a total volume of 125 cc. and boil for a minute (Note 3).

If a white precipitate of the oxychloride separates during dilution (which
should not occur if the directions are followed), it is best to discard the

determination and to start anew.

Carefully neutralize most of the acid with ammonium hydroxide solution
(sp. gr. 0.96), but leave it distinctly acid (Note 4). Dissolve 3 grams of
sodium bicarbonate in 200 cc. of water in a 500 cc. beaker, and pour the
cold solution of the antimony chloride into this, avoiding loss by
effervescence. Make sure that the solution contains an excess of the
bicarbonate, and then add 1 cc. or 2 cc. of starch solution and titrate with
iodine solution to the appearance of the blue, avoiding excess (Notes 5 and
6).

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

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

[Note 2: Antimony salts, such as the chloride, are readily hydrolyzed, and
compounds such as SbOCl are formed which are often relatively insoluble;
but in the presence of tartaric acid compounds with complex ions are
formed, and these are soluble. An excess of hydrochloric acid also prevents
precipitation of the oxychloride because the H^{+} ions from the acid
lessen the dissociation of the water and thus prevent any considerable
hydrolysis.]

[Note 3: The action of hydrochloric acid upon the sulphide sets free
sulphureted hydrogen, a part of which is held in solution by the acid. This is
usually expelled by the heating upon the water bath; but if it is not wholly
driven out, a point is reached during dilution at which the antimony
sulphide, being no longer held in solution by the acid, separates. If the
dilution is immediately stopped and the solution warmed, this sulphide is

again brought into solution and at the same time more of the sulphureted
hydrogen is expelled. This procedure must be continued until the
sulphureted hydrogen is all removed, since it reacts with iodine. If no
precipitation of the sulphide occurs, it is an indication that the sulphureted
hydrogen was all expelled on solution of the stibnite.]

[Note 4: Ammonium hydroxide is added to neutralize most of the acid,
thus lessening the amount of sodium bicarbonate to be added. The ammonia
should not neutralize all of the acid.]

[Note 5: The reaction which takes place during titration may be
expressed thus:

Na_{3}SbO_{3} + 2NaHCO_{3} + I_{2} —> Na_{3}SbO_{4} + 2NaI
+ H_{2}O + 2CO_{2}.]

[Note 6: If the end-point is not permanent, that is, if the blue of the iodo-
starch is discharged after standing a few moments, the cause may be an
insufficient quantity of sodium bicarbonate, leaving the solution slightly
acid, or a very slight precipitation of an antimony compound which is
slowly acted upon by the iodine when the latter is momentarily present in
excess. In either case it is better to discard the analysis and to repeat the
process, using greater care in the amounts of reagents employed.]

CHLORIMETRY

The processes included under the term !chlorimetry! comprise those
employed to determine chlorine, hypochlorites, bromine, and hypobromites.
The reagent employed is sodium arsenite in the presence of sodium
bicarbonate. The reaction in the case of the hypochlorites is

NaClO + Na_{3}AsO_{3} —> Na_{3}AsO_{4} + NaCl.

The sodium arsenite may be prepared from pure arsenious oxide, as
described below, and is stable for considerable periods; but commercial
oxide requires resublimation to remove arsenic sulphide, which may be
present in small quantity. To prepare the solution, dissolve about 5 grams of
the powdered oxide, accurately weighed, in 10 cc. of a concentrated sodium
hydroxide solution, dilute the solution to 300 cc., and make it faintly acid
with dilute hydrochloric acid. Add 30 grams of sodium bicarbonate
dissolved in a little water, and dilute the solution to exactly 1000 cc. in a
measuring flask. Transfer the solution to a dry liter bottle and mix
thoroughly.

It is possible to dissolve the arsenious oxide directly in a solution of
sodium bicarbonate, with gentle warming, but solution in sodium hydroxide
takes place much more rapidly, and the excess of the hydroxide is readily
neutralized by hydrochloric acid, with subsequent addition of the
bicarbonate to maintain neutrality during the titration.

The indicator required for this process is made by dipping strips of filter
paper in a starch solution prepared as described on page 76, to which 1
gram of potassium iodide has been added. These strips are allowed to drain
and spread upon a watch-glass until dry. When touched by a drop of the
solution the paper turns blue until the hypochlorite has all been reduced and
an excess of the arsenite has been added.

DETERMINATION OF THE AVAILABLE
CHLORINE IN BLEACHING POWDER

Bleaching powder consists mainly of a calcium compound which is a
derivative of both hydrochloric and hypochlorous acids. Its formula is
CaClOCl. Its use as a bleaching or disinfecting agent, or as a source of
chlorine, depends upon the amount of hypochlorous acid which it yields
when treated with a stronger acid. It is customary to express the value of
bleaching powder in terms of "available chlorine," by which is meant the
chlorine present as hypochlorite, but not the chlorine present as chloride.

PROCEDURE.—Weigh out from a stoppered test tube into a porcelain
mortar about 3.5 grams of bleaching powder (Note 1). Triturate the powder
in the mortar with successive portions of water until it is well ground and
wash the contents into a 500 cc. measuring flask (Note 2). Fill the flask to
the mark with water and shake thoroughly. Measure off 25 cc. of this semi-
solution in a measuring flask, or pipette, observing the precaution that the
liquid removed shall contain approximately its proportion of suspended
matter.

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

From the volume of solution required to react with the bleaching powder,
calculate the percentage of available chlorine in the latter, assuming the
titration reaction to be that between chlorine and arsenious oxide:

As_{4}O_{6} + 4Cl_{2} + 4H_{2}O —> 2As_{2}O_{5} + 8HCl

Note that only one twentieth of the original weight of bleaching powder
enters into the reaction.

[Note 1: The powder must be triturated until it is fine, otherwise the
lumps will inclose calcium hypochlorite, which will fail to react with the
arsenious acid. The clear supernatant liquid gives percentages which are
below, and the sediment percentages which are above, the average. The
liquid measured off should, therefore, carry with it its proper proportion of
the sediment, so far as that can be brought about by shaking the solution
just before removal of the aliquot part for titration.]

[Note 2: Bleaching powder is easily acted upon by the carbonic acid in
the air, which liberates the weak hypochlorous acid. This, of course, results
in a loss of available chlorine. The original material for analysis should be
kept in a closed container and protected form the air as far as possible. It is
difficult to obtain analytical samples which are accurately representative of
a large quantity of the bleaching powder. The procedure, as outlined, will
yield results which are sufficiently exact for technical purposes.]

III. PRECIPITATION METHODS

DETERMINATION OF SILVER BY THE
THIOCYANATE PROCESS

The addition of a solution of potassium or ammonium thiocyanate to one of
silver in nitric acid causes a deposition of silver thiocyanate as a white,
curdy precipitate. If ferric nitrate is also present, the slightest excess of the
thiocyanate over that required to combine with the silver is indicated by the
deep red which is characteristic of the thiocyanate test for iron.

The reactions involved are:

AgNO_{3} + KSCN —> AgSCN + KNO_{3}, 3KSCN + Fe(NO_{3})
{3} —> Fe(SCN){3} + 3KNO_{3}.

The ferric thiocyanate differs from the great majority of salts in that it is
but very little dissociated in aqueous solutions, and the characteristic color
appears to be occasioned by the formation of the un-ionized ferric salt.

The normal solution of potassium thiocyanate should contain an amount
of the salt per liter of solution which would yield sufficient (CNS)^{-} to
combine with one gram of hydrogen to form HCNS, i.e., a gram-molecular
weight of the salt or 97.17 grams. If the ammonium thiocyanate is used, the
amount is 76.08 grams. To prepare the solution for this determination,
which should be approximately 0.05 N, dissolve about 5 grams of
potassium thiocyanate, or 4 grams of ammonium thiocyanate, in a small

amount of water; dilute this solution to 1000 cc. in a liter bottle and mix as
usual.

Prepare 20 cc. of a saturated solution of ferric alum and add 5 cc. of
dilute nitric acid (sp. gr. 1.20). About 5 cc. of this solution should be used as
an indicator.

STANDARDIZATION

PROCEDURE.—Crush a small quantity of silver nitrate crystals in a mortar
(Note 1). Transfer them to a watch-glass and dry them for an hour at
110°C., protecting them from dust or other organic matter (Note 2). Weigh
out two portions of about 0.5 gram each and dissolve them in 50 cc. of
water. Add 10 cc. of dilute nitric acid which has been recently boiled to
expel the lower oxides of nitrogen, if any, and then add 5 cc. of the
indicator solution. Run in the thiocyanate solution from a burette, with
constant stirring, allowing the precipitate to settle occasionally to obtain an
exact recognition of the end-point, until a faint red tinge can be detected in
the solution.

From the data obtained, calculate the relation of the thiocyanate solution
to the normal.

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

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

It is plain that a standard solution of silver nitrate (made by weighing out
the crystals) is convenient or necessary if many titrations of this nature are
to be made. In the absence of such a solution the liability of passing the
end-point is lessened by setting aside a small fraction of the silver solution,
to be added near the close of the titration.]

DETERMINATION OF SILVER IN COIN

PROCEDURE.— Weigh out two portions of the coin of about 0.5 gram
each. Dissolve them in 15 cc. of dilute nitric acid (sp. gr. 1.2) and boil until
all the nitrous compounds are expelled (Note 1). Cool the solution, dilute to
50 cc., and add 5 cc. of the indicator solution, and titrate with the
thiocyanate to the appearance of the faint red coloration (Note 2).

From the corrected volume of the thiocyanate solution required, calculate
the percentage of silver in the coin.

[Note 1: The reaction with silver may be carried out in nitric acid
solutions and in the presence of copper, if the latter does not exceed 70 per
cent. Above that percentage it is necessary to add silver in known quantity
to the solution. The liquid must be cold at the time of titration and entirely
free from nitrous compounds, as these sometimes cause a reddening of the
indicator solution. All utensils, distilled water, the nitric acid and the
beakers must be free from chlorides, as the presence of these will cause
precipitation of silver chloride, thereby introducing an error.]

[Note 2: The solution containing the silver precipitate, as well as those
from the standardization, should be placed in the receptacle for "silver
residues" as a matter of economy.]

PART III

GRAVIMETRIC ANALYSIS

GENERAL DIRECTIONS

Gravimetric analyses involve the following principal steps: first, the
weighing of the sample; second, the solution of the sample; third, the
separation of some substance from solution containing, or bearing a definite
relation to, the constituent to be measured, under conditions which render
this separation as complete as possible; and finally, the segregation of that
substance, commonly by filtration, and the determination of its weight, or
that of some stable product formed from it on ignition. For example, the
gravimetric determination of aluminium is accomplished by solution of the
sample, by precipitation in the form of hydroxide, collection of the
hydroxide upon a filter, complete removal by washing of all foreign soluble
matter, and the burning of the filter and ignition of the precipitate to
aluminium oxide, in which condition it is weighed.

Among the operations which are common to nearly all gravimetric
analyses are precipitation, washing of precipitates, ignition of precipitates,
and the use of desiccators. In order to avoid burdensome repetitions in the
descriptions of the various gravimetric procedures which follow, certain
general instructions are introduced at this point. These instructions must,
therefore, be considered to be as much a part of all subsequent procedures
as the description of apparatus, reagents, or manipulations.

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

PRECIPITATION

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

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

Amorphous substances, such as ferric hydroxide, aluminium hydroxide,
or silicic acid, separate in a gelatinous form and are relatively difficult to
filter and wash. Substances of this class also exhibit a tendency to form,
with pure water, what are known as colloidal solutions. To prevent this as
far as possible, they are washed with solutions of volatile salts, as will be
described in some of the following procedures.

In all precipitations the reagent should be added slowly, with constant
stirring, and should be hot when circumstances permit. The slow addition is

less likely to occasion contamination of the precipitate by the inclosure of
other substances which may be in the solution, or of the reagent itself.

FUNNELS AND FILTERS

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

The glass funnels selected for use with paper filters should have an angle
as near 60° as possible, and a narrow stem about six inches in length. The
filters employed should be washed filters, i.e., those which have been
treated with hydrochloric and hydrofluoric acids, and which on incineration
leave a very small and definitely known weight of ash, generally about
.00003 gram. Such filters are readily obtainable on the market.

The filter should be carefully folded to fit the funnel according to either
of the two well-established methods described in the Appendix. It should
always be placed so that the upper edge of the paper is about one fourth
inch below the top of the funnel. Under no circumstances should the filter
extend above the edge of the funnel, as it is then utterly impossible to effect
complete washing.

To test the efficiency of the filter, fill it with distilled water. This water
should soon fill the stem completely, forming a continuous column of liquid
which, by its hydrostatic pressure, produces a gentle suction, thus materially
promoting the rapidity of filtration. Unless the filter allows free passage of

water under these conditions, it is likely to give much trouble when a
precipitate is placed upon it.

The use of a suction pump to promote filtration is rarely altogether
advantageous in quantitative analysis, if paper filters are employed. The
tendency of the filter to break, unless the point of the filter paper is
supported by a perforated porcelain cone or a small "hardened filter" of
parchment, and the tendency of the precipitates to pass through the pores of
the filter, more than compensate for the possible gain in time. On the other
hand, filtration by suction may be useful in the case of precipitates which
do not require ignition before weighing, or in the case of precipitates which
are to be discarded without weighing. This is best accomplished with the
aid of the special apparatus called a Gooch filter referred to above.

FILTRATION AND WASHING OF PRECIPITATES

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

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

The vessels which contain the initial filtrate should !always! be replaced
by clean ones, properly labeled, before the washing of a precipitate begins.
In many instances a finely divided precipitate which shows no tendency to
pass through the filter at first, while the solution is relatively dense, appears
at once in the washings. Under such conditions the advantages accruing
from the removal of the first filtrate are obvious, both as regards the

diminished volume requiring refiltration, and also the smaller number of
washings subsequently required.

Much time may often be saved by washing precipitates by decantation,
i.e., by pouring over them, while still in the original vessel, considerable
volumes of wash-water and allowing them to settle. The supernatant, clear
wash-water is then decanted through the filter, so far as practicable without
disturbing the precipitate, and a new portion of wash-water is added. This
procedure can be employed to special advantage with gelatinous
precipitates, which fill up the pores of the filter paper. As the medium from
which the precipitate is to settle becomes less dense it subsides less readily,
and it ultimately becomes necessary to transfer it to the filter and complete
the washing there.

A precipitate should never completely fill a filter. The wash-water should
be applied at the top of the filter, above the precipitate. It may be shown
mathematically that the washing is most !rapidly! accomplished by filling
the filter well to the top with wash-water each time, and allowing it to drain
completely after each addition; but that when a precipitate is to be washed
with the !least possible volume! of liquid the latter should be applied in
repeated !small! quantities.

Gelatinous precipitates should not be allowed to dry before complete
removal of foreign matter is effected. They are likely to shrink and crack,
and subsequent additions of wash-water pass through these channels only.

All filtrates and wash-waters without exception must be properly tested.
!This lies at the foundation of accurate work!, and the student should clearly
understand that it is only by the invariable application of this rule that
assurance of ultimate reliability can be secured. Every original filtrate must
be tested to prove complete precipitation of the compound to be separated,
and the wash-waters must also be tested to assure complete removal of

foreign material. In testing the latter, the amount first taken should be but a
few drops if the filtrate contains material which is to be subsequently
determined. When, however, the washing of the filter and precipitate is
nearly completed the amount should be increased, and for the final test not
less than 3 cc. should be used.

It is impossible to trust to one's judgment with regard to the washing of
precipitates; the washings from !each precipitate! of a series simultaneously
treated must be tested, since the rate of washing will often differ materially
under apparently similar conditions, !No exception can ever be made to this
rule!.

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

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

For the complete removal of precipitates from containing vessels, it is
often necessary to rub the sides of these vessels to loosen the adhering
particles. This can best be done by slipping over the end of a stirring rod a
soft rubber device sometimes called a "policeman."

DESICCATORS

Desiccators should be filled with fused, anhydrous calcium chloride, over
which is placed a clay triangle, or an iron triangle covered with silica tubes,

to support the crucible or other utensils. The cover of the desiccator should
be made air-tight by the use of a thin coating of vaseline.

Pumice moistened with concentrated sulphuric acid may be used in place
of the calcium chloride, and is essential in special cases; but for most
purposes the calcium chloride, if renewed occasionally and not allowed to
cake together, is practically efficient and does not slop about when the
desiccator is moved.

Desiccators should never remain uncovered for any length of time. The
dehydrating agents rapidly lose their efficiency on exposure to the air.

CRUCIBLES

It is often necessary in quantitative analysis to employ fluxes to bring into
solution substances which are not dissolved by acids. The fluxes in most
common use are sodium carbonate and sodium or potassium acid sulphate.
In gravimetric analysis it is usually necessary to ignite the separated
substance after filtration and washing, in order to remove moisture, or to
convert it through physical or chemical changes into some definite and
stable form for weighing. Crucibles to be used in fusion processes must be
made of materials which will withstand the action of the fluxes employed,
and crucibles to be used for ignitions must be made of material which will
not undergo any permanent change during the ignition, since the initial
weight of the crucible must be deducted from the final weight of the
crucible and product to obtain the weight of the ignited substance. The three
materials which satisfy these conditions, in general, are platinum, porcelain,
and silica.

Platinum crucibles have the advantage that they can be employed at high
temperatures, but, on the other hand, these crucibles can never be used

when there is a possibility of the reduction to the metallic state of metals
like lead, copper, silver, or gold, which would alloy with and ruin the
crucible. When platinum crucibles are used with compounds of arsenic or
phosphorus, special precautions are necessary to prevent damage. This
statement applies to both fusions and ignitions.

Fusions with sodium carbonate can be made only in platinum, since
porcelain or silica crucibles are attacked by this reagent. Acid sulphate
fusions, which require comparatively low temperatures, can sometimes be
made in platinum, although platinum is slightly attacked by the flux.
Porcelain or silica crucibles may be used with acid fluxes.

Silica crucibles are less likely to crack on heating than porcelain
crucibles on account of their smaller coefficient of expansion. Ignition of
substances not requiring too high a temperature may be made in porcelain
or silica crucibles.

Iron, nickel or silver crucibles are used in special cases.

In general, platinum crucibles should be used whenever such use is
practicable, and this is the custom in private, research or commercial
laboratories. Platinum has, however, become so valuable that it is liable to
theft unless constantly under the protection of the user. As constant
protection is often difficult in instructional laboratories, it is advisable, in
order to avoid serious monetary losses, to use porcelain or silica crucibles
whenever these will give satisfactory service. When platinum utensils are
used the danger of theft should always be kept in mind.

PREPARATION OF CRUCIBLES FOR USE

All crucibles, of whatever material, must always be cleaned, ignited and
allowed to cool in a desiccator before weighing, since all bodies exposed to
the air condense on their surfaces a layer of moisture which increases their
weight. The amount and weight of this moisture varies with the humidity of
the atmosphere, and the latter may change from hour to hour. The air in the
desiccator (see above) is kept at a constant and low humidity by the drying
agent which it contains. Bodies which remain in a desiccator for a sufficient
time (usually 20-30 minutes) retain, therefore, on their surfaces a constant
weight of moisture which is the same day after day, thus insuring constant
conditions.

Hot objects, such as ignited crucibles, should be allowed to cool in the air
until, when held near the skin, but little heat is noticeable. If this precaution
is not taken, the air within the desiccator is strongly heated and expands
before the desiccator is covered. As the temperature falls, the air contracts,
causing a reduction of air pressure within the covered vessel. When the
cover is removed (which is often rendered difficult) the inrush of air from
the outside may sweep light particles out of a crucible, thus ruining an
entire analysis.

Constant heating of platinum causes a slight crystallization of the surface
which, if not removed, penetrates into the crucible. Gentle polishing of the
surface destroys the crystalline structure and prevents further damage. If sea
sand is used for this purpose, great care is necessary to keep it from the
desk, since beakers are easily scratched by it, and subsequently crack on
heating.

Platinum crucibles stained in use may often be cleaned by the fusion in
them of potassium or sodium acid sulphate, or by heating with ammonium
chloride. If the former is used, care should be taken not to heat so strongly
as to expel all of the sulphuric acid, since the normal sulphates sometimes

expand so rapidly on cooling as to split the crucible. The fused material
should be poured out, while hot, on to a !dry! tile or iron surface.

IGNITION OF PRECIPITATES

Most precipitates may, if proper precautions are taken, be ignited without
previous drying. If, however, such precipitates can be dried without loss of
time to the analyst (as, for example, over night), it is well to submit them to
this process. It should, nevertheless, be remembered that a partially dried
precipitate often requires more care during ignition than a thoroughly moist
one.

The details of the ignition of precipitates vary so much with the character
of the precipitate, its moisture content, and temperature to which it is to be
heated, that these details will be given under the various procedures which
follow.

DETERMINATION OF CHLORINE IN
SODIUM CHLORIDE

!Method A. With the Use of a Gooch Filter!

PROCEDURE.—Carefully clean a weighing-tube containing the sodium
chloride, handling it as little as possible with the moist fingers, and weigh it
accurately to 0.0001 gram, recording the weight at once in the notebook
(see Appendix). Hold the tube over the top of a beaker (200-300 cc.), and
cautiously remove the stopper, noting carefully that no particles fall from it,
or from the tube, elsewhere than into the beaker. Pour out a small portion of
the chloride, replace the stopper, and determine by approximate weighing
how much has been removed. Continue this procedure until 0.25-0.30 gram
has been taken from the tube, then weigh accurately and record the weight
beneath the first in the notebook. The difference of the two weights
represents the weight of the chloride taken for analysis. Again weigh a
second portion of 0.25-0.30 gram into a second beaker of the same size as
the first. The beakers should be plainly marked to correspond with the
entries in the notebook. Dissolve each portion of the chloride in 150 cc. of
distilled water and add about ten drops of dilute nitric acid (sp. gr. 1.20)
(Note 2). Calculate the volume of silver nitrate solution required to effect
complete precipitation in each case, and add slowly about 5 cc. in excess of
that amount, with constant stirring. Heat the solutions cautiously to boiling,
stirring occasionally, and continue the heating and stirring until the

precipitates settle promptly, leaving a nearly clear supernatant liquid (Note
3). This heating should not take place in direct sunlight (Note 4). The
beaker should be covered with a watch-glass, and both boiling and stirring
so regulated as to preclude any possibility of loss of material. Add to the
clear liquid one or two drops of silver nitrate solution, to make sure that an
excess of the reagent is present. If a precipitate, or cloudiness, appears as
the drops fall into the solution, heat again, and stir until the whole
precipitate has coagulated. The solution is then ready for filtration.

Prepare a Gooch filter as follows: Fold over the top of a Gooch funnel
(Fig. 2) a piece of rubber-band tubing, such as is known as "bill-tie" tubing,
and fit into the mouth of the funnel a perforated porcelain crucible (Gooch
crucible), making sure that when the crucible is gently forced into the
mouth of the funnel an airtight joint results. (A small 1 or 1-1/4-inch glass
funnel may be used, in which case the rubber tubing is stretched over the
top of the funnel and then drawn up over the side of the crucible until an
air-tight joint is secured.)

[ILLUSTRATION: FIG. 2]

Fit the funnel into the stopper of a filter bottle, and connect the filter
bottle with the suction pump. Suspend some finely divided asbestos, which
has been washed with acid, in 20 to 30 cc. of water (Note 1); allow this to
settle, pour off the very fine particles, and then pour some of the mixture
cautiously into the crucible until an even felt of asbestos, not over 1/32 inch
in thickness, is formed. A gentle suction must be applied while preparing
this felt. Wash the felt thoroughly by passing through it distilled water until
all fine or loose particles are removed, increasing the suction at the last until
no more water can be drawn out of it; place on top of the felt the small,
perforated porcelain disc and hold it in place by pouring a very thin layer of
asbestos over it, washing the whole carefully; then place the crucible in a

small beaker, and place both in a drying closet at 100-110°C. for thirty to
forty minutes. Cool the crucible in a desiccator, and weigh. Heat again for
twenty to thirty minutes, cool, and again weigh, repeating this until the
weight is constant within 0.0003 gram. The filter is then ready for use.

Place the crucible in the funnel, and apply a gentle suction, !after which!
the solution to be filtered may be poured in without disturbing the asbestos
felt. When pouring liquid onto a Gooch filter hold the stirring-rod at first
well down in the crucible, so that the liquid does not fall with any force
upon the asbestos, and afterward keep the crucible will filled with the
solution.

Pour the liquid above the silver chloride slowly onto the filter, leaving
the precipitate in the beaker as far as possible. Wash the precipitate twice by
decantation with warm water; then transfer it to the filter with the aid of a
stirring-rod with a rubber tip and a stream from the wash-bottle.

Examine the first portions of the filtrate which pass through the filter
with great care for asbestos fibers, which are most likely to be lost at this
point. Refilter the liquid if any fibers are visible. Finally, wash the
precipitate thoroughly with warm water until free from soluble silver salts.
To test the washings, disconnect the suction at the flask and remove the
funnel or filter tube from the suction flask. Hold the end of the tube over the
mouth of a small test tube and add from a wash-bottle 2-3 cc. of water.
Allow the water to drip through into the test tube and add a drop of dilute
hydrochloric acid. No precipitate or cloud should form in the wash-water
(Note 16). Dry the filter and contents at 100-110°C. until the weight is
constant within 0.0003 gram, as described for the preparation of the filter.
Deduct the weight of the dry crucible from the final weight, and from the
weight of silver chloride thus obtained calculate the percentage of chlorine
in the sample of sodium chloride.

[Note 1: The washed asbestos for this type of filter is prepared by
digesting in concentrated hydrochloric acid, long-fibered asbestos which
has been cut in pieces of about 0.5 cm. in length. After digestion, the
asbestos is filtered off on a filter plate and washed with hot, distilled water
until free from chlorides. A small portion of the asbestos is shaken with
water, forming a thin suspension, which is bottled and kept for use.]

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

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

[Note 4: The precipitate should not be exposed to strong sunlight, since
under those conditions a reduction of the silver chloride ensues which is
accompanied by a loss of chlorine. The superficial alteration which the
chloride undergoes in diffused daylight is not sufficient to materially affect
the accuracy of the determination. It should be noted, however, that a slight
error does result from the effect of light upon the silver chloride precipitate
and in cases in which the greatest obtainable accuracy is required, the
procedure described under "Method B" should be followed, in which this
slight reduction of the silver chloride is corrected by subsequent treatment
with nitric and hydrochloric acids.]

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

[Note 6: The precipitate must be washed with warm water until it is
absolutely free from silver and sodium nitrates. It may be assumed that the
sodium salt is completely removed when the wash-water shows no evidence
of silver. It must be borne in mind that silver chloride is somewhat soluble
in hydrochloric acid, and only a single drop should be added. The washing
should be continued until no cloudiness whatever can be detected in 3 cc. of
the washings.

Silver chloride is but slightly soluble in water. The solubility varies with
its physical condition within small limits, and is about 0.0018 gram per liter
at 18°C. for the curdy variety usually precipitated. The chloride is also
somewhat soluble in solutions of many chlorides, in solutions of silver
nitrate, and in concentrated nitric acid.

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

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

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

PROCEDURE.—Weigh out two portions of sodium chloride of about
0.25-0.3 gram each and proceed with the precipitation of the silver chloride
as described under Method A above. When the chloride is ready for
filtration prepare two 9 cm. washed paper filters (see Appendix). Pour the
liquid above the precipitates through the filters, wash twice by decantation
and transfer the precipitates to the filters, finally washing them until free
from silver solution as described. The funnel should then be covered with a
moistened filter paper by stretching it over the top and edges, to which it
will adhere on drying. It should be properly labeled with the student's name
and desk number, and then placed in a drying closet, at a temperature of
about 100-110°C., until completely dry.

The perfectly dry filter is then opened over a circular piece of clean,
smooth, glazed paper about six inches in diameter, placed upon a larger
piece about twelve inches in diameter. The precipitate is removed from the
filter as completely as possible by rubbing the sides gently together, or by
scraping them cautiously with a feather which has been cut close to the quill
and is slightly stiff (Note 1). In either case, care must be taken not to rub off
any considerable quantity of the paper, nor to lose silver chloride in the
form of dust. Cover the precipitate on the glazed paper with a watch-glass
to prevent loss of fine particles and to protect it from dust from the air. Fold
the filter paper carefully, roll it into a small cone, and wind loosely around
!the top! a piece of small platinum wire (Note 2). Hold the filter by the wire
over a small porcelain crucible (which has been cleaned, ignited, cooled in
a desiccator, and weighed), ignite it, and allow the ash to fall into the
crucible. Place the crucible upon a clean clay triangle, on its side, and
ignite, with a low flame well at its base, until all the carbon of the filter has
been consumed. Allow the crucible to cool, add two drops of concentrated
nitric acid and one drop of concentrated hydrochloric acid, and heat !very

cautiously!, to avoid spattering, until the acids have been expelled; then
transfer the main portion of the precipitate from the glazed paper to the
cooled crucible, placing the latter on the larger piece of glazed paper and
brushing the precipitate from the smaller piece into it, sweeping off all
particles belonging to the determination.

Moisten the precipitate with two drops of concentrated nitric acid and
one drop of concentrated hydrochloric acid, and again heat with great
caution until the acids are expelled and the precipitate is white, when the
temperature is slowly raised until the silver chloride just begins to fuse at
the edges (Note 3). The crucible is then cooled in a desiccator and weighed,
after which the heating (without the addition of acids) is repeated, and it is
again weighed. This must be continued until the weight is constant within
0.0003 gram in two consecutive weighings. Deduct the weight of the
crucible, and calculate the percentage of chlorine in the sample of sodium
chloride taken for analysis.

[Note 1: The separation of the silver chloride from the filter is essential,
since the burning carbon of the paper would reduce a considerable quantity
of the precipitate to metallic silver, and its complete reconversion to the
chloride within the crucible, by means of acids, would be accompanied by
some difficulty. The small amount of silver reduced from the chloride
adhering to the filter paper after separating the bulk of the precipitate, and
igniting the paper as prescribed, can be dissolved in nitric acid, and
completely reconverted to chloride by hydrochloric acid. The subsequent
addition of the two acids to the main portion of the precipitate restores the
chlorine to any chloride which may have been partially reduced by the
sunlight. The excess of the acids is volatilized by heating.]

[Note 2: The platinum wire is wrapped around the top of the filter during
its incineration to avoid contact with any reduced silver from the reduction

of the precipitate. If the wire were placed nearer the apex, such contact
could hardly be avoided.]

[Note 3: Silver chloride should not be heated to complete fusion, since a
slight loss by volatilization is possible at high temperatures. The
temperature of fusion is not always sufficient to destroy filter shreds; hence
these should not be allowed to contaminate the precipitate.]

DETERMINATION OF IRON AND OF
SULPHUR IN FERROUS AMMONIUM
SULPHATE,

FESO_{4}.(NH_{4}){2}SO{4}.6H_{2}O

DETERMINATION OF IRON

PROCEDURE.—Weigh out into beakers (200-250 cc.) two portions of the
sample (Note 1) of about 1 gram each and dissolve these in 50 cc. of water,
to which 1 cc. of dilute hydrochloric acid (sp. gr. 1.12) has been added
(Note 2). Heat the solution to boiling, and while at the boiling point add
concentrated nitric acid (sp. gr. 1.42), !drop by drop! (noting the volume
used), until the brown coloration, which appears after the addition of a part
of the nitric acid, gives place to a yellow or red (Note 3). Avoid a large
excess of nitric acid, but be sure that the action is complete. Pour this
solution cautiously into about 200 cc. of water, containing a slight excess of
ammonia. Calculate for this purpose the amount of aqueous ammonia
required to neutralize the hydrochloric and nitric acids added (see Appendix
for data), and also to precipitate the iron as ferric hydroxide from the weight
of the ferrous ammonium sulphate taken for analysis, assuming it to be pure
(Note 4). The volume thus calculated will be in excess of that actually
required for precipitation, since the acids are in part consumed in the
oxidation process, or are volatilized. Heat the solution to boiling, and allow

the precipitated ferric hydroxide to settle. Decant the clear liquid through a
washed filter (9 cm.), keeping as much of the precipitate in the beaker as
possible. Wash twice by decantation with 100 cc. of hot water. Reserve the
filtrate. Dissolve the iron from the filter with hot, dilute hydrochloric acid
(sp. gr. 1.12), adding it in small portions, using as little as possible and
noting the volume used. Collect the solution in the beaker in which
precipitation took place. Add 1 cc. of nitric acid (sp. gr. 1.42), boil for a few
moments, and again pour into a calculated excess of ammonia.

Wash the precipitate twice by decantation, and finally transfer it to the
original filter. Wash continuously with hot water until finally 3 cc. of the
washings, acidified with nitric acid (Note 5), show no evidences of the
presence of chlorides when tested with silver nitrate. The filtrate and
washings are combined with those from the first precipitation and treated
for the determination of sulphur, as prescribed on page 112.

[Note 1: If a selection of pure material for analysis is to be made, crystals
which are cloudy are to be avoided on account of loss of water of
crystallization; and also those which are red, indicating the presence of
ferric iron. If, on the other hand, the value of an average sample of material
is desired, it is preferable to grind the whole together, mix thoroughly, and
take a sample from the mixture for analysis.]

[Note 2: When aqueous solutions of ferrous compounds are heated in the
air, oxidation of the Fe^{++} ions to Fe^{+++} ions readily occurs in the
absence of free acid. The H^{+} and OH^{-} ions from water are involved
in the oxidation process and the result is, in effect, the formation of some
ferric hydroxide which tends to separate. Moreover, at the boiling
temperature, the ferric sulphate produced by the oxidation hydrolyzes in
part with the formation of a basic ferric sulphate, which also tends to
separate from solution. The addition of the hydrochloric acid prevents the

formation of ferric hydroxide, and so far reduces the ionization of the water
that the hydrolysis of the ferric sulphate is also prevented, and no
precipitation occurs on heating.]

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

The ionic changes which are involved in this oxidation are perhaps most
simply expressed by the equation

3Fe^{++} + NO_{3}^{-}+ 4H^{+} —> 3Fe^{+++} + 2H_{2}O + NO,

the H^{+} ions coming from the acid in the solution, in this case either
the nitric or the hydrochloric acid. The full equation on which this is based
may be written thus:

6FeSO_{4} + 2HNO_{3} + 6HCl —> 2Fe_{2}(SO_{4}){3} + 2FeCl{3}
+ 2NO + 4H_{2}O,

assuming that only enough nitric acid is added to complete the
oxidation.]

[Note 4: The ferric hydroxide precipitate tends to carry down some
sulphuric acid in the form of basic ferric sulphate. This tendency is lessened
if the solution of the iron is added to an excess of OH^{-} ions from the
ammonium hydroxide, since under these conditions immediate and

complete precipitation of the ferric hydroxide ensues. A gradual
neutralization with ammonia would result in the local formation of a neutral
solution within the liquid, and subsequent deposition of a basic sulphate as
a consequence of a local deficiency of OH^{-} ions from the NH_{4}OH
and a partial hydrolysis of the ferric salt. Even with this precaution the
entire absence of sulphates from the first iron precipitate is not assured. It
is, therefore, redissolved and again thrown down by ammonia. The organic
matter of the filter paper may occasion a partial reduction of the iron during
solution, with consequent possibility of incomplete subsequent precipitation
with ammonia. The nitric acid is added to reoxidize this iron.

To avoid errors arising from the solvent action of ammoniacal liquids
upon glass, the iron precipitate should be filtered without unnecessary
delay.]

[Note 5: The washings from the ferric hydroxide are acidified with nitric
acid, before testing with silver nitrate, to destroy the ammonia which is a
solvent of silver chloride.

The use of suction to promote filtration and washing is permissible,
though not prescribed. The precipitate should not be allowed to dry during
the washing.]

!Ignition of the Iron Precipitate!

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

Fold the top of the filter paper over the moist precipitate of ferric
hydroxide and transfer it cautiously to the crucible. Wipe the inside of the

funnel with a small fragment of washed filter paper, if necessary, and place
the paper in the crucible.

Incline the crucible on its side, on a triangle supported on a ring-stand,
and stand the cover on edge at the mouth of the crucible. Place a burner
below the front edge of the crucible, using a low flame and protecting it
from drafts of air by means of a chimney. The heat from the burner is thus
reflected into the crucible and dries the precipitate without danger of loss as
the result of a sudden generation of steam within the mass of ferric
hydroxide. As the drying progresses the burner may be gradually moved
toward the base of the crucible and the flame increased until the paper of
the filter begins to char and finally to smoke, as the volatile matter is
expelled. This is known as "smoking off" a filter, and the temperature
should not be raised sufficiently high during this process to cause the paper
to ignite, as the air currents produced by the flame of the blazing paper may
carry away particles of the precipitate.

When the paper is fully charred, move the burner to the base of the
crucible and raise the temperature to the full heat of the burner for fifteen
minutes, with the crucible still inclined on its side, but without the cover
(Note 1). Finally set the crucible upright in the triangle, cover it, and heat at
the full temperature of a blast lamp or other high temperature burner. Cool
and weigh in the usual manner (Note 2). Repeat the strong heating until the
weight is constant within 0.0003 gram.

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

[Note 1: These directions for the ignition of the precipitate must be
closely followed. A ready access of atmospheric oxygen is of special
importance to insure the reoxidation to ferric oxide of any iron which may
be reduced to magnetic oxide (Fe_{3}O_{4}) during the combustion of the

filter. The final heating over the blast lamp is essential for the complete
expulsion of the last traces of water from the hydroxide.]

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

[Note 3: The gravimetric determination of aluminium or chromium is
comparable with that of iron just described, with the additional precaution
that the solution must be boiled until it contains but a very slight excess of
ammonia, since the hydroxides of aluminium and chromium are more
soluble than ferric hydroxide.

The most important properties of these hydroxides, from a quantitative
standpoint, other than those mentioned, are the following: All are
precipitable by the hydroxides of sodium and potassium, but always inclose
some of the precipitant, and should be reprecipitated with ammonium
hydroxide before ignition to oxides. Chromium and aluminium hydroxides
dissolve in an excess of the caustic alkalies and form anions, probably of
the formula AlO_2^{-} and CrO_{2}^{-}. Chromium hydroxide is
reprecipitated from this solution on boiling. When first precipitated the
hydroxides are all readily soluble in acids, but aluminium hydroxide
dissolves with considerable difficulty after standing or boiling for some
time. The precipitation of the hydroxides is promoted by the presence of
ammonium chloride, but is partially or entirely prevented by the presence of
tartaric or citric acids, glycerine, sugars, and some other forms of soluble
organic matter. The hydroxides yield on ignition an oxide suitable for
weighing (Al_{2}O_{3}, Cr_{2}O_{3}, Fe_{2}O_{3}).]

DETERMINATION OF SULPHUR

PROCEDURE.—Add to the combined filtrates from the ferric hydroxide
about 0.6 gram of anhydrous sodium carbonate; cover the beaker, and then
add dilute hydrochloric acid (sp. gr. 1.12) in moderate excess and evaporate
to dryness on the water bath. Add 10 cc. of concentrated hydrochloric acid
(sp. gr. 1.20) to the residue, and again evaporate to dryness on the bath.
Dissolve the residue in water, filter if not clear, transfer to a 700 cc. beaker,
dilute to about 400 cc., and cautiously add hydrochloric acid until the
solution shows a distinctly acid reaction (Note 1). Heat the solution to
boiling, and add !very slowly! and with constant stirring, 20 cc. in excess of
the calculated amount of a hot barium chloride solution, containing about
20 grams BaCl_{2}.2H_{2}O per liter (Notes 2 and 3). Continue the
boiling for about two minutes, allow the precipitate to settle, and decant the
liquid at the end of half an hour (Note 4). Replace the beaker containing the
original filtrate by a clean beaker, wash the precipitated sulphate by
decantation with hot water, and subsequently upon the filter until it is freed
from chlorides, testing the washings as described in the determination of
iron. The filter is then transferred to a platinum or porcelain crucible and
ignited, as described above, until the weight is constant (Note 5). From the
weight of barium sulphate (BaSO_{4}) obtained, calculate the percentage
of sulphur (S) in the sample.

[Note 1: Barium sulphate is slightly soluble in hydrochloric acid, even
dilute, probably as a result of the reduction in the degree of dissociation of
sulphuric acid in the presence of the H^{+} ions of the hydrochloric acid,
and possibly because of the formation of a complex anion made up of
barium and chlorine; hence only the smallest excess should be added over
the amount required to acidify the solution.]

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

Ba^{++} + SO_{4}^{—} —> [BaSO_{4}]

This case affords one of the best illustrations of the effect of an excess of
a precipitant in decreasing the solubility of a precipitate. If the conditions
are considered which exist at the moment when just enough of the Ba^{++}
ions have been added to correspond to the SO_{4}^{—} ions in the
solution, it will be seen that nearly all of the barium sulphate has been
precipitated, and that the small amount which then remains in the solution
which is in contact with the precipitate must represent a saturated solution
for the existing temperature, and that this solution is comparable with a
solution of sugar to which more sugar has been added than will dissolve. It
should be borne in mind that the quantity of barium sulphate in this
!saturated solution is a constant quantity! for the existing conditions. The
dissolved barium sulphate, like any electrolyte, is dissociated, and the
equilibrium conditions may be expressed thus:

(!Conc'n Ba^{++} x Conc'n SO_{4}^{—})/(Conc'n BaSO_{4}) =
Const.!,

and since !Conc'n BaSO_{4}! for the saturated solution has a constant
value (which is very small), it may be eliminated, when the expression
becomes !Conc'n Ba^{++} x Conc'n SO_{4}^{—} = Const.!, which is the

"solubility product" of BaSO_{4}. If, now, an excess of the precipitant, a
soluble barium salt, is added in the form of a relatively concentrated
solution (the slight change of volume of a few cubic centimeters may be
disregarded for the present discussion) the concentration of the Ba^{++}
ions is much increased, and as a consequence the !Conc'n SO_{4}! must
decrease in proportion if the value of the expression is to remain constant,
which is a requisite condition if the law of mass action upon which our
argument depends holds true. In other words, SO_{4}^{—} ions must
combine with some of the added Ba^{++} ions to form [BaSO_{4}]; but it
will be recalled that the solution is already saturated with BaSO_{4}, and
this freshly formed quantity must, therefore, separate and add itself to the
precipitate. This is exactly what is desired in order to insure more complete
precipitation and greater accuracy, and leads to the conclusion that the
larger the excess of the precipitant added the more successful the analysis;
but a practical limit is placed upon the quantity of the precipitant which
may be properly added by other conditions, as stated in the following note.]

[Note 3: Barium sulphate, in a larger measure than most compounds,
tends to carry down other substances which are present in the solution from
which it separates, even when these other substances are relatively soluble,
and including the barium chloride used as the precipitant. This is also
notably true in the case of nitrates and chlorates of the alkalies, and of ferric
compounds; and, since in this analysis ammonium nitrate has resulted from
the neutralization of the excess of the nitric acid added to oxidize the iron, it
is essential that this should be destroyed by repeated evaporation with a
relatively large quantity of hydrochloric acid. During evaporation a mutual
decomposition of the two acids takes place, and the nitric acid is finally
decomposed and expelled by the excess of hydrochloric acid.

Iron is usually found in the precipitate of barium sulphate when thrown
down from hot solutions in the presence of ferric salts. This, according to

Kuster and Thiel (!Zeit. anorg. Chem.!, 22, 424), is due to the formation of
a complex ion (Fe(SO_{4})_{2}) which precipitates with the Ba^{++} ion,
while Richards (!Zeit. anorg. Chem.!, 23, 383) ascribes it to hydrolytic
action, which causes the formation of a basic ferric complex which is
occluded in the barium precipitate. Whatever the character of the compound
may be, it has been shown that it loses sulphuric anhydride upon ignition,
causing low results, even though the precipitate contains iron.

The contamination of the barium sulphate by iron is much less in the
presence of ferrous than ferric salts. If, therefore, the sulphur alone were to
be determined in the ferrous ammonium sulphate, the precipitation by
barium might be made directly from an aqueous solution of the salt, which
had been made slightly acid with hydrochloric acid.]

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

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

[Note 5: A reduction of barium sulphate to the sulphide may very readily
be caused by the reducing action of the burning carbon of the filter, and
much care should be taken to prevent any considerable reduction from this
cause. Subsequent ignition, with ready access of air, reconverts the sulphide

to sulphate unless a considerable reduction has occurred. In the latter case it
is expedient to add one or two drops of sulphuric acid and to heat cautiously
until the excess of acid is expelled.]

[Note 6: Barium sulphate requires about 400,000 parts of water for its
solution. It is not decomposed at a red heat but suffers loss, probably of
sulphur trioxide, at a temperature above 900°C.]

DETERMINATION OF SULPHUR IN BARIUM
SULPHATE

PROCEDURE.—Weigh out, into platinum crucibles, two portions of about
0.5 gram of the sulphate. Mix each in the crucible with five to six times its
weight of anhydrous sodium carbonate. This can best be done by placing
the crucible on a piece of glazed paper and stirring the mixture with a clean,
dry stirring-rod, which may finally be wiped off with a small fragment of
filter paper, the latter being placed in the crucible. Cover the crucible and
heat until a quiet, liquid fusion ensues. Remove the burner, and tip the
crucible until the fused mass flows nearly to its mouth. Hold it in that
position until the mass has solidified. When cold, the material may usually
be detached in a lump by tapping the crucible or gently pressing it near its
upper edge. If it still adheres, a cubic centimeter or so of water may be
placed in the cold crucible and cautiously brought to boiling, when the cake
will become loosened and may be removed and placed in about 250 cc. of
hot, distilled water to dissolve. Clean the crucible completely, rubbing the
sides with a rubber-covered stirring-rod, if need be.

When the fused mass has completely disintegrated and nothing further
will dissolve, decant the solution from the residue of barium carbonate
(Note 1). Pour over the residue 20 cc. of a solution of sodium carbonate and
10 cc. of water and heat to gentle boiling for about three minutes (Note 2).
Filter off the carbonate and wash it with hot water, testing the slightly

acidified washings for sulphate and preserving any precipitates which
appear in these tests. Acidify the filtrate with hydrochloric acid until just
acid, bring to boiling, and slowly add hot barium chloride solution, as in the
preceding determination. Add also any tests from the washings in which
precipitates have appeared. Filter, wash, ignite, and weigh.

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

[Note 1: This alkaline fusion is much employed to disintegrate substances
ordinarily insoluble in acids into two components, one of which is water
soluble and the other acid soluble. The reaction involved is:

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

As the sodium sulphate is soluble in water, and the barium carbonate
insoluble, a separation between them is possible and the sulphur can be
determined in the water-soluble portion.

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

[Note 2: The barium carbonate is boiled with sodium carbonate solution
before filtration because the reaction above is reversible; and it is only by
keeping the sodium carbonate present in excess until nearly all of the
sodium sulphate solution has been removed by filtration that the reversion
of some of the barium carbonate to barium sulphate is prevented. This is an
application of the principle of mass action, in which the concentration of the
reagent (the carbonate ion) is kept as high as practicable and that of the

sulphate ion as low as possible, in order to force the reaction in the desired
direction (see Appendix).]

DETERMINATION OF PHOSPHORIC
ANHYDRIDE IN APATITE

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

For the purpose of gravimetric determination, phosphoric acid is usually
precipitated from ammoniacal solutions in the form of magnesium
ammonium phosphate which, on ignition, is converted into magnesium
pyrophosphate. Since the calcium phosphate of the apatite is also insoluble
in ammoniacal solutions, this procedure cannot be applied directly. The
separation of the phosphoric acid from the calcium must first be
accomplished by precipitation in the form of ammonium phosphomolybdate
in nitric acid solution, using ammonium molybdate as the precipitant. The
"yellow precipitate," as it is often called, is not always of a definite
composition, and therefore not suitable for direct weighing, but may be
dissolved in ammonia, and the phosphoric acid thrown out as magnesium
ammonium phosphate from the solution.

Of the substances likely to occur in apatite, silicic acid alone interferes
with the precipitation of the phosphoric acid in nitric acid solution.

PRECIPITATION OF AMMONIUM PHOSPHOMOLYBDATE

PROCEDURE.—Grind the mineral in an agate mortar until no grit is
perceptible. Transfer the substance to a weighing-tube, and weigh out two
portions, not exceeding 0.20 gram each (Note 1) into two beakers of about
200 cc. capacity. Pour over them 20 cc. of dilute nitric acid (sp. gr. 1.2) and
warm gently until solvent action has apparently ceased. Evaporate the
solution cautiously to dryness, heat the residue for about an hour at 100-
110°C., and treat it again with nitric acid as described above; separate the
residue of silica by filtration on a small filter (7 cm.) and wash with warm
water, using as little as possible (Note 2). Receive the filtrate in a beaker
(200-500 cc.). Test the washings with ammonia for calcium phosphate, but
add all such tests in which a precipitate appears to the original nitrate (Note
3). The filtrate and washings must be kept as small as possible and should
not exceed 100 cc. in volume. Add aqueous ammonia (sp. gr. 0.96) until the
precipitate of calcium phosphate first produced just fails to redissolve, and
then add a few drops of nitric acid until this is again brought into solution
(Note 4). Warm the solution until it cannot be comfortably held in the hand
(about 60°C.) and, after removal of the burner, add 75 cc. of ammonium
molybdate solution which has been !gently! warmed, but which must be
perfectly clear. Allow the mixture to stand at a temperature of about 50 or
60°C. for twelve hours (Notes 5 and 6). Filter off the yellow precipitate on a
9 cm. filter, and wash by decantation with a solution of ammonium nitrate
made acid with nitric acid.[1] Allow the precipitate to remain in the beaker
as far as possible. Test the washings for calcium with ammonia and
ammonium oxalate (Note 3).

[Footnote 1: This solution is prepared as follows: Mix 100 cc. of
ammonia solution (sp. gr. 0.96) with 325 cc. of nitric acid (sp. gr. 1.2) and
dilute with 100 cc. of water.]

Add 10 cc. of molybdate solution to the nitrate, and leave it for a few
hours. It should then be carefully examined for a !yellow! precipitate; a

white precipitate may be neglected.

[Note 1: Magnesium ammonium phosphate, as noted below, is slightly
soluble under the conditions of operation. Consequently the unavoidable
errors of analysis are greater in this determination than in those which have
preceded it, and some divergence may be expected in duplicate analyses. It
is obvious that the larger the amount of substance taken for analysis the less
will be the relative loss or gain due to unavoidable experimental errors; but,
in this instance, a check is placed upon the amount of material which may
be taken both by the bulk of the resulting precipitate of ammonium
phosphomolybdate and by the excessive amount of ammonium molybdate
required to effect complete separation of the phosphoric acid, since a liberal
excess above the theoretical quantity is demanded. Molybdic acid is one of
the more expensive reagents.]

[Note 2: Soluble silicic acid would, if present, partially separate with the
phosphomolybdate, although not in combination with molybdenum. Its
previous removal by dehydration is therefore necessary.]

[Note 3: When washing the siliceous residue the filtrate may be tested for
calcium by adding ammonia, since that reagent neutralizes the acid which
holds the calcium phosphate in solution and causes precipitation; but after
the removal of the phosphoric acid in combination with the molybdenum,
the addition of an oxalate is required to show the presence of calcium.]

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

[Note 5: The precipitation of the phosphomolybdate takes place more
promptly in warm than in cold solutions, but the temperature should not
exceed 60°C. during precipitation; a higher temperature tends to separate

molybdic acid from the solution. This acid is nearly white, and its
deposition in the filtrate on long standing should not be mistaken for a
second precipitation of the yellow precipitate. The addition of 75 cc. of
ammonium molybdate solution insures the presence of a liberal excess of
the reagent, but the filtrate should be tested as in all quantitative procedures.

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

[Note 6: The composition of the "yellow precipitate" undoubtedly varies
slightly with varying conditions at the time of its formation. Its composition
may probably fairly be represented by the formula, (NH_{4})
{3}PO{4}.12MoO_{3}.H_{2}O, when precipitated under the conditions
prescribed in the procedure. Whatever other variations may occur in its
composition, the ratio of 12 MoO_{3}:1 P seems to hold, and this fact is
utilized in volumetric processes for the determination of phosphorus, in
which the molybdenum is reduced to a lower oxide and reoxidized by a
standard solution of potassium permanganate. In principle, the procedure is
comparable with that described for the determination of iron by
permanganate.]

PRECIPITATION OF MAGNESIUM AMMONIUM PHOSPHATE

PROCEDURE.—Dissolve the precipitate of phosphomolybdate upon the
filter by pouring through it dilute aqueous ammonia (one volume of dilute
ammonia (sp. gr. 0.96) and three volumes of water, which should be
carefully measured), and receive the solution in the beaker containing the
bulk of the precipitate. The total volume of nitrate and washings should not
much exceed 100 cc. Acidify the solution with dilute hydrochloric acid, and

heat it nearly to boiling. Calculate the volume of magnesium ammonium
chloride solution ("magnesia mixture") required to precipitate the
phosphoric acid, assuming 40 per cent P_{2}O_{5} in the apatite. Measure
out about 5 cc. in excess of this amount, and pour it into the acid solution.
Then add slowly dilute ammonium hydroxide (1 volume of strong ammonia
(sp. gr. 0.90) and 9 volumes of water), stirring constantly until a precipitate
forms. Then add a volume of filtered, concentrated ammonia (sp. gr. 0.90)
equal to one third of the volume of liquid in the beaker (Note 1). Allow the
whole to cool. The precipitated magnesium ammonium phosphate should
then be definitely crystalline in appearance (Note 2). (If it is desired to
hasten the precipitation, the solution may be cooled, first in cold and then in
ice-water, and stirred !constantly! for half an hour, when precipitation will
usually be complete.)

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

[Note 1: Magnesium ammonium phosphate is not a wholly insoluble
substance, even under the most favorable analytical conditions. It is least
soluble in a liquid containing one fourth of its volume of concentrated
aqueous ammonia (sp. gr. 0.90) and this proportion should be carefully
maintained as prescribed in the procedure. On account of this slight
solubility the volume of solutions should be kept as small as possible and
the amount of wash-water limited to that absolutely required.

A large excess of the magnesium solution tends both to throw out
magnesium hydroxide (shown by a persistently flocculent precipitate) and
to cause the phosphate to carry down molybdic acid. The tendency of the

magnesium precipitate to carry down molybdic acid is also increased if the
solution is too concentrated. The volume should not be less than 90 cc., nor
more than 125 cc., at the time of the first precipitation with the magnesia
mixture.]

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

REPRECIPITATION AND IGNITION OF MAGNESIUM
AMMONIUM PHOSPHATE

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

PROCEDURE.—Dissolve the precipitate from the filter in a little dilute
hydrochloric acid (sp. gr. 1.12), allowing the acid solution to run into the
beaker in which the original precipitation was made (Note 1). Wash the
filter with water until the wash-water shows no test for chlorides, but avoid
an unnecessary amount of wash-water. Add to the solution 2 cc. (not more)
of magnesia mixture, and then dilute ammonium hydroxide solution (sp. gr.
0.96), drop by drop, with constant stirring, until the liquid smells distinctly

of ammonia. Stir for a few moments and then add a volume of strong
ammonia (sp. gr. 0.90), equal to one third of the volume of the solution.
Allow the solution to stand for some hours, and then filter off the
magnesium ammonium phosphate, which should be distinctly crystalline in
character. Wash the precipitate with dilute ammonia water, as prescribed
above, until, finally, 3 cc. of the washings, after acidifying with nitric acid,
show no evidence of chlorides. Test both filtrates for complete precipitation
by adding a few cubic centimeters of magnesia mixture and allowing them
to stand for some time.

Transfer the moist precipitate to a weighed porcelain or platinum crucible
and ignite, using great care to raise the temperature slowly while drying the
filter in the crucible, and to insure the ready access of oxygen during the
combustion of the filter paper, thus guarding against a possible reduction of
the phosphate, which would result in disastrous consequences both to the
crucible, if of platinum, and the analysis. Do not raise the temperature
above moderate redness until the precipitate is white. (Keep this precaution
well in mind.) Ignite finally at the highest temperature of the Tirrill burner,
and repeat the heating until the weight is constant. If the ignited precipitate
is persistently discolored by particles of unburned carbon, moisten the mass
with a drop or two of concentrated nitric acid and heat cautiously, finally
igniting strongly. The acid will dissolve magnesium pyrophosphate from the
surface of the particles of carbon, which will then burn away. Nitric acid
also aids as an oxidizing agent in supplying oxygen for the combustion of
the carbon.

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

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

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

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

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

that is, the PO_{4}^{—} ions and the H^{+} ions lose their identity in
the formation of the new ion, HPO_{4}^{—}, and this continues until the
magnesium ammonium phosphate is entirely dissolved.]

[Note 2: During ignition the magnesium ammonium phosphate loses
ammonia and water and is converted into magnesium pyrophosphate:

2MgNH_{4}PO_{4} —> Mg_{2}P_{2}O_{7} + 2NH_{3} + H_{2}O.

The precautions mentioned on pages 111 and 123 must be observed with
great care during the ignition of this precipitate. The danger here lies in a
possible reduction of the phosphate by the carbon of the filter paper, or by
the ammonia evolved, which may act as a reducing agent. The phosphorus

then attacks and injures a platinum crucible, and the determination is
valueless.]

ANALYSIS OF LIMESTONE

Limestones vary widely in composition from a nearly pure marble through
the dolomitic limestones, containing varying amounts of magnesium, to the
impure varieties, which contain also ferrous and manganous carbonates and
siliceous compounds in variable proportions. Many other minerals may be
inclosed in limestones in small quantities, and an exact qualitative analysis
will often show the presence of sulphides or sulphates, phosphates, and
titanates, and the alkali or even the heavy metals. No attempt is made in the
following procedures to provide a complete quantitative scheme which
would take into account all of these constituents. Such a scheme for a
complete analysis of a limestone may be found in Bulletin No. 700 of the
United States Geological Survey. It is assumed that, for these practice
determinations, a limestone is selected which contains only the more
common constituents first enumerated above.

DETERMINATION OF MOISTURE

The determination of the amount of moisture in minerals or ores is often of
great importance. Ores which have been exposed to the weather during
shipment may have absorbed enough moisture to appreciably affect the
results of analysis. Since it is essential that the seller and buyer should make
their analyses upon comparable material, it is customary for each analyst to
determine the moisture in the sample examined, and then to calculate the

percentages of the various constituents with reference to a sample dried in
the air, or at a temperature a little above 100°C., which, unless the ore has
undergone chemical change because of the wetting, should be the same
before and after shipment.

PROCEDURE.—Spread 25 grams of the powdered sample on a weighed
watch-glass; weigh to the nearest 10 milligrams only and heat at 105°C.;
weigh at intervals of an hour, after cooling in a desiccator, until the loss of
weight after an hour's heating does not exceed 10 milligrams. It should be
noted that a variation in weight of 10 milligrams in a total weight of 25
grams is no greater relatively than a variation of 0.1 milligram when the
sample taken weighs 0.25 gram

DETERMINATION OF THE INSOLUBLE MATTER AND SILICA

PROCEDURE.—Weigh out two portions of the original powdered
sample (not the dried sample), of about 5 grams each, into 250 cc.
casseroles, and cover each with a watch-glass (Note 1). Pour over the
powder 25 cc. of water, and then add 50 cc. of dilute hydrochloric acid (sp.
gr. 1.12) in small portions, warming gently, until nothing further appears to
dissolve (Note 2). Evaporate to dryness on the water bath. Pour over the
residue a mixture of 5 cc. of water and 5 cc. of concentrated hydrochloric
acid (sp. gr. 1.2) and again evaporate to dryness, and finally heat for at least
an hour at a temperature of 110°C. Pour over this residue 50 cc. of dilute
hydrochloric acid (one volume acid (sp. gr. 1.12) to five volumes water),
and boil for about five minutes; then filter and wash twice with the dilute
hydrochloric acid, and then with hot water until free from chlorides.
Transfer the filter and contents to a porcelain crucible, dry carefully over a
low flame, and ignite to constant weight. The residue represents the
insoluble matter and the silica from any soluble silicates (Note 3).

Calculate the combined percentage of these in the limestone.

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

[Note 2: It is plain that the amount of the insoluble residue and also its
character will often depend upon the strength of acid used for solution of
the limestone. It cannot, therefore, be regarded as representing any well-
defined constituent, and its determination is essentially empirical.]

[Note 3: It is probable that some of the silicates present are wholly or
partly decomposed by the acid, and the soluble silicic acid must be
converted by evaporation to dryness, and heating, into white, insoluble
silica. This change is not complete after one evaporation. The heating at a
temperature somewhat higher than that of the water bath for a short time
tends to leave the silica in the form of a powder, which promotes
subsequent filtration. The siliceous residue is washed first with dilute acid
to prevent hydrolytic changes, which would result in the formation of
appreciable quantities of insoluble basic iron or aluminium salts on the filter
when washing with hot water.

If it is desired to determine the percentage of silica separately, the ignited
residue should be mixed in a platinum crucible with about six times its
weight of anhydrous sodium carbonate, and the procedure given on page
151 should be followed. The filtrate from the silica is then added to the
main filtrate from the insoluble residue.]

DETERMINATION OF FERRIC OXIDE AND
ALUMINIUM OXIDE (WITH MANGANESE)

PROCEDURE.—To the filtrate from the insoluble residue add ammonium
hydroxide until the solution just smells distinctly of ammonia, but do not
add an excess. Then add 5 cc. of saturated bromine water (Note 1), and boil
for five minutes. If the smell of ammonia has disappeared, again add
ammonium hydroxide in slight excess, and 3 cc. of bromine water, and heat
again for a few minutes. Finally add 10 cc. of ammonium chloride solution
and keep the solution warm until it barely smells of ammonia; then filter
promptly (Note 2). Wash the filter twice with hot water, then (after
replacing the receiving beaker) pour through it 25 cc. of hot, dilute
hydrochloric acid (one volume dilute HCl [sp. gr. 1.12] to five volumes
water). A brown residue insoluble in the acid may be allowed to remain on
the filter. Wash the filter five times with hot water, add to the filtrate
ammonium hydroxide and bromine water as described above, and repeat the
precipitation. Collect the precipitate on the filter already used, wash it free
from chlorides with hot water, and ignite and weigh as described for ferric
hydroxide on page 110. The residue after ignition consists of ferric oxide,
alumina, and mangano-manganic oxide (Mn_{3}O_{4}), if manganese is
present. These are commonly determined together (Note 3).

Calculate the percentage of the combined oxides in the limestone.

[Note 1: The addition of bromine water to the ammoniacal solutions
serves to oxidize any ferrous hydroxide to ferric hydroxide and to
precipitate manganese as MnO(OH)_{2}. The solution must contain not
more than a bare excess of hydroxyl ions (ammonium hydroxide) when it is
filtered, on account of the tendency of the aluminium hydroxide to
redissolve.

The solution should not be strongly ammoniacal when the bromine is
added, as strong ammonia reacts with the bromine, with the evolution of
nitrogen.]

[Note 2: The precipitate produced by ammonium hydroxide and bromine
should be filtered off promptly, since the alkaline solution absorbs carbon
dioxide from the air, with consequent partial precipitation of the calcium as
carbonate. This is possible even under the most favorable conditions, and
for this reason the iron precipitate is redissolved and again precipitated to
free it from calcium. When the precipitate is small, this reprecipitation may
be omitted.]

[Note 3: In the absence of significant amounts of manganese the iron and
aluminium may be separately determined by fusion of the mixed ignited
precipitate, after weighing, with about ten times its weight of acid
potassium sulphate, solution of the cold fused mass in water, and
volumetric determination of the iron, as described on page 66. The
aluminium is then determined by difference, after subtracting the weight of
ferric oxide corresponding to the amount of iron found.

If a separate determination of the iron, aluminium, and manganese is
desired, the mixed precipitate may be dissolved in acid before ignition, and
the separation effected by special methods (see, for example, Fay,
!Quantitative Analyses!, First Edition, pp. 15-19 and 23-27).]

DETERMINATION OF CALCIUM

PROCEDURE.—To the combined filtrates from the double precipitation of
the hydroxides just described, add 5 cc. of dilute ammonium hydroxide (sp.
gr. 0.96), and transfer the liquid to a 500 cc. graduated flask, washing out
the beaker carefully. Cool to laboratory temperature, and fill the flask with
distilled water until the lowest point of the meniscus is exactly level with
the mark on the neck of the flask. Carefully remove any drops of water
which are on the inside of the neck of the flask above the graduation by
means of a strip of filter paper, make the solution uniform by pouring it out
into a dry beaker and back into the flask several times. Measure off one
fifth of this solution as follows (Note 1): Pour into a 100 cc. graduated flask
about 10 cc. of the solution, shake the liquid thoroughly over the inner
surface of the small flask, and pour it out. Repeat the same operation. Fill
the 100 cc. flask until the lowest point of the meniscus is exactly level with
the mark on its neck, remove any drops of solution from the upper part of
the neck with filter paper, and pour the solution into a beaker (400-500 cc.).
Wash out the flask with small quantities of water until it is clean, adding
these to the 100 cc. of solution. When the duplicate portion of 100 cc. is
measured out from the solution, remember that the flask must be rinsed out
twice with that solution, as prescribed above, before the measurement is
made. (A 100 cc. pipette may be used to measure out the aliquot portions, if
preferred.)

Dilute each of the measured portions to 250 cc. with distilled water, heat
the whole to boiling, and add ammonium oxalate solution slowly in
moderate excess, stirring well. Boil for two minutes; allow the precipitated
calcium oxalate to settle for a half-hour, and decant through a filter. Test the
filtrate for complete precipitation by adding a few cubic centimeters of the
precipitant, allowing it to stand for fifteen minutes. If no precipitate forms,
make the solution slightly acid with hydrochloric acid (Note 2); see that it is
properly labeled and reserve it to be combined with the filtrate from the
second calcium oxalate precipitation (Notes 3 and 4).

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

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

The precipitate may, if preferred, be placed in a weighted porcelain
crucible. After burning off the filter and heating for ten minutes the calcium

precipitate may be converted into calcium sulphate by placing 2 cc. of
dilute sulphuric acid in the crucible (cold), heating the covered crucible
very cautiously over a low flame to drive off the excess of acid, and finally
at redness to constant weight (Note 7).

From the weight of the oxide or sulphate, calculate the percentage of the
calcium (Ca) in the limestone, remembering that only one fifth of the total
solution is used for this determination.

[Note 1: If the calcium were precipitated from the entire solution, the
quantity of the precipitate would be greater than could be properly treated.
The solution is, therefore, diluted to a definite volume (500 cc.), and exactly
one fifth (100 cc.) is measured off in a graduated flask or by means of a
pipette.]

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

[Note 3: The accurate quantitative separation of calcium and magnesium
as oxalates requires considerable care. The calcium precipitate usually
carries down with it some magnesium, and this can best be removed by
redissolving the precipitate after filtration, and reprecipitation in the
presence of only the small amount of magnesium which was included in the
first precipitate. When, however, the proportion of magnesium is not very
large, the second precipitation of the calcium can usually be avoided by
precipitating it from a rather dilute solution (800 cc. or so) and in the
presence of a considerable excess of the precipitant, that is, rather more
than enough to convert both the magnesium and calcium into oxalates.]

[Note 4: The ionic changes involved in the precipitation of calcium as
oxalate are exceedingly simple, and the principles discussed in connection

with the barium sulphate precipitation on page 113 also apply here. The
reaction is

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

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

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

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

[Note 5: The small quantity of ammonium oxalate solution is added
before the second precipitation of the calcium oxalate to insure the presence

of a slight excess of the reagent, which promotes the separation of the
calcium compound.]

[Note 6: On ignition the calcium oxalate loses carbon dioxide and carbon
monoxide, leaving calcium oxide:

CaC_{2}O_{4}.H_{2}O —> CaO + CO_{2} + CO + H_{2}O.

For small weights of the oxalate (0.6 gram or less) this reaction may be
brought about in a platinum crucible at the highest temperature of a Tirrill
burner, but it is well to ignite larger quantities than this over the blast lamp
until the weight is constant.]

[Note 7: The heat required to burn the filter, and that subsequently
applied as described, will convert most of the calcium oxalate to calcium
carbonate, which is changed to sulphate by the sulphuric acid. The reactions
involved are

CaC_{2}O_{4} —> CaCO_{3} + CO,
CaCO_{3} + H_{2}SO_{4} —> CaSO_{4} + H_{2}O + CO_{2}.

If a porcelain crucible is employed for ignition, this conversion to
sulphate is to be preferred, as a complete conversion to oxide is difficult to
accomplish.]

[Note 8: The determination of the calcium may be completed
volumetrically by washing the calcium oxalate precipitate from the filter
into dilute sulphuric acid, warming, and titrating the liberated oxalic acid
with a standard solution of potassium permanganate as described on page
72. When a considerable number of analyses are to be made, this procedure
will save much of the time otherwise required for ignition and weighing.]

DETERMINATION OF MAGNESIUM

PROCEDURE.—Evaporate the acidified filtrates from the calcium
precipitates until the salts begin to crystallize, but do !not! evaporate to
dryness (Note 1). Dilute the solution cautiously until the salts are brought
into solution, adding a little acid if the solution has evaporated to very small
volume. The solution should be carefully examined at this point and must
be filtered if a precipitate has appeared. Heat the clear solution to boiling;
remove the burner and add 25 cc. of a solution of disodium phosphate. Then
add slowly dilute ammonia (1 volume strong ammonia (sp. gr. 0.90) and 9
volumes water) as long as a precipitate continues to form. Finally, add a
volume of concentrated ammonia (sp. gr. 0.90) equal to one third of the
volume of the solution, and allow the whole to stand for about twelve hours.

Decant the solution through a filter, wash it with dilute ammonia water,
proceeding as prescribed for the determination of phosphoric anhydride on
page 122, including; the reprecipitation (Note 2), except that 3 cc. of
disodium phosphate solution are added before the reprecipitation of the
magnesium ammonium phosphate instead of the magnesia mixture there
prescribed. From the weight of the pyrophosphate, calculate the percentage
of magnesium oxide (MgO) in the sample of limestone. Remember that the
pyrophosphate finally obtained is from one fifth of the original sample.

[Note 1: The precipitation of the magnesium should be made in as small
volume as possible, and the ratio of ammonia to the total volume of solution

should be carefully provided for, on account of the relative solubility of the
magnesium ammonium phosphate. This matter has been fully discussed in
connection with the phosphoric anhydride determination.]

[Note 2: The first magnesium ammonium phosphate precipitate is rarely
wholly crystalline, as it should be, and is not always of the proper
composition when precipitated in the presence of such large amounts of
ammonium salts. The difficulty can best be remedied by filtering the
precipitate and (without washing it) redissolving in a small quantity of
hydrochloric acid, from which it may be again thrown down by ammonia
after adding a little disodium phosphate solution. If the flocculent character
was occasioned by the presence of magnesium hydroxide, the second
precipitation, in a smaller volume containing fewer salts, will often result
more favorably.

The removal of iron or alumina from a contaminated precipitate is a
matter involving a long procedure, and a redetermination of the magnesium
from a new sample, with additional precautions, is usually to be preferred.]

DETERMINATION OF CARBON DIOXIDE

!Absorption Apparatus!

[Illustration: Fig. 3]

The apparatus required for the determination of the carbon dioxide
should be arranged as shown in the cut (Fig. 3). The flask (A) is an ordinary
wash bottle, which should be nearly filled with dilute hydrochloric acid
(100 cc. acid (sp. gr. 1.12) and 200 cc. of water). The flask is connected by
rubber tubing (a) with the glass tube (b) leading nearly to the bottom of the
evolution flask (B) and having its lower end bent upward and drawn out to
small bore, so that the carbon dioxide evolved from the limestone cannot
bubble back into (b). The evolution flask should preferably be a wide-
mouthed Soxhlet extraction flask of about 150 cc. capacity because of the
ease with which tubes and stoppers may be fitted into the neck of a flask of
this type. The flask should be fitted with a two-hole rubber stopper. The
condenser (C) may consist of a tube with two or three large bulbs blown in
it, for use as an air-cooled condenser, or it may be a small water-jacketed
condenser. The latter is to be preferred if a number of determinations are to
be made in succession.

A glass delivery tube (c) leads from the condenser to the small U-tube
(D) containing some glass beads or small pieces of glass rod and 3 cc. of a
saturated solution of silver sulphate, with 3 cc. of concentrated sulphuric

acid (sp. gr. 1.84). The short rubber tubing (d) connects the first U-tube to a
second U-tube (E) which is filled with small dust-free lumps of dry calcium
chloride, with a small, loose plug of cotton at the top of each arm. Both
tubes should be closed by cork stoppers, the tops of which are cut off level
with, or preferably forced a little below, the top of the U-tube, and then
neatly sealed with sealing wax.

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

The Geissler bulb (F'), if used, should be filled with potassium hydroxide
solution (1 part of solid potassium hydroxide dissolved in two parts of
water) until each small bulb is about two thirds full (Note 4). A small tube
containing calcium chloride is connected with the Geissler bulb proper by a
ground joint and should be wired to the bulb for safety. This is designed to
retain any moisture from the hydroxide solution. A piece of clean, fine
copper wire is so attached to the bulb that it can be hung from the hook
above a balance pan, or other support.

The small bottle (G) with concentrated sulphuric acid (sp. gr. 1.84) is so
arranged that the tube (f) barely dips below the surface. This will prevent
the absorption of water vapor by (F) or (F') and serves as an aid in
regulating the flow of air through the apparatus. (H) is an aspirator bottle of
about four liters capacity, filled with water; (k) is a safety tube and a means

of refilling (H); (h) is a screw clamp, and (K) a U-tube filled with soda
lime.

[Note 1: The air current, which is subsequently drawn through the
apparatus, to sweep all of the carbon dioxide into the absorption apparatus,
is likely to carry with it some hydrochloric acid from the evolution flask.
This acid is retained by the silver sulphate solution. The addition of
concentrated sulphuric acid to this solution reduces its vapor pressure so far
that very little water is carried on by the air current, and this slight amount
is absorbed by the calcium chloride in (E). As the calcium chloride
frequently contains a small amount of a basic material which would absorb
carbon dioxide, it is necessary to pass carbon dioxide through (E) for a
short time and then drive all the gas out with a dry air current for thirty
minutes before use.]

[Note 2: Soda-lime absorption tubes are to be preferred if a satisfactory
quality of soda lime is available and the number of determinations to be
made successively is small. The potash bulbs will usually permit of a larger
number of successive determinations without refilling, but they require
greater care in handling and in the analytical procedure.]

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

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

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

After the potassium hydroxide is approximately half consumed in the
first bulb of the absorption apparatus, potassium bicarbonate is formed, and
as it is much less soluble than the carbonate, it often precipitates. Its
formation is a warning that the absorbing power of the hydroxide is much
diminished.]

!The Analysis!

PROCEDURE.— Weigh out into the flask (B) about 1 gram of
limestone. Cover it with 15 cc. of water. Weigh the absorption apparatus (F)
or (F') accurately after allowing it to stand for 30 minutes in the balance
case, and wiping it carefully with a lintless cloth, taking care to handle it as
little as possible after wiping (Note 1). Connect the absorption apparatus
with (e) and (f). If a soda-lime tube is used, be sure that the arm containing
the soda lime is next the tube (E) and that the glass stopcocks are open.

To be sure that the whole apparatus is airtight, disconnect the rubber tube
from the flask (A), making sure that the tubes (a) and (b) do not contain any
hydrochloric acid, close the pinchcocks (a) and (k) and open (h). No
bubbles should pass through (D) or (G) after a few seconds. When assured
that the fittings are tight, close (h) and open (a) cautiously to admit air to
restore atmospheric pressure. This precaution is essential, as a sudden
inrush of air will project liquid from (D) or (F'). Reconnect the rubber tube

with the flask (A). Open the pinchcocks (a) and (k) and blow over about 10
cc. of the hydrochloric acid from (A) into (B). When the action of the acid
slackens, blow over (slowly) another 10 cc.

The rate of gas evolution should not exceed for more than a few seconds
that at which about two bubbles per second pass through (G) (Note 2).
Repeat the addition of acid in small portions until the action upon the
limestone seems to be at an end, taking care to close (a) after each addition
of acid (Note 3). Disconnect (A) and connect the rubber tubing with the
soda-lime tube (K) and open (a). Then close (k) and open (h), regulating the
flow of water from (H) in such a way that about two bubbles per second
pass through (G). Place a small flame under (B) and !slowly! raise the
contents to boiling and boil for three minutes. Then remove the burner from
under (B) and continue to draw air through the apparatus for 20-30 minutes,
or until (H) is emptied (Note 4). Remove the absorption apparatus, closing
the stopcocks on (F) or stoppering the open ends of (F'), leave the apparatus
in the balance case for at least thirty minutes, wipe it carefully and weigh,
after opening the stopcocks (or removing plugs). The increase in weight is
due to absorption of CO_{2}, from which its percentage in the sample may
be calculated.

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

[Note 1: The absorption tubes or bulbs have large surfaces on which
moisture may collect. By allowing them to remain in the balance case for
some time before weighing, the amount of moisture absorbed on the surface
is as nearly constant as practicable during two weighings, and a uniform
temperature is also assured. The stopcocks of the U-tube should be opened,
or the plugs used to close the openings of the Geissler bulb should be

removed before weighing in order that the air contents shall always be at
atmospheric pressure.]

[Note 2: If the gas passes too rapidly into the absorption apparatus, some
carbon dioxide may be carried through, not being completely retained by
the absorbents.]

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

[Note 4: Carbon dioxide is dissolved by cold water, but the gas is
expelled by boiling, and, together with that which is distributed through the
apparatus, is swept out into the absorption bulb by the current of air. This
air is purified by drawing it through the tube (K) containing soda lime,
which removes any carbon dioxide which may be in it.]

DETERMINATION OF LEAD, COPPER, IRON,
AND ZINC IN BRASS

ELECTROLYTIC SEPARATIONS

!General Discussion!

When a direct current of electricity passes from one electrode to another
through solutions of electrolytes, the individual ions present in these
solutions tend to move toward the electrode of opposite electrical charge to
that which each ion bears, and to be discharged by that electrode. Whether
or not such discharge actually occurs in the case of any particular ion
depends upon the potential (voltage) of the current which is passing through
the solution, since for each ion there is, under definite conditions, a
minimum potential below which the discharge of the ion cannot be effected.
By taking advantage of differences in discharge-potentials, it is possible to
effect separations of a number of the metallic ions by electrolysis, and at the
same time to deposit the metals in forms which admit of direct weighing. In
this way the slower procedures of precipitation and filtration may
frequently be avoided. The following paragraphs present a brief statement
of the fundamental principles and conditions underlying electro-analysis.

The total energy of an electric current as it passes through a solution is
distributed among three factors, first, its potential, which is measured in

volts, and corresponds to what is called "head" in a stream of water; second,
current strength, which is measured in amperes, and corresponds to the
volume of water passing a cross-section of a stream in a given time interval;
and third, the resistance of the conducting medium, which is measured in
ohms. The relation between these three factors is expressed by Ohm's law,
namely, that !I = E/R!, when I is current strength, E potential, and R
resistance. It is plain that, for a constant resistance, the strength of the
current and its potential are mutually and directly interdependent.

As already stated, the applied electrical potential determines whether or
not deposition of a metal upon an electrode actually occurs. The current
strength determines the rate of deposition and the physical characteristics of
the deposit. The resistance of the solution is generally so small as to fall out
of practical consideration.

Approximate deposition-potentials have been determined for a number of
the metallic elements, and also for hydrogen and some of the acid-forming
radicals. The values given below are those required for deposition from
normal solutions at ordinary temperatures with reference to a hydrogen
electrode. They must be regarded as approximate, since several disturbing
factors and some secondary reactions render difficult their exact application
under the conditions of analysis. They are:

Zn Cd Fe Ni Pb H Cu Sb Hg Ag SO_{4} +0.77 +0.42 +0.34 +0.33
+0.13 0 -0.34 -0.67 -0.76 -0.79 +1.90

From these data it is evident that in order to deposit copper from a
normal solution of copper sulphate a minimum potential equal to the
algebraic sum of the deposition-potentials of copper ions and sulphate ions
must be applied, that is, +1.56 volts. The deposition of zinc from a solution
of zinc sulphate would require +2.67 volts, but, since the deposition of
hydrogen from sulphuric acid solution requires only +1.90 volts, the

quantitative deposition of zinc by electrolysis from a sulphuric acid solution
of a zinc salt is not practicable. On the other hand, silver, if present in a
solution of copper sulphate, would deposit with the copper.

The foregoing examples suffice to illustrate the application of the
principle of deposition potentials, but it must further be noted that the
values stated apply to normal solutions of the compounds in question, that
is, to solutions of considerable concentrations. As the concentration of the
ions diminishes, and hence fewer ions approach the electrodes, somewhat
higher voltages are required to attract and discharge them. From this it
follows that the concentrations should be kept as high as possible to effect
complete deposition in the least practicable time, or else the potentials
applied must be progressively increased as deposition proceeds. In practice,
the desired result is obtained by starting with small volumes of solution,
using as large an electrode surface as possible, and by stirring the solution
to bring the ions in contact with the electrodes. This is, in general, a more
convenient procedure than that of increasing the potential of the current
during electrolysis, although that method is also used.

As already stated, those ions in a solution of electrolytes will first be
discharged which have the lowest deposition potentials, and so long as these
ions are present around the electrode in considerable concentration they,
almost alone, are discharged, but, as their concentration diminishes, other
ions whose deposition potentials are higher but still within that of the
current applied, will also begin to separate. For example, from a nitric acid
solution of copper nitrate, the copper ions will first be discharged at the
cathode, but as they diminish in concentration hydrogen ions from the acid
(or water) will be also discharged. Since the hydrogen thus liberated is a
reducing agent, the nitric acid in the solution is slowly reduced to ammonia,
and it may happen that if the current is passed through for a long time, such
a solution will become alkaline. Oxygen is liberated at the anode, but, since

there is no oxidizable substance present around that electrode, it escapes as
oxygen gas. It should be noted that, in general, the changes occurring at the
cathode are reductions, while those at the anode are oxidations.

For analytical purposes, solutions of nitrates or sulphates of the metals
are preferable to those of the chlorides, since liberated chlorine attacks the
electrodes. In some cases, as for example, that of silver, solution of salts
forming complex ions, like that of the double cyanide of silver and
potassium, yield better metallic deposits.

Most metals are deposited as such upon the cathode; a few, notably lead
and manganese, separate in the form of dioxides upon the anode. It is
evidently important that the deposited material should be so firmly adherent
that it can be washed, dried, and weighed without loss in handling. To
secure these conditions it is essential that the current density (that is, the
amount of current per unit of area of the electrodes) shall not be too high. In
prescribing analytical conditions it is customary to state the current strength
in "normal densities" expressed in amperes per 100 sq. cm. of electrode
surface, as, for example, "N.D_{100} = 2 amps."

If deposition occurs too rapidly, the deposit is likely to be spongy or
loosely adherent and falls off on subsequent treatment. This places a
practical limit to the current density to be employed, for a given electrode
surface. The cause of the unsatisfactory character of the deposit is
apparently sometimes to be found in the coincident liberation of
considerable hydrogen and sometimes in the failure of the rapidly deposited
material to form a continuous adherent surface. The effect of rotating
electrodes upon the character of the deposit is referred to below.

The negative ions of an electrolyte are attracted to the anode and are
discharged on contact with it. Anions such as the chloride ion yield chlorine
atoms, from which gaseous chlorine molecules are formed and escape. The

radicals which compose such ions as NO_{3}^{-} or SO_{4}^{—} are not
capable of independent existence after discharge, and break down into
oxygen and N_{2}O_{5} and SO_{3} respectively. The oxygen escapes
and the anhydrides, reacting with water, re-form nitric and sulphuric acids.

The law of Faraday expresses the relation between current strength and
the quantities of the decomposition products which, under constant
conditions, appear at the electrodes, namely, that a given quantity of
electricity, acting for a given time, causes the separation of chemically
equivalent quantities of the various elements or radicals. For example, since
107.94 grams of silver is equivalent to 1.008 grams of hydrogen, and that in
turn to 8 grams of oxygen, or 31.78 grams of copper, the quantity of
electricity which will cause the deposit of 107.94 grams of silver in a given
time will also separate the weights just indicated of the other substances.
Experiments show that a current of one ampere passing for one second, i.e.,
a coulomb of electricity, causes the deposition of 0.001118 gram of silver
from a normal solution of a silver salt. The number of coulombs required to
deposit 107.94 grams is 107.94/0.001118 or 96,550 and the same number of
coulombs will also cause the separation of 1.008 grams of hydrogen, 8
grams of oxygen or 31.78 grams of copper. While it might at first appear
that Faraday's law could thus be used as a basis for the calculation of the
time required for the deposition of a given quantity of an electrolyte from
solution, it must be remembered that the law expresses what occurs when
the concentration of the ions in the solution is kept constant, as, for
example, when the anode in a silver salt solution is a plate of metallic silver.
Under the conditions of electro-analysis the concentration of the ions is
constantly diminishing as deposition proceeds and the time actually
required for complete deposition of a given weight of material by a current
of constant strength is, therefore, greater than that calculated on the basis of
the law as stated above.

The electrodes employed in electro-analysis are almost exclusively of
platinum, since that metal alone satisfactorily resists chemical action of the
electrolytes, and can be dried and weighed without change in composition.
The platinum electrodes may be used in the form of dishes, foil or gauze.
The last, on account of the ease of circulation of the electrolyte, its
relatively large surface in proportion to its weight and the readiness with
which it can be washed and dried, is generally preferred.

Many devices have been described by the use of which the electrode
upon which deposition occurs can be mechanically rotated. This has an
effect parallel to that of greatly increasing the electrode surface and also
provides a most efficient means of stirring the solution. With such an
apparatus the amperage may be increased to 5 or even 10 amperes and
depositions completed with great rapidity and accuracy. It is desirable,
whenever practicable, to provide a rotating or stirring device, since, for
example, the time consumed in the deposition of the amount of copper
usually found in analysis may be reduced from the 20 to 24 hours required
with stationary electrodes, and unstirred solutions, to about one half hour.

DETERMINATION OF COPPER AND LEAD

PROCEDURE.—Weigh out two portions of about 0.5 gram each (Note 1)
into tall, slender lipless beakers of about 100 cc. capacity. Dissolve the
metal in a solution of 5 cc. of dilute nitric acid (sp. gr. 1.20) and 5 cc. of
water, heating gently, and keeping the beaker covered. When the sample has
all dissolved (Note 2), wash down the sides of the beaker and the bottom of
the watch-glass with water and dilute the solution to about 50 cc. Carefully
heat to boiling and boil for a minute or two to expel nitrous fumes.

Meanwhile, four platinum electrodes, two anodes and two cathodes,
should be cleaned by dipping in dilute nitric acid, washing with water and
finally with 95 per cent alcohol (Note 3). The alcohol may be ignited and
burned off. The electrodes are then cooled in a desiccator and weighed.
Connect the electrodes with the binding posts (or other device for
connection with the electric circuit) in such a way that the copper will be
deposited upon the electrode with the larger surface, which is made the
cathode. The beaker containing the solution should then be raised into place
from below the electrodes until the latter reach nearly to the bottom of the
beaker. The support for the beaker must be so arranged that it can be easily
raised or lowered.

If the electrolytic apparatus is provided with a mechanism for the rotation
of the electrode or stirring of the electrolyte, proceed as follows: Arrange
the resistance in the circuit to provide a direct current of about one ampere.

Pass this current through the solution to be electrolyzed, and start the
rotating mechanism. Keep the beaker covered as completely as possible,
using a split watch-glass (or other device) to avoid loss by spattering. When
the solution is colorless, which is usually the case after about 35 minutes,
rinse off the cover glass, wash down the sides of the beaker, add about 0.30
gram of urea and continue the electrolysis for another five minutes (Notes 4
and 5).

If stationary electrodes are employed, the current strength should be
about 0.1 ampere, which may, after 12 to 15 hours, be increased to 0.2
ampere. The time required for complete deposition is usually from 20 to 24
hours. It is advisable to add 5 cc. of nitric acid (sp. gr. 1.2) if the electrolysis
extends over this length of time. No urea is added in this case.

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

Test the solution in the beaker for copper as follows, remembering that it
is to be used for subsequent determinations of iron and zinc: Remove about
5 cc. and add a slight excess of ammonia. Compare the mixture with some
distilled water, holding both above a white surface. The solution should not
show any tinge of blue. If the presence of copper is indicated, add the test

portion to the main solution, evaporate the whole to a volume of about 100
cc., and again electrolyze with clean electrodes (Note 7).

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

[Note 1: It is obvious that the brass taken for analysis should be
untarnished, which can be easily assured, when wire is used, by scouring
with emery. If chips or borings are used, they should be well mixed, and the
sample for analysis taken from different parts of the mixture.]

[Note 2: If a white residue remains upon treatment of the alloy with nitric
acid, it indicates the presence of tin. The material is not, therefore, a true
brass. This may be treated as follows: Evaporate the solution to dryness,
moisten the residue with 5 cc. of dilute nitric acid (sp. gr. 1.2) and add 50
cc. of hot water. Filter off the meta-stannic acid, wash, ignite in porcelain
and weigh as SnO_{2}. This oxide is never wholly free from copper and
must be purified for an exact determination. If it does not exceed 2 per cent
of the alloy, the quantity of copper which it contains may usually be
neglected.]

[Note 3: The electrodes should be freed from all greasy matter before
using, and those portions upon which the metal will deposit should not be
touched with the fingers after cleaning.]

[Note 4: Of the ions in solution, the H^{+}, Cu^{++}, Zn^{++}, and
Fe^{+++} ions tend to move toward the cathode. The NO_{3}^{-} ions and
the lead, probably in the form of PbO_{2}^{—} ions, move toward the
anode. At the cathode the Cu^{++} ions are discharged and plate out as
metallic copper. This alone occurs while the solution is relatively
concentrated. Later on, H^{+} ions are also discharged. In the presence of

considerable quantities of H^{+} ions, as in this acid solution, no Zn^{++}
or Fe^{+++} ions are discharged because of their greater deposition
potentials. At the anode the lead is deposited as PbO_{2} and oxygen is
evolved.

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

[Note 5: Urea reacts with nitrous acid, which may be formed in the
solution as a result of the reducing action of the liberated hydrogen. Its
removal promotes the complete precipitation of the copper. The reaction is

CO(NH_{2}){2} + 2HNO{2} —> CO_{2} + 2N_{2} + 3H_{2}O.]

[Note 6: The electrodes must be washed nearly or quite free from the
nitric acid solution before the circuit is broken to prevent re-solution of the
copper.

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

[Note 7: The electrodes upon which the copper has been deposited may
be cleaned by immersion in warm nitric acid. To remove the lead dioxide,
add a few crystals of oxalic acid to the nitric acid.]

DETERMINATION OF IRON

Most brasses contain small percentages of iron (usually not over 0.1 per
cent) which, unless removed, is precipitated as phosphate and weighed with
the zinc.

PROCEDURE.—To the solution from the precipitation of copper and
lead by electrolysis, add dilute ammonia (sp. gr. 0.96) until the precipitate
of zinc hydroxide which first forms re-dissolves, leaving only a slight red
precipitate of ferric hydroxide. Filter off the iron precipitate, using a washed
filter, and wash five times with hot water. Test a portion of the last washing
with a dilute solution of ammonium sulphide to assure complete removal of
the zinc.

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

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

DETERMINATION OF ZINC

PROCEDURE.—Acidify the filtrate from the iron determination with dilute
nitric acid. Concentrate it to 150 cc. Add to the cold solution dilute
ammonia (sp. gr. 0.96) cautiously until it barely smells of ammonia; then
add !one drop! of a dilute solution of litmus (Note 1), and drop in, with the
aid of a dropper, dilute nitric acid until the blue of the litmus just changes to
red. It is important that this point should not be overstepped. Heat the
solution nearly to boiling and pour into it slowly a filtered solution of di-
ammonium hydrogen phosphate[1] containing a weight of the phosphate
about equal to twelve times that of the zinc to be precipitated. (For this
calculation the approximate percentage of zinc is that found by subtracting
the sum of the percentages of the copper, lead and iron from 100 per cent.)
Keep the solution just below boiling for fifteen minutes, stirring frequently
(Note 2). If at the end of this time the amorphous precipitate has become
crystalline, allow the solution to cool for about four hours, although a
longer time does no harm (Note 3), and filter upon an asbestos filter in a
porcelain Gooch crucible. The filter is prepared as described on page 103,
and should be dried to constant weight at 105°C.

[Footnote 1: The ammonium phosphate which is commonly obtainable
contains some mono-ammonium salt, and this is not satisfactory as a
precipitant. It is advisable, therefore, to weigh out the amount of the salt
required, dissolve it in a small volume of water, add a drop of

phenolphthalein solution, and finally add dilute ammonium hydroxide
solution cautiously until the solution just becomes pink, but do not add an
excess.]

Wash the precipitate until free from sulphates with a warm 1 per cent
solution of the di-ammonium phosphate, and then five times with 50 per
cent alcohol (Note 4). Dry the crucible and precipitate for an hour at
105°C., and finally to constant weight (Note 5). The filtrate should be made
alkaline with ammonia and tested for zinc with a few drops of ammonium
sulphide, allowing it to stand (Notes 6, 7 and 8).

From the weight of the zinc ammonium phosphate (ZnNH_{4}PO_{4})
calculate the percentage of the zinc (Zn) in the brass.

[Note 1: The zinc ammonium phosphate is soluble both in acids and in
ammonia. It is, therefore, necessary to precipitate the zinc in a nearly
neutral solution, which is more accurately obtained by adding a drop of a
litmus solution to the liquid than by the use of litmus paper.]

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

[Note 3: In a carefully neutralized solution containing a considerable
excess of the precipitant, and also ammonium salts, the separation of the
zinc is complete after standing four hours. The ionic changes connected
with the precipitation of the zinc as zinc ammonium phosphate are similar
to those described for magnesium ammonium phosphate, except that the
zinc precipitate is soluble in an excess of ammonium hydroxide, probably

as a result of the formation of complex ions of the general character
Zn(NH_{3})_{4}^{++}.]

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

[Note 5: The precipitate may be ignited and weighed as
Zn_{2}P_{2}O_{7}, by cautiously heating the porcelain Gooch crucible
within a nickel or iron crucible, used as a radiator. The heating must be very
slow at first, as the escaping ammonia may reduce the precipitate if it is
heated too quickly.]

[Note 6: If the ammonium sulphide produced a distinct precipitate, this
should be collected on a small filter, dissolved in a few cubic centimeters of
dilute nitric acid, and the zinc reprecipitated as phosphate, filtered off,
dried, and weighed, and the weight added to that of the main precipitate.]

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

[Note 8. The details of this method of precipitation of zinc are fully
discussed in an article by Dakin, !Ztschr. Anal. Chem.!, 39 (1900), 273.]

DETERMINATION OF SILICA IN SILICATES

Of the natural silicates, or artificial silicates such as slags and some of the
cements, a comparatively few can be completely decomposed by treatment
with acids, but by far the larger number require fusion with an alkaline flux
to effect decomposition and solution for analysis. The procedure given
below applies to silicates undecomposable by acids, of which the mineral
feldspar is taken as a typical example. Modifications of the procedure,
which are applicable to silicates which are completely or partially
decomposable by acids, are given in the Notes on page 155.

PREPARATION OF THE SAMPLE

Grind about 3 grams of the mineral in an agate mortar (Note 1) until no
grittiness is to be detected, or, better, until it will entirely pass through a
sieve made of fine silk bolting cloth. The sieve may be made by placing a
piece of the bolting cloth over the top of a small beaker in which the ground
mineral is placed, holding the cloth in place by means of a rubber band
below the lip of the beaker. By inverting the beaker over clean paper and
gently tapping it, the fine particles pass through the sieve, leaving the
coarser particles within the beaker. These must be returned to the mortar
and ground, and the process of sifting and grinding repeated until the entire
sample passes through the sieve.

[Note 1: If the sample of feldspar for analysis is in the massive or
crystalline form, it should be crushed in an iron mortar until the pieces are
about half the size of a pea, and then transferred to a steel mortar, in which
they are reduced to a coarse powder. A wooden mallet should always be
used to strike the pestle of the steel mortar, and the blows should not be
sharp.

It is plain that final grinding in an agate mortar must be continued until
the whole of the portion of the mineral originally taken has been ground so
that it will pass the bolting cloth, otherwise the sifted portion does not
represent an average sample, the softer ingredients, if foreign matter is
present, being first reduced to powder. For this reason it is best to start with
not more than the quantity of the feldspar needed for analysis. The mineral
must be thoroughly mixed after the grinding.]

FUSION AND SOLUTION

PROCEDURE.—Weigh into platinum crucibles two portions of the ground
feldspar of about 0.8 gram each. Weigh on rough balances two portions of
anhydrous sodium carbonate, each amounting to about six times the weight
of the feldspar taken for analysis (Note 1). Pour about three fourths of the
sodium carbonate into the crucible, place the latter on a piece of clean,
glazed paper, and thoroughly mix the substance and the flux by carefully
stirring for several minutes with a dry glass rod, the end of which has been
recently heated and rounded in a flame and slowly cooled. The rod may be
wiped off with a small fragment of filter paper, which may be placed in the
crucible. Place the remaining fourth of the carbonate on the top of the
mixture. Cover the crucible, heat it to dull redness for five minutes, and
then gradually increase the heat to the full capacity of a Bunsen or Tirrill
burner for twenty minutes, or until a quiet, liquid fusion is obtained (Note

2). Finally, heat the sides and cover strongly until any material which may
have collected upon them is also brought to fusion.

Allow the crucible to cool, and remove the fused mass as directed on
page 116. Disintegrate the mass by placing it in a previously prepared
mixture of 100 cc. of water and 50 cc. of dilute hydrochloric acid (sp. gr.
1.12) in a covered casserole (Note 3). Clean the crucible and lid by means
of a little hydrochloric acid, adding this acid to the main solution (Notes 4
and 5).

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

[Note 2: During the fusion the feldspar, which, when pure, is a silicate of
aluminium and either sodium or potassium, but usually contains some iron,
calcium, and magnesium, is decomposed by the alkaline flux. The sodium
of the latter combines with the silicic acid of the silicate, with the evolution
of carbon dioxide, while about two thirds of the aluminium forms sodium
aluminate and the remainder is converted into basic carbonate, or the oxide.
The calcium and magnesium, if present, are changed to carbonates or
oxides.

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

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

Whether the resulting solution of the silicate contains ortho-silicic acid
(H_{4}SiO_{4}) or whether it is a colloidal solution of some other less
hydrated acid, such as meta-silicic acid (H_{2}SiO_{3}), is a matter that is
still debatable. It is certain, however, that the gelatinous material which
readily separates from such solutions is of the nature of a hydrogel, that is, a
colloid which is insoluble in water. This substance when heated to 100°C.,
or higher, is completely dehydrated, leaving only the anhydride, SiO_{2}.
The changes may be represented by the equation:

SiO_{3}^{—} + 2H^{+} —> [H_{2}SiO_{3}] —> H_{2}O +
SiO_{2}.]

[Note 4: A portion of the fused mass is usually projected upward by the
escaping carbon dioxide during the fusion. The crucible must therefore be
kept covered as much as possible and the lid carefully cleaned.]

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

DEHYDRATION AND FILTRATION

PROCEDURE.—Evaporate the solution of the fusion to dryness, stirring
frequently until the residue is a dry powder. Moisten the residue with 5 cc.
of strong hydrochloric acid (sp. gr. 1.20) and evaporate again to dryness.
Heat the residue for at least one hour at a temperature of 110°C. (Note 1).
Again moisten the residue with concentrated hydrochloric acid, warm
gently, making sure that the acid comes into contact with the whole of the

residue, dilute to about 200 cc. and bring to boiling. Filter off the silica
without much delay (Note 2), and wash five times with warm dilute
hydrochloric acid (one part dilute acid (1.12 sp. gr.) to three parts of water).
Allow the filter to drain for a few moments, then place a clean beaker below
the funnel and wash with water until free from chlorides, discarding these
washings. Evaporate the original filtrate to dryness, dehydrate at 110°C. for
one hour (Note 3), and proceed as before, using a second filter to collect the
silica after the second dehydration. Wash this filter with warm, dilute
hydrochloric acid (Note 4), and finally with hot water until free from
chlorides.

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

A flocculent residue will often remain after the decomposition of the
fused mass is effected. This is usually partially dehydrated silicic acid and
does not require further treatment at this point. The progress of the
dehydration is indicated by the behavior of the solution, which as
evaporation proceeds usually gelatinizes. On this account it is necessary to
allow the solution to evaporate on a steam bath, or to stir it vigorously, to
avoid loss by spattering.]

[Note 2: To obtain an approximately pure silica, the residue after
evaporation must be thoroughly extracted by warming with hydrochloric
acid, and the solution freely diluted to prevent, as far as possible, the
inclosure of the residual salts in the particles of silica. The filtration should

take place without delay, as the dehydrated silica slowly dissolves in
hydrochloric acid on standing.]

[Note 3: It has been shown by Hillebrand that silicic acid cannot be
completely dehydrated by a single evaporation and heating, nor by several
such treatments, unless an intermediate filtration of the silica occurs. If,
however, the silica is removed and the filtrates are again evaporated and the
residue heated, the amount of silica remaining in solution is usually
negligible, although several evaporations and filtrations are required with
some silicates to insure absolute accuracy.

It is probable that temperatures above 100°C. are not absolutely
necessary to dehydrate the silica; but it is recommended, as tending to leave
the silica in a better condition for filtration than when the lower temperature
of the water bath is used. This, and many other points in the analysis of
silicates, are fully discussed by Dr. Hillebrand in the admirable monograph
on "The Analysis of Silicate and Carbonate Rocks," Bulletin No. 700 of the
United States Geological Survey.

The double evaporation and filtration spoken of above are essential
because of the relatively large amount of alkali salts (sodium chloride)
present after evaporation. For the highest accuracy in the determination of
silica, or of iron and alumina, it is also necessary to examine for silica the
precipitate produced in the filtrate by ammonium hydroxide by fusing it
with acid potassium sulphate and solution of the fused mass in water. The
insoluble silica is filtered, washed, and weighed, and the weight added to
the weight of silica previously obtained.]

[Note 4: Aluminium and iron are likely to be thrown down as basic salts
from hot, very dilute solutions of their chlorides, as a result of hydrolysis. If
the silica were washed only with hot water, the solution of these chlorides
remaining in the filter after the passage of the original filtrate would

gradually become so dilute as to throw down basic salts within the pores of
the filter, which would remain with the silica. To avoid this, an acid wash-
water is used until the aluminium and iron are practically removed. The
acid is then removed by water.]

IGNITION AND TESTING OF SILICA

PROCEDURE.—Transfer the two washed filters belonging to each
determination to a platinum crucible, which need not be previously
weighed, and burn off the filter (Note 1). Ignite for thirty minutes over the
blast lamp with the cover on the crucible, and then for periods of ten
minutes, until the weight is constant.

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

If the precipitate has dissolved in this quantity of acid, add two drops of
concentrated sulphuric acid, and heat very slowly (always under the hood)
until all the liquid has evaporated, finally igniting to redness. Cool in a
desiccator, and weigh the crucible and residue. Deduct this weight from the
previous weight of crucible and impure silica, and from the difference
calculate the percentage of silica in the sample (Note 3).

[Note 1: The silica undergoes no change during the ignition beyond the
removal of all traces of water; but Hillebrand (!loc. cit.!) has shown that the
silica holds moisture so tenaciously that prolonged ignition over the blast
lamp is necessary to remove it entirely. This finely divided, ignited silica
tends to absorb moisture, and should be weighed quickly.]

[Note 2: Notwithstanding all precautions, the ignited precipitate of silica
is rarely wholly pure. It is tested by volatilisation of the silica as silicon
fluoride after solution in hydrofluoric acid, and, if the analysis has been
properly conducted, the residue, after treatment with the acids and ignition,
should not exceed 1 mg.

The acid produces ulceration if brought into contact with the skin, and its
fumes are excessively harmful if inhaled.]

[Note 3: The impurities are probably weighed with the original
precipitate in the form of oxides. The addition of the sulphuric acid
displaces the hydrofluoric acid, and it may be assumed that the resulting
sulphates (usually of iron or aluminium) are converted to oxides by the final
ignition.

It is obvious that unless the sulphuric and hydrofluoric acids used are
known to leave no residue on evaporation, a quantity equal to that
employed in the analysis must be evaporated and a correction applied for
any residue found.]

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

A silicate which gelatinizes on treatment with acids should be mixed first
with a little water, and the strong acid added in small portions with stirring,
otherwise the gelatinous silicic acid incloses particles of the original silicate
and prevents decomposition. The water, by separating the particles and
slightly lessening the rapidity of action, prevents this difficulty. This

procedure is one which applies in general to the solution of fine mineral
powders in acids.

If a small residue remains undecomposed by the treatment of the silicate
with acid, this may be filtered, washed, ignited and fused with sodium
carbonate and a solution of the fused mass added to the original acid
solution. This double procedure has an advantage, in that it avoids adding
so large a quantity of sodium salts as is required for disintegration of the
whole of the silicate by the fusion method.]

PART IV

STOICHIOMETRY

The problems with which the analytical chemist has to deal are not, as a
matter of actual fact, difficult either to solve or to understand. That they
appear difficult to many students is due to the fact that, instead of
understanding the principles which underlie each of the small number of
types into which these problems may be grouped, each problem is
approached as an individual puzzle, unrelated to others already solved or
explained. This attitude of mind should be carefully avoided.

It is obvious that ability to make the calculations necessary for the
interpretation of analytical data is no less important than the manipulative
skill required to obtain them, and that a moderate time spent in the careful
study of the solutions of the typical problems which follow may save much
later embarrassment.

1. It is often necessary to calculate what is known as a "chemical factor,"
or its equivalent logarithmic value called a "log factor," for the conversion
of the weight of a given chemical substance into an equivalent weight of
another substance. This is, in reality, a very simple problem in proportion,
making use of the atomic or molecular weights of the substances in
question which are chemically equivalent to each other. One of the simplest
cases of this sort is the following: What is the factor for the conversion of a

given weight of barium sulphate (BaSO_{4}) into an equivalent weight of
sulphur (S)? The molecular weight of BaSO_{4} is 233.5. There is one
atom of S in the molecule and the atomic weight of S is 32.1. The chemical
factor is, therefore, 32.1/233.5, or 0.1375 and the weight of S corresponding
to a given weight of BaSO_{4} is found by multiplying the weight of
BaSO_{4} by this factor. If the problem takes the form, "What is the factor
for the conversion of a given weight of ferric oxide (Fe_{2}O_{3}) into
ferrous oxide (FeO), or of a given weight of mangano-manganic oxide
(Mn_{3}O_{4}) into manganese (Mn)?" the principle involved is the same,
but it must then be noted that, in the first instance, each molecule of
Fe_{2}O_{3} will be equivalent to two molecules of FeO, and in the
second instance that each molecule of Mn_{3}O_{4} is equivalent to three
atoms of Mn. The respective factors then become

(2FeO/Fe_{2}O_{3}) or (143.6/159.6) and (3Mn/Mn_{3}O_{4}) or
(164.7/228.7).

It is obvious that the arithmetical processes involved in this type of
problem are extremely simple. It is only necessary to observe carefully the
chemical equivalents. It is plainly incorrect to express the ratio of ferrous to
ferric oxide as (FeO/Fe_{2}O_{3}), since each molecule of the ferric oxide
will yield two molecules of the ferrous oxide. Mistakes of this sort are
easily made and constitute one of the most frequent sources of error.

2. A type of problem which is slightly more complicated in appearance,
but exactly comparable in principle, is the following: "What is the factor for
the conversion of a given weight of ferrous sulphate (FeSO_{4}), used as a
reducing agent against potassium permanganate, into the equivalent weight
of sodium oxalate (Na_{2}C_{2}O_{4})?" To determine the chemical
equivalents in such an instance it is necessary to inspect the chemical
reactions involved. These are:

10FeSO_{4} + 2KMnO_{4} + 8H_{2}SO_{4} —> 5Fe_{2}(SO_{4})
{3} + K{2}SO_{4} + 2MnSO_{4} + 8H_{2}O,

5Na_{2}C_{2}O_{4} + 2KMnO_{4} + 8H_{2}SO_{4} —>
5Na_{2}SO_{4} + 10CO_{2} + K_{2}SO_{4} + 2MnSO_{4} +
8H_{2}O.

It is evident that 10FeSO_{4} in the one case, and 5Na_{2}C_{2}O_{4}
in the other, each react with 2KMnO_{4}. These molecular quantities are
therefore equivalent, and the factor becomes
(10FeSO_{4}/5Na_{2}C_{2}O_{4}) or
(2FeSO_{4}/Na_{2}C_{2}O_{4}) or (303.8/134).

Again, let it be assumed that it is desired to determine the factor required
for the conversion of a given weight of potassium permanganate
(KMnO_{4}) into an equivalent weight of potassium bichromate
(K_{2}Cr_{2}O_{7}), each acting as an oxidizing agent against ferrous
sulphate. The reactions involved are:

10FeSO_{4} + 2KMnO_{4} + 8H_{2}SO_{4} —> 5Fe_{2}(SO_{4})
{3} + K{2}SO_{4} + 2MnSO_{4} + 8H_{2}O,

6FeSO_{4} + K_{2}Cr_{2}O_{7} + 7H_{2}SO_{4} —> 3Fe_{2}
(SO_{3}){3} + K{2}SO_{4} + Cr_{2}(SO_{4}){3} + 7H{2}O.

An inspection of these equations shows that 2KMO_{4} react with
10FeSO_{4}, while K_{2}Cr_{2}O_{7} reacts with 6FeSO_{4}. These are
not equivalent, but if the first equation is multiplied by 3 and the second by
5 the number of molecules of FeSO_{4} is then the same in both, and the
number of molecules of KMnO_{4} and K_{2}Cr_{2}O_{7} reacting with
these 30 molecules become 6 and 5 respectively. These are obviously

chemically equivalent and the desired factor is expressed by the fraction
(6KMnO_{4}/5K_{2}Cr_{2}O_{7}) or (948.0/1471.0).

3. It is sometimes necessary to calculate the value of solutions according
to the principles just explained, when several successive reactions are
involved. Such problems may be solved by a series of proportions, but it is
usually possible to eliminate the common factors and solve but a single one.
For example, the amount of MnO_{2} in a sample of the mineral pyrolusite
may be determined by dissolving the mineral in hydrochloric acid,
absorbing the evolved chlorine in a solution of potassium iodide, and
measuring the liberated iodine by titration with a standard solution of
sodium thiosulphate. The reactions involved are:

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

Assuming that the weight of thiosulphate corresponding to the volume of
sodium thiosulphate solution used is known, what is the corresponding
weight of manganese dioxide? From the reactions given above, the
following proportions may be stated:

2Na_{2}S_{2}O_{3}:I_{2} = 316.4:253.9,

I_{2}:Cl_{2} = 253.9:71,

Cl_{2}:MnO_{2} = 71:86.9.

After canceling the common factors, there remains
2Na_{2}S_{2}O_{3}:MnO_{2} = 316.4:86.9, and the factor for the
conversion of thiosulphate into an equivalent of manganese dioxide is
86.9/316.4.

4. To calculate the volume of a reagent required for a specific operation,
it is necessary to know the exact reaction which is to be brought about, and,
as with the calculation of factors, to keep in mind the molecular relations
between the reagent and the substance reacted upon. For example, to
estimate the weight of barium chloride necessary to precipitate the sulphur
from 0.1 gram of pure pyrite (FeS_{2}), the proportion should read

488. 120.0 2(BaCl_{2}.2H_{2}O):FeS_{2} = x:0.1,

where !x! represents the weight of the chloride required. Each of the two
atoms of sulphur will form upon oxidation a molecule of sulphuric acid or a
sulphate, which, in turn, will require a molecule of the barium chloride for
precipitation. To determine the quantity of the barium chloride required, it
is necessary to include in its molecular weight the water of crystallization,
since this is inseparable from the chloride when it is weighed. This applies
equally to other similar instances.

If the strength of an acid is expressed in percentage by weight, due regard
must be paid to its specific gravity. For example, hydrochloric acid (sp. gr.
1.12) contains 23.8 per cent HCl !by weight!; that is, 0.2666 gram HCl in
each cubic centimeter.

5. It is sometimes desirable to avoid the manipulation required for the
separation of the constituents of a mixture of substances by making what is
called an "indirect analysis." For example, in the analysis of silicate rocks,
the sodium and potassium present may be obtained in the form of their
chlorides and weighed together. If the weight of such a mixture is known,
and also the percentage of chlorine present, it is possible to calculate the
amount of each chloride in the mixture. Let it be assumed that the weight of
the mixed chlorides is 0.15 gram, and that it contains 53 per cent of
chlorine.

The simplest solution of such a problem is reached through algebraic
methods. The weight of chlorine is evidently 0.15 x 0.53, or 0.0795 gram.
Let x represent the weight of sodium chloride present and y that of
potassium chloride. The molecular weight of NaCl is 58.5 and that of KCl
is 74.6. The atomic weight of chlorine is 35.5. Then

x + y = 0.15 (35.5/58.5)x + (35.5/74.6)y = 0.00795

Solving these equations for x shows the weight of NaCl to be 0.0625
gram. The weight of KCl is found by subtracting this from 0.15.

The above is one of the most common types of indirect analyses. Others
are more complex but they can be reduced to algebraic expressions and
solved by their aid. It should, however, be noted that the results obtained by
these indirect methods cannot be depended upon for high accuracy, since
slight errors in the determination of the common constituent, as chlorine in
the above mixture, will cause considerable variations in the values found
for the components. They should not be employed when direct methods are
applicable, if accuracy is essential.

PROBLEMS

(The reactions necessary for the solution of these problems are either stated
with the problem or may be found in the earlier text. In the calculations
from which the answers are derived, the atomic weights given on page 195
have been employed, using, however, only the first decimal but increasing
this by 1 when the second decimal is 5 or above. Thus, 39.1 has been taken
as the atomic weight of potassium, 32.1 for sulphur, etc. This has been done
merely to secure uniformity of treatment, and the student should remember
that it is always well to take into account the degree of accuracy desired in a
particular instance in determining the number of decimal places to retain.
Four-place logarithms were employed in the calculations. Where four
figures are given in the answer, the last figure may vary by one or (rarely)
by two units, according to the method by which the problem is solved.)

VOLUMETRIC ANALYSIS

1. How many grams of pure potassium hydroxide are required for exactly 1
liter of normal alkali solution?

!Answer!: 56.1 grams.

2. Calculate the equivalent in grams (a) of sulphuric acid as an acid; (b)
of hydrochloric acid as an acid; (c) of oxalic acid as an acid; (d) of nitric

acid as an acid.

!Answers!: (a) 49.05; (b) 36.5; (c) 63; (d) 63.

3. Calculate the equivalent in grams of (a) potassium hydroxide; (b) of
sodium carbonate; (c) of barium hydroxide; (d) of sodium bicarbonate when
titrated with an acid.

!Answers!: (a) 56.1; (b) 53.8; (c) 85.7; (d) 84.

4. What is the equivalent in grams of Na_{2}HPO_{4} (a) as a
phosphate; (b) as a sodium salt?

!Answers!: (a) 47.33; (b) 71.0.

5. A sample of aqueous hydrochloric acid has a specific gravity of 1.12
and contains 23.81 per cent hydrochloric acid by weight. Calculate the
grams and the milliequivalents of hydrochloric acid (HCl) in each cubic
centimeter of the aqueous acid.

!Answers!: 0.2667 gram; 7.307 milliequivalents.

6. How many cubic centimeters of hydrochloric acid (sp. gr. 1.20
containing 39.80 per cent HCl by weight) are required to furnish 36.45
grams of the gaseous compound?

!Answer!: 76.33 cc.

7. A given solution contains 0.1063 equivalents of hydrochloric acid in
976 cc. What is its normal value?

!Answer!: 0.1089 N.

8. In standardizing a hydrochloric acid solution it is found that 47.26 cc.
of hydrochloric acid are exactly equivalent to 1.216 grams of pure sodium
carbonate, using methyl orange as an indicator. What is the normal value of
the hydrochloric acid?

!Answer!: 0.4855 N.

9. Convert 42.75 cc. of 0.5162 normal hydrochloric acid to the equivalent
volume of normal hydrochloric acid.

!Answer!: 22.07 cc.

10. A solution containing 25.27 cc. of 0.1065 normal hydrochloric acid is
added to one containing 92.21 cc. of 0.5431 normal sulphuric acid and 50
cc. of exactly normal potassium hydroxide added from a pipette. Is the
solution acid or alkaline? How many cubic centimeters of 0.1 normal acid
or alkali must be added to exactly neutralize the solution?

!Answer!: 27.6 cc. alkali (solution is acid).

11. By experiment the normal value of a sulphuric acid solution is found
to be 0.5172. Of this acid 39.65 cc. are exactly equivalent to 21.74 cc. of a
standard alkali solution. What is the normal value of the alkali?

!Answer!: 0.9432 N.

12. A solution of sulphuric acid is standardized against a sample of
calcium carbonate which has been previously accurately analyzed and
found to contain 92.44% CaCO_{3} and no other basic material. The
sample weighing 0.7423 gram was titrated by adding an excess of acid
(42.42 cc.) and titrating the excess with sodium hydroxide solution (11.22
cc.). 1 cc. of acid is equivalent to 0.9976 cc. of sodium hydroxide. Calculate
the normal value of each.

!Answers!: Acid 0.4398 N; alkali 0.4409 N.

13. Given five 10 cc. portions of 0.1 normal hydrochloric acid, (a) how
many grams of silver chloride will be precipitated by a portion when an
excess of silver nitrate is added? (b) how many grams of pure anhydrous
sodium carbonate (Na_{2}CO_{3}) will be neutralized by a portion of it?
(c) how many grams of silver will there be in the silver chloride formed
when an excess of silver nitrate is added to a portion? (d) how many grams
of iron will be dissolved to FeCl_{2} by a portion of it? (e) how many
grams of magnesium chloride will be formed and how many grams of
carbon dioxide liberated when an excess of magnesium carbonate is treated
with a portion of the acid?

!Answers!: (a) 0.1434; (b) 0.053; (c) 0.1079; (d) 0.0279; (e) 0.04765, and
0.022.

14. If 30.00 grams of potassium tetroxalate
(KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O) are dissolved and the
solution diluted to exactly 1 liter, and 40 cc. are neutralized with 20 cc. of a
potassium carbonate solution, what is the normal value of the carbonate
solution?

!Answer!: 0.7084 N.

15. How many cubic centimeters of 0.3 normal sulphuric acid will be
required to neutralize (a) 30 cc. of 0.5 normal potassium hydroxide; (b) to
neutralize 30 cc. of 0.5 normal barium hydroxide; (c) to neutralize 20 cc. of
a solution containing 10.02 grams of potassium bicarbonate per 100 cc.; (d)
to give a precipitate of barium sulphate weighing 0.4320 gram?

!Answers!: (a) 50 cc.; (b) 50 cc.; (c) 66.73 cc.; (d) 12.33 cc.

16. It is desired to dilute a solution of sulphuric acid of which 1 cc. is
equivalent to 0.1027 gram of pure sodium carbonate to make it exactly
1.250 normal. 700 cc. of the solution are available. To what volume must it
be diluted?

!Answer!: 1084 cc.

17. Given the following data: 1 cc. of NaOH = 1.117 cc. HCl. The HCl is
0.4876 N. How much water must be added to 100 cc. of the alkali to make
it exactly 0.5 N.?

!Answer!: 9.0 cc.

18. What is the normal value of a sulphuric acid solution which has a
specific gravity of 1.839 and contains 95% H_{2}SO_{4} by weight?

!Answer!: 35.61 N.

19. A sample of Rochelle Salt (KNaC_{4}H_{4}O_{6}.4H_{2}O), after
ignition in platinum to convert it to the double carbonate, is titrated with
sulphuric acid, using methyl orange as an indicator. From the following data
calculate the percentage purity of the sample:

Wt. sample = 0.9500 gram
H_{2}SO_{4} used = 43.65 cc.
NaOH used = 1.72 cc.
1 cc. H_{2}SO_{4} = 1.064 cc. NaOH
Normal value NaOH = 0.1321 N.

!Answer!: 87.72 cc.

20. One gram of a mixture of 50% sodium carbonate and 50% potassium
carbonate is dissolved in water, and 17.36 cc. of 1.075 N acid is added. Is

the resulting solution acid or alkaline? How many cubic centimeters of
1.075 N acid or alkali will have to be added to make the solution exactly
neutral?

!Answers!: Acid; 1.86 cc. alkali.

21. In preparing an alkaline solution for use in volumetric work, an
analyst, because of shortage of chemicals, mixed exactly 46.32 grams of
pure KOH and 27.64 grams of pure NaOH, and after dissolving in water,
diluted the solution to exactly one liter. How many cubic centimeters of
1.022 N hydrochloric acid are necessary to neutralize 50 cc. of the basic
solution?

!Answer!: 74.18 cc.

22. One gram of crude ammonium salt is treated with strong potassium
hydroxide solution. The ammonia liberated is distilled and collected in 50
cc. of 0.5 N acid and the excess titrated with 1.55 cc. of 0.5 N sodium
hydroxide. Calculate the percentage of NH_{3} in the sample.

!Answer!: 41.17%.

23. In titrating solutions of alkali carbonates in the presence of
phenolphthalein, the color change takes place when the carbonate has been
converted to bicarbonate. In the presence of methyl orange, the color
change takes place only when the carbonate has been completely
neutralized. From the following data, calculate the percentages of
Na_{2}CO_{3} and NaOH in an impure mixture. Weight of sample, 1.500
grams; HCl (0.5 N) required for phenolphthalein end-point, 28.85 cc.; HCl
(0.5 N) required to complete the titration after adding methyl orange, 23.85
cc.

!Answers!: 6.67% NaOH; 84.28% Na_{2}CO_{3}.

24. A sample of sodium carbonate containing sodium hydroxide weighs
1.179 grams. It is titrated with 0.30 N hydrochloric acid, using
phenolphthalein in cold solution as an indicator and becomes colorless after
the addition of 48.16 cc. Methyl orange is added and 24.08 cc. are needed
for complete neutralization. What is the percentage of NaOH and
Na_{2}CO_{3}?

!Answers!: 24.50% NaOH; 64.92% Na_{2}CO_{3}.

25. From the following data, calculate the percentages of Na_{2}CO_{3}
and NaHCO_{3} in an impure mixture. Weight of sample 1.000 gram;
volume of 0.25 N hydrochloric acid required for phenolphthalein end-point,
26.40 cc.; after adding an excess of acid and boiling out the carbon dioxide,
the total volume of 0.25 N hydrochloric acid required for phenolphthalein
end-point, 67.10 cc.

!Answer!: 69.95% Na_{2}CO_{3}; 30.02% NaHCO_{3}.

26. In the analysis of a one-gram sample of soda ash, what must be the
normality of the acid in order that the number of cubic centimeters of acid
used shall represent the percentage of carbon dioxide present?

!Answer!: 0.4544 gram.

27. What weight of pearl ash must be taken for analysis in order that the
number of cubic centimeters of 0.5 N acid used may be equal to one third
the percentage of K_{2}CO_{3}?

!Answer!: 1.152 grams.

28. What weight of cream of tartar must have been taken for analysis in
order to have obtained 97.60% KHC_{4}H_{4}O_{6} in an analysis
involving the following data: NaOH used = 30.06 cc.; H_{2}SO_{4}
solution used = 0.50 cc.; 1 cc. H_{2}SO_{4} sol. = 0.0255 gram
CaCO_{3}; 1 cc. H_{2}SO_{4} sol. = 1.02 cc. NaOH sol.?

!Answer!: 2.846 grams.

29. Calculate the percentage of potassium oxide in an impure sample of
potassium carbonate from the following data: Weight of sample = 1.00
gram; HCl sol. used = 55.90 cc.; NaOH sol. used = 0.42 cc.; 1 cc. NaOH
sol. = 0.008473 gram of KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O; 2
cc. HCl sol. = 5 cc. NaOH sol.

!Answer!: 65.68%.

30. Calculate the percentage purity of a sample of calcite (CaCO_{3})
from the following data: (Standardization); Weight of
H_{2}C_{2}O_{4}.2H_{2}O = 0.2460 gram; NaOH solution used = 41.03
cc.; HCl solution used = 0.63; 1 cc. NaOH solution = 1.190 cc. HCl
solution. (Analysis); Weight of sample 0.1200 gram; HCl used = 36.38 cc.;
NaOH used = 6.20 cc.

!Answer!: 97.97%.

31. It is desired to dilute a solution of hydrochloric acid to exactly 0.05
N. The following data are given: 44.97 cc. of the hydrochloric acid are
equivalent to 43.76 cc. of the NaOH solution. The NaOH is standardized
against a pure potassium tetroxalate
(KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O) weighing 0.2162 gram
and requires 49.14 cc. How many cc. of water must be added to 1000 cc. of
the aqueous hydrochloric acid?

!Answer!: 11 cc.

32. How many cubic centimeters of 3 N phosphoric acid must be added
to 300 cc. of 0.4 N phosphoric acid in order that the resulting solution may
be 0.6 N?

!Answer!: 25 cc.

33. To oxidize the iron in 1 gram of FeSO_{4}(NH_{4})
{2}SO{4}.6H_{2}O (mol. wgt. 392) requires 3 cc. of a given solution of
HNO_{3}. What is the normality of the nitric acid when used as an acid?
6FeSO_{4} + 2HNO_{3} + 2H_{2}SO_{4} = 3Fe_{2}(SO_{4}){3} +
2NO + 4H{2}O.

!Answer!: 0.2835 N.

34. The same volume of carbon dioxide at the same temperature and the
same pressure is liberated from a 1 gram sample of dolomite, by adding an
excess of hydrochloric acid, as can be liberated by the addition of 35 cc. of
0.5 N hydrochloric acid to an excess of any pure or impure carbonate.
Calculate the percentage of CO_{2} in the dolomite.

!Answer!: 38.5%.

35. How many cubic centimeters of sulphuric acid (sp. gr. 1.84,
containing 96% H_{2}SO_{4} by weight) will be required to displace the
chloride in the calcium chloride formed by the action of 100 cc. of 0.1072
N hydrochloric acid on an excess of calcium carbonate, and how many
grams of CaSO_{4} will be formed?

!Answers!: 0.298 cc.; 0.7300 gram.

36. Potassium hydroxide which has been exposed to the air is found on
analysis to contain 7.62% water, 2.38% K_{2}CO_{3}. and 90% KOH.
What weight of residue will be obtained if one gram of this sample is added
to 46 cc. of normal hydrochloric acid and the resulting solution, after exact
neutralization with 1.070 N potassium hydroxide solution, is evaporated to
dryness?

!Answer!: 3.47 grams.

37. A chemist received four different solutions, with the statement that
they contained either pure NaOH; pure Na_{2}CO_{3}; pure NaHCO_{3},
or mixtures of these substances. From the following data identify them:

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

Sample II. On titrating with standard acid, it required 15.26 cc. for a
change in color, using phenolphthalein, and 17.90 cc. additional, using
methyl orange as an indicator.

Sample III. The sample was titrated with hydrochloric acid until the pink
of phenolphthalein disappeared, and on the addition of methyl orange the
solution was colored pink.

Sample IV. On titrating with hydrochloric acid, using phenolphthalein,
15.00 cc. were required. A new sample of the same weight required exactly
30 cc. of the same acid for neutralization, using methyl orange.

!Answers!: (a) NaHCO_{3}; (b) NaHCO_{3}+Na_{2}CO_{3};
(c)NaOH; (d)
Na_{2}CO_{3}.

38. In the analysis of a sample of KHC_{4}H_{4}O_{6} the following
data are obtained: Weight sample = 0.4732 gram. NaOH solution used =
24.97 cc. 3.00 cc. NaOH = 1 cc. of H_{3}PO_{4} solution of which 1 cc.
will precipitate 0.01227 gram of magnesium as MgNH_{4}PO_{4}.
Calculate the percentage of KHC_{4}H_{4}O_{6}.

!Answer!: 88.67%.

39. A one-gram sample of sodium hydroxide which has been exposed to
the air for some time, is dissolved in water and diluted to exactly 500 cc.
One hundred cubic centimeters of the solution, when titrated with 0.1062 N
hydrochloric acid, using methyl orange as an indicator, requires 38.60 cc.
for complete neutralization. Barium chloride in excess is added to a second
portion of 100 cc. of the solution, which is diluted to exactly 250 cc.,
allowed to stand and filtered. Two hundred cubic centimeters of this filtrate
require 29.62 cc. of 0.1062 N hydrochloric acid for neutralization, using
phenolphthalein as an indicator. Calculate percentage of NaOH,
Na_{2}CO_{3}, and H_{2}O.

!Answers!: 78.63% NaOH; 4.45% Na_{2}CO_{3}; 16.92% H_{2}O.

40. A sodium hydroxide solution (made from solid NaOH which has been
exposed to the air) was titrated against a standard acid using methyl orange
as an indicator, and was found to be exactly 0.1 N. This solution was used
in the analysis of a material sold at 2 cents per pound per cent of an acid
constituent A, and always mixed so that it was supposed to contain 15% of
A, on the basis of the analyst's report. Owing to the carelessness of the
analyst's assistant, the sodium hydroxide solution was used with
phenolphthalein as an indicator in cold solution in making the analyses. The
concern manufacturing this material sells 600 tons per year, and when the
mistake was discovered it was estimated that at the end of a year the error in
the use of indicators would either cost them or their customers $6000. Who

would lose and why? Assuming the impure NaOH used originally in
making the titrating solution consisted of NaOH and Na_{2}CO_{3} only,
what per cent of each was present?

!Answers!: Customer lost; 3.94% Na_{2}CO_{3}; 96.06% NaOH.

41. In the standardization of a K_{2}Cr_{2}O_{7} solution against iron
wire, 99.85% pure, 42.42 cc. of the solution were added. The weight of the
wire used was 0.22 gram. 3.27 cc. of a ferrous sulphate solution having a
normal value as a reducing agent of 0.1011 were added to complete the
titration. Calculate the normal value of the K_{2}Cr_{2}O_{7}.

!Answer!: 0.1006 N.

42. What weight of iron ore containing 56.2% Fe should be taken to
standardize an approximately 0.1 N oxidizing solution, if not more than 47
cc. are to be used?

!Answer!: 0.4667 gram.

43. One tenth gram of iron wire, 99.78% pure, is dissolved in
hydrochloric acid and the iron oxidized completely with bromine water.
How many grams of stannous chloride are there in a liter of solution if it
requires 9.47 cc. to just reduce the iron in the above? What is the normal
value of the stannous chloride solution as a reducing agent?

!Answer!: 17.92 grams; 0.1888 N.

44. One gram of an oxide of iron is fused with potassium acid sulphate
and the fusion dissolved in acid. The iron is reduced with stannous chloride,
mercuric chloride is added, and the iron titrated with a normal
K_{2}Cr_{2}O_{7} solution. 12.94 cc. were used. What is the formula of
the oxide, FeO, Fe_{2}O_{3}, or Fe_{3}O_{4}?

!Answer!: Fe_{3}O_{4}.

45. If an element has 98 for its atomic weight, and after reduction with
stannous chloride could be oxidized by bichromate to a state corresponding
to an XO_{4}^{-} anion, compute the oxide, or valence, corresponding to
the reduced state from the following data: 0.3266 gram of the pure element,
after being dissolved, was reduced with stannous chloride and oxidized by
40 cc. of K_{2}Cr_{2}O_{7}, of which one cc. = 0.1960 gram of
FeSO_{4}(NH_{4}){2}SO{4}.6H_{2}O.

!Answer!: Monovalent.

46. Determine the percentage of iron in a sample of limonite from the
following data: Sample = 0.5000 gram. KMnO_{4} used = 50 cc. 1 cc.
KMnO_{4} = 0.005317 gram Fe. FeSO_{4} used = 6 cc. 1 cc. FeSO_{4} =
0.009200 gram FeO.

!Answer!: 44.60%.

47. If 1 gram of a silicate yields 0.5000 gram of Fe_{2}O_{3} and
Al_{2}O_{3} and the iron present requires 25 cc. of 0.2 N KMnO_{4},
calculate the percentage of FeO and Al_{2}O_{3} in the sample.

!Answer!: 35.89% FeO; 10.03% Al_{2}O_{3}.

48. A sample of magnesia limestone has the following composition:
Silica, 3.00%; ferric oxide and alumina, 0.20%; calcium oxide, 33.10%;
magnesium oxide, 20.70%; carbon dioxide, 43.00%. In manufacturing lime
from the above the carbon dioxide is reduced to 3.00%. How many cubic
centimeters of normal KMnO_{4} will be required to determine the
calcium oxide volumetrically in a 1 gram sample of the lime?

!Answer!: 20.08 cc.

49. If 100 cc. of potassium bichromate solution (10 gram
K_{2}Cr_{2}O_{7} per liter), 5 cc. of 6 N sulphuric acid, and 75 cc. of
ferrous sulphate solution (80 grams FeSO_{4}.7H_{2}O per liter) are
mixed, and the resulting solution titrated with 0.2121 N KMnO_{4}, how
many cubic centimeters of the KMnO_{4} solution will be required to
oxidize the iron?

!Answer!: 5.70 cc.

50. If a 0.5000 gram sample of limonite containing 59.50 per cent
Fe_{2}O_{3} requires 40 cc. of KMnO_{4} to oxidize the iron, what is the
value of 1 cc. of the permanganate in terms of (a) Fe, (b)
H_{2}C_{2}O_{4}.2H_{2}O?

!Answers!: (a) 0.005189 gram; (b) 0.005859 gram.

51. A sample of pyrolusite weighing 0.6000 gram is treated with 0.9000
gram of oxalic acid. The excess oxalic acid requires 23.95 cc. of
permanganate (1 cc. = 0.03038 gram FeSO_{4}.7H_{2}O). What is the
percentage of MnO_{2}, in the sample?

!Answer!: 84.47%.

52. A solution contains 50 grams of
KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O per liter. What is the
normal value of the solution (a) as an acid, and (b) as a reducing agent?

!Answers!: (a) 0.5903 N; (b) 0.7872 N.

53. In the analysis of an iron ore containing 60% Fe_{2}O_{3}, a sample
weighing 0.5000 gram is taken and the iron is reduced with sulphurous
acid. On account of failure to boil out all the excess SO_{2}, 38.60 cubic
centimeters of 0.1 N KMnO_{4} were required to titrate the solution. What

was the error, percentage error, and what weight of sulphur dioxide was in
the solution?

!Answers!: (a) 1.60%; (b) 2.67%; (c) 0.00322 gram.

54. From the following data, calculate the ratio of the nitric acid as an
oxidizing agent to the tetroxalate solution as a reducing agent: 1 cc.
HNO_{3} = 1.246 cc. NaOH solution; 1 cc. NaOH = 1.743 cc.
KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O solution; Normal value
NaOH = 0.12.

!Answer!: 4.885.

55. Given the following data: 25 cc. of a hydrochloric acid, when
standardized gravimetrically as silver chloride, yields a precipitate weighing
0.5465 gram. 24.35 cc. of the hydrochloric acid are exactly equivalent to
30.17 cc. of KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O solution. How
much water must be added to a liter of the oxalate solution to make it
exactly 0.025 N as a reducing agent?

!Answer!: 5564 cc.

56. Ten grams of a mixture of pure potassium tetroxalate
(KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O) and pure oxalic acid
(H_{2}C_{2}O_{4}.2H_{2}O) are dissolved in water and diluted to
exactly 1000 cc. The normal value of the oxalate solution when used as an
acid is 0.1315. Calculate the ratio of tetroxalate to oxalate used in making
up the solution and the normal value of the solution as a reducing agent.

!Answers!: 2:1; 0.1577 N.

57. A student standardized a solution of NaOH and one of KMnO_{4}
against pure KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O and found the

former to be 0.07500 N as an alkali and the latter exactly 0.1 N as an
oxidizing agent. By coincidence, exactly 47.26 cc. were used in each
standardization. Find the ratio of the oxalate used in the NaOH
standardization to the oxalate used in the permanganate standardization.

!Answer!: 1:1.

58. A sample of apatite weighing 0.60 gram is analyzed for its phosphoric
anhydride content. If the phosphate is precipitated as (NH_{4})
{3}PO{4}.12MoO_{3}, and the precipitate (after solution and reduction of
the MoO_{3} to Mo_{24}O_{37}), requires 100 cc. of normal KMnO_{4}
to oxidize it back to MoO_{3}, what is the percentage of P_{2}O_{5}?

!Answer!: 33.81%.

59. In the analysis of a sample of steel weighing 1.881 grams the
phosphorus was precipitated with ammonium molybdate and the yellow
precipitate was dissolved, reduced and titrated with KMnO_{4}. If the
sample contained 0.025 per cent P and 6.01 cc. of KMnO_{4} were used, to
what oxide was the molybdenum reduced? 1 cc. KMnO_{4} = 0.007188
gram Na_{2}C_{2}O_{4}.

!Answer!: Mo_{4}O_{5}.

60. What is the value of 1 cc. of an iodine solution (1 cc. equivalent to
0.0300 gram Na_{2}S_{2}O_{3}) in terms of As_{2}O_{3}?

!Answer!: 0.009385 gram.

61. 48 cc. of a solution of sodium thiosulphate are required to titrate the
iodine liberated from an excess of potassium iodide solution by 0.3000
gram of pure KIO_{3}. (KIO_{3} + 5KI + 3H_{2}SO_{4} =

3K_{2}SO_{4} + 3I_{2} + 3H_{2}O.) What is the normal strength of the
sodium thiosulphate and the value of 1 cc. of it in terms of iodine?

!Answers!: 0.1753 N; 0.02224 gram.

62. One thousand cubic centimeters of 0.1079 N sodium thiosulphate
solution is allowed to stand. One per cent by weight of the thiosulphate is
decomposed by the carbonic acid present in the solution. To what volume
must the solution be diluted to make it exactly 0.1 N as a reducing agent?
(Na_{2}S_{2}O_{3} + 2H_{2}CO_{3} = H_{2}SO_{3} + 2NaHCO_{3}
+ S.)

!Answer!: 1090 cc.

63. An analyzed sample of stibnite containing 70.05% Sb is given for
analysis. A student titrates it with a solution of iodine of which 1 cc. is
equivalent to 0.004950 gram of As_{2}O_{3}. Due to an error on his part
in standardization, the student's analysis shows the sample to contain
70.32% Sb. Calculate the true normal value of the iodine solution, and the
percentage error in the analysis.

!Answers!: 0.1000 N; 0.39%.

64. A sample of pyrolusite weighing 0.5000 gram is treated with an
excess of hydrochloric acid, the liberated chlorine is passed into potassium
iodide and the liberated iodine is titrated with sodium thiosulphate solution
(49.66 grams of pure Na_{2}S_{2}O_{3}.5H_{2}O per liter). If 38.72 cc.
are required, what volume of 0.25 normal permanganate solution will be
required in an indirect determination in which a similar sample is reduced
with 0.9012 gram H_{2}C_{2}O_{4}.2H_{2}O and the excess oxalic acid
titrated?

!Answer!: 26.22 cc.

65. In the determination of sulphur in steel by evolving the sulphur as
hydrogen sulphide, precipitating cadmium sulphide by passing the liberated
hydrogen sulphide through ammoniacal cadmium chloride solution, and
decomposing the CdS with acid in the presence of a measured amount of
standard iodine, the following data are obtained: Sample, 5.027 grams; cc.
Na_{2}S_{2}O_{3} sol. = 12.68; cc. Iodine sol. = 15.59; 1 cc. Iodine sol. =
1.086 cc. Na_{2}S_{2}O_{3} sol.; 1 cc. Na_{2}S_{2}O_{3}= 0.005044
gram Cu. Calculate the percentage of sulphur. (H_{2}S + I_{2} = 2HI + S.)

!Answer!: 0.107%.

66. Given the following data, calculate the percentage of iron in a sample
of crude ferric chloride weighing 1.000 gram. The iodine liberated by the
reaction 2FeCl_{3}+ 2HI = 2HCl + 2FeCl_{2} + I_{2} is reduced by the
addition of 50 cc. of sodium thiosulphate solution and the excess
thiosulphate is titrated with standard iodine and requires 7.85 cc. 45 cc.
I_{2} solution = 45.95 cc. Na_{2}S_{2}O_{3} solution; 45 cc.
As_{2}O_{3} solution = 45.27 cc. I_{2} solution. 1 cc. arsenite solution =
0.005160 gram As_{2}O_{3}.

!Answer!: 23.77%.

67. Sulphide sulphur was determined in a sample of reduced barium
sulphate by the evolution method, in which the sulphur was evolved as
hydrogen sulphide and was passed into CdCl_{2} solution, the acidified
precipitate being titrated with iodine and thiosulphate. Sample, 5.076
grams; cc. I_{2} = 20.83; cc. Na_{2}S_{2}O_{3} = 12.37; 43.45 cc.
Na_{2}S_{2}O_{3} = 43.42 cc. I_{2}; 8.06 cc. KMnO_{4} = 44.66 cc.
Na_{2}S_{2}O_{3}; 28.87 cc. KMnO_{4} = 0.2004 gram

Na_{2}C_{2}O_{4}. Calculate the percentage of sulphide sulphur in the
sample.

!Answer!: 0.050%.

68. What weight of pyrolusite containing 89.21% MnO_{2} will oxidize
the same amount of oxalic acid as 37.12 cc. of a permanganate solution, of
which 1 cc. will liberate 0.0175 gram of I_{2} from KI?

!Answer!: 0.2493 gram.

69. A sample of pyrolusite weighs 0.2400 gram and is 92.50% pure
MnO_{2}. The iodine liberated from KI by the manganese dioxide is
sufficient to react with 46.24 cc. of Na_{2}S_{2}O_{3} sol. What is the
normal value of the thiosulphate?

!Answer!:: 0.1105 N.

70. In the volumetric analysis of silver coin (90% Ag), using a 0.5000
gram sample, what is the least normal value that a potassium thiocyanate
solution may have and not require more than 50 cc. of solution in the
analysis?

!Answer!: 0.08339 N.

71. A mixture of pure lithium chloride and barium bromide weighing 0.6
gram is treated with 45.15 cubic centimeters of 0.2017 N silver nitrate, and
the excess titrated with 25 cc. of 0.1 N KSCN solution, using ferric alum as
an indicator. Calculate the percentage of bromine in the sample.

!Answer!: 40.11%.

72. A mixture of the chlorides of sodium and potassium from 0.5000
gram of a feldspar weighs 0.1500 gram, and after solution in water requires
22.71 cc. of 0.1012 N silver nitrate for the precipitation of the chloride ions.
What are the percentages of Na_{2}O and K_{2}O in the feldspar?

!Answer!: 8.24% Na_{2}O; 9.14% K_{2}O.

GRAVIMETRIC ANALYSIS

73. Calculate (a) the grams of silver in one gram of silver chloride; (b) the
grams of carbon dioxide liberated by the addition of an excess of acid to
one gram of calcium carbonate; (c) the grams of MgCl_{2} necessary to
precipitate 1 gram of MgNH_{4}PO_{4}.

!Answers!: (a) 0.7526; (b) 0.4397; (c) 0.6940.

74. Calculate the chemical factor for (a) Sn in SnO_{2}; (b) MgO in
Mg_{2}P_{2}O_{7}; (c) P_{2}O_{5} in Mg_{2}P_{2}O_{7}; (d) Fe in
Fe_{2}O_{3}; (e) SO_{4} in BaSO_{4}.

!Answers!: (a) 0.7879; (b) 0.3620; (c) 0.6378; (d) 0.6990; (e) 0.4115.

75. Calculate the log factor for (a) Pb in PbCrO_{4}; (b) Cr_{2}O_{3}
in PbCrO_{4}; (c) Pb in PbO_{2} and (d) CaO in CaC_{2}O_{4}.

!Answers!: (a) 9.8069-10, (b) 9.3713-10; (c) 9.9376-10; (d) 9.6415-10.

76. How many grams of Mn_{3}O_{4} can be obtained from 1 gram of
MnO_{2}?

!Answer!: 0.8774 gram.

77. If a sample of silver coin weighing 0.2500 gram gives a precipitate of
AgCl weighing 0.2991 gram, what weight of AgI could have been obtained
from the same weight of sample, and what is the percentage of silver in the
coin?

!Answers!: 0.4898 gr.; 90.05%.

78. How many cubic centimeters of hydrochloric acid (sp. gr. 1.13
containing 25.75% HCl by weight) are required to exactly neutralize 25 cc.
of ammonium hydroxide (sp. gr. .90 containing 28.33% NH_{3} by
weight)?

!Answer!: 47.03 cc.

79. How many cubic centimeters of ammonium hydroxide solution (sp.
gr. 0.96 containing 9.91% NH_{3} by weight) are required to precipitate
the aluminium as aluminium hydroxide from a two-gram sample of alum
(KAl(SO_{4}){2}.12H{2}O)? What will be the weight of the ignited
precipitate?

!Answers!: 2.26 cc.; 0.2154 gram.

80. What volume of nitric acid (sp. gr. 1.05 containing 9.0% HNO_{3}
by weight) is required to oxidize the iron in one gram of
FeSO_{4}.7H_{2}O in the presence of sulphuric acid? 6FeSO_{4} +
2HNO_{3} + 3H_{2}SO_{4} = 3Fe_{2}(SO_{4}){3} + 2NO + 4H{2}O.

!Answer!: 0.80 cc.

81. If 0.7530 gram of ferric nitrate (Fe(NO_{3}){3}.9H{2}O) is
dissolved in water and 1.37 cc. of HCl (sp. gr. 1.11 containing 21.92% HCl
by weight) is added, how many cubic centimeters of ammonia (sp. gr. 0.96

containing 9.91% NH_{3} by weight) are required to neutralize the acid
and precipitate the iron as ferric hydroxide?

!Answer!: 2.63 cc.

82. To a suspension of 0.3100 gram of Al(OH){3} in water are added
13.00 cc. of aqueous ammonia (sp. gr. 0.90 containing 28.4% NH{3} by
weight). How many cubic centimeters of sulphuric acid (sp. gr. 1.18
containing 24.7% H_{2}SO_{4} by weight) must be added to the mixture
in order to bring the aluminium into solution?

!Answer!: 34.8 cc.

83. How many cubic centimeters of sulphurous acid (sp. gr. 1.04
containing 75 grams SO_{2} per liter) are required to reduce the iron in 1
gram of ferric alum (KFe(SO_{4}){2}.12H{2}O)? Fe_{2}(SO_{4}){3} +
SO{2} + 2H_{2}O = 2FeSO_{4} + 2H_{2}SO_{4}.

!Answer!: 0.85 cc.

84. How many cubic centimeters of a solution of potassium bichromate
containing 26.30 grams of K_{2}Cr_{2}O_{7} per liter must be taken in
order to yield 0.6033 gram of Cr_{2}O_{3} after reduction and
precipitation of the chromium?

K_{2}Cr_{2}O_{7} + 3SO_{2} + H_{2}SO_{4} = K_{2}SO_{4} +
Cr_{2}(SO_{4}){3} + H{2}O.

!Answer!: 44.39 cc.

85. How many cubic centimeters of ammonium hydroxide (sp. gr. 0.946
containing 13.88% NH_{3} by weight) are required to precipitate the iron
as Fe(OH){3} from a sample of pure FeSO{4}.(NH_{4})

{2}SO{4}.6H_{2}O, which requires 0.34 cc. of nitric acid (sp. gr. 1.350
containing 55.79% HNO_{3} by weight) for oxidation of the iron? (See
problem No. 80 for reaction.)

!Answer!: 4.74 cc.

86. In the analysis of an iron ore by solution, oxidation and precipitation
of the iron as Fe(OH){3}, what weight of sample must be taken for analysis
so that each one hundredth of a gram of the ignited precipitate of
Fe{2}O_{3} shall represent one tenth of one per cent of iron?

!Answer!: 6.99 grams.

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

!Answer!: 0.6870 gram.

88. What weight of magnetite must be taken for analysis in order that,
after precipitating and igniting all the iron to Fe_{2}O_{3}, the percentage
of Fe_{2}O_{4} in the sample may be found by multiplying the weight in
grams of the ignited precipitate by 100?

!Answer!: 0.9665 gram.

89. After oxidizing the arsenic in 0.5000 gram of pure As_{2}S_{3} to
arsenic acid, it is precipitated with "magnesia mixture" (MgCl_{2} +
2NH_{4}Cl). If exactly 12.6 cc. of the mixture are required, how many
grams of MgCl_{2} per liter does the solution contain? H_{3}AsO_{4} +
MgCl_{2} + 3NH_{4}OH = MgNH_{4}AsO_{4} + 2NH_{4}Cl +
3H_{2}O.

!Answer!: 30.71 grams.

90. A sample is prepared for student analysis by mixing pure apatite
(Ca_{3}(PO_{4}){2}.CaCl{2}) with an inert material. If 1 gram of the
sample gives 0.4013 gram of Mg_{2}P_{2}O_{7}, how many cubic
centimeters of ammonium oxalate solution (containing 40 grams of
(NH_{4}){2}C{2}O_{4}.H_{2}O per liter) would be required to
precipitate the calcium from the same weight of sample?

!Answer!: 25.60 cc.

91. If 0.6742 gram of a mixture of pure magnesium carbonate and pure
calcium carbonate, when treated with an excess of hydrochloric acid, yields
0.3117 gram of carbon dioxide, calculate the percentage of magnesium
oxide and of calcium oxide in the sample.

!Answers!: 13.22% MgO; 40.54% CaO. 92. The calcium in a sample of
dolomite weighing 0.9380 gram is precipitated as calcium oxalate and
ignited to calcium oxide. What volume of gas, measured over water at
20°C. and 765 mm. pressure, is given off during ignition, if the resulting
oxide weighs 0.2606 gram? (G.M.V. = 22.4 liters; V.P. water at 20°C. = 17.4
mm.)

!Answer!: 227 cc.

93. A limestone is found to contain 93.05% CaCO_{3}, and 5.16 %
MgCO_{3}. Calculate the weight of CaO obtainable from 3 tons of the
limestone, assuming complete conversion to oxide. What weight of
Mg_{2}P_{2}O_{7} could be obtained from a 3-gram sample of the
limestone?

!Answers!: 1.565 tons; 0.2044 gram.

94. A sample of dolomite is analyzed for calcium by precipitating as the
oxalate and igniting the precipitate. The ignited product is assumed to be
CaO and the analyst reports 29.50% Ca in the sample. Owing to insufficient
ignition, the product actually contained 8% of its weight of CaCO_{3}.
What is the correct percentage of calcium in the sample, and what is the
percentage error?

!Answers!: 28.46%; 3.65% error.

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

!Answer!: 0.2359 gram.

96. How many cubic centimeters of HNO_{3} (sp. gr. 1.13 containing
21.0% HNO_{3} by weight) are required to dissolve 5 grams of brass,
containing 0.61% Pb, 24.39% Zn, and 75% Cu, assuming reduction of the
nitric acid to NO by each constituent? What fraction of this volume of acid
is used for oxidation?

!Answers!: 55.06 cc.; 25%.

97. What weight of metallic copper will be deposited from a cupric salt
solution by a current of 1.5 amperes during a period of 45 minutes,
assuming 100% current efficiency? (1 Faraday = 96,500 coulombs.)

!Answer!: 1.335 grams.

98. In the electrolysis of a 0.8000 gram sample of brass, there is obtained
0.0030 gram of PbO_{2}, and a deposit of metallic copper exactly equal in
weight to the ignited precipitate of Zn_{2}P_{2}O_{7} subsequently

obtained from the solution. What is the percentage composition of the
brass?

!Answers!: 69.75% Cu; 29.92% Zn; 0.33% Pb.

99. A sample of brass (68.90% Cu; 1.10% Pb and 30.00% Zn) weighing
0.9400 gram is dissolved in nitric acid. The lead is determined by weighing
as PbSO_{4}, the copper by electrolysis and the zinc by precipitation with
(NH_{4}){2}HPO{4} in a neutral solution.

(a) Calculate the cubic centimeters of nitric acid (sp. gr. 1.42 containing
69.90% HNO_{3} by weight) required to just dissolve the brass, assuming
reduction to NO.

!Answer!: 2.48 cc.

(b) Calculate the cubic centimeters of sulphuric acid (sp. gr. 1.84
containing 94% H_{2}SO_{4} by weight) to displace the nitric acid.

!Answer!: 0.83 cc.

(c) Calculate the weight of PbSO_{4}.

!Answer!: 0.0152 gram.

(d) The clean electrode weighs 10.9640 grams. Calculate the weight after
the copper has been deposited.

!Answer!: 11.6116 grams.

(e) Calculate the grams of (NH_{4}){2}HPO{4} required to precipitate
the zinc as ZnNH_{4}PO_{4}.

!Answer!: 0.5705 gram.

(f) Calculate the weight of ignited Zn_{2}P_{2}O_{7}.

!Answer!: 0.6573 gram.

100. If in the analysis of a brass containing 28.00% zinc an error is made
in weighing a 2.5 gram portion by which 0.001 gram too much is weighed
out, what percentage error in the zinc determination would result? What
volume of a solution of sodium hydrogen phosphate, containing 90 grams
of Na_{2}HPO_{4}.12H_{2}O per liter, would be required to precipitate
the zinc as ZnNH_{4}PO_{4} and what weight of precipitate would be
obtained?

!Answers!: (a) 0.04% error; (b) 39.97 cc.; (c) 1.909 grams.

101. A sample of magnesium carbonate, contaminated with SiO_{2} as
its only impurity, weighs 0.5000 gram and loses 0.1000 gram on ignition.
What volume of disodium phosphate solution (containing 90 grams
Na_{2}HPO_{4}.12H_{2}O per liter) will be required to precipitate the
magnesium as magnesium ammonium phosphate?

!Answer!: 9.07 cc.

102. 2.62 cubic centimeters of nitric acid (sp. gr. 1.42 containing 69.80%
HNO_{2} by weight) are required to just dissolve a sample of brass
containing 69.27% Cu; 0.05% Pb; 0.07% Fe; and 30.61% Zn. Assuming the
acid used as oxidizing agent was reduced to NO in every case, calculate the
weight of the brass and the cubic centimeters of acid used as acid.

!Answer!: 0.992 gram; 1.97 cc.

103. One gram of a mixture of silver chloride and silver bromide is found
to contain 0.6635 gram of silver. What is the percentage of bromine?

!Answer!: 21.30%.

104. A precipitate of silver chloride and silver bromide weighs 0.8132
gram. On heating in a current of chlorine, the silver bromide is converted to
silver chloride, and the mixture loses 0.1450 gram in weight. Calculate the
percentage of chlorine in the original precipitate.

!Answer!: 6.13%.

105. A sample of feldspar weighing 1.000 gram is fused and the silica
determined. The weight of silica is 0.6460 gram. This is fused with 4 grams
of sodium carbonate. How many grams of the carbonate actually combined
with the silica in fusion, and what was the loss in weight due to carbon
dioxide during the fusion?

!Answers!: 1.135 grams; 0.4715 gram.

106. A mixture of barium oxide and calcium oxide weighing 2.2120
grams is transformed into mixed sulphates, weighing 5.023 grams.
Calculate the grams of calcium oxide and barium oxide in the mixture.

!Answers!: 1.824 grams CaO; 0.3877 gram BaO.

APPENDIX

ELECTROLYTIC DISSOCIATION THEORY

The following brief statements concerning the ionic theory and a few of its
applications are intended for reference in connection with the explanations
which are given in the Notes accompanying the various procedures. The
reader who desires a more extended discussion of the fundamental theory
and its uses is referred to such books as Talbot and Blanchard's !Electrolytic
Dissociation Theory! (Macmillan Company), or Alexander Smith's
!Introduction to General Inorganic Chemistry! (Century Company).

The !electrolytic dissociation theory!, as propounded by Arrhenius in
1887, assumes that acids, bases, and salts (that is, electrolytes) in aqueous
solution are dissociated to a greater or less extent into !ions!. These ions are
assumed to be electrically charged atoms or groups of atoms, as, for
example, H^{+} and Br^{-} from hydrobromic acid, Na^{+} and OH^{-}
from sodium hydroxide, 2NH_{4}^{+} and SO_{4}^{—} from ammonium
sulphate. The unit charge is that which is dissociated with a hydrogen ion.
Those upon other ions vary in sign and number according to the chemical
character and valence of the atoms or radicals of which the ions are
composed. In any solution the aggregate of the positive charges upon the
positive ions (!cations!) must always balance the aggregate negative
charges upon the negative ions (!anions!).

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

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

The chemical changes upon which the procedures of analytical chemistry
depend are almost exclusively those in which the reacting substances are
electrolytes, and analytical chemistry is, therefore, essentially the chemistry
of the ions. The percentage dissociation of the same electrolyte tends to
increase with increasing dilution of its solution, although not in direct
proportion. The percentage dissociation of different electrolytes in solutions
of equivalent concentrations (such, for example, as normal solutions) varies
widely, as is indicated in the following tables, in which approximate figures
are given for tenth-normal solutions at a temperature of about 18°C.

ACIDS
======================================================
===================
|
SUBSTANCE | PERCENTAGE DISSOCIATION IN

| 0.1 EQUIVALENT SOLUTION
_____________________________________________|_______________
____________
|
HCl, HBr, HI, HNO_{3} | 90
|
HClO_{3}, HClO_{4}, HMnO_{4} | 90
|
H_{2}SO_{4} <—> H^{+} + HSO_{4}^{-} | 90
|
H_{2}C_{2}O_{4} <—> H^{+} + HC_{2}O_{4}^{-} | 50
|
H_{2}SO_{3} <—> H^{+} + HSO{}3^{-} | 20
|
H{3}PO_{4} <—> H^{+} + H_{2}PO_{4}^{-} | 27
|
H_{2}PO_{4}^{-} <—> H^{+} + HPO_{4}^{—} | 0.2
|
H_{3}AsO_{4} <—> H^{+} + H_{2}AsO_{4}^{-} | 20
|
HF | 9
|
HC_{2}H_{3}O_{2} | 1.4
|
H_{2}CO_{3} <—> H^{+} + HCO_{3}^{-} | 0.12
|
H_{2}S <—> H^{+} + HS^{-} | 0.05
|
HCN | 0.01
|

======================================================
===================

BASES
======================================================
===================
|
SUBSTANCE | PERCENTAGE DISSOCIATION IN
| 0.1 EQUIVALENT SOLUTION
_____________________________________________|_______________
____________
|
KOH, NaOH | 86
|
Ba(OH){2} | 75
|
NH{4}OH | 1.4
|
======================================================
===================

SALTS
======================================================
===================
|
TYPE OF SALT | PERCENTAGE DISSOCIATION IN
| 0.1 EQUIVALENT SOLUTION
_____________________________________________|_______________
____________
|

R^{+}R^{-} | 86
|
R^{++}(R^{-}){2} | 72
|
(R^{+}){2}R^{—} | 72
|
R^{++}R^{—} | 45
|
======================================================
===================

The percentage dissociation is determined by studying the electrical
conductivity of the solutions and by other physico-chemical methods, and
the following general statements summarize the results:

!Salts!, as a class, are largely dissociated in aqueous solution.

!Acids! yield H^{+} ions in water solution, and the comparative
!strength!, that is, the activity, of acids is proportional to the concentration
of the H^{+} ions and is measured by the percentage dissociation in
solutions of equivalent concentration. The common mineral acids are
largely dissociated and therefore give a relatively high concentration of
H^{+} ions, and are commonly known as "strong acids." The organic acids,
on the other hand, belong generally to the group of "weak acids."

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

Ionic reactions are all, to a greater or less degree, !reversible reactions!. A
typical example of an easily reversible reaction is that representing the
changes in ionization which an electrolyte such as acetic acid undergoes on
dilution or concentration of its solutions, !i.e.!, HC_{2}H_{3}O_{2} <—>
H^{+} + C_{2}H_{3}O_{2}^{-}. As was stated above, the ionization
increases with dilution, the reaction then proceeding from left to right,
while concentration of the solution occasions a partial reassociation of the
ions, and the reaction proceeds from right to left. To understand the
principle underlying these changes it is necessary to consider first the
conditions which prevail when a solution of acetic acid, which has been
stirred until it is of uniform concentration throughout, has come to a
constant temperature. A careful study of such solutions has shown that there
is a definite state of equilibrium between the constituents of the solution;
that is, there is a definite relation between the undissociated acetic acid and
its ions, which is characteristic for the prevailing conditions. It is not,
however, assumed that this is a condition of static equilibrium, but rather
that there is continual dissociation and association, as represented by the
opposing reactions, the apparent condition of rest resulting from the fact
that the amount of change in one direction during a given time is exactly
equal to that in the opposite direction. A quantitative study of the amount of
undissociated acid, and of H^{+} ions and C_{2}H_{3}O_{2}^{-} ions
actually to be found in a large number of solutions of acetic acid of varying
dilution (assuming them to be in a condition of equilibrium at a common
temperature), has shown that there is always a definite relation between
these three quantities which may be expressed thus:

(!Conc'n H^{+} x Conc'n C_{2}H_{3}O_{2}^{-})/Conc'n
HC_{2}H_{3}O_{2} =
Constant!.

In other words, there is always a definite and constant ratio between the
product of the concentrations of the ions and the concentration of the
undissociated acid when conditions of equilibrium prevail.

It has been found, further, that a similar statement may be made
regarding all reversible reactions, which may be expressed in general terms
thus: The rate of chemical change is proportional to the product of the
concentrations of the substances taking part in the reaction; or, if conditions
of equilibrium are considered in which, as stated, the rate of change in
opposite directions is assumed to be equal, then the product of the
concentrations of the substances entering into the reaction stands in a
constant ratio to the product of the concentrations of the resulting
substances, as given in the expression above for the solutions of acetic acid.
This principle is called the !Law of Mass Action!.

It should be borne in mind that the expression above for acetic acid
applies to a wide range of dilutions, provided the temperature remains
constant. If the temperature changes the value of the constant changes
somewhat, but is again uniform for different dilutions at that temperature.
The following data are given for temperatures of about 18°C.[1]

====================================================
======================
||||
MOLAL | FRACTION | MOLAL CONCENTRA- | MOLAL
CONCENTRA- | VALUE OF
CONCENTRATION | IONIZED | TION OF H^{+} AND| TION OF
UNDIS- | CONSTANT
CONSTANT | | ACETATE^{-} IONS | SOCIATED ACID |
______________|__________|__________________|_________________
_|__________

||||
1.0 | .004 | .004 | .996 | .0000161
||||
0.1 | .013 | .0013 | .0987 | .0000171
||||
0.01 | .0407 | .000407 | .009593 | .0000172
||||
======================================================
=====================

[Footnote 1: Alexander Smith, !General Inorganic Chemistry!, p. 579.]

The molal concentrations given in the table refer to fractions of a gram-
molecule per liter of the undissociated acid, and to fractions of the
corresponding quantities of H^{+} and C_{2}H_{3}O_{2}^{-} ions per
liter which would result from the complete dissociation of a gram-molecule
of acetic acid. The values calculated for the constant are subject to some
variation on account of experimental errors in determining the percentage
ionized in each case, but the approximate agreement between the values
found for molal and centimolal (one hundredfold dilution) is significant.

The figures given also illustrate the general principle, that the !relative!
ionization of an electrolyte increases with the dilution of its solution. If we
consider what happens during the (usually) brief period of dilution of the
solution from molal to 0.1 molal, for example, it will be seen that on the
addition of water the conditions of concentration which led to equality in
the rate of change, and hence to equilibrium in the molal solution, cease to
exist; and since the dissociating tendency increases with dilution, as just
stated, it is true at the first instant after the addition of water that the
concentration of the undissociated acid is too great to be permanent under
the new conditions of dilution, and the reaction, HC_{2}H_{3}O_{2} <—>

H^{+} + C_{2}H_{3}O_{2}^{-}, will proceed from left to right with great
rapidity until the respective concentrations adjust themselves to the new
conditions.

That which is true of this reaction is also true of all reversible reactions,
namely, that any change of conditions which occasions an increase or a
decrease in concentration of one or more of the components causes the
reaction to proceed in one direction or the other until a new state of
equilibrium is established. This principle is constantly applied throughout
the discussion of the applications of the ionic theory in analytical chemistry,
and it should be clearly understood that whenever an existing state of
equilibrium is disturbed as a result of changes of dilution or temperature, or
as a consequence of chemical changes which bring into action any of the
constituents of the solution, thus altering their concentrations, there is
always a tendency to re-establish this equilibrium in accordance with the
law. Thus, if a base is added to the solution of acetic acid the H^{+} ions
then unite with the OH^{-} ions from the base to form undissociated water.
The concentration of the H^{+} ions is thus diminished, and more of the
acid dissociates in an attempt to restore equilbrium, until finally practically
all the acid is dissociated and neutralized.

Similar conditions prevail when, for example, silver ions react with
chloride ions, or barium ions react with sulphate ions. In the former case the
dissociation reaction of the silver nitrate is AgNO_{3} <—> Ag^{+} +
NO_{3}^{-}, and as soon as the Ag^{+} ions unite with the Cl^{-} ions the
concentration of the former is diminished, more of the AgNO_{3}
dissociates, and this process goes on until the Ag^{+} ions are practically
all removed from the solution, if the Cl^{-} ions are present in sufficient
quantity.

For the sake of accuracy it should be stated that the mass law cannot be
rigidly applied to solutions of those electrolytes which are largely
dissociated. While the explanation of the deviation from quantitative
exactness in these cases is not known, the law is still of marked service in
developing analytical methods along more logical lines than was formerly
practicable. It has not seemed wise to qualify each statement made in the
Notes to indicate this lack of quantitative exactness. The student should
recognize its existence, however, and will realize its significance better as
his knowledge of physical chemistry increases.

If we apply the mass law to the case of a substance of small solubility,
such as the compounds usually precipitated in quantitative analysis, we
derive what is known as the !solubility product!, as follows: Taking silver
chloride as an example, and remembering that it is not absolutely insoluble
in water, the equilibrium expression for its solution is:

(!Conc'n Ag^{+} x Conc'n Cl^{-})/Conc'n AgCl = Constant!.

But such a solution of silver chloride which is in contact with the solid
precipitate must be saturated for the existing temperature, and the quantity
of undissociated AgCl in the solution is definite and constant for that
temperature. Since it is a constant, it may be eliminated, and the expression
becomes !Conc'n Ag^{+} x Conc'n Cl^{-} = Constant!, and this is known
as the solubility product. No precipitation of a specific substance will occur
until the product of the concentrations of its ions in a solution exceeds the
solubility product for that substance; whenever that product is exceeded
precipitation must follow.

It will readily be seen that if a substance which yields an ion in common
with the precipitated compound is added to such a solution as has just been
described, the concentration of that ion is increased, and as a result the
concentration of the other ion must proportionately decrease, which can

only occur through the formation of some of the undissociated compound
which must separate from the already saturated solution. This explains why
the addition of an excess of the precipitant is often advantageous in
quantitative procedures. Such a case is discussed at length in Note 2 on
page 113.

Similarly, the ionization of a specific substance in solution tends to
diminish on the addition of another substance with a common ion, as, for
instance, the addition of hydrochloric acid to a solution of hydrogen
sulphide. Hydrogen sulphide is a weak acid, and the concentration of the
hydrogen ions in its aqueous solutions is very small. The equilibrium in
such a solution may be represented as:

(!(Conc'n H^{+})^{2} x Conc'n S^{—})/Conc'n H_{2}S = Constant!,
and a marked increase in the concentration of the H^{+} ions, such as
would result from the addition of even a small amount of the highly ionized
hydrochloric acid, displaces the point of equilibrium and some of the
S^{—} ions unite with H^{+} ions to form undissociated H_{2}S. This is
of much importance in studying the reactions in which hydrogen sulphide is
employed, as in qualitative analysis. By a parallel course of reasoning it
will be seen that the addition of a salt of a weak acid or base to solutions of
that acid or base make it, in effect, still weaker because they decrease its
percentage ionization.

To understand the changes which occur when solids are dissolved where
chemical action is involved, it should be remembered that no substance is
completely insoluble in water, and that those products of a chemical change
which are least dissociated will first form. Consider, for example, the action
of hydrochloric acid upon magnesium hydroxide. The minute quantity of
dissolved hydroxide dissociates thus: Mg(OH){2} <—> Mg^{++} +
2OH^{-}. When the acid is introduced, the H^{+} ions of the acid unite with

the OH^{-} ions to form undissociated water. The concentration of the
OH^{-} ions is thus diminished, more Mg(OH){2} dissociates, the solution
is no longer saturated with the undissociated compound, and more of the
solid dissolves. This process repeats itself with great rapidity until, if
sufficient acid is present, the solid passes completely into solution.

Exactly the same sort of process takes place if calcium oxalate, for
example, is dissolved in hydrochloric acid. The C_{2}O_{4}^{—} ions
unite with the H^{+} ions to form undissociated oxalic acid, the acid being
less dissociated than normally in the presence of the H^{+} ions from the
hydrochloric acid (see statements regarding hydrogen sulphide above). As
the undissociated oxalic acid forms, the concentration of the
C_{2}O_{4}^{—} ions lessens and more CaC_{2}O_{4} dissolves, as
described for the Mg(OH)_{2} above. Numerous instances of the
applications of these principles are given in the Notes.

Water itself is slightly dissociated, and although the resulting H^{+} and
OH^{-} ions are present only in minute concentrations (1 mol. of
dissociated water in 10^{7} liters), yet under some conditions they may
give rise to important consequences. The term !hydrolysis! is applied to the
changes which result from the reaction of these ions. Any salt which is
derived from a weak base or a weak acid (or both) is subject to hydrolytic
action. Potassium cyanide, for example, when dissolved in water gives an
alkaline solution because some of the H^{+} ions from the water unite with
CN^{-} ions to form (HCN), which is a very weak acid, and is but very
slightly dissociated. Potassium hydroxide, which might form from the
OH^{-} ions, is so largely dissociated that the OH^{-} ions remain as such
in the solution. The union of the H^{+} ions with the CN^{-} ions to form
the undissociated HCN diminishes the concentration of the H^{+} ions, and
more water dissociates (H_{2}O <—> H^{+} + OH^{-}) to restore the
equilibrium. It is clear, however, that there must be a gradual accumulation

of OH^{-} ions in the solution as a result of these changes, causing the
solution to exhibit an alkaline reaction, and also that ultimately the further
dissociation of the water will be checked by the presence of these ions, just
as the dissociation of the H_{2}S was lessened by the addition of HCl.

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

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

The impelling force which causes an element to pass from the atomic to
the ionic condition is termed !electrolytic solution pressure!, or ionization
tension. This force may be measured in terms of electrical potential, and the
table below shows the relative values for a number of elements.

In general, an element with a greater solution pressure tends to cause the
deposition of an element of less solution pressure when placed in a solution
of its salt, as, for instance, when a strip of zinc or iron is placed in a solution
of a copper salt, with the resulting precipitation of metallic copper.

Hydrogen is included in the table, and its position should be noted with
reference to the other common elements. For a more extended discussion of
this topic the student should refer to other treatises.

POTENTIAL SERIES OF THE METALS

___________________________________________________________
_______ | | | | POTENTIAL | | POTENTIAL | IN VOLTS | | IN VOLTS
_____________________|___________|____________________|_______
____ | | | Sodium Na^{+} | +2.44 | Lead Pb^{++} | -0.13 Calcium Ca^{++} |
| Hydrogen H^{+} | -0.28 Magnesium Mg^{++} | | Bismuth Bi^{+++}|
Aluminum A1^{+++} | +1.00 | Antimony | -0.75 Manganese Mn^{++} | |
Arsenic | Zinc Zn^{++} | +0.49 | Copper Cu^{++} | -0.61 Cadmium
Cd^{++} | +0.14 | Mercury Hg^{+} | -1.03 Iron Fe^{++} | +0.063 | Silver
Ag^{+} | -1.05 Cobalt Co^{++} | -0.045 | Platinum | Nickel Ni^{++} |
-0.049 | Gold | Tin Sn^{++} | -0.085(?) | |
_____________________|___________|____________________|_______
___

THE FOLDING OF A FILTER PAPER

If a filter paper is folded along its diameter, and again folded along the
radius at right angles to the original fold, a cone is formed on opening, the
angle of which is 60°. Funnels for analytical use are supposed to have the
same angle, but are rarely accurate. It is possible, however, with care, to fit
a filter thus folded into a funnel in such a way as to prevent air from passing
down between the paper and the funnel to break the column of liquid in the
stem, which aids greatly, by its gentle suction, in promoting the rate of
filtration.

Such a filter has, however, the disadvantage that there are three
thicknesses of paper back of half of its filtering surface, as a consequence of
which one half of a precipitate washes or drains more slowly. Much time
may be saved in the aggregate by learning to fold a filter in such a way as to
improve its effective filtering surface. The directions which follow, though

apparently complicated on first reading, are easily applied and easily
remembered. Use a 6-inch filter for practice. Place four dots on the filter,
two each on diameters which are at right angles to each other. Then proceed
as follows: (1) Fold the filter evenly across one of the diameters, creasing it
carefully; (2) open the paper, turn it over, rotate it 90° to the right, bring the
edges together and crease along the other diameter; (3) open, and rotate 45°
to the right, bring edges together, and crease evenly; (4) open, and rotate
90° to the right, and crease evenly; (5) open, turn the filter over, rotate 22-
(1/2)° to the right, and crease evenly; (6) open, rotate 45° to the right and
crease evenly; (7) open, rotate 45° to the right and crease evenly; (8) open,
rotate 45° to the right and crease evenly; (9) open the filter, and, starting
with one of the dots between thumb and forefinger of the right hand, fold
the second crease to the left over on it, and do the same with each of the
other dots. Place it, thus folded, in the funnel, moisten it, and fit to the side
of the funnel. The filter will then have four short segments where there are
three thicknesses and four where there is one thickness, but the latter are
evenly distributed around its circumference, thus greatly aiding the passage
of liquids through the paper and hastening both filtration and washing of the
whole contents of the filter.

!SAMPLE PAGES FOR LABORATORY RECORDS!

!Page A!

Date

CALIBRATION OF BURETTE No.

___________________________________________________________
________________
||||

BURETTE | DIFFERENCE | OBSERVED | DIFFERENCE |
CALCULATED
READINGS | | WEIGHTS | | CORRECTION
_______________|______________|______________|______________|__
____________
0.02 | | 16.27 | |
10.12 | 10.10 | 26.35 | 10.08 | -.02
20.09 | 9.97 | 36.26 | 9.91 | -.06
30.16 | 10.07 | 46.34 | 10.08 | +.01
40.19 | 10.03 | 56.31 | 9.97 | -.06
50.00 | 9.81 | 66.17 | 9.86 | +.05
_______________|______________|______________|______________|__
____________

These data to be obtained in duplicate for each burette.

!Page B!

Date

DETERMINATION OF COMPARATIVE STRENGTH HCl vs. NaOH

___________________________________________________________
________________
||
DETERMINATION | I | II
_________________________|________________________|___________
_____________
||
| Corrected | Corrected

Final Reading HCl | 48.17 48.08 | 43.20 43.14
Initial Reading HCl | 0.12 .12 | .17 .17
| ——- ——- | ——- ——-
| 47.96 | 42.97
||
| Corrected | Corrected
Final Reading HCl | 46.36 46.29 | 40.51 40.37
Initial Reading HCl | 1.75 1.75 | .50 .50
| ——- ——- | ——- ——-
| 44.54 | 39.87
||
log cc. NaOH | 1.6468 | 1.6008
colog cc. HCl | 8.3192 | 8.3668
| ——— | ———
| 9.9680 - 10 | 9.9676 - 10
1 cc. HCl | .9290 cc. NaOH | .9282 cc. NaOH
Mean | .9286 |
_________________________|________________________|___________
_____________

Signed

!Page C!
Date

STANDARDIZATION OF HYDROCHLORIC ACID
======================================================
===============
||
Weight sample and tube| 9.1793 | 8.1731

| 8.1731 | 6.9187
| ——— | ———
Weight sample | 1.0062 | 1.2544
||
Final Reading HCl | 39.97 39.83 | 49.90 49.77
Initial Reading HCl | .00 .00 | .04 .04
| ——- ——- | ——- ——-
| 39.83 | 49.73
||
Final Reading NaOH | .26 .26 | .67 .67
Initial Reading NaOH | .12 .12 | .36 .36
| —- —- | —- —-
| .14 | .31
||
| .14 | .31
Corrected cc. HCl | 39.83 - ——- = 39.68 | 49.73 - ——- = 49.40
| .9286 | .9286
||
log sample | 0.0025 | 0.0983
colog cc | 8.4014 - 10 | 8.3063 - 10
colog milli equivalent| 1.2757 | 1.2757
| ——— | ———
| 9.6796 - 10 | 9.6803 - 10
||
Normal value HCl | .4782 | .4789
Mean | .4786 |
||
======================================================
===============

Signed

!Page D!
Date

DETERMINATION OF CHLORINE IN CHLORIDE, SAMPLE No.
======================================================
===============
||
Weight sample and tube| 16.1721 | 15.9976
| 15.9976 | 15.7117
| ———- | ———-
Weight sample | .1745 | .2859
||
Weight crucible | |
+ precipitate | 14.4496 | 15.6915
Constant weights | 14.4487 | 15.6915
| 14.4485 |
||
Weight crucible | 14.2216 | 15.3196
Constant weight | 14.2216 | 15.3194
||
Weight AgCl | .2269 | .3721
||
log Cl | 1.5496 | 1.5496
log weight AgCl | 9.3558 - 10 | 9.5706 - 10
log 100 | 2.0000 | 2.0000
colog AgCl | 7.8438 - 10 | 7.7438 - 10
colog sample | 0.7583 | 0.5438
| ———- | ———-
| 1.5075 | 1.5078
||

Cl in sample No. | 32.18% | 32.20%
||
======================================================
===============

Signed

STRENGTH OF REAGENTS

The concentrations given in this table are those suggested for use in the
procedures described in the foregoing pages. It is obvious, however, that an
exact adherence to these quantities is not essential.

Approx. Approx.
Grams relation relation
per to normal to molal
liter. solution solution

Ammonium oxalate, (NH_{4}){2}C{2}O_{4}.H_{2}O 40 0.5N 0.25
Barium chloride, BaCl_{2}.2H_{2}O 25 0.2N 0.1
Magnesium ammonium chloride (of MgCl_{2}) 71 1.5N 0.75
Mercuric chloride, HgCl_{2} 45 0.33N 0.66
Potassium hydroxide, KOH (sp. gr. 1.27) 480
Potassium thiocyanate, KSCN 5 0.05N 0.55
Silver nitrate, AgNO_{3} 21 0.125N 0.125
Sodium hydroxide, NaOH 100 2.5N 2.5
Sodium carbonate. Na_{2}CO_{3} 159 3N 1.5
Sodium phosphate, Na_{2}HPO_{4}.12H_{2}O 90 0.5N or 0.75N 0.25

Stannous chloride, SnCl_{2}, made by saturating hydrochloric acid (sp.
gr. 1.2) with tin, diluting with an equal volume of water, and adding a slight

excess of acid from time to time. A strip of metallic tin is kept in the bottle.

A solution of ammonium molybdate is best prepared as follows: Stir 100
grams of molybdic acid (MoO_{3}) into 400 cc. of cold, distilled water.
Add 80 cc. of concentrated ammonium hydroxide (sp. gr. 0.90). Filter, and
pour the filtrate slowly, with constant stirring, into a mixture of 400 cc.
concentrated nitric acid (sp. gr. 1.42) and 600 cc. of water. Add to the
mixture about 0.05 gram of microcosmic salt. Filter, after allowing the
whole to stand for 24 hours.

The following data regarding the common acids and aqueous ammonia
are based upon percentages given in the Standard Tables of the
Manufacturing Chemists' Association of the United States [!J.S.C.I.!, 24
(1905), 787-790]. All gravities are taken at 15.5°C. and compared with
water at the same temperature.

Aqueous ammonia (sp. gr. 0.96) contains 9.91 per cent NH_{3} by
weight, and corresponds to a 5.6 N and 5.6 molal solution.

Aqueous ammonia (sp. gr. 0.90) contains 28.52 per cent NH_{3} by
weight, and corresponds to a 5.6 N and 5.6 molal solution.

Hydrochloric acid (sp. gr. 1.12) contains 23.81 per cent HCl by weight,
and corresponds to a 7.3 N and 7.3 molal solution.

Hydrochloric acid (sp. gr. 1.20) contains 39.80 per cent HCl by weight,
and corresponds to a 13.1 N and 13.1 molal solution.

Nitric acid (sp. gr. 1.20) contains 32.25 per cent HNO_{3} by weight,
and corresponds to a 6.1 N and 6.1 molal solution:

Nitric acid (sp. gr. 1.42) contains 69.96 per cent HNO_{3} by weight,
and corresponds to a 15.8 N and 15.8 molal solution.

Sulphuric acid (sp. gr. 1.8354) contains 93.19 per cent H_{2}SO_{4} by
weight, and corresponds to a 34.8 N or 17.4 molal solution.

Sulphuric acid (sp. gr. 1.18) contains 24.74 per cent H_{2}SO_{4} by
weight, and corresponds to a 5.9 N or 2.95 molal solution.

The term !normal! (N), as used above, has the same significance as in
volumetric analyses. The molal solution is assumed to contain one
molecular weight in grams in a liter of solution.

DENSITIES AND VOLUMES OF WATER AT TEMPERATURES FROM 15-30°C.

Temperature Density. Volume.
Centigrade.

4° 1.000000 1.000000 15° 0.999126 1.000874 16° 0.998970
1.001031 17° 0.998801 1.001200 18° 0.998622 1.001380 19°
0.998432 1.001571 20° 0.998230 1.001773 21° 0.998019
1.001985 22° 0.997797 1.002208 23° 0.997565 1.002441 24°
0.997323 1.002685 25° 0.997071 1.002938 26° 0.996810
1.003201 27° 0.996539 1.003473 28° 0.996259 1.003755 29°
0.995971 1.004046 30° 0.995673 1.004346

Authority: Landolt, Börnstein, and Meyerhoffer's !Tabellen!, third
edition.

CORRECTIONS FOR CHANGE OF TEMPERATURE OF
STANDARD SOLUTIONS

The values below are average values computed from data relating to a
considerable number of solutions. They are sufficiently accurate for use in
chemical analyses, except in the comparatively few cases where the highest

attainable accuracy is demanded in chemical investigations. The expansion
coefficients should then be carefully determined for the solutions employed.
For a compilation of the existing data, consult Landolt, Börnstein, and
Meyerhoffer's !Tabellen!, third edition.

Corrections for 1 cc.
Concentration. of solution between
15° and 35°C.

Normal .00029
0.5 Normal .00025
0.1 Normal or more dilute solutions .00020

The volume of solution used should be multiplied by the values given,
and that product multiplied by the number of degrees which the temperature
of the solution varies from the standard temperature selected for the
laboratory. The total correction thus found is subtracted from the observed
burette reading if the temperature is higher than the standard, or added, if it
is lower. Corrections are not usually necessary for variations of temperature
of 2°C. or less.

INTERNATIONAL ATOMIC WEIGHTS

======================================================
====
|||
| 1920 | | 1920
_________________|_________|___________________|__________
|||
Aluminium Al | 27.1 | Molybdenum Mo | 96.0
Antimony Sb | 120.2 | Neodymium Nd | 144.3
Argon A | 39.9 | Neon Ne | 20.2
Arsenic As | 74.96 | Nickel Ni | 58.68
Barium Ba | 137.37 | Nitrogen N | 14.008
Bismuth Bi | 208.0 | Osmium Os | 190.9
Boron B | 11.0 | Oxygen O | 16.00
Bromine Br | 79.92 | Palladium Pd | 106.7
Cadmium Cd | 112.40 | Phosphorus P | 31.04
Caesium Cs | 132.81 | Platinum Pt | 195.2
Calcium Ca | 40.07 | Potassium K | 39.10
Carbon C | 12.005 | Praseodymium Pr | 140.9
Cerium Ce | 140.25 | Radium Ra | 226.0
Chlorine Cl | 35.46 | Rhodium Rh | 102.9
Chromium Cr | 52.0 | Rubidium Rb | 85.45
Cobalt Co | 58.97 | Ruthenium Ru | 101.7
Columbium Cb | 93.1 | Samarium Sm | 150.4
Copper Cu | 63.57 | Scandium Sc | 44.1
Dysprosium Dy | 162.5 | Selenium Se | 79.2
Erbium Er | 167.7 | Silicon Si | 28.3

Europium Eu | 152.0 | Silver Ag | 107.88
Fluorine Fl | 19.0 | Sodium Na | 23.00
Gadolinium Gd | 157.3 | Strontium Sr | 87.63
Gallium Ga | 69.9 | Sulphur S | 32.06
Germanium Ge | 72.5 | Tantalum Ta | 181.5
Glucinum Gl | 9.1 | Tellurium Te | 127.5
Gold Au | 197.2 | Terbium Tb | 159.2
Helium He | 4.00 | Thallium Tl | 204.0
Hydrogen H | 1.008 | Thorium Th | 232.4
Indium In | 114.8 | Thulium Tm | 168.5
Iodine I | 126.92 | Tin Sn | 118.7
Iridium Ir | 193.1 | Titanium Ti | 48.1
Iron Fe | 55.84 | Tungsten W | 184.0
Krypton Kr | 82.92 | Uranium U | 238.2
Lanthanum La | 139.0 | Vanadium V | 51.0
Lead Pb | 207.2 | Xenon Xe | 130.2
Lithium Li | 6.94 | Ytterbium Yb | 173.5
Lutecium Lu | 175.0 | Yttrium Y | 88.7
Magnesium Mg | 24.32 | Zinc Zn | 65.37
Manganese Mn | 54.93 | Zirconium Zr | 90.6
Mercury Hg | 200.6 | |
======================================================
====

INDEX

Acidimetry
Acid solutions, normal
standard
Acids, definition of
Acids, weak, action of other acids on
action of salts on
Accuracy demanded
Alkalimetry
Alkali solutions, normal
standard
Alumina, determination of in stibnite
Ammonium nitrate, acid
Analytical chemistry, subdivisions of
Antimony, determination of, in stibnite
Apatite, analysis of
Asbestos filters
Atomic weights, table of

Balances, essential features of use and care of Barium sulphate,
determination of sulphur in Bases, definition of Bichromate process for iron
Bleaching powder, analysis of Brass, analysis of Burette, description of
calibration of cleaning of reading of

Calcium, determination of, in limestone
Calibration, definition of
of burettes
of flasks
Carbon dioxide, determination of, in limestone
Chlorimetry
Chlorine, gravimetric determination of
Chrome iron ore, analysis of
Coin, determination of silver in
Colloidal solution of precipitates
Colorimetric analyses, definition of
Copper, determination of, in brass
determination of in copper ores
Crucibles, use of
Crystalline precipitates

Densities of water
Deposition potentials
Desiccators
Direct methods
Dissociation, degree of

Economy of time
Electrolytic dissociation, theory of
Electrolytic separations, principles of
End-point, definition of
Equilibrium, chemical
Evaporation of liquids

Faraday's law
Feldspar, analysis of

Ferrous ammonium sulphate, analysis of
Filters, folding of
how fitted
Filtrates, testing of
Filtration
Flasks, graduation of
Funnels
Fusions, removal of from crucibles

General directions for gravimetric analysis
volumetric analysis
Gooch filter
Gravimetric analysis, definition of

Hydrochloric acid, standardization of
Hydrolysis

Ignition of precipitates Indicators, definition of for acidimetry
preparation of Indirect methods Insoluble matter, determination of in
limestone Integrity Iodimetry Ions, definition of Iron, gravimetric
determination of volumetric determination of

Jones reductor

Lead, determination of in brass
Limestone, analysis of
Limonite, determination of iron in
Liquids, evaporation of
transfer of
Litmus
Logarithms

Magnesium, determination of
Mass action, law of
Measuring instruments
Methyl orange
Moisture, determination of in limestone

Neutralization methods Normal solutions, acid and alkali oxidizing
agents reducing agents Notebooks, sample pages of

Oxalic acid, determination of strength of
Oxidation processes
Oxidizing power of pyrolusite

Permanganate process for iron Phenolphthalein Phosphoric anhydride,
determination of Pipette, calibration of description of Platinum crucibles,
care of Precipitates, colloidal crystalline ignition of separation from filter
washing of Precipitation Precipitation methods (volumetric) Problems
Pyrolusite, oxidizing power of

Quantitative Analyses, subdivisions of

Reagents, strength of
Reducing solution, normal
Reductor, Jones
Reversible reactions

Silica, determination of, in limestone
determination of, in silicates
purification of
Silicic acid, dehydration of
Silver, determination of in coin
Soda ash, alkaline strength of

Sodium chloride, determination of chlorine in
Solubility product
Solution pressure
Solutions, normal
standard
Standardization, definition of
Standard solutions, acidimetry and alkalimetry
chlorimetry
iodimetry
oxidizing and reducing agents
thiocyanate
Starch solutions
Stibnite, determination of antimony in
Stirring rods
Stoichiometry
Strength of reagents
Suction, use of
Sulphur, determination of in ferrous ammonium sulphate
in barium sulphate

Temperature, corrections for
Testing of washings
Theory of electrolytic dissociation
Thiocyanate process for silver
Titration, definition of
Transfer of liquids

Volumetric analysis, definition of
general directions

Wash-bottles
Washed filters
Washing of precipitates
Washings, testing of
Water, ionization of
densities of
Weights, care of

Zimmermann-Reinhardt method for iron
Zinc, determination of, in brass

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