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Title: An Introductory Course of Quantitative Chemical Analysis - With Explanatory Notes
Author: Talbot, Henry P.
Language: English
As this book started as an ASCII text book there are no pictures available.

*** Start of this LibraryBlog Digital Book "An Introductory Course of Quantitative Chemical Analysis - With Explanatory Notes" ***

[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.]












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.


!Massachusetts Institute of Technology, September, 1921!.




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


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

!I. Neutralization Methods!

  Preparation and Standardization of Solutions; Indicators

!II. Oxidation Processes!


!III. Precipitation Methods!



  Precipitation; Funnels and Filters; Filtration and Washing of
  Precipitates; Desiccators; Crucibles and their Preparation
  for Use; Ignition of Precipitates
  Determination of Moisture; Insoluble Matter and Silica; Ferric
  Oxide and Alumina; Calcium; Magnesium; Carbon Dioxide
  Electrolytic Separations; Determination of Lead, Copper, Iron
  and Zinc.









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

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.


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.


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 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.


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 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


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.


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.



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

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.



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

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

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

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

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.


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.


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.


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.]


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.]


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.


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

[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

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.]


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.


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.]


!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.


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.]


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.


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!.


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

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

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.




!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

!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

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.


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.]


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

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

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

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),

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

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.]


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.


!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

[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


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

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

[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.]



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

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.


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

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

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.]


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.]


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.]



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.


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

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


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.


!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.


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.


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

[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.]


!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.


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

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


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.]


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.]


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

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.


!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.]


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.


!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.]


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

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

!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.]


!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.]



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

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

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

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.]


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.


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.


!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.]


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

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.]


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


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.]


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

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

[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.]


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

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

[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 +

[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.]


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

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

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.


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

[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.]



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.


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

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.]


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.]




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.


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.


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.


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

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!

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 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


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

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

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.


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

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

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.


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.


!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.)


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

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

[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.]




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

[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

[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

[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

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

[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

[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}).]


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

[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.]


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


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

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


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

[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.]


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

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


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

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

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


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.


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


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

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.]


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

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


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

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.]


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

[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.]


!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

[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.]



!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

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

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.


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

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

[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

[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

[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.]


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.


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

[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),


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.


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.]


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

[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.]


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

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.]


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



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

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

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

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

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

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

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.


(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


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

!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

!Answers!: (a) NaHCO_{3}; (b) NaHCO_{3}+Na_{2}CO_{3}; (c)NaOH; (d)

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

!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

!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)

!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 =

!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

!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

!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.


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

!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

!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

!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

!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

!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

!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

!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

!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.



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.

                 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

                 SUBSTANCE                   | PERCENTAGE DISSOCIATION IN
                                             |  0.1 EQUIVALENT SOLUTION
KOH, NaOH                                    |            86
Ba(OH)_{2}                                   |            75
NH_{4}OH                                     |             1.4

               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} =

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

              |          |                  |                  |
  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 |                    | 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(?) |                    |


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

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.


!Page A!



               |              |              |              |
    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!



                         |                        |
       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      |


!Page C!

                      |                       |
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   |
                      |                       |


!Page D!

                      |                       |
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%
                      |                       |



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.


Temperature           Density.        Volume.

     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


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.


                 |         |                   |
                 |  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   |                   |


Acid solutions, normal
Acids, definition of
Acids, weak, action of other acids on
  action of salts on
Accuracy demanded
Alkali solutions, normal
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
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
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
Flasks, graduation of
Fusions, removal of from crucibles

General directions for gravimetric analysis
  volumetric analysis
Gooch filter
Gravimetric analysis, definition of

Hydrochloric acid, standardization of

Ignition of precipitates
Indicators, definition of
  for acidimetry
  preparation of
Indirect methods
Insoluble matter, determination of in limestone
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

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
Phosphoric anhydride, determination of
Pipette, calibration of
  description of
Platinum crucibles, care of
Precipitates, colloidal
  ignition of
  separation from filter
  washing of
Precipitation methods (volumetric)
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
Standardization, definition of
Standard solutions, acidimetry and alkalimetry
  oxidizing and reducing agents
Starch solutions
Stibnite, determination of antimony in
Stirring rods
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

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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