An Introductory Course of Quantitative Chemical Analysis
by Henry Paul Talbot · English · Gutenberg #12787 · 420,953 chars · public domain (US)
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Title: An Introductory Course of Quantitative Chemical Analysis
Author: Henry Paul Talbot
Release date: June 30, 2004 [eBook #12787]
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Language: English
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[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
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AN INTRODUCTORY COURSE
OF
QUANTITATIVE
CHEMICAL ANALYSIS
WITH
EXPLANATORY NOTES
BY
HENRY P. TALBOT
PROFESSOR OF INORGANIC CHEMISTRY AT THE MASSACHUSETTS INSTITUTE OF
TECHNOLOGY
SIXTH EDITION, COMPLETELY REWRITTEN
PREFACE
This Introductory Course of Quantitative Analysis has been prepared
to meet the needs of students who are just entering upon the subject,
after a course of qualitative analysis. It is primarily intended to
enable the student to work successfully and intelligently without the
necessity for a larger measure of personal assistance and supervision
than can reasonably be given to each member of a large class. To this
end the directions are given in such detail that there is very little
opportunity for the student to go astray; but the manual is not, the
author believes, on this account less adapted for use with small
classes, where the instructor, by greater personal influence, can
stimulate independent thought on the part of the pupil.
The method of presentation of the subject is that suggested by
Professor A.A. Noyes' excellent manual of Qualitative Analysis. For
each analysis the procedure is given in considerable detail, and
this is accompanied by explanatory notes, which are believed to be
sufficiently expanded to enable the student to understand fully the
underlying reason for each step prescribed. The use of the book
should, nevertheless, be supplemented by classroom instruction, mainly
of the character of recitations, and the student should be taught to
consult larger works. The general directions are intended to emphasize
those matters upon which the beginner in quantitative analysis must
bestow special care, and to offer helpful suggestions. The student
can hardly be expected to appreciate the force of all the statements
contained in these directions, or, indeed, to retain them all in
the memory after a single reading; but the instructor, by frequent
reference to special paragraphs, as suitable occasion presents itself,
can soon render them familiar to the student.
The analyses selected for practice are those comprised in the first
course of quantitative analysis at the Massachusetts Institute of
Technology, and have been chosen, after an experience of years,
as affording the best preparation for more advanced work, and as
satisfactory types of gravimetric and volumetric methods. From the
latter point of view, they also seem to furnish the best insight into
quantitative analysis for those students who can devote but a limited
time to the subject, and who may never extend their study beyond the
field covered by this manual. The author has had opportunity to test
the efficiency of the course for use with such students, and has found
the results satisfactory.
In place of the usual custom of selecting simple salts as material for
preliminary practice, it has been found advantageous to substitute, in
most instances, approximately pure samples of appropriate minerals or
industrial products. The difficulties are not greatly enhanced, while
the student gains in practical experience.
The analytical procedures described in the following pages have been
selected chiefly with reference to their usefulness in teaching the
subject, and with the purpose of affording as wide a variety of
processes as is practicable within an introductory course of this
character. The scope of the manual precludes any extended attempt to
indicate alternative procedures, except through general references to
larger works on analytical chemistry. The author is indebted to the
standard works for many suggestions for which it is impracticable to
make specific acknowledgment; no considerable credit is claimed by him
for originality of procedure.
For many years, as a matter of convenience, the classes for which this
text was originally prepared were divided, one part beginning with
gravimetric processes and the other with volumetric analyses. After a
careful review of the experience thus gained the conclusion has been
reached that volumetric analysis offers the better approach to the
subject. Accordingly the arrangement of the present (the sixth)
edition of this manual has been changed to introduce volumetric
procedures first. Teachers who are familiar with earlier editions
will, however, find that the order of presentation of the material
under the various divisions is nearly the same as that previously
followed, and those who may still prefer to begin the course of
instruction with gravimetric processes will, it is believed, be able
to follow that order without difficulty.
Procedures for the determination of sulphur in insoluble sulphates,
for the determination of copper in copper ores by iodometric methods,
for the determination of iron by permanganate in hydrochloric acid
solutions, and for the standardization of potassium permanganate
solutions using sodium oxalate as a standard, and of thiosulphate
solutions using copper as a standard, have been added. The
determination of silica in silicates decomposable by acids, as a
separate procedure, has been omitted.
The explanatory notes have been rearranged to bring them into closer
association with the procedures to which they relate. The number of
problems has been considerably increased.
The author wishes to renew his expressions of appreciation of the
kindly reception accorded the earlier editions of this manual. He has
received helpful suggestions from so many of his colleagues within the
Institute, and friends elsewhere, that his sense of obligation must
be expressed to them collectively. He is under special obligations
to Professor L.F. Hamilton for assistance in the preparation of the
present edition.
HENRY P. TALBOT
!Massachusetts Institute of Technology, September, 1921!.
CONTENTS
PART I. INTRODUCTION
SUBDIVISIONS OF ANALYTICAL CHEMISTRY
GENERAL DIRECTIONS
Accuracy and Economy of Time; Notebooks; Reagents; Wash-bottles;
Transfer of Liquids
PART II. VOLUMETRIC ANALYSIS
GENERAL DISCUSSION
Subdivisions; The Analytical Balance; Weights; Burettes;
Calibration of Measuring Devices
GENERAL DIRECTIONS
Standard and Normal Solutions
!I. Neutralization Methods!
ALKALIMETRY AND ACIDIMETRY
Preparation and Standardization of Solutions; Indicators
STANDARDIZATION OF HYDROCHLORIC ACID
DETERMINATION OF TOTAL ALKALINE STRENGTH OF SODA ASH
DETERMINATION OF ACID STRENGTH OF OXALIC ACID
!II. Oxidation Processes!
GENERAL DISCUSSION
BICHROMATE PROCESS FOR THE DETERMINATION OF IRON
DETERMINATION OF IRON IN LIMONITE BY THE BICHROMATE PROCESS
DETERMINATION OF CHROMIUM IN CHROME IRON ORE
PERMANGANATE PROCESS FOR THE DETERMINATION OF IRON
DETERMINATION OF IRON IN LIMONITE BY THE PERMANGANATE PROCESS
DETERMINATION OF IRON IN LIMONITE BY THE ZIMMERMANN-REINHARDT PROCESS
DETERMINATION OF THE OXIDIZING POWER OF PYROLUSITE
IODIMETRY
DETERMINATION OF COPPER IN ORES
DETERMINATION OF ANTIMONY IN STIBNITE
CHLORIMETRY
DETERMINATION OF AVAILABLE CHLORINE IN BLEACHING POWDER
!III. Precipitation Methods!
DETERMINATION OF SILVER BY THE THIOCYANATE PROCESS
PART III. GRAVIMETRIC ANALYSIS
GENERAL DIRECTIONS
Precipitation; Funnels and Filters; Filtration and Washing of
Precipitates; Desiccators; Crucibles and their Preparation
for Use; Ignition of Precipitates
DETERMINATION OF CHLORINE IN SODIUM CHLORIDE
DETERMINATION OF IRON AND OF SULPHUR IN FERROUS AMMONIUM SULPHATE
DETERMINATION OF SULPHUR IN BARIUM SULPHATE
DETERMINATION OF PHOSPHORIC ANHYDRIDE IN APATITE
ANALYSIS OF LIMESTONE
Determination of Moisture; Insoluble Matter and Silica; Ferric
Oxide and Alumina; Calcium; Magnesium; Carbon Dioxide
ANALYSIS OF BRASS
Electrolytic Separations; Determination of Lead, Copper, Iron
and Zinc.
DETERMINATION OF SILICA IN SILICATES
PART IV. STOICHIOMETRY
SOLUTIONS OF TYPICAL PROBLEMS
PROBLEMS
APPENDIX
ELECTROLYTIC DISSOCIATION THEORY
FOLDING OF A FILTER PAPER
SAMPLE NOTEBOOK PAGES
STRENGTH OF REAGENTS
DENSITIES AND VOLUMES OF WATER
CORRECTIONS FOR CHANGE OF TEMPERATURE OF STANDARD SOLUTIONS
ATOMIC WEIGHTS
LOGARITHM TABLES
QUANTITATIVE CHEMICAL ANALYSIS
PART I
INTRODUCTION
SUBDIVISIONS OF ANALYTICAL CHEMISTRY
A complete chemical analysis of a body of unknown composition involves
the recognition of its component parts by the methods of !qualitative
analysis!, and the determination of the proportions in which these
components are present by the processes of !quantitative analysis!.
A preliminary qualitative examination is generally indispensable, if
intelligent and proper provisions are to be made for the separation of
the various constituents under such conditions as will insure accurate
quantitative estimations.
It is assumed that the operations of qualitative analysis are familiar
to the student, who will find that the reactions made use of in
quantitative processes are frequently the same as those employed in
qualitative analyses with respect to both precipitation and systematic
separation from interfering substances; but it should be noted that
the conditions must now be regulated with greater care, and in such
a manner as to insure the most complete separation possible. For
example, in the qualitative detection of sulphates by precipitation
as barium sulphate from acid solution it is not necessary, in most
instances, to take into account the solubility of the sulphate
in hydrochloric acid, while in the quantitative determination of
sulphates by this reaction this solubility becomes an important
consideration. The operations of qualitative analysis are, therefore,
the more accurate the nearer they are made to conform to quantitative
conditions.
The methods of quantitative analysis are subdivided, according
to their nature, into those of !gravimetric analysis, volumetric
analysis!, and !colorimetric analysis!. In !gravimetric! processes the
constituent to be determined is sometimes isolated in elementary
form, but more commonly in the form of some compound possessing a
well-established and definite composition, which can be readily and
completely separated, and weighed either directly or after ignition.
From the weight of this substance and its known composition, the
amount of the constituent in question is determined.
In !volumetric! analysis, instead of the final weighing of a definite
body, a well-defined reaction is caused to take place, wherein the
reagent is added from an apparatus so designed that the volume of the
solution employed to complete the reaction can be accurately measured.
The strength of this solution (and hence its value for the reaction
in question) is accurately known, and the volume employed serves,
therefore, as a measure of the substance acted upon. An example will
make clear the distinction between these two types of analysis.
The percentage of chlorine in a sample of sodium chloride may be
determined by dissolving a weighed amount of the chloride in water
and precipitating the chloride ions as silver chloride, which is
then separated by filtration, ignited, and weighed (a !gravimetric!
process); or the sodium chloride may be dissolved in water, and a
solution of silver nitrate containing an accurately known amount of
the silver salt in each cubic centimeter may be cautiously added from
a measuring device called a burette until precipitation is complete,
when the amount of chlorine may be calculated from the number of cubic
centimeters of the silver nitrate solution involved in the reaction.
This is a !volumetric! process, and is equivalent to weighing without
the use of a balance.
Volumetric methods are generally more rapid, require less apparatus,
and are frequently capable of greater accuracy than gravimetric
methods. They are particularly useful when many determinations of the
same sort are required.
In !colorimetric! analyses the substance to be determined is converted
into some compound which imparts to its solutions a distinct color,
the intensity of which must vary in direct proportion to the amount of
the compound in the solution. Such solutions are compared with respect
to depth of color with standard solutions containing known amounts of
the colored compound, or of other similar color-producing substance
which has been found acceptable as a color standard. Colorimetric
methods are, in general, restricted to the determinations of very
small quantities, since only in dilute solutions are accurate
comparisons of color possible.
GENERAL DIRECTIONS
The following paragraphs should be read carefully and thoughtfully. A
prime essential for success as an analyst is attention to details and
the avoidance of all conditions which could destroy, or even lessen,
confidence in the analyses when completed. The suggestions here given
are the outcome of much experience, and their adoption will tend to
insure permanently work of a high grade, while neglect of them will
often lead to disappointment and loss of time.
ACCURACY AND ECONOMY OF TIME
The fundamental conception of quantitative analysis implies a
necessity for all possible care in guarding against loss of material
or the introduction of foreign matter. The laboratory desk, and all
apparatus, should be scrupulously neat and clean at all times. A
sponge should always be ready at hand, and desk and filter-stands
should be kept dry and in good order. Funnels should never be allowed
to drip upon the base of the stand. Glassware should always be
wiped with a clean, lintless towel just before use. All filters and
solutions should be covered to protect them from dust, just as far as
is practicable, and every drop of solution or particle of precipitate
must be regarded as invaluable for the success of the analysis.
An economical use of laboratory hours is best secured by acquiring
a thorough knowledge of the character of the work to be done before
undertaking it, and then by so arranging the work that no time shall
be wasted during the evaporation of liquids and like time-consuming
operations. To this end the student should read thoughtfully not only
the !entire! procedure, but the explanatory notes as well, before
any step is taken in the analysis. The explanatory notes furnish, in
general, the reasons for particular steps or precautions, but they
also occasionally contain details of manipulation not incorporated,
for various reasons, in the procedure. These notes follow the
procedures at frequent intervals, and the exact points to which they
apply are indicated by references. The student should realize that a
!failure to study the notes will inevitably lead to mistakes, loss of
time, and an inadequate understanding of the subject!.
All analyses should be made in duplicate, and in general a close
agreement of results should be expected. It should, however, be
remembered that a close concordance of results in "check analyses" is
not conclusive evidence of the accuracy of those results, although the
probability of their accuracy is, of course, considerably enhanced.
The satisfaction in obtaining "check results" in such analyses must
never be allowed to interfere with the critical examination of the
procedure employed, nor must they ever be regarded as in any measure a
substitute for absolute truth and accuracy.
In this connection it must also be emphasized that only the operator
himself can know the whole history of an analysis, and only he can
know whether his work is worthy of full confidence. No work should be
continued for a moment after such confidence is lost, but should
be resolutely discarded as soon as a cause for distrust is fully
established. The student should, however, determine to put forth his
best efforts in each analysis; it is well not to be too ready to
condone failures and to "begin again," as much time is lost in these
fruitless attempts. Nothing less than !absolute integrity! is or can
be demanded of a quantitative analyst, and any disregard of this
principle, however slight, is as fatal to success as lack of chemical
knowledge or inaptitude in manipulation can possibly be.
NOTEBOOKS
Notebooks should contain, beside the record of observations,
descriptive notes. All records of weights should be placed upon the
right-hand page, while that on the left is reserved for the notes,
calculations of factors, or the amount of reagents required.
The neat and systematic arrangement of the records of analyses is
of the first importance, and is an evidence of careful work and an
excellent credential. Of two notebooks in which the results may be,
in fact, of equal value as legal evidence, that one which is neatly
arranged will carry with it greater weight.
All records should be dated, and all observations should be recorded
at once in the notebook. The making of records upon loose paper is a
practice to be deprecated, as is also that of copying original entries
into a second notebook. The student should accustom himself to orderly
entries at the time of observation. Several sample pages of systematic
records are to be found in the Appendix. These are based upon
experience; but other arrangements, if clear and orderly, may prove
equally serviceable. The student is advised to follow the sample pages
until he is in a position to plan out a system of his own.
REAGENTS
The habit of carefully testing reagents, including distilled water,
cannot be too early acquired or too constantly practiced; for, in
spite of all reasonable precautionary measures, inferior chemicals
will occasionally find their way into the stock room, or errors will
be made in filling reagent bottles. The student should remember that
while there may be others who share the responsibility for the purity
of materials in the laboratory of an institution, the responsibility
will later be one which he must individually assume.
The stoppers of reagent bottles should never be laid upon the desk,
unless upon a clean watch-glass or paper. The neck and mouth of all
such bottles should be kept scrupulously clean, and care taken that no
confusion of stoppers occurs.
WASH-BOTTLES
Wash-bottles for distilled water should be made from flasks of about
750 cc. capacity and be provided with gracefully bent tubes, which
should not be too long. The jet should be connected with the tube
entering the wash-bottle by a short piece of rubber tubing in such
a way as to be flexible, and should deliver a stream about one
millimeter in diameter. The neck of the flask may be wound with cord,
or covered with wash-leather, for greater comfort when hot water is
used. It is well to provide several small wash-bottles for liquids
other than distilled water, which should invariably be clearly
labeled.
TRANSFER OF LIQUIDS
Liquids should never be transferred from one vessel to another, nor to
a filter, without the aid of a stirring rod held firmly against the
side or lip of the vessel. When the vessel is provided with a lip it
is not usually necessary to use other means to prevent the loss of
liquid by running down the side; whenever loss seems imminent a !very
thin! layer of vaseline, applied with the finger to the edge of the
vessel, will prevent it. The stirring rod down which the liquid runs
should never be drawn upward in such a way as to allow the solution to
collect on the under side of the rim or lip of a vessel.
The number of transfers of liquids from one vessel to another during
an analysis should be as small as possible to avoid the risk of slight
losses. Each vessel must, of course, be completely washed to insure
the transfer of all material; but it should be remembered that this
can be accomplished better by the use of successive small portions of
wash-water (perhaps 5-10 cc.), if each wash-water is allowed to drain
away for a few seconds, than by the addition of large amounts which
unnecessarily increase the volume of the solutions, causing loss of
time in subsequent filtrations or evaporations.
All stirring rods employed in quantitative analyses should be rounded
at the ends by holding them in the flame of a burner until they begin
to soften. If this is not done, the rods will scratch the inner
surface of beakers, causing them to crack on subsequent heating.
EVAPORATION OF LIQUIDS
The greatest care must be taken to prevent loss of solutions during
processes of evaporation, either from too violent ebullition, from
evaporation to dryness and spattering, or from the evolution of gas
during the heating. In general, evaporation upon the steam bath is to
be preferred to other methods on account of the impossibility of
loss by spattering. If the steam baths are well protected from dust,
solutions should be left without covers during evaporation; but
solutions which are boiled upon the hot plate, or from which gases are
escaping, should invariably be covered. In any case a watch-glass may
be supported above the vessel by means of a glass triangle, or other
similar device, and the danger of loss of material or contamination by
dust thus be avoided. It is obvious that evaporation is promoted by
the use of vessels which admit of the exposure of a broad surface to
the air.
Liquids which contain suspended matter (precipitates) should always
be cautiously heated, since the presence of the solid matter is
frequently the occasion of violent "bumping," with consequent risk to
apparatus and analysis.
PART II
VOLUMETRIC ANALYSIS
The processes of volumetric analysis are, in general, simpler than
those of gravimetric analysis and accordingly serve best as an
introduction to the practice of quantitative analysis. For their
execution there are required, first, an accurate balance with which
to weigh the material for analysis; second, graduated instruments in
which to measure the volume of the solutions employed; third, standard
solutions, that is, solutions the value of which is accurately known;
and fourth, indicators, which will furnish accurate evidence of the
point at which the desired reaction is completed. The nature of the
indicators employed will be explained in connection with the different
analyses.
The process whereby a !standard solution! is brought into reaction is
called !titration!, and the point at which the reaction is exactly
completed is called the !end-point!. The !indicator! should show the
!end-point! of the !titration!. The volume of the standard solution
used then furnishes the measure of the substance to be determined as
truly as if that substance had been separated and weighed.
The processes of volumetric analysis are easily classified, according
to their character, into:
I. NEUTRALIZATION METHODS; such, for example, as those of acidimetry
and alkalimetry.
II. OXIDATION PROCESSES; as exemplified in the determination of
ferrous iron by its oxidation with potassium bichromate.
III. PRECIPITATION METHODS; of which the titration for silver with
potassium thiocyanate solution is an illustration.
From a somewhat different standpoint the methods in each case may
be subdivided into (a) DIRECT METHODS, in which the substance to be
measured is directly determined by titration to an end-point with a
standard solution; and (b) INDIRECT METHODS, in which the substance
itself is not measured, but a quantity of reagent is added which is
known to be an excess with respect to a specific reaction, and the
unused excess determined by titration. Examples of the latter class
will be pointed out as they occur in the procedures.
MEASURING INSTRUMENTS
THE ANALYTICAL BALANCE
For a complete discussion of the physical principles underlying the
construction and use of balances, and the various methods of weighing,
the student is referred to larger manuals of Quantitative Analysis,
such as those of Fresenius, or Treadwell-Hall, and particularly to
the admirable discussion of this topic in Morse's !Exercises in
Quantitative Chemistry!.
The statements and rules of procedure which follow are sufficient
for the intelligent use of an analytical balance in connection with
processes prescribed in this introductory manual. It is, however,
imperative that the student should make himself familiar with these
essential features of the balance, and its use. He should fully
realize that the analytical balance is a delicate instrument which
will render excellent service under careful treatment, but such
treatment is an essential condition if its accuracy is to be depended
upon. He should also understand that no set of rules, however
complete, can do away with the necessity for a sense of personal
responsibility, since by carelessness he can render inaccurate not
only his own analyses, but those of all other students using the same
balance.
Before making any weighings the student should seat himself before a
balance and observe the following details of construction:
1. The balance case is mounted on three brass legs, which should
preferably rest in glass cups, backed with rubber to prevent slipping.
The front legs are adjustable as to height and are used to level the
balance case; the rear leg is of permanent length.
2. The front of the case may be raised to give access to the balance.
In some makes doors are provided also at the ends of the balance case.
3. The balance beam is mounted upon an upright in the center of the
case on the top of which is an inlaid agate plate. To the center of
the beam there is attached a steel or agate knife-edge on which the
beam oscillates when it rests on the agate plate.
4. The balance beam, extending to the right and left, is graduated
along its upper edge, usually on both sides, and has at its
extremities two agate or steel knife-edges from which are suspended
stirrups. Each of these stirrups has an agate plate which, when the
balance is in action, rests upon the corresponding knife-edge of the
beam. The balance pans are suspended from the stirrups.
5. A pointer is attached to the center of the beam, and as the beam
oscillates this pointer moves in front of a scale near the base of the
post.
6. At the base of the post, usually in the rear, is a spirit-level.
7. Within the upright is a mechanism, controlled by a knob at the
front of the balance case, which is so arranged as to raise the entire
beam slightly above the level at which the knife-edges are in contact
with the agate plates. When the balance is not in use the beam must
be supported by this device since, otherwise, the constant jarring
to which a balance is inevitably subjected, will soon dull the
knife-edges, and lessen the sensitiveness of the balance.
8. A small weight, or bob, is attached to the pointer (or sometimes
to the beam) by which the center of gravity of the beam and its
attachments may be regulated. The center of gravity must lie very
slightly below the level of the agate plates to secure the desired
sensitiveness of the balance. This is provided for when the balance is
set up and very rarely requires alteration. The student should never
attempt to change this adjustment.
9. Below the balance pans are two pan-arrests operated by a button
from the front of the case. These arrests exert a very slight upward
pressure upon the pans and minimize the displacement of the beam when
objects or weights are being placed upon the pans.
10. A movable rod, operated from one end of the balance case, extends
over the balance beam and carries a small wire weight, called a rider.
By means of this rod the rider can be placed upon any desired division
of the scale on the balance beam. Each numbered division on the beam
corresponds to one milligram, and the use of the rider obviates the
placing of very small fractional weights on the balance pan.
If a new rider is purchased, or an old one replaced, care must be
taken that its weight corresponds to the graduations on the beam of
the balance on which it is to be used. The weight of the rider in
milligrams must be equal to the number of large divisions (5, 6, 10,
or 12) between the central knife-edge and the knife-edge at the end of
the beam. It should be noted that on some balances the last division
bears no number. Each new rider should be tested against a 5 or
10-milligram weight.
In some of the most recent forms of the balance a chain device
replaces the smaller weights and the use of the rider as just
described.
Before using a balance, it is always best to test its adjustment. This
is absolutely necessary if the balance is used by several workers; it
is always a wise precaution under any conditions. For this purpose,
brush off the balance pans with a soft camel's hair brush. Then note
(1) whether the balance is level; (2) that the mechanism for raising
and lowering the beams works smoothly; (3) that the pan-arrests touch
the pans when the beam is lowered; and (4) that the needle swings
equal distances on either side of the zero-point when set in motion
without any load on the pans. If the latter condition is not
fulfilled, the balance should be adjusted by means of the adjusting
screw at the end of the beam unless the variation is not more than one
division on the scale; it is often better to make a proper allowance
for this small zero error than to disturb the balance by an attempt at
correction. Unless a student thoroughly understands the construction
of a balance he should never attempt to make adjustments, but should
apply to the instructor in charge.
The object to be weighed should be placed on the left-hand balance pan
and the weights upon the right-hand pan. Every substance which
could attack the metal of the balance pan should be weighed upon a
watch-glass, and all objects must be dry and cold. A warm body gives
rise to air currents which vitiate the accuracy of the weighing.
The weights should be applied in the order in which they occur in the
weight-box (not at haphazard), beginning with the largest weight which
is apparently required. After a weight has been placed upon the pan
the beam should be lowered upon its knife-edges, and, if necessary,
the pan-arrests depressed. The movement of the pointer will then
indicate whether the weight applied is too great or too small. When
the weight has been ascertained, by the successive addition of small
weights, to the nearest 5 or 10 milligrams, the weighing is completed
by the use of the rider. The correct weight is that which causes the
pointer to swing an equal number of divisions to the right and left
of the zero-point, when the pointer traverses not less than five
divisions on either side.
The balance case should always be closed during the final weighing,
while the rider is being used, to protect the pans from the effect of
air currents.
Before the final determination of an exact weight the beam should
always be lifted from the knife-edges and again lowered into place,
as it frequently happens that the scale pans are, in spite of the
pan-arrests, slightly twisted by the impact of the weights, the beam
being thereby virtually lengthened or shortened. Lifting the beam
restores the proper alignment.
The beam should never be set in motion by lowering it forcibly upon
the knife-edges, nor by touching the pans, but rather by lifting the
rider (unless the balance be provided with some of the newer devices
for the purpose), and the swing should be arrested only when the
needle approaches zero on the scale, otherwise the knife-edges become
dull. For the same reason the beam should never be left upon its
knife-edges, nor should weights be removed from or placed on the
pans without supporting the beam, except in the case of the small
fractional weights.
When the process of weighing has been completed, the weight should
be recorded in the notebook by first noting the vacant spaces in the
weight-box, and then checking the weight by again noting the weights
as they are removed from the pan. This practice will often detect and
avoid errors. It is obvious that the weights should always be returned
to their proper places in the box, and be handled only with pincers.
It should be borne in mind that if the mechanism of a balance is
deranged or if any substance is spilled upon the pans or in the
balance case, the damage should be reported at once. In many instances
serious harm can be averted by prompt action when delay might ruin the
balance.
Samples for analysis are commonly weighed in small tubes with cork
stoppers. Since the stoppers are likely to change in weight from
the varying amounts of moisture absorbed from the atmosphere, it is
necessary to confirm the recorded weight of a tube which has been
unused for some time before weighing out a new portion of substance
from it.
WEIGHTS
The sets of weights commonly used in analytical chemistry range from
20 grams to 5 milligrams. The weights from 20 grams to 1 gram are
usually of brass, lacquered or gold plated. The fractional weights
are of German silver, gold, platinum or aluminium. The rider is of
platinum or aluminium wire.
The sets of weights purchased from reputable dealers are usually
sufficiently accurate for analytical work. It is not necessary that
such a set should be strictly exact in comparison with the absolute
standard of weight, provided they are relatively correct among
themselves, and provided the same set of weights is used in all
weighings made during a given analysis. The analyst should assure
himself that the weights in a set previously unfamiliar to him are
relatively correct by a few simple tests. For example, he should make
sure that in his set two weights of the same denomination (i.e., two
10-gram weights, or the two 100-milligram weights) are actually equal
and interchangeable, or that the 500-milligram weight is equal to
the sum of the 200, 100, 100, 50, 20, 20 and 10-milligram weights
combined, and so on. If discrepancies of more than a few tenths of a
milligram (depending upon the total weight involved) are found, the
weights should be returned for correction. The rider should also be
compared with a 5 or 10-milligram weight.
In an instructional laboratory appreciable errors should be reported
to the instructor in charge for his consideration.
When the highest accuracy is desired, the weights may be calibrated
and corrections applied. A calibration procedure is described in a
paper by T.W. Richards, !J. Am. Chem. Soc.!, 22, 144, and in many
large text-books.
Weights are inevitably subject to corrosion if not properly protected
at all times, and are liable to damage unless handled with great care.
It is obvious that anything which alters the weight of a single piece
in an analytical set will introduce an error in every weighing made
in which that piece is used. This source of error is often extremely
obscure and difficult to detect. The only safeguard against such
errors is to be found in scrupulous care in handling and protection
on the part of the analyst, and an equal insistence that if several
analysts use the same set of weights, each shall realize his
responsibility for the work of others as well as his own.
BURETTES
A burette is made from a glass tube which is as uniformly cylindrical
as possible, and of such a bore that the divisions which are etched
upon its surface shall correspond closely to actual contents.
The tube is contracted at one extremity, and terminates in either a
glass stopcock and delivery-tube, or in such a manner that a piece of
rubber tubing may be firmly attached, connecting a delivery-tube of
glass. The rubber tubing is closed by means of a glass bead. Burettes
of the latter type will be referred to as "plain burettes."
The graduations are usually numbered in cubic centimeters, and the
latter are subdivided into tenths.
One burette of each type is desirable for the analytical procedures
which follow.
PREPARATION OF A BURETTE FOR USE
The inner surface of a burette must be thoroughly cleaned in order
that the liquid as drawn out may drain away completely, without
leaving drops upon the sides. This is best accomplished by treating
the inside of the burette with a warm solution of chromic acid in
concentrated sulphuric acid, applied as follows: If the burette is of
the "plain" type, first remove the rubber tip and force the lower
end of the burette into a medium-sized cork stopper. Nearly fill the
burette with the chromic acid solution, close the upper end with a
cork stopper and tip the burette backward and forward in such a way
as to bring the solution into contact with the entire inner surface.
Remove the stopper and pour the solution into a stock bottle to be
kept for further use, and rinse out the burette with water several
times. Unless the water then runs freely from the burette without
leaving drops adhering to the sides, the process must be repeated
(Note 1).
If the burette has a glass stopcock, this should be removed after
the cleaning and wiped, and also the inside of the ground joint. The
surface of the stopcock should then be smeared with a thin coating of
vaseline and replaced. It should be attached to the burette by means
of a wire, or elastic band, to lessen the danger of breakage.
Fill the burettes with distilled water, and allow the water to run out
through the stopcock or rubber tip until convinced that no air
bubbles are inclosed (Note 2). Fill the burette to a point above the
zero-point and draw off the water until the meniscus is just below
that mark. It is then ready for calibration.
[Note 1: The inner surface of the burette must be absolutely clean if
the liquid is to run off freely. Chromic acid in sulphuric acid is
usually found to be the best cleansing agent, but the mixture must be
warm and concentrated. The solution can be prepared by pouring over a
few crystals of potassium bichromate a little water and then adding
concentrated sulphuric acid.]
[Note 2: It is always necessary to insure the absence of air bubbles
in the tips or stopcocks. The treatment described above will usually
accomplish this, but, in the case of plain burettes it is sometimes
better to allow a little of the liquid to flow out of the tip while it
is bent upwards. Any air which may be entrapped then rises with the
liquid and escapes.
If air bubbles escape during subsequent calibration or titration, an
error is introduced which vitiates the results.]
READING OF A BURETTE
All liquids when placed in a burette form what is called a meniscus at
their upper surfaces. In the case of liquids such as water or
aqueous solutions this meniscus is concave, and when the liquids are
transparent accurate readings are best obtained by observing the
position on the graduated scales of the lowest point of the meniscus.
This can best be done as follows: Wrap around the burette a piece of
colored paper, the straight, smooth edges of which are held evenly
together with the colored side next to the burette (Note 1). Hold the
paper about two small divisions below the meniscus and raise or lower
the level of the eyes until the edge of the paper at the back of the
burette is just hidden from the eye by that in front (Note 2). Note
the position of the lowest point of the curve of the meniscus,
estimating the tenths of the small divisions, thus reading its
position to hundredths of a cubic centimeter.
[Note 1: The ends of the colored paper used as an aid to accurate
readings may be fastened together by means of a gummed label. The
paper may then remain on the burette and be ready for immediate use by
sliding it up or down, as required.]
[Note 2: To obtain an accurate reading the eye must be very nearly on
a level with the meniscus. This is secured by the use of the paper
as described. The student should observe by trial how a reading is
affected when the meniscus is viewed from above or below.
The eye soon becomes accustomed to estimating the tenths of the
divisions. If the paper is held as directed, two divisions below the
meniscus, one whole division is visible to correct the judgment. It is
not well to attempt to bring the meniscus exactly to a division mark
on the burette. Such readings are usually less accurate than those in
which the tenths of a division are estimated.]
CALIBRATION OF GLASS MEASURING DEVICES
If accuracy of results is to be attained, the correctness of all
measuring instruments must be tested. None of the apparatus offered
for sale can be implicitly relied upon except those more expensive
instruments which are accompanied by a certificate from the !National
Bureau of Standards! at Washington, or other equally authentic source.
The bore of burettes is subject to accidental variations, and since
the graduations are applied by machine without regard to such
variations of bore, local errors result.
The process of testing these instruments is called !calibration!.
It is usually accomplished by comparing the actual weight of water
contained in the instrument with its apparent volume.
There is, unfortunately, no uniform standard of volume which has been
adopted for general use in all laboratories. It has been variously
proposed to consider the volume of 1000 grams of water at 4°, 15.5°,
16°, 17.5°, and even 20°C., as a liter for practical purposes, and to
consider the cubic centimeter to be one one-thousandth of that volume.
The true liter is the volume of 1000 grams of water at 4°C.; but
this is obviously a lower temperature than that commonly found in
laboratories, and involves the constant use of corrections if taken as
a laboratory standard. Many laboratories use 15.5°C. (60° F.) as the
working standard. It is plain that any temperature which is deemed
most convenient might be chosen for a particular laboratory, but it
cannot be too strongly emphasized that all measuring instruments,
including burettes, pipettes, and flasks, should be calibrated at that
temperature in order that the contents of each burette, pipette,
etc., shall be comparable with that of every other instrument, thus
permitting general interchange and substitution. For example, it is
obvious that if it is desired to remove exactly 50 cc. from a solution
which has been diluted to 500 cc. in a graduated flask, the 50 cc.
flask or pipette used to remove the fractional portion must give
a correct reading at the same temperature as the 500 cc. flask.
Similarly, a burette used for the titration of the 50 cc. of solution
removed should be calibrated under the same conditions as the
measuring flasks or pipettes employed with it.
The student should also keep constantly in mind the fact that all
volumetric operations, to be exact, should be carried out as nearly at
a constant temperature as is practicable. The spot selected for
such work should therefore be subject to a minimum of temperature
variations, and should have as nearly the average temperature of
the laboratory as is possible. In all work, whether of calibration,
standardization, or analysis, the temperature of the liquids employed
must be taken into account, and if the temperature of these liquids
varies more than 3° or 4° from the standard temperature chosen for the
laboratory, corrections must be applied for errors due to expansion or
contraction, since volumes of a liquid measured at different times are
comparable only under like conditions as to temperature. Data to be
used for this purpose are given in the Appendix. Neglect of this
correction is frequently an avoidable source of error and annoyance in
otherwise excellent work. The temperature of all solutions at the time
of standardization should be recorded to facilitate the application of
temperature corrections, if such are necessary at any later time.
CALIBRATION OF THE BURETTES
Two burettes, one at least of which should have a glass stopper, are
required throughout the volumetric work. Both burettes should be
calibrated by the student to whom they are assigned.
PROCEDURE.--Weigh a 50 cc., flat-bottomed flask (preferably a
light-weight flask), which must be dry on the outside, to the nearest
centigram. Record the weight in the notebook. (See Appendix for
suggestions as to records.) Place the flask under the burette and draw
out into it about 10 cc. of water, removing any drop on the tip by
touching it against the inside of the neck of the flask. Do not
attempt to stop exactly at the 10 cc. mark, but do not vary more than
0.1 cc. from it. Note the time, and at the expiration of three minutes
(or longer) read the burette accurately, and record the reading in the
notebook (Note 1). Meanwhile weigh the flask and water to centigrams
and record its weight (Note 2). Draw off the liquid from 10 cc. to
about 20 cc. into the same flask without emptying it; weigh, and at
the expiration of three minutes take the reading, and so on throughout
the length of the burette. When it is completed, refill the burette
and check the first calibration.
The differences in readings represent the apparent volumes, the
differences in weights the true volumes. For example, if an apparent
volume of 10.05 cc. is found to weigh 10.03 grams, it may be assumed
with sufficient accuracy that the error in that 10 cc. amounts to
-0.02 cc., or -0.002 for each cubic centimeter (Note 3).
In the calculation of corrections the temperature of the water must be
taken into account, if this varies more than 4°C. from the laboratory
standard temperature, consulting the table of densities of water in
the Appendix.
From the final data, plot the corrections to be applied so that they
may be easily read for each cubic centimeter throughout the burette.
The total correction at each 10 cc. may also be written on the burette
with a diamond, or etching ink, for permanence of record.
[Note 1: A small quantity of liquid at first adheres to the side of
even a clean burette. This slowly unites with the main body of liquid,
but requires an appreciable time. Three minutes is a sufficient
interval, but not too long, and should be adopted in every instance
throughout the whole volumetric practice before final readings are
recorded.]
[Note 2: A comparatively rough balance, capable of weighing to
centigrams, is sufficiently accurate for use in calibrations, for a
moment's reflection will show that it would be useless to weigh the
water with an accuracy greater than that of the readings taken on
the burette. The latter cannot exceed 0.01 cc. in accuracy, which
corresponds to 0.01 gram.
The student should clearly understand that !all other weighings!,
except those for calibration, should be made accurately to 0.0001
gram, unless special directions are given to the contrary.
Corrections for temperature variations of less than 4°C. are
negligible, as they amount to less than 0.01 gram for each 10 grams of
water withdrawn.]
[Note 3: Should the error discovered in any interval of 10 cc. on the
burette exceed 0.10 cc., it is advisable to weigh small portions (even
1 cc.) to locate the position of the variation of bore in the
tube rather than to distribute the correction uniformly over the
corresponding 10 cc. The latter is the usual course for small
corrections, and it is convenient to calculate the correction
corresponding to each cubic centimeter and to record it in the form
of a table or calibration card, or to plot a curve representing the
values.
Burettes may also be calibrated by drawing off the liquid in
successive portions through a 5 cc. pipette which has been accurately
calibrated, as a substitute for weighing. If many burettes are to be
tested, this is a more rapid method.]
PIPETTES
A !pipette! may consist of a narrow tube, in the middle of which is
blown a bulb of a capacity a little less than that which it is desired
to measure by the pipette; or it may be a miniature burette, without
the stopcock or rubber tip at the lower extremity. In either case, the
flow of liquid is regulated by the pressure of the finger on the top,
which governs the admission of the air.
Pipettes are usually already graduated when purchased, but they
require calibration for accurate work.
CALIBRATION OF PIPETTES
PROCEDURE.--Clean the pipette. Draw distilled water into it by sucking
at the upper end until the water is well above the graduation mark.
Quickly place the forefinger over the top of the tube, thus preventing
the entrance of air and holding the water in the pipette. Cautiously
admit a little air by releasing the pressure of the finger, and allow
the level of the water to fall until the lowest point of the meniscus
is level with the graduation. Hold the water at that point by pressure
of the finger and then allow the water to run out from the pipette
into a small tared, or weighed, beaker or flask. After a definite time
interval, usually two to three minutes, touch the end of the pipette
against the side of the beaker or flask to remove any liquid adhering
to it (Note 1). The increase in weight of the flask in grams
represents the volume of the water in cubic centimeters delivered by
the pipette. Calculate the necessary correction.
[Note 1: A definite interval must be allowed for draining, and a
definite practice adopted with respect to the removal of the liquid
which collects at the end of the tube, if the pipette is designed to
deliver a specific volume when emptied. This liquid may be removed
at the end of a definite interval either by touching the side of the
vessel or by gently blowing out the last drops. Either practice, when
adopted, must be uniformly adhered to.]
FLASKS
!Graduated or measuring flasks! are similar to the ordinary
flat-bottomed flasks, but are provided with long, narrow necks in
order that slight variations in the position of the meniscus with
respect to the graduation shall represent a minimum volume of liquid.
The flasks must be of such a capacity that, when filled with the
specified volume, the liquid rises well into the neck.
GRADUATION OF FLASKS
It is a general custom to purchase the flasks ungraduated and to
graduate them for use under standard conditions selected for the
laboratory in question. They may be graduated for "contents" or
"delivery." When graduated for "contents" they contain a specified
volume when filled to the graduation at a specified temperature, and
require to be washed out in order to remove all of the solution from
the flask. Flasks graduated for "delivery" will deliver the specified
volume of a liquid without rinsing. A flask may, of course, be
graduated for both contents and delivery by placing two graduation
marks upon it.
PROCEDURE.--To calibrate a flask for !contents!, proceed as follows:
Clean the flask, using a chromic acid solution, and dry it carefully
outside and inside. Tare it accurately; pour water into the flask
until the weight of the latter counterbalances weights on the opposite
pan which equal in grams the number of cubic centimeters of water
which the flask is to contain. Remove any excess of water with the aid
of filter paper (Note 1). Take the flask from the balance, stopper
it, place it in a bath at the desired temperature, usually 15.5°
or 17.5°C., and after an hour mark on the neck with a diamond the
location of the lowest point of the meniscus (Note 2). The mark may
be etched upon the flask by hydrofluoric acid, or by the use of an
etching ink now commonly sold on the market.
To graduate a flask which is designed to !deliver! a specified volume,
proceed as follows: Clean the flask as usual and wipe all moisture
from the outside. Fill it with distilled water. Pour out the water
and allow the water to drain from the flask for three minutes.
Counterbalance the flask with weights to the nearest centigram.
Add weights corresponding in grams to the volume desired, and add
distilled water to counterbalance these weights. An excess of water,
or water adhering to the neck of the flask, may be removed by means of
a strip of clean filter paper. Stopper the flask, place it in a bath
at 15.5°C. or 17.5°C. and, after an hour, mark the location of the
lowest point of the meniscus, as described above.
[Note 1: The allowable error in counterbalancing the water and
weights varies with the volume of the flask. It should not exceed one
ten-thousandth of the weight of water.]
[Note 2: Other methods are employed which involve the use of
calibrated apparatus from which the desired volume of water may be run
into the dry flask and the position of the meniscus marked directly
upon it. For a description of a procedure which is most convenient
when many flasks are to be calibrated, the student is referred to the
!Am. Chem J.!, 16, 479.]
GENERAL DIRECTIONS FOR VOLUMETRIC ANALYSES
It cannot be too strongly emphasized that for the success of analyses
uniformity of practice must prevail throughout all volumetric work
with respect to those factors which can influence the accuracy of the
measurement of liquids. For example, whatever conditions are imposed
during the calibration of a burette, pipette, or flask (notably the
time allowed for draining), must also prevail whenever the flask or
burette is used.
The student should also be constantly watchful to insure parallel
conditions during both standardization and analyst with respect to the
final volume of liquid in which a titration takes place. The value
of a standard solution is only accurate under the conditions which
prevailed when it was standardized. It is plain that the standard
solutions must be scrupulously protected from concentration or
dilution, after their value has been established. Accordingly, great
care must be taken to thoroughly rinse out all burettes, flasks, etc.,
with the solutions which they are to contain, in order to remove all
traces of water or other liquid which could act as a diluent. It is
best to wash out a burette at least three times with small portions of
a solution, allowing each to run out through the tip before assuming
that the burette is in a condition to be filled and used. It is, of
course, possible to dry measuring instruments in a hot closet, but
this is tedious and unnecessary.
To the same end, all solutions should be kept stoppered and away from
direct sunlight or heat. The bottles should be shaken before use to
collect any liquid which may have distilled from the solution and
condensed on the sides.
The student is again reminded that variations in temperature of
volumetric solutions must be carefully noted, and care should always
be taken that no source of heat is sufficiently near the solutions to
raise the temperature during use.
Much time may be saved by estimating the approximate volume of a
standard solution which will be required for a titration (if the data
are obtainable) before beginning the operation. It is then possible to
run in rapidly approximately the required amount, after which it is
only necessary to determine the end-point slowly and with accuracy.
In such cases, however, the knowledge of the approximate amount to be
required should never be allowed to influence the judgment regarding
the actual end-point.
STANDARD SOLUTIONS
The strength or value of a solution for a specific reaction is
determined by a procedure called !Standardization!, in which the
solution is brought into reaction with a definite weight of a
substance of known purity. For example, a definite weight of pure
sodium carbonate may be dissolved in water, and the volume of a
solution of hydrochloric acid necessary to exactly neutralize the
carbonate accurately determined. From these data the strength or value
of the acid is known. It is then a !standard solution!.
NORMAL SOLUTIONS
Standard solutions may be made of a purely empirical strength dictated
solely by convenience of manipulation, or the concentration may
be chosen with reference to a system which is applicable to all
solutions, and based upon chemical equivalents. Such solutions are
called !Normal Solutions! and contain such an amount of the reacting
substance per liter as is equivalent in its chemical action to one
gram of hydrogen, or eight grams of oxygen. Solutions containing one
half, one tenth, or one one-hundredth of this quantity per liter are
called, respectively, half-normal, tenth-normal, or hundredth-normal
solutions.
Since normal solutions of various reagents are all referred to a
common standard, they have an advantage not possessed by empirical
solutions, namely, that they are exactly equivalent to each other.
Thus, a liter of a normal solution of an acid will exactly neutralize
a liter of a normal alkali solution, and a liter of a normal oxidizing
solution will exactly react with a liter of a normal reducing
solution, and so on.
Beside the advantage of uniformity, the use of normal solutions
simplifies the calculations of the results of analyses. This is
particularly true if, in connection with the normal solution, the
weight of substance for analysis is chosen with reference to the
atomic or molecular weight of the constituent to be determined. (See
problem 26.)
The preparation of an !exactly! normal, half-normal, or tenth-normal
solution requires considerable time and care. It is usually carried
out only when a large number of analyses are to be made, or when the
analyst has some other specific purpose in view. It is, however, a
comparatively easy matter to prepare standard solutions which differ
but slightly from the normal or half-normal solution, and these have
the advantage of practical equality; that is, two approximately
half-normal solutions are more convenient to work with than two which
are widely different in strength. It is, however, true that some of
the advantage which pertains to the use of normal solutions as regards
simplicity of calculations is lost when using these approximate
solutions.
The application of these general statements will be made clear in
connection with the use of normal solutions in the various types of
volumetric processes which follow.
I. NEUTRALIZATION METHODS
ALKALIMETRY AND ACIDIMETRY
GENERAL DISCUSSION
!Standard Acid Solutions! may be prepared from either hydrochloric,
sulphuric, or oxalic acid. Hydrochloric acid has the advantage of
forming soluble compounds with the alkaline earths, but its solutions
cannot be boiled without danger of loss of strength; sulphuric acid
solutions may be boiled without loss, but the acid forms insoluble
sulphates with three of the alkaline earths; oxalic acid can be
accurately weighed for the preparation of solutions, and its solutions
may be boiled without loss, but it forms insoluble oxalates with
three of the alkaline earths and cannot be used with certain of the
indicators.
!Standard Alkali Solutions! may be prepared from sodium or potassium
hydroxide, sodium carbonate, barium hydroxide, or ammonia. Of sodium
and potassium hydroxide, it may be said that they can be used with all
indicators, and their solutions may be boiled, but they absorb carbon
dioxide readily and attack the glass of bottles, thereby losing
strength; sodium carbonate may be weighed directly if its purity is
assured, but the presence of carbonic acid from the carbonate is a
disadvantage with many indicators; barium hydroxide solutions may
be prepared which are entirely free from carbon dioxide, and such
solutions immediately show by precipitation any contamination from
absorption, but the hydroxide is not freely soluble in water; ammonia
does not absorb carbon dioxide as readily as the caustic alkalies,
but its solutions cannot be boiled nor can they be used with all
indicators. The choice of a solution must depend upon the nature of
the work in hand.
A !normal acid solution! should contain in one liter that quantity of
the reagent which represents 1 gram of hydrogen replaceable by a base.
For example, the normal solution of hydrochloric acid (HCl) should
contain 36.46 grams of gaseous hydrogen chloride, since that amount
furnishes the requisite 1 gram of replaceable hydrogen. On the other
hand, the normal solution of sulphuric acid (H_{2}SO_{4}) should
contain only 49.03 grams, i.e., one half of its molecular weight in
grams.
A !normal alkali solution! should contain sufficient alkali in a liter
to replace 1 gram of hydrogen in an acid. This quantity is represented
by the molecular weight in grams (40.01) of sodium hydroxide (NaOH),
while a sodium carbonate solution (Na_{2}CO_{3}) should contain but
one half the molecular weight in grams (i.e., 53.0 grams) in a liter
of normal solution.
Half-normal or tenth-normal solutions are employed in most analyses
(except in the case of the less soluble barium hydroxide). Solutions
of the latter strength yield more accurate results when small
percentages of acid or alkali are to be determined.
INDICATORS
It has already been pointed out that the purpose of an indicator is to
mark (usually by a change of color) the point at which just enough of
the titrating solution has been added to complete the chemical change
which it is intended to bring about. In the neutralization processes
which are employed in the measurement of alkalies (!alkalimetry!)
or acids (!acidimetry!) the end-point of the reaction should, in
principle, be that of complete neutrality. Expressed in terms of ionic
reactions, it should be the point at which the H^{+} ions from an
acid[Note 1] unite with a corresponding number of OH^{-} ions from a
base to form water molecules, as in the equation
H^{+}, Cl^{-} + Na^{+}, OH^{-} --> Na^{+}, Cl^{-} + (H_{2}O).
It is not usually possible to realize this condition of exact
neutrality, but it is possible to approach it with sufficient
exactness for analytical purposes, since substances are known which,
in solution, undergo a sharp change of color as soon as even a minute
excess of H^{+} or OH^{-} ions are present. Some, as will be seen,
react sharply in the presence of H^{+} ions, and others with OH^{-}
ions. These substances employed as indicators are usually organic
compounds of complex structure and are closely allied to the dyestuffs
in character.
[Note 1: A knowledge on the part of the student of the ionic theory
as applied to aqueous solutions of electrolytes is assumed. A brief
outline of the more important applications of the theory is given in
the Appendix.]
BEHAVIOR OF ORGANIC INDICATORS
The indicators in most common use for acid and alkali titrations are
methyl orange, litmus, and phenolphthalein.
In the following discussion of the principles underlying the behavior
of the indicators as a class, methyl orange and phenolphthalein will
be taken as types. It has just been pointed out that indicators are
bodies of complicated structure. In the case of the two indicators
named, the changes which they undergo have been carefully studied by
Stieglitz (!J. Am. Chem. Soc.!, 25, 1112) and others, and it appears
that the changes involved are of two sorts: First, a rearrangement
of the atoms within the molecule, such as often occurs in organic
compounds; and, second, ionic changes. The intermolecular changes
cannot appropriately be discussed here, as they involve a somewhat
detailed knowledge of the classification and general behavior of
organic compounds; they will, therefore, be merely alluded to, and
only the ionic changes followed.
Methyl orange is a representative of the group of indicators which,
in aqueous solutions, behave as weak bases. The yellow color which it
imparts to solutions is ascribed to the presence of the undissociated
base. If an acid, such as HCl, is added to such a solution, the acid
reacts with the indicator (neutralizes it) and a salt is formed, as
indicated by the equation:
(M.o.)^{+}, OH^{-} + H^{+}, Cl^{-} --> (M.o.)^{+} Cl^{-} + (H_{2}O).
This salt ionizes into (M.o.)^{+} (using this abbreviation for the
positive complex) and Cl^{-}; but simultaneously with this ionization
there appears to be an internal rearrangement of the atoms which
results in the production of a cation which may be designated as
(M'.o'.)^{+}, and it is this which imparts a characteristic red color
to the solution. As these changes occur in the presence of even a
very small excess of acid (that is, of H^{+} ions), it serves as the
desired index of their presence in the solution. If, now, an alkali,
such as NaOH, is added to this reddened solution, the reverse
series of changes takes place. As soon as the free acid present is
neutralized, the slightest excess of sodium hydroxide, acting as
a strong base, sets free the weak, little-dissociated base of the
indicator, and at the moment of its formation it reverts, because of
the rearrangement of the atoms, to the yellow form:
OH^{-} + (M'.o'.)^{+} --> [M'.o'.OH] --> [M.o.OH].
Phenolphthalein, on the other hand, is a very weak, little-dissociated
acid, which is colorless in neutral aqueous solution or in the
presence of free H^{+} ions. When an alkali is added to such a
solution, even in slight excess, the anion of the salt which has
formed from the acid of the indicator undergoes a rearrangement of the
atoms, and a new ion, (Ph')^{+}, is formed, which imparts a pink color
to the solution:
H^{+}, (Ph)^{-} + Na^{+}, OH^{-} --> (H_{2}O) + Na^{+}, (Ph)^{-}
--> Na^{+}, (Ph')^{-}
The addition of the slightest excess of an acid to this solution, on
the other hand, occasions first the reversion to the colorless ion and
then the setting free of the undissociated acid of the indicator:
H^{+}, (Ph')^{-} --> H^{+}, (Ph)^{-} --> (HPh).
Of the common indicators methyl orange is the most sensitive toward
alkalies and phenolphthalein toward acids; the others occupy
intermediate positions. That methyl orange should be most sensitive
toward alkalies is evident from the following considerations: Methyl
orange is a weak base and, therefore, but little dissociated. It
should, then, be formed in the undissociated condition as soon as even
a slight excess of OH^{-} ions is present in the solution, and there
should be a prompt change from red to yellow as outlined above. On the
other hand, it should be an unsatisfactory indicator for use with weak
acids (acetic acid, for example) because the salts which it forms
with such acids are, like all salts of that type, hydrolyzed to a
considerable extent. This hydrolytic change is illustrated by the
equation:
(M.o.)^{+} C_{2}H_{3}O_{2}^{-} + H^{+}, OH^{-} --> [M.o.OH] + H^{+},
C_{2}H_{3}O_{2}^{-}.
Comparison of this equation with that on page 30 will make it plain
that hydrolysis is just the reverse of neutralization and must,
accordingly, interfere with it. Salts of methyl orange with weak acids
are so far hydrolyzed that the end-point is uncertain, and methyl
orange cannot be used in the titration of such acids, while with
the very weak acids, such as carbonic acid or hydrogen sulphide
(hydrosulphuric acid), the salts formed with methyl orange are, in
effect, completely hydrolyzed (i.e., no neutralization occurs), and
methyl orange is accordingly scarcely affected by these acids. This
explains its usefulness, as referred to later, for the titration of
strong acids, such as hydrochloric acid, even in the presence of
carbonates or sulphides in solution.
Phenolphthalein, on the other hand, should be, as it is, the best of
the common indicators for use with weak acids. For, since it is
itself a weak acid, it is very little dissociated, and its nearly
undissociated, colorless molecules are promptly formed as soon as
there is any free acid (that is, free H^{+} ions) in the solution.
This indicator cannot, however, be successfully used with weak bases,
even ammonium hydroxide; for, since it is weak acid, the salts
which it forms with weak alkalies are easily hydrolyzed, and as a
consequence of this hydrolysis the change of color is not sharp.
This indicator can, however, be successfully used with strong bases,
because the salts which it forms with such bases are much less
hydrolyzed and because the excess of OH^{-} ions from these bases also
diminishes the hydrolytic action of water.
This indicator is affected by even so weak an acid as carbonic acid,
which must be removed by boiling the solution before titration. It is
the indicator most generally employed for the titration of organic
acids.
In general, it may be stated that when a strong acid, such as
hydrochloric, sulphuric or nitric acid, is titrated against a strong
base, such as sodium hydroxide, potassium hydroxide, or barium
hydroxide, any of these indicators may be used, since very little
hydrolysis ensues. It has been noted above that the color change does
not occur exactly at theoretical neutrality, from which it follows
that no two indicators will show exactly the same end-point when acids
and alkalis are brought together. It is plain, therefore, that the
same indicator must be employed for both standardization and analysis,
and that, if this is done, accurate results are obtainable.
The following table (Note 1) illustrates the variations in the volume
of an alkali solution (tenth-normal sodium hydroxide) required to
produce an alkaline end-point when run into 10 cc. of tenth-normal
sulphuric acid, diluted with 50 cc. of water, using five drops of each
of the different indicator solutions.
====================================================================
| | | |
INDICATOR | N/10 | N/10 |COLOR IN ACID|COLOR IN ALKA-
| H_{2}SO_{4}| NaOH |SOLUTION |LINE SOLUTION
_______________|____________|__________|_____________|______________
| cc. | cc. | cc. |
Methyl orange | 10 | 9.90 | Red | Yellow
Lacmoid | 10 | 10.00 | Red | Blue
Litmus | 10 | 10.00 | Red | Blue
Rosalic acid | 10 | 10.07 | Yellow | Pink
Phenolphthalein| 10 | 10.10 | Colorless | Pink
====================================================================
It should also be stated that there are occasionally secondary
changes, other than those outlined above, which depend upon the
temperature and concentration of the solutions in which the indicators
are used. These changes may influence the sensitiveness of an
indicator. It is important, therefore, to take pains to use
approximately the same volume of solution when standardizing that is
likely to be employed in analysis; and when it is necessary, as is
often the case, to titrate the solution at boiling temperature, the
standardization should take place under the same conditions. It is
also obvious that since some acid or alkali is required to react with
the indicator itself, the amount of indicator used should be uniform
and not excessive. Usually a few drops of solution will suffice.
The foregoing statements with respect to the behavior of indicators
present the subject in its simplest terms. Many substances other than
those named may be employed, and they have been carefully studied to
determine the exact concentration of H^{+} ions at which the color
change of each occurs. It is thus possible to select an indicator
for a particular purpose with considerable accuracy. As data of this
nature do not belong in an introductory manual, reference is made to
the following papers or books in which a more extended treatment of
the subject may be found:
Washburn, E.W., Principles of Physical Chemistry (McGraw-Hill Book
Co.), (Second Edition, 1921), pp. 380-387.
Prideaux, E.B.R., The Theory and Use of Indicators (Constable & Co.,
Ltd.), (1917).
Salm, E., A Study of Indicators, !Z. physik. Chem.!, 57 (1906),
471-501.
Stieglitz, J., Theories of Indicators, !J. Am. Chem. Soc.!, 25 (1903),
1112-1127.
Noyes, A.A., Quantitative Applications of the Theory of Indicators to
Volumetric Analysis, !J. Am. Chem. Soc.!, 32 (1911), 815-861.
Bjerrum, N., General Discussion, !Z. Anal. Chem.!, 66 (1917), 13-28
and 81-95.
Ostwald, W., Colloid Chemistry of Indicators, !Z. Chem. Ind.
Kolloide!, 10 (1912), 132-146.
[Note 1: Glaser, !Indikatoren der Acidimetrie und Alkalimetrie!.
Wiesbaden, 1901.]
PREPARATION OF INDICATOR SOLUTIONS
A !methyl orange solution! for use as an indicator is commonly made by
dissolving 0.05-0.1 gram of the compound (also known as Orange III) in
a few cubic centimeters of alcohol and diluting with water to 100 cc.
A good grade of material should be secured. It can be successfully
used for the titration of hydrochloric, nitric, sulphuric, phosphoric,
and sulphurous acids, and is particularly useful in the determination
of bases, such as sodium, potassium, barium, calcium, and ammonium
hydroxides, and even many of the weak organic bases. It can also be
used for the determination, by titration with a standard solution of
a strong acid, of the salts of very weak acids, such as carbonates,
sulphides, arsenites, borates, and silicates, because the weak acids
which are liberated do not affect the indicator, and the reddening of
the solution does not take place until an excess of the strong acid
is added. It should be used in cold, not too dilute, solutions. Its
sensitiveness is lessened in the presence of considerable quantities
of the salts of the alkalies.
A !phenolphthalein solution! is prepared by dissolving 1 gram of the
pure compound in 100 cc. of 95 per cent alcohol. This indicator is
particularly valuable in the determination of weak acids, especially
organic acids. It cannot be used with weak bases, even ammonia. It
is affected by carbonic acid, which must, therefore, be removed by
boiling when other acids are to be measured. It can be used in hot
solutions. Some care is necessary to keep the volume of the solutions
to be titrated approximately uniform in standardization and in
analysis, and this volume should not in general exceed 125-150 cc. for
the best results, since the compounds formed by the indicator undergo
changes in very dilute solution which lessen its sensitiveness.
The preparation of a !solution of litmus! which is suitable for use
as an indicator involves the separation from the commercial litmus of
azolithmine, the true coloring principle. Soluble litmus tablets are
often obtainable, but the litmus as commonly supplied to the market is
mixed with calcium carbonate or sulphate and compressed into lumps. To
prepare a solution, these are powdered and treated two or three times
with alcohol, which dissolves out certain constituents which cause a
troublesome intermediate color if not removed. The alcohol is decanted
and drained off, after which the litmus is extracted with hot water
until exhausted. The solution is allowed to settle for some time, the
clear liquid siphoned off, concentrated to one-third its volume and
acetic acid added in slight excess. It is then concentrated to a
sirup, and a large excess of 95 per cent. alcohol added to it. This
precipitates the blue coloring matter, which is filtered off, washed
with alcohol, and finally dissolved in a small volume of water and
diluted until about three drops of the solution added to 50 cc. of
water just produce a distinct color. This solution must be kept in an
unstoppered bottle. It should be protected from dust by a loose plug
of absorbent cotton. If kept in a closed bottle it soon undergoes a
reduction and loses its color, which, however, is often restored by
exposure to the air.
Litmus can be employed successfully with the strong acids and bases,
and also with ammonium hydroxide, although the salts of the latter
influence the indicator unfavorably if present in considerable
concentration. It may be employed with some of the stronger organic
acids, but the use of phenolphthalein is to be preferred.
PREPARATION OF STANDARD SOLUTIONS
!Hydrochloric Acid and Sodium Hydroxide. Approximate Strength!, 0.5 N
PROCEDURE.--Measure out 40 cc. of concentrated, pure hydrochloric
acid into a clean liter bottle, and dilute with distilled water to an
approximate volume of 1000 cc. Shake the solution vigorously for a
full minute to insure uniformity. Be sure that the bottle is not too
full to permit of a thorough mixing, since lack of care at this point
will be the cause of much wasted time (Note 1).
Weigh out, upon a rough balance, 23 grams of sodium hydroxide (Note
2). Dissolve the hydroxide in water in a beaker. Pour the solution
into a liter bottle and dilute, as above, to approximately 1000 cc.
This bottle should preferably have a rubber stopper, as the hydroxide
solution attacks the glass of the ground joint of a glass stopper, and
may cement the stopper to the bottle. Shake the solution as described
above.
[Note 1: The original solutions are prepared of a strength greater
than 0.5 N, as they are more readily diluted than strengthened if
later adjustment is desired.
Too much care cannot be taken to insure perfect uniformity of
solutions before standardization, and thoroughness in this respect
will, as stated, often avoid much waste of time. A solution once
thoroughly mixed remains uniform.]
[Note 2: Commercial sodium hydroxide is usually impure and always
contains more or less carbonate; an allowance is therefore made for
this impurity by placing the weight taken at 23 grams per liter. If
the hydroxide is known to be pure, a lesser amount (say 21 grams) will
suffice.]
COMPARISON OF ACID AND ALKALI SOLUTIONS
PROCEDURE.--Rinse a previously calibrated burette three times with the
hydrochloric acid solution, using 10 cc. each time, and allowing the
liquid to run out through the tip to displace all water and air
from that part of the burette. Then fill the burette with the acid
solution. Carry out the same procedure with a second burette, using
the sodium hydroxide solution.
The acid solution may be placed in a plain or in a glass-stoppered
burette as may be more convenient, but the alkaline solution should
never be allowed to remain long in a glass-stoppered burette, as it
tends to cement the stopper to the burette, rendering it useless. It
is preferable to use a plain burette for this solution.
When the burettes are ready for use and all air bubbles displaced from
the tip (see Note 2, page 17) note the exact position of the liquid in
each, and record the readings in the notebook. (Consult page 188.) Run
out from the burette into a beaker about 40 cc. of the acid and add
two drops of a solution of methyl orange; dilute the acid to about
80 cc. and run out alkali solution from the other burette, stirring
constantly, until the pink has given place to a yellow. Wash down the
sides of the beaker with a little distilled water if the solution has
spattered upon them, return the beaker to the acid burette, and add
acid to restore the pink; continue these alternations until the point
is accurately fixed at which a single drop of either solutions served
to produce a distinct change of color. Select as the final end-point
the appearance of the faintest pink tinge which can be recognized, or
the disappearance of this tinge, leaving a pure yellow; but always
titrate to the same point (Note 1). If the titration has occupied more
than the three minutes required for draining the sides of the burette,
the final reading may be taken immediately and recorded in the
notebook.
Refill the burettes and repeat the titration. From the records of
calibration already obtained, correct the burette readings and make
corrections for temperature, if necessary. Obtain the ratio of the
sodium hydroxide solution to that of hydrochloric acid by dividing
the number of cubic centimeters of acid used by the number of cubic
centimeters of alkali required for neutralization. The check results
of the two titrations should not vary by more than two parts in one
thousand (Note 2). If the variation in results is greater than this,
refill the burettes and repeat the titration until satisfactory values
are obtained. Use a new page in the notebook for each titration.
Inaccurate values should not be erased or discarded. They should be
retained and marked "correct" or "incorrect," as indicated by the
final outcome of the titrations. This custom should be rigidly
followed in all analytical work.
[Note 1: The end-point should be chosen exactly at the point of
change; any darker tint is unsatisfactory, since it is impossible to
carry shades of color in the memory and to duplicate them from day to
day.]
[Note 2: While variation of two parts in one thousand in the values
obtained by an inexperienced analyst is not excessive, the idea must
be carefully avoided that this is a standard for accurate work to be
!generally applied!. In many cases, after experience is gained, the
allowable error is less than this proportion. In a few cases a
larger variation is permissible, but these are rare and can only
be recognized by an experienced analyst. It is essential that the
beginner should acquire at least the degree of accuracy indicated if
he is to become a successful analyst.]
STANDARDIZATION OF HYDROCHLORIC ACID
SELECTION AND PREPARATION OF STANDARD
The selection of the best substance to be used as a standard for acid
solutions has been the subject of much controversy. The work of Lunge
(!Ztschr. angew. Chem.! (1904), 8, 231), Ferguson (!J. Soc. Chem.
Ind.! (1905), 24, 784), and others, seems to indicate that the best
standard is sodium carbonate prepared from sodium bicarbonate by
heating the latter at temperature between 270° and 300°C. The
bicarbonate is easily prepared in a pure state, and at the
temperatures named the decomposition takes place according to the
equation
2HNaCO_{3} --> Na_{2}CO_{3} + H_{2}O + CO_{2}
and without loss of any carbon dioxide from the sodium carbonate, such
as may occur at higher temperatures. The process is carried out as
described below.
PROCEDURE.--Place in a porcelain crucible about 6 grams (roughly
weighed) of the purest sodium bicarbonate obtainable. Rest the
crucible upon a triangle of iron or copper wire so placed within a
large crucible that there is an open air space of about three eighths
of an inch between them. The larger crucible may be of iron, nickel or
porcelain, as may be most convenient. Insert the bulb of a thermometer
reading to 350°C. in the bicarbonate, supporting it with a clamp so
that the bulb does not rest on the bottom of the crucible. Heat
the outside crucible, using a rather small flame, and raise the
temperature of the bicarbonate fairly rapidly to 270°C. Then regulate
the heat in such a way that the temperature rises !slowly! to 300°C.
in the course of a half-hour. The bicarbonate should be frequently
stirred with a clean, dry, glass rod, and after stirring, should be
heaped up around the bulb of the thermometer in such a way as to cover
it. This will require attention during most of the heating, as the
temperature should not be permitted to rise above 310°C. for any
length of time. At the end of the half-hour remove the thermometer and
transfer the porcelain crucible, which now contains sodium carbonate,
to a desiccator. When it is cold, transfer the carbonate to a
stoppered weighing tube or weighing-bottle.
STANDARDIZATION
PROCEDURE.--Clean carefully the outside of a weighing-tube, or
weighing-bottle, containing the pure sodium carbonate, taking care
to handle it as little as possible after wiping. Weigh the tube
accurately to 0.0001 gram, and record the weight in the notebook. Hold
the tube over the top of a beaker (200-300 cc.) and cautiously remove
the stopper, making sure that no particles fall from it or from the
tube elsewhere than in the beaker. Pour out from the tube a portion
of the carbonate, replace the stopper and determine approximately how
much has been removed. Continue this procedure until 1.00 to 1.10
grams has been taken from the tube. Then weigh the tube accurately
and record the weight under the first weight in the notebook.
The difference in the two weights is the weight of the carbonate
transferred to the beaker. Proceed in the same way to transfer a
second portion of the carbonate from the tube to another beaker of
about the same size as the first. The beakers should be labeled and
plainly marked to correspond with the entries in the notebook.
Pour over the carbonate in each beaker about 80 cc. of water, stir
until solution is complete, and add two drops of methyl orange
solution. Fill the burettes with the standard acid and alkali
solutions, noting the initial readings of the burettes and temperature
of the solutions. Run in acid from the burette, stirring and avoiding
loss by effervescence, until the solution has become pink. Wash down
the sides of the beaker with a !little! water from a wash-bottle, and
then run in alkali from the other burette until the pink is replaced
by yellow; then finish the titration as described on page 37. Note the
readings of the burettes after the proper interval, and record them in
the notebook. Repeat the procedure, using the second portion of sodium
carbonate. Apply the necessary calibration corrections to the volumes
of the solutions used, and correct for temperature if necessary.
From the data obtained, calculate the volume of the hydrochloric
acid solution which is equivalent to the volume of sodium hydroxide
solution used in this titration. Subtract this volume from the volume
of hydrochloric acid. The difference represents the volume of acid
used to react with the sodium carbonate. Divide the weight of sodium
carbonate by this volume in cubic centimeters, thus obtaining the
weight of sodium carbonate equivalent to each cubic centimeter of the
acid.
From this weight it is possible to calculate the corresponding weight
of HCl in each cubic centimeter of the acid, and in turn the relation
of the acid to the normal.
If, however, it is recalled that normal solutions are equivalent to
each other, it will be seen that the same result may be more readily
reached by dividing the weight in grams of sodium carbonate per cubic
centimeter just found by titration by the weight which would be
contained in the same volume of a normal solution of sodium carbonate.
A normal solution of sodium carbonate contains 53.0 grams per liter,
or 0.0530 gram per cc. (see page 29). The relation of the acid
solution to the normal is, therefore, calculated by dividing the
weight of the carbonate to which each cubic centimeter of the acid is
equivalent by 0.0530. The standardization must be repeated until the
values obtained agree within, at most, two parts in one thousand.
When the standard of the acid solution has been determined, calculate,
from the known ratio of the two solutions, the relation of the sodium
hydroxide solution to a normal solution (Notes 1 and 2).
[Note 1: In the foregoing procedure the acid solution is standardized
and the alkali solution referred to this standard by calculation. It
is equally possible, if preferred, to standardize the alkali solution.
The standards in a common use for this purpose are purified
oxalic acid (H_{2}C_{2}O_{4}.2H_{2}O), potassium acid oxalate
(KHC_{2}O_{4}.H_{2}O or KHC_{2}O_{4}), potassium tetroxalate
(KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O), or potassium acid tartrate
(KHC_{4}O_{6}), with the use of a suitable indicator. The oxalic acid
and the oxalates should be specially prepared to insure purity,
the main difficulty lying in the preservation of the water of
crystallization.
It should be noted that the acid oxalate and the acid tartrate each
contain one hydrogen atom replaceable by a base, while the tetroxalate
contains three such atoms and the oxalic acid two. Each of the two
salts first named behave, therefore, as monobasic acids, and the
tetroxalate as a tribasic acid.]
[Note 2: It is also possible to standardize a hydrochloric acid
solution by precipitating the chloride ions as silver chloride and
weighing the precipitate, as prescribed under the analysis of sodium
chloride to be described later. Sulphuric acid solutions may be
standardized by precipitation of the sulphate ions as barium sulphate
and weighing the ignited precipitate, but the results are not above
criticism on account of the difficulty in obtaining large precipitates
of barium sulphate which are uncontaminated by inclosures or are not
reduced on ignition.]
DETERMINATION OF THE TOTAL ALKALINE STRENGTH OF SODA ASH
Soda ash is crude sodium carbonate. If made by the ammonia process it
may contain also sodium chloride, sulphate, and hydroxide; when made
by the Le Blanc process it may contain sodium sulphide, silicate, and
aluminate, and other impurities. Some of these, notably the hydroxide,
combine with acids and contribute to the total alkaline strength,
but it is customary to calculate this strength in terms of sodium
carbonate; i.e., as though no other alkali were present.
PROCEDURE.--In order to secure a sample which shall represent the
average value of the ash, it is well to take at least 5 grams. As this
is too large a quantity for convenient titration, an aliquot portion
of the solution is measured off, representing one fifth of the entire
quantity. This is accomplished as follows: Weigh out on an analytical
balance two samples of soda ash of about 5 grams each into beakers
of about 500 cc. capacity. (The weighings need be made to centigrams
only.) Dissolve the ash in 75 cc. of water, warming gently, and filter
off the insoluble residue; wash the filter by filling it at least
three times with distilled water, and allowing it to drain, adding the
washings to the main filtrate. Cool the filtrate to approximately the
standard temperature of the laboratory, and transfer it to a 250 cc.
measuring flask, washing out the beaker thoroughly. Add distilled
water of laboratory temperature until the lowest point of the meniscus
is level with the graduation on the neck of the flask and remove any
drops of water that may be on the neck above the graduation by means
of a strip of filter paper; make the solution thoroughly uniform by
pouring it out into a dry beaker and back into the flask several
times. Measure off 50 cc. of the solution in a measuring flask, or
pipette, either of which before use should, unless they are dry on the
inside, be rinsed out with at least two small portions of the soda ash
solution to displace any water.
If a flask is used, fill it to the graduation with the soda ash
solution and remove any liquid from the neck above the graduation with
filter paper. Empty it into a beaker, and wash out the small flask,
unless it is graduated for !delivery!, using small quantities of
water, which are added to the liquid in the beaker. A second 50 cc.
portion from the main solution should be measured off into a second
beaker. Dilute the solutions in each beaker to 100 cc., add two drops
of a solution of methyl orange (Note 1) and titrate for the alkali
with the standard hydrochloric acid solution, using the alkali
solution to complete the titration as already prescribed.
From the volumes of acid and alkali employed, corrected for burette
errors and temperature changes, and the data derived from the
standardization, calculate the percentage of alkali present, assuming
it all to be present as sodium carbonate (Note 2).
[Note 1: The hydrochloric acid sets free carbonic acid which is
unstable and breaks down into water and carbon dioxide, most of which
escapes from the solution. Carbonic acid is a weak acid and, as such,
does not yield a sufficient concentration of H^{+} ions to cause the
indicator to change to a pink (see page 32).
The chemical changes involved may be summarized as follows:
2H^{+}, 2Cl^{-} + 2Na^{+}, CO_{3}^{--} --> 2Na^{+}, 2Cl^{-} +
[H_{2}CO_{3}] --> H_{2}O + CO_{2}]
[Note 2: A determination of the alkali present as hydroxide in soda
ash may be determined by precipitating the carbonate by the addition
of barium chloride, removing the barium carbonate by filtration, and
titrating the alkali in the filtrate.
The caustic alkali may also be determined by first using
phenolphthalein as an indicator, which will show by its change from
pink to colorless the point at which the caustic alkali has been
neutralized and the carbonate has been converted to bicarbonate, and
then adding methyl orange and completing the titration. The amount of
acid necessary to change the methyl orange to pink is a measure of one
half of the carbonate present. The results of the double titration
furnish the data necessary for the determination of the caustic alkali
and of the carbonate in the sample.]
DETERMINATION OF THE ACID STRENGTH OF OXALIC ACID
PROCEDURE.--Weigh out two portions of the acid of about 1 gram
each. Dissolve these in 50 cc. of warm water. Add two drops of
phenolphthalein solution, and run in alkali from the burette until the
solution is pink; add acid from the other burette until the pink is
just destroyed, and then add 0.3 cc. (not more) in excess. Heat the
solution to boiling for three minutes. If the pink returns during the
boiling, discharge it with acid and again add 0.3 cc. in excess and
repeat the boiling (Note 1). If the color does not then reappear, add
alkali until it does, and a !drop or two! of acid in excess and boil
again for one minute (Note 2). If no color reappears during this time,
complete the titration in the hot solution. The end-point should be
the faintest visible shade of color (or its disappearance), as the
same difficulty would exist here as with methyl orange if an attempt
were made to match shades of pink.
From the corrected volume of alkali required to react with the
oxalic acid, calculate the percentage of the crystallized acid
(H_{2}C_{2}O_{4}.2H_{2}O) in the sample (Note 3).
[Note 1: All commercial caustic soda such as that from which the
standard solution was made contains some sodium carbonate. This reacts
with the oxalic acid, setting free carbonic acid, which, in turn,
forms sodium bicarbonate with the remaining carbonate:
H_{2}CO_{3} + Na_{2}CO_{3} --> 2HNaCO_{3}.
This compound does not hydrolyze sufficiently to furnish enough OH^{-}
ions to cause phenolphthalein to remain pink; hence, the color of
the indicator is discharged in cold solutions at the point at which
bicarbonate is formed. If, however, the solution is heated to boiling,
the bicarbonate loses carbon dioxide and water, and reverts to sodium
carbonate, which causes the indicator to become again pink:
2HNaCO_{3} --> H_{2}O + CO_{2} + Na_{2}CO_{3}.
By adding successive portions of hydrochloric acid and boiling, the
carbonate is ultimately all brought into reaction.
The student should make sure that the difference in behavior of the
two indicators, methyl orange and phenolphthalein, is understood.]
[Note 2: Hydrochloric acid is volatilized from aqueous solutions,
except such as are very dilute. If the directions in the procedure
are strictly followed, no loss of acid need be feared, but the amount
added in excess should not be greater than 0.3-0.4 cc.]
[Note 3: Attention has already been called to the fact that the color
changes in the different indicators occur at varying concentrations
of H^{+} or OH^{-} ions. They do not indicate exact theoretical
neutrality, but a particular indicator always shows its color change
at a particular concentration of H^{+} or OH^{-} ions. The results
of titration with a given indicator are, therefore, comparable. As a
matter of fact, a small error is involved in the procedure as outlined
above. The comparison of the acid and alkali solutions was made, using
methyl orange as an indicator, while the titration of the oxalic acid
is made with the use of phenolphthalein. For our present purposes the
small error may be neglected but, if time permits, the student is
recommended to standardize the alkali solution against one of the
substances named in Note 1, page 41, and also to ascertain
the comparative value of the acid and alkali solutions, using
phenolphthalein as indicator throughout, and conducting the titrations
as described above. This will insure complete accuracy.]
II. OXIDATION PROCESSES
GENERAL DISCUSSION
In the oxidation processes of volumetric analysis standard solutions
of oxidizing agents and of reducing agents take the place of the acid
and alkali solutions of the neutralization processes already studied.
Just as an acid solution was the principal reagent in alkalimetry, and
the alkali solution used only to make certain of the end-point, the
solution of the oxidizing agent is the principal reagent for the
titration of substances exerting a reducing action. It is, in general,
true that oxidizable substances are determined by !direct! titration,
while oxidizing substances are determined by !indirect! titration.
The important oxidizing agents employed in volumetric solutions are
potassium bichromate, potassium permangenate, potassium ferricyanide,
iodine, ferric chloride, and sodium hypochlorite.
The important reducing agents which are used in the form of standard
solutions are ferrous sulphate (or ferrous ammonium sulphate), oxalic
acid, sodium thiosulphate, stannous chloride, arsenious acid, and
potassium cyanide. Other reducing agents, as sulphurous acid,
sulphureted hydrogen, and zinc (nascent hydrogen), may take part in
the processes, but not as standard solutions.
The most important combinations among the foregoing are: Potassium
bichromate and ferrous salts; potassium permanganate and ferrous
salts; potassium permanganate and oxalic acid, or its derivatives;
iodine and sodium thiosulphate; hypochlorites and arsenious acid.
BICHROMATE PROCESS FOR THE DETERMINATION OF IRON
Ferrous salts may be promptly and completely oxidized to ferric salts,
even in cold solution, by the addition of potassium bichromate,
provided sufficient acid is present to hold in solution the ferric and
chromic compounds which are formed.
The acid may be either hydrochloric or sulphuric, but the former is
usually preferred, since it is by far the best solvent for iron and
its compounds. The reaction in the presence of hydrochloric acid is as
follows:
6FeCl_{2} + K_{2}Cr_{2}O_{7} + 14HCl --> 6FeCl_{3} + 2CrCl_{3} + 2KCl
+ 7H_{2}O.
NORMAL SOLUTIONS OF OXIDIZING OR REDUCING AGENTS
It will be recalled that the system of normal solutions is based upon
the equivalence of the reagents which they contain to 8 grams of
oxygen or 1 gram of hydrogen. A normal solution of an oxidizing agent
should, therefore, contain that amount per liter which is equivalent
in oxidizing power to 8 grams of oxygen; a normal reducing solution
must be equivalent in reducing power to 1 gram of hydrogen. In order
to determine what the amount per liter will be it is necessary to know
how the reagents enter into reaction. The two solutions to be employed
in the process under consideration are those of potassium bichromate
and ferrous sulphate. The reaction between them, in the presence of an
excess of sulphuric acid, may be expressed as follows:
6FeSO_{4} + K_{2}Cr_{2}O_{7} + 7H_{2}SO_{4} --> 3Fe_{2}(SO_{4})_{3} +
K_{2}SO_{4} + Cr_{2}(SO_{4})_{3} + 7H_{2}O.
If the compounds of iron and chromium, with which alone we are now
concerned, be written in such a way as to show the oxides of these
elements in each, they would appear as follows: On the left-hand side
of the equation 6(FeO.SO_{3}) and K_{2}O.2CrO_{3}; on the right-hand
side, 3(Fe_{2}O_{3}.3SO_{3}) and Cr_{2}O_{3}.3SO_{3}. A careful
inspection shows that there are three less oxygen atoms associated
with chromium atoms on the right-hand side of the equation than on the
left-hand, but there are three more oxygen atoms associated with iron
atoms on the right than on the left. In other words, a molecule of
potassium bichromate has given up three atoms of oxygen for oxidation
purposes; i.e., a molecular weight in grams of the bichromate (294.2)
will furnish 3 X 16 or 48 grams of oxygen for oxidation purposes.
As this 48 grams is six times 8 grams, the basis of the system, the
normal solution of potassium bichromate should contain per liter one
sixth of 294.2 grams or 49.03 grams.
A further inspection of the dissected compounds above shows that six
molecules of FeO.SO_{3} were required to react with the three atoms of
oxygen from the bichromate. From the two equations
3H_{2} + 3O --> 3H_{2}O
6(FeO.SO_{3}) + 3O --> 3(Fe_{2}O_{3}.3SO_{3})
it is plain that one molecule of ferrous sulphate is equivalent to one
atom of hydrogen in reducing power; therefore one molecular weight in
grams of ferrous sulphate (151.9) is equivalent to 1 gram of
hydrogen. Since the ferrous sulphate crystalline form has the formula
FeSO_{4}.7H_{2}O, a normal reducing solution of this crystalline salt
should contain 277.9 grams per liter.
PREPARATION OF SOLUTIONS
!Approximate Strength 0.1 N!
It is possible to purify commercial potassium bichromate by
recrystallization from hot water. It must then be dried and cautiously
heated to fusion to expel the last traces of moisture, but not
sufficiently high to expel any oxygen. The pure salt thus prepared,
may be weighed out directly, dissolved, and the solution diluted in a
graduated flask to a definite volume. In this case no standardization
is made, as the normal value can be calculated directly. It is,
however, more generally customary to standardize a solution of
the commercial salt by comparison with some substance of definite
composition, as described below.
PROCEDURE.--Pulverize about 5 grams of potassium bichromate of good
quality. Dissolve the bichromate in distilled water, transfer the
solution to a liter bottle, and dilute to approximately 1000 cc. Shake
thoroughly until the solution is uniform.
To prepare the solution of the reducing agent, pulverize about 28
grams of ferrous sulphate (FeSO_{4}.7H_{2}O) or about 40 grams of
ferrous ammonium sulphate (FeSO_{4}.(NH_{4})_{2}SO_{4}.6H_{2}O) and
dissolve in distilled water containing 5 cc. of concentrated sulphuric
acid. Transfer the solution to a liter bottle, add 5 cc. concentrated
sulphuric acid, make up to about 1000 cc. and shake vigorously to
insure uniformity.
INDICATOR SOLUTION
No indicator is known which, like methyl orange, can be used within
the solution, to show when the oxidation process is complete. Instead,
an outside indicator solution is employed to which drops of the
titrated solution are transferred for testing. The reagent used is
potassium ferricyanide, which produces a blue precipitate (or color)
with ferrous compounds as long as there are unoxidized ferrous ions in
the titrated solution. Drops of the indicator solution are placed upon
a glazed porcelain tile, or upon white cardboard which has been coated
with paraffin to render it waterproof, and drops of the titrated
solution are transferred to the indicator on the end of a stirring
rod. When the oxidation is nearly completed only very small amounts
of the ferrous compounds remain unoxidized and the reaction with the
indicator is no longer instantaneous. It is necessary to allow a brief
time to elapse before determining that no blue color is formed. Thirty
seconds is a sufficient interval, and should be adopted throughout the
analytical procedure. If left too long, the combined effect of light
and dust from the air will cause a reduction of the ferric compounds
already formed and a resultant blue will appear which misleads the
observer with respect to the true end-point.
The indicator solution must be highly diluted, otherwise its own color
interferes with accurate observation. Prepare a fresh solution, as
needed each day, by dissolving a crystal of potassium ferricyanide
about the size of a pin's head in 25 cc. of distilled water. The salt
should be carefully tested with ferric chloride for the presence of
ferrocyanides, which give a blue color with ferric salts.
In case of need, the ferricyanide can be purified by adding to its
solution a little bromine water and recrystallizing the compound.
COMPARISON OF OXIDIZING AND REDUCING SOLUTIONS
PROCEDURE.--Fill one burette with each of the solutions, observing
the general procedure with respect to cleaning and rinsing already
prescribed. The bichromate solution is preferably to be placed in a
glass-stoppered burette.
Run out from a burette into a beaker of about 300 cc. capacity nearly
40 cc. of the ferrous solution, add 15 cc. of dilute hydrochloric acid
(sp. gr. 1.12) and 150 cc. of water and run in the bichromate
solution from another burette. Since both solutions are approximately
tenth-normal, 35 cc. of the bichromate solution may be added without
testing. Test at that point by removing a very small drop of the
iron solution on the end of a stirring rod, mixing it with a drop of
indicator on the tile (Note 1). If a blue precipitate appears at once,
0.5 cc. of the bichromate solution may be added before testing again.
The stirring rod which has touched the indicator should be dipped in
distilled water before returning it to the iron solution. As soon as
the blue appears to be less intense, add the bichromate solution in
small portions, finally a single drop at a time, until the point is
reached at which no blue color appears after the lapse of thirty
seconds from the time of mixing solution and indicator. At the close
of the titration a large drop of the iron solution should be taken for
the test. To determine the end-point beyond any question, as soon as
the thirty seconds have elapsed remove another drop of the solution
of the same size as that last taken and mix it with the indicator,
placing it beside the last previous test. If this last previous test
shows a blue tint in comparison with the fresh mixture, the end-point
has not been reached; if no difference can be noted the reaction is
complete. Should the end-point be overstepped, a little more of the
ferrous solution may be added and the end-point definitely fixed.
From the volumes of the solutions used, after applying corrections for
burette readings, and, if need be, for the temperature of solutions,
calculate the value of the ferrous solution in terms of the oxidizing
solution.
[Note 1: The accuracy of the work may be much impaired by the removal
of unnecessarily large quantities of solution for the tests. At the
beginning of the titration, while much ferrous iron is still present,
the end of the stirring rod need only be moist with the solution; but
at the close of the titration drops of considerable size may properly
be taken for the final tests. The stirring rod should be washed to
prevent transfer of indicator to the main solution. This cautious
removal of solution does not seriously affect the accuracy of the
determination, as it will be noted that the volume of the titrated
solution is about 200 cc. and the portions removed are very
small. Moreover, if the procedure is followed as prescribed, the
concentration of unoxidized iron decreases very rapidly as the
titration is carried out so that when the final tests are made, though
large drops may be taken, the amount of ferrous iron is not sufficient
to produce any appreciable error in results.
If the end-point is determined as prescribed, it can be as accurately
fixed as that of other methods; and if a ferrous solution is at
hand, the titration need consume hardly more time than that of the
permanganate process to be described later on.]
STANDARDIZATION OF POTASSIUM BICHROMATE SOLUTIONS
!Selection of a Standard!
A substance which will serve satisfactorily as a standard for
oxidizing solutions must possess certain specific properties: It must
be of accurately known composition and definite in its behavior as a
reducing agent, and it must be permanent against oxidation in the air,
at least for considerable periods. Such standards may take the form of
pure crystalline salts, such as ferrous ammonium sulphate, or may be
in the form of iron wire or an iron ore of known iron content. It is
not necessary that the standard should be of 100 per cent purity,
provided the content of the active reducing agent is known and no
interfering substances are present.
The two substances most commonly used as standards for a bichromate
solution are ferrous ammonium sulphate and iron wire. A standard wire
is to be purchased in the market which answers the purpose well, and
its iron content may be determined for each lot purchased by a number
of gravimetric determinations. It may best be preserved in jars
containing calcium chloride, but this must not be allowed to come
into contact with the wire. It should, however, even then be examined
carefully for rust before use.
If pure ferrous ammonium sulphate is used as the standard, clear
crystals only should be selected. It is perhaps even better to
determine by gravimetric methods once for all the iron content of a
large commercial sample which has been ground and well mixed. This
salt is permanent over long periods if kept in stoppered containers.
STANDARDIZATION
PROCEDURE.--Weigh out two portions of iron wire of about 0.24-0.26
gram each, examining the wire carefully for rust. It should be handled
and wiped with filter paper (not touched by the fingers), should
be weighed on a watch-glass, and be bent in such a way as not to
interfere with the movement of the balance.
Place 30 cc. of hydrochloric acid (sp. gr. 1.12) in each of two 300
cc. Erlenmeyer flasks, cover them with watch-glasses, and bring the
acid just to boiling. Remove them from the flame and drop in the
portions of wire, taking great care to avoid loss of liquid during
solution. Boil for two or three minutes, keeping the flasks covered
(Note 1), then wash the sides of the flasks and the watch-glass with
a little water and add stannous chloride solution to the hot liquid
!from a dropper! until the solution is colorless, but avoid more than
a drop or two in excess (Note 2). Dilute with 150 cc. of water and
cool !completely!. When cold, add rapidly about 30 cc. of mercuric
chloride solution. Allow the solutions to stand about three minutes
and then titrate without further delay (Note 3), add about 35 cc. of
the standard solution at once and finish the titration as prescribed
above, making use of the ferrous solution if the end-point should be
passed.
From the corrected volumes of the bichromate solution required to
oxidize the iron actually know to be present in the wire, calculate
the relation of the standard solution to the normal.
Repeat the standardization until the results are concordant within at
least two parts in one thousand.
[Note 1: The hydrochloric acid is added to the ferrous solution
to insure the presence of at least sufficient free acid for the
titration, as required by the equation on page 48.
The solution of the wire in hot acid and the short boiling insure the
removal of compounds of hydrogen and carbon which are formed from the
small amount of carbon in the iron. These might be acted upon by the
bichromate if not expelled.]
[Note 2: It is plain that all the iron must be reduced to the ferrous
condition before the titration begins, as some oxidation may have
occurred from the oxygen of the air during solution. It is also
evident that any excess of the agent used to reduce the iron must be
removed; otherwise it will react with the bichromate added later.
The reagents available for the reduction of iron are stannous
chloride, sulphurous acid, sulphureted hydrogen, and zinc; of these
stannous chloride acts most readily, the completion of the reaction
is most easily noted, and the excess of the reagent is most readily
removed. The latter object is accomplished by oxidation to stannic
chloride by means of mercuric chloride added in excess, as the
mercuric salts have no effect upon ferrous iron or the bichromate. The
reactions involved are:
2FeCl_{3} + SnCl_{2} --> 2FeCl_{2} + SnCl_{4}
SnCl_{2} + 2HgCl_{2} --> SnCl_{4} + 2HgCl
The mercurous chloride is precipitated.
It is essential that the solution should be cold and that the stannous
chloride should not be present in great excess, otherwise a secondary
reaction takes place, resulting in the reduction of the mercurous
chloride to metallic mercury:
SnCl_{2} + 2HgCl --> SnCl_{4} + 2Hg.
The occurrence of this secondary reaction is indicated by the
darkening of the precipitate; and, since potassium bichromate oxidizes
this mercury slowly, solutions in which it has been precipitated are
worthless as iron determinations.]
[Note 3: The solution should be allowed to stand about three minutes
after the addition of mercuric chloride to permit the complete
deposition of mercurous chloride. It should then be titrated without
delay to avoid possible reoxidation of the iron by the oxygen of the
air.]
DETERMINATION OF IRON IN LIMONITE
PROCEDURE.--Grind the mineral (Note 1) to a fine powder. Weigh out
accurately two portions of about 0.5 gram (Note 2) into porcelain
crucibles; heat these crucibles to dull redness for ten minutes,
allow them to cool, and place them, with their contents, in beakers
containing 30 cc. of dilute hydrochloric acid (sp. gr. 1.12). Heat
at a temperature just below boiling until the undissolved residue is
white or until solvent action has ceased. If the residue is white,
or known to be free from iron, it may be neglected and need not be
removed by filtration. If a dark residue remains, collect it on a
filter, wash free from hydrochloric acid, and ignite the filter in a
platinum crucible (Note 3). Mix the ash with five times its weight of
sodium carbonate and heat to fusion; cool, and disintegrate the fused
mass with boiling water in the crucible. Unite this solution and
precipitate (if any) with the acid solution, taking care to avoid loss
by effervescence. Wash out the crucible, heat the acid solution
to boiling, add stannous chloride solution until it is colorless,
avoiding a large excess (Note 4); cool, and when !cold!, add 40 cc. of
mercuric chloride solution, dilute to 200 cc., and proceed with the
titration as already described.
From the standardization data already obtained, and the known weight
of the sample, calculate the percentage of iron (Fe) in the limonite.
[Note 1: Limonite is selected as a representative of iron ores in
general. It is a native, hydrated oxide of iron. It frequently occurs
in or near peat beds and contains more or less organic matter which,
if brought into solution, would be acted upon by the potassium
bichromate. This organic matter is destroyed by roasting. Since a high
temperature tends to lessen the solubility of ferric oxide, the heat
should not be raised above low redness.]
[Note 2: It is sometimes advantageous to dissolve a large portion--say
5 grams--and to take one tenth of it for titration. The sample will
then represent more closely the average value of the ore.]
[Note 3: A platinum crucible may be used for the roasting of the
limonite and must be used for the fusion of the residue. When used, it
must not be allowed to remain in the acid solution of ferric chloride
for any length of time, since the platinum is attacked and dissolved,
and the platinic chloride is later reduced by the stannous chloride,
and in the reduced condition reacts with the bichromate, thus
introducing an error. It should also be noted that copper and antimony
interfere with the determination of iron by the bichromate process.]
[Note 4: The quantity of stannous chloride required for the reduction
of the iron in the limonite will be much larger than that added to the
solution of iron wire, in which the iron was mainly already in the
ferrous condition. It should, however, be added from a dropper to
avoid an unnecessary excess.]
DETERMINATION OF CHROMIUM IN CHROME IRON ORE
PROCEDURE.--Grind the chrome iron ore (Note 1) in an agate mortar
until no grit is perceptible under the pestle. Weigh out two portions
of 0.5 gram each into iron crucibles which have been scoured inside
until bright (Note 2). Weigh out on a watch-glass (Note 3), using the
rough balances, 5 grams of dry sodium peroxide for each portion, and
pour about three quarters of the peroxide upon the ore. Mix ore and
flux by thorough stirring with a dry glass rod. Then cover the mixture
with the remainder of the peroxide. Place the crucible on a triangle
and raise the temperature !slowly! to the melting point of the flux,
using a low flame, and holding the lamp in the hand (Note 4). Maintain
the fusion for five minutes, and stir constantly with a stout iron
wire, but do not raise the temperature above moderate redness (Notes 5
and 6).
Allow the crucible to cool until it can be comfortably handled (Note
7) and then place it in a 300 cc. beaker, and cover it with distilled
water (Note 8). The beaker must be carefully covered to avoid loss
during the disintegration of the fused mass. When the evolution of
gas ceases, rinse off and remove the crucible; then heat the solution
!while still alkaline! to boiling for fifteen minutes. Allow the
liquid to cool for a few minutes; then acidify with dilute sulphuric
acid (1:5), adding 10 cc. in excess of the amount necessary to
dissolve the ferric hydroxide (Note 9). Dilute to 200 cc., cool, add
from a burette an excess of a standard ferrous solution, and titrate
for the excess with a standard solution of potassium bichromate, using
the outside indicator (Note 10).
From the corrected volumes of the two standard solutions, and their
relations to normal solutions, calculate the percentage of chromium in
the ore.
[Note 1: Chrome iron ore is essentially a ferrous chromite, or
combination of FeO and Cr_{2}O_{3}. It must be reduced to a state of
fine subdivision to ensure a prompt reaction with the flux.]
[Note 2: The scouring of the iron crucible is rendered much easier if
it is first heated to bright redness and plunged into cold water. In
this process oily matter is burned off and adhering scale is caused to
chip off when the hot crucible contracts rapidly in the cold water.]
[Note 3: Sodium peroxide must be kept off of balance pans and should
not be weighed out on paper, as is the usual practice in the rough
weighing of chemicals. If paper to which the peroxide is adhering is
exposed to moist air it is likely to take fire as a result of
the absorption of moisture, and consequent evolution of heat and
liberation of oxygen.]
[Note 4: The lamp should never be allowed to remain under the
crucible, as this will raise the temperature to a point at which the
crucible itself is rapidly attacked by the flux and burned through.]
[Note 5: The sodium peroxide acts as both a flux and an oxidizing
agent. The chromic oxide is dissolved by the flux and oxidized to
chromic anhydride (CrO_{3}) which combines with the alkali to form
sodium chromate. The iron is oxidized to ferric oxide.]
[Note 6: The sodium peroxide cannot be used in porcelain, platinum, or
silver crucibles. It attacks iron and nickel as well; but crucibles
made from these metals may be used if care is exercised to keep the
temperature as low as possible. Preference is here given to iron
crucibles, because the resulting ferric hydroxide is more readily
brought into solution than the nickelic oxide from a nickel crucible.
The peroxide must be dry, and must be protected from any admixture of
dust, paper, or of organic matter of any kind, otherwise explosions
may ensue.]
[Note 7: When an iron crucible is employed it is desirable to allow
the fusion to become nearly cold before it is placed in water,
otherwise scales of magnetic iron oxide may separate from the
crucible, which by slowly dissolving in acid form ferrous sulphate,
which reduces the chromate.]
[Note 8: Upon treatment with water the chromate passes into solution,
the ferric hydroxide remains undissolved, and the excess of peroxide
is decomposed with the evolution of oxygen. The subsequent boiling
insures the complete decomposition of the peroxide. Unless this is
complete, hydrogen peroxide is formed when the solution is acidified,
and this reacts with the bichromate, reducing it and introducing a
serious error.]
[Note 9: The addition of the sulphuric acid converts the sodium
chromate to bichromate, which behaves exactly like potassium
bichromate in acid solution.]
[Note 10: If a standard solution of a ferrous salt is not at hand, a
weight of iron wire somewhat in excess of the amount which would be
required if the chromite were pure FeO.Cr_{2}O_{3} may be weighed out
and dissolved in sulphuric acid; after reduction of all the iron by
stannous chloride and the addition of mercuric chloride, this solution
may be poured into the chromate solution and the excess of iron
determined by titration with standard bichromate solution.]
PERMANGANATE PROCESS FOR THE DETERMINATION OF IRON
Potassium permanganate oxidizes ferrous salts in cold, acid solution
promptly and completely to the ferric condition, while in hot acid
solution it also enters into a definite reaction with oxalic acid, by
which the latter is oxidized to carbon dioxide and water.
The reactions involved are these:
10FeSO_{4} + 2KMnO_{4} + 8H_{2}S_{4} --> 5Fe_{2}(SO_{4})_{3} +
K_{2}SO_{4} + 2MnSO_{4} + 8H_{2}O
5C_{2}H_{2}O_{4}(2H_{2}O) + 2KMnO_{4} +3H_{2}SO_{4} --> K_{2}SO_{4} +
2MnSO_{4} + 10CO_{2} + 1 H_{2}O.
These are the fundamental reactions upon which the extensive use of
potassium permanganate depends; but besides iron and oxalic acid the
permanganate enters into reaction with antimony, tin, copper, mercury,
and manganese (the latter only in neutral solution), by which these
metals are changed from a lower to a higher state of oxidation; and it
also reacts with sulphurous acid, sulphureted hydrogen, nitrous acid,
ferrocyanides, and most soluble organic bodies. It should be noted,
however, that very few of these organic compounds react quantitatively
with the permanganate, as is the case with oxalic acid and the
oxalates.
Potassium permanganate is acted upon by hydrochloric acid; the action
is rapid in hot or concentrated solution (particularly in the presence
of iron salts, which appear to act as catalyzers, increasing the
velocity of the reaction), but slow in cold, dilute solutions.
However, the greater solubility of iron compounds in hydrochloric acid
makes it desirable to use this acid as a solvent, and experiments made
with this end in view have shown that in cold, dilute hydrochloric
acid solution, to which considerable quantities of manganous sulphate
and an excess of phosphoric acid have been added, it is possible to
obtain satisfactory results.
It is also possible to replace the hydrochloric acid by evaporating
the solutions with an excess of sulphuric acid until the latter fumes.
This procedure is somewhat more time-consuming, but the end-point of
the permanganate titration is more permanent. Both procedures are
described below.
Potassium permanganate has an intense coloring power, and since the
solution resulting from the oxidation of the iron and the reduction of
the permanganate is colorless, the latter becomes its own indicator.
The slightest excess is indicated with great accuracy by the pink
color of the solution.
PREPARATION OF A STANDARD SOLUTION
!Approximate Strength 0.1 N!
A study of the reactions given above which represent the oxidation of
ferrous compounds by potassium permanganate, shows that there are 2
molecules of KMnO_{4} and 10 molecules of FeSO_{4} on the
left-hand side, and 2 molecules of MnSO_{4} and 5 molecules of
Fe_{2}(SO_{4})_{5} on the right-hand side. Considering only these
compounds, and writing the formulas in such a way as to show the
oxides of the elements in each, the equation becomes:
K_{2}O.Mn_{2}O_{7} + 10(FeO.SO_{3}) --> K_{2}O.SO_{3} + 2(MnO.SO_{3})
+ 5(Fe_{2}O_{3}.3SO_{3}).
From this it appears that two molecules of KMnO_{4} (or 316.0 grams)
have given up five atoms (or 80 grams) of oxygen to oxidize the
ferrous compound. Since 8 grams of oxygen is the basis of normal
oxidizing solutions and 80 grams of oxygen are supplied by 316.0 grams
of KMnO_{4}, the normal solution of the permanganate should contain,
per liter, 316.0/10 grams, or 31.60 grams (Note 1).
The preparation of an approximately tenth-normal solution of the
reagent may be carried out as follows:
PROCEDURE.--Dissolve about 3.25 grams of potassium permanganate
crystals in approximately 1000 cc. of distilled water in a large
beaker, or casserole. Heat slowly and when the crystals have
dissolved, boil the solution for 10-15 minutes. Cover the solution
with a watch-glass; allow it to stand until cool, or preferably over
night. Filter the solution through a layer of asbestos. Transfer the
filtrate to a liter bottle and mix thoroughly (Note 2).
[Note 1: The reactions given on page 61 are those which take place in
the presence of an excess of acid. In neutral solutions the reduction
of the permanganate is less complete, and, under these conditions,
two gram-molecular weights of KMnO_{4} will furnish only 48 grams
of oxygen. A normal solution for use under these conditions should,
therefore, contain 316.0/6 grams, or 52.66 grams.]
[Note 2: Potassium permanganate solutions are not usually stable for
long periods, and change more rapidly when first prepared than after
standing some days. This change is probably caused by interaction
with the organic matter contained in all distilled water, except that
redistilled from an alkaline permanganate solution. The solutions
should be protected from light and heat as far as possible, since both
induce decomposition with a deposition of manganese dioxide, and it
has been shown that decomposition proceeds with considerable rapidity,
with the evolution of oxygen, after the dioxide has begun to form. As
commercial samples of the permanganate are likely to be contaminated
by the dioxide, it is advisable to boil and filter solutions through
asbestos before standardization, as prescribed above. Such solutions
are relatively stable.]
COMPARISON OF PERMANGANATE AND FERROUS SOLUTIONS
PROCEDURE.--Fill a glass-stoppered burette with the permanganate
solution, observing the usual precautions, and fill a second burette
with the ferrous sulphate solution prepared for use with the potassium
bichromate. The permanganate solution cannot be used in burettes with
rubber tips, as a reduction takes place upon contact with the rubber.
The solution has so deep a color that the lower line of the meniscus
cannot be detected; readings must therefore be made from the upper
edge. Run out into a beaker about 40 cc. of the ferrous solution,
dilute to about 100 cc., add 10 cc. of dilute sulphuric acid, and run
in the permanganate solution to a slight permanent pink. Repeat, until
the ratio of the two solutions is satisfactorily established.
STANDARDIZATION OF A POTASSIUM PERMANGANATE SOLUTION
!Selection of a Standard!
Commercial potassium permanganate is rarely sufficiently pure to admit
of its direct weighing as a standard. On this account, and because
of the uncertainties as to the permanence of its solutions, it is
advisable to standardize them against substances of known value. Those
in most common use are iron wire, ferrous ammonium sulphate, sodium
oxalate, oxalic acid, and some other derivatives of oxalic acid.
With the exception of sodium oxalate, these all contain water of
crystallization which may be lost on standing. They should, therefore,
be freshly prepared, and with great care. At present, sodium oxalate
is considered to be one of the most satisfactory standards.
!Method A!
!Iron Standards!
The standardization processes employed when iron or its compounds are
selected as standards differ from those applicable in connection with
oxalate standards. The procedure which immediately follows is that in
use with iron standards.
As in the case of the bichromate process, it is necessary to reduce
the iron completely to the ferrous condition before titration. The
reducing agents available are zinc, sulphurous acid, or sulphureted
hydrogen. Stannous chloride may also be used when the titration is
made in the presence of hydrochloric acid. Since the excess of both
the gaseous reducing agents can only be expelled by boiling, with
consequent uncertainty regarding both the removal of the excess and
the reoxidation of the iron, zinc or stannous chlorides are the most
satisfactory agents. For prompt and complete reduction it is essential
that the iron solution should be brought into ultimate contact with
the zinc. This is brought about by the use of a modified Jones
reductor, as shown in Figure 1. This reductor is a standard apparatus
and is used in other quantitative processes.
[Illustration: Fig. 1]
The tube A has an inside diameter of 18 mm. and is 300 mm. long; the
small tube has an inside diameter of 6 mm. and extends 100 mm. below
the stopcock. At the base of the tube A are placed some pieces of
broken glass or porcelain, covered by a plug of glass wool about 8 mm.
thick, and upon this is placed a thin layer of asbestos, such as is
used for Gooch filters, 1 mm. thick. The tube is then filled with the
amalgamated zinc (Note 1) to within 50 mm. of the top, and on the zinc
is placed a plug of glass wool. If the top of the tube is not already
shaped like the mouth of a thistle-tube (B), a 60 mm. funnel is fitted
into the tube with a rubber stopper and the reductor is connected
with a suction bottle, F. The bottle D is a safety bottle to
prevent contamination of the solution by water from the pump. After
preparation for use, or when left standing, the tube A should be
filled with water, to prevent clogging of the zinc.
[Note 1: The use of fine zinc in the reductor is not necessary and
tends to clog the tube. Particles which will pass a 10-mesh sieve, but
are retained by one of 20 meshes to the inch, are most satisfactory.
The zinc can be amalgamated by stirring or shaking it in a mixture of
25 cc. of normal mercuric chloride solution, 25 cc. of hydrochloric
acid (sp. gr. 1.12) and 250 cc. of water for two minutes. The solution
should then be poured off and the zinc thoroughly washed. It is then
ready for bottling and preservation under water. A small quantity of
glass wool is placed in the neck of the funnel to hold back foreign
material when the reductor is in use.]
STANDARDIZATION
PROCEDURE.--Weigh out into Erlenmeyer flasks two portions of iron wire
of about 0.25 gram each. Dissolve these in hot dilute sulphuric acid
(5 cc. of concentrated acid and 100 cc. of water), using a covered
flask to avoid loss by spattering. Boil the solution for two or
three minutes after the iron has dissolved to remove any volatile
hydrocarbons. Meanwhile prepare the reductor for use as follows:
Connect the vacuum bottle with the suction pump and pour into the
funnel at the top warm, dilute sulphuric acid, prepared by adding 5
cc. of concentrated sulphuric acid to 100 cc. of distilled water. See
that the stopcock (C) is open far enough to allow the acid to run
through slowly. Continue to pour in acid until 200 cc. have passed
through, then close the stopcock !while a small quantity of liquid
is still left in the funnel!. Discard the filtrate, and again
pass through 100 cc. of the warm, dilute acid. Test this with the
permanganate solution. A single drop should color it permanently; if
it does not, repeat the washing, until assured that the zinc is not
contaminated with appreciable quantities of reducing substances. Be
sure that no air enters the reductor (Note 1).
Pour the iron solution while hot (but not boiling) through the
reductor at a rate not exceeding 50 cc. per minute (Notes 2 and 3).
Wash out the beaker with dilute sulphuric acid, and follow the iron
solution without interruption with 175 cc. of the warm acid and
finally with 75 cc. of distilled water, leaving the funnel partially
filled. Remove the filter bottle and cool the solution quickly under
the water tap (Note 4), avoiding unnecessary exposure to the oxygen of
the air. Add 10 cc. of dilute sulphuric acid and titrate to a faint
pink with the permanganate solution, adding it directly to the
contents of the vacuum flask. Should the end-point be overstepped, the
ferrous sulphate solution may be added.
From the volume of the solution required to oxidize the iron in
the wire, calculate the relation to the normal of the permanganate
solution. The duplicate results should be concordant within two parts
in one thousand.
[Note 1: The funnel of the reductor must never be allowed to empty.
If it is left partially filled with water the reductor is ready for
subsequent use after a very little washing; but a preliminary test is
always necessary to safeguard against error.
If more than a small drop of permanganate solution is required to
color 100 cc. of the dilute acid after the reductor is well washed, an
allowance must be made for the iron in the zinc. !Great care! must be
used to prevent the access of air to the reductor after it has been
washed out ready for use. If air enters, hydrogen peroxide forms,
which reacts with the permanganate, and the results are worthless.]
[Note 2: The iron is reduced to the ferrous condition by contact with
the zinc. The active agent may be considered to be !nascent! hydrogen,
and it must be borne in mind that the visible bubbles are produced by
molecular hydrogen, which is without appreciable effect upon ferric
iron.
The rate at which the iron solution passes through the zinc should not
exceed that prescribed, but the rate may be increased somewhat when
the wash-water is added. It is well to allow the iron solution to run
nearly, but not entirely, out of the funnel before the wash-water
is added. If it is necessary to interrupt the process, the complete
emptying of the funnel can always be avoided by closing the stopcock.
It is also possible to reduce the iron by treatment with zinc in a
flask from which air is excluded. The zinc must be present in excess
of the quantity necessary to reduce the iron and is finally completely
dissolved. This method is, however, less convenient and more tedious
than the use of the reductor.]
[Note 3: The dilute sulphuric acid for washing must be warmed ready
for use before the reduction of the iron begins, and it is of the
first importance that the volume of acid and of wash-water should
be measured, and the volume used should always be the same in the
standardizations and all subsequent analyses.]
[Note 4: The end-point is more permanent in cold than hot solutions,
possibly because of a slight action of the permanganate upon the
manganous sulphate formed during titration. If the solution turns
brown, it is an evidence of insufficient acid, and more should be
immediately added. The results are likely to be less accurate in this
case, however, as a consequence of secondary reactions between the
ferrous iron and the manganese dioxide thrown down. It is wiser to
discard such results and repeat the process.]
[Note 5: The potassium permanganate may, of course, be diluted and
brought to an exactly 0.1 N solution from the data here obtained. The
percentage of iron in the iron wire must be taken into account in all
calculations.]
!Method B!
!Oxalate Standards!
PROCEDURE.--Weigh out two portions of pure sodium oxalate of 0.25-0.3
gram each into beakers of about 600 cc. capacity. Add about 400 cc. of
boiling water and 20 cc. of manganous sulphate solution (Note 1).
When the solution of the oxalate is complete, heat the liquid, if
necessary, until near its boiling point (70-90°C.) and run in the
standard permanganate solution drop by drop from a burette, stirring
constantly until an end-point is reached (Note 2). Make a blank test
with 20 cc. of manganous sulphate solution and a volume of distilled
water equal to that of the titrated solution to determine the volume
of the permanganate solution required to produce a very slight pink.
Deduct this volume from the amount of permanganate solution used in
the titration.
From the data obtained, calculate the relation of the permanganate
solution to the normal. The reaction involved is:
5Na_{2}C_{2}O_{4} + 2KMnO_{4} + 8H_{2}SO_{4} --> 5Na_{2}SO_{4} +
K_{2}SO_{4} + 2MnSO_{4} + 10CO_{2} + 8H_{2}O
[Note 1: The manganous sulphate titrating solution is made by
dissolving 20 grams of MnSO_{4} in 200 cubic centimeters of water and
adding 40 cc. of concentrated sulphuric acid (sp. gr. 1.84) and 40 cc.
or phosphoric acid (85%).]
[Note 2: The reaction between oxalates and permanganates takes place
quantitatively only in hot acid solutions. The temperatures must not
fall below 70°C.]
DETERMINATION OF IRON IN LIMONITE
!Method A!
The procedures, as here prescribed, are applicable to iron ores in
general, provided these ores contain no constituents which are reduced
by zinc or stannous chloride and reoxidized by permanganates. Many
iron ores contain titanium, and this element among others does
interfere with the determination of iron by the process described.
If, however, the solutions of such ores are treated with sulphureted
hydrogen or sulphurous acid, instead of zinc or stannous chloride to
reduce the iron, and the excess reducing agent removed by boiling, an
accurate determination of the iron can be made.
PROCEDURE.--Grind the mineral to a fine powder. Weigh out two portions
of about 0.5 gram each into small porcelain crucibles. Roast the ore
at dull redness for ten minutes (Note 1), allow the crucibles to cool,
and place them and their contents in casseroles containing 30 cc. of
dilute hydrochloric acid (sp. gr. 1.12).
Proceed with the solution of the ore, and the treatment of the
residue, if necessary, exactly as described for the bichromate process
on page 56. When solution is complete, add 6 cc. of concentrated
sulphuric acid to each casserole, and evaporate on the steam bath
until the solution is nearly colorless (Note 2). Cover the casseroles
and heat over the flame of the burner, holding the casserole in
the hand and rotating it slowly to hasten evaporation and prevent
spattering, until the heavy white fumes of sulphuric anhydride are
freely evolved (Note 3). Cool the casseroles, add 100 cc. of water
(measured), and boil gently until the ferric sulphate is dissolved;
pour the warm solution through the reductor which has been previously
washed; proceed as described under standardization, taking pains
to use the same volume and strength of acid and the same volume of
wash-water as there prescribed, and titrate with the permanganate
solution in the reductor flask, using the ferrous sulphate solution if
the end-point should be overstepped.
From the corrected volume of permanganate solution used, calculate the
percentage of iron (Fe) in the limonite.
[Note 1: The preliminary roasting is usually necessary because, even
though the sulphuric acid would subsequently char the carbonaceous
matter, certain nitrogenous bodies are not thereby rendered insoluble
in the acid, and would be oxidized by the permanganate.]
[Note 2: The temperature of the steam bath is not sufficient to
volatilize sulphuric acid. Solutions may, therefore, be left to
evaporate overnight without danger of evaporation to dryness.]
[Note 3: The hydrochloric acid, both free and combined, is displaced
by the less volatile sulphuric acid at its boiling point. Ferric
sulphate separates at this point, since there is no water to hold
it in solution and care is required to prevent bumping. The ferric
sulphate usually has a silky appearance and is easily distinguished
from the flocculent silica which often remains undissolved.]
!Zimmermann-Reinhardt Procedure!
!Method (B)!
PROCEDURE.--Grind the mineral to a fine powder. Weigh out two portions
of about 0.5 gram each into small porcelain crucibles. Proceed with
the solution of the ore, treat the residue, if necessary, and reduce
the iron by the addition of stannous chloride, followed by mercuric
chloride, as described for the bichromate process on page 56. Dilute
the solution to about 400 cc. with cold water, add 10 cc. of the
manganous sulphate titrating solution (Note 1, page 68) and titrate
with the standard potassium permanganate solution to a faint pink
(Note 1).
From the standardization data already obtained calculate the
percentage of iron (Fe) in the limonite.
[Note 1: It has already been noted that hydrochloric acid reacts
slowly in cold solutions with potassium permanganate. It is, however,
possible to obtain a satisfactory, although somewhat fugitive
end-point in the presence of manganous sulphate and phosphoric acid.
The explanation of the part played by these reagents is somewhat
obscure as yet. It is possible that an intermediate manganic compound
is formed which reacts rapidly with the ferrous compounds--thus in
effect catalyzing the oxidizing process.
While an excess of hydrochloric acid is necessary for the successful
reduction of the iron by stannous chloride, too large an amount
should be avoided in order to lessen the chance of reduction of the
permanganate by the acid during titration.]
DETERMINATION OF THE OXIDIZING POWER OF PYROLUSITE
INDIRECT OXIDATION
Pyrolusite, when pure, consists of manganese dioxide. Its value as an
oxidizing agent, and for the production of chlorine, depends upon the
percentage of MnO_{2} in the sample. This percentage is determined
by an indirect method, in which the manganese dioxide is reduced and
dissolved by an excess of ferrous sulphate or oxalic acid in the
presence of sulphuric acid, and the unused excess determined by
titration with standard permanganate solution.
PROCEDURE.--Grind the mineral in an agate mortar until no grit
whatever can be detected under the pestle (Note 1). Transfer it to a
stoppered weighing-tube, and weigh out two portions of about 0.5 gram
into beakers (400-500 cc.) Read Note 2, and then calculate in each
case the weight of oxalic acid (H_{2}C_{2}O_{4}.2H_{2}O) required to
react with the weights of pyrolusite taken. The reaction involved is
MnO_{2} + H_{2}C_{2}O_{4}(2H_{2}O) + H_{2}SO_{4} --> MnSO_{4} +
2CO_{2} + 4H_{2}O.
Weigh out about 0.2 gram in excess of this quantity of !pure! oxalic
acid into the corresponding beakers, weighing the acid accurately and
recording the weight in the notebook. Pour into each beaker 25 cc. of
water and 50 cc. of dilute sulphuric acid (1:5), cover and warm the
beaker and its contents gently until the evolution of carbon dioxide
ceases (Note 3). If a residue remains which is sufficiently colored to
obscure the end-reaction of the permanganate, it must be removed by
filtration.
Finally, dilute the solution to 200-300 cc., heat the solution to a
temperature just below boiling, add 15 cc. of a manganese sulphate
solution and while hot, titrate for the excess of the oxalic acid with
standard permanganate solution (Notes 4 and 5).
From the corrected volume of the solution required, calculate the
amount of oxalic acid undecomposed by the pyrolusite; subtract this
from the total quantity of acid used, and calculate the weight of
manganese dioxide which would react with the balance of the acid, and
from this the percentage in the sample.
[Note 1: The success of the analysis is largely dependent upon the
fineness of the powdered mineral. If properly ground, solution should
be complete in fifteen minutes or less.]
[Note 2: A moderate excess of oxalic acid above that required to react
with the pyrolusite is necessary to promote solution; otherwise the
residual quantity of oxalic acid would be so small that the last
particles of the mineral would scarcely dissolve. It is also desirable
that a sufficient excess of the acid should be present to react with a
considerable volume of the permanganate solution during the titration,
thus increasing the accuracy of the process. On the other hand, the
excess of oxalic acid should not be so large as to react with more of
the permanganate solution than is contained in a 50 cc. burette. If
the pyrolusite under examination is known to be of high grade, say 80
per cent pure, or above the calculation of the oxalic acid needed may
be based upon an assumption that the mineral is all MnO_{2}. If the
quality of the mineral is unknown, it is better to weigh out three
portions instead of two and to add to one of these the amount of
oxalic prescribed, assuming complete purity of the mineral. Then run
in the permanganate solution from a pipette or burette to determine
roughly the amount required. If the volume exceeds the contents of a
burette, the amount of oxalic acid added to the other two portions is
reduced accordingly.]
[Note 3: Care should be taken that the sides of the beaker are not
overheated, as oxalic acid would be decomposed by heat alone if
crystallization should occur on the sides of the vessel. Strong
sulphuric acid also decomposes the oxalic acid. The dilute acid
should, therefore, be prepared before it is poured into the beaker.]
[Note 4: Ferrous ammonium sulphate, ferrous sulphate, or iron wire
may be substituted for the oxalic acid. The reaction is then the
following:
2 FeSO_{4} + MnO_{2} + 2H_{2}SO_{4} --> Fe_{2}(SO_{4})_{3} + 2H_{2}O
The excess of ferrous iron may also be determined by titration with
potassium bichromate, if desired. Care is required to prevent the
oxidation of the iron by the air, if ferrous salts are employed.]
[Note 5: The oxidizing power of pyrolusite may be determined by other
volumetric processes, one of which is outlined in the following
reactions:
MnO_{2} + 4HCl --> MnCl_{2} + Cl_{2} + 2H_{2}O
Cl_{2} + 2KI --> I_{2} + 2KCl
I_{2} + 2Na_{2}S_{2}O_{3} --> Na_{2}S_{4}O_{6} + 2NaI.
The chlorine generated by the pyrolusite is passed into a solution of
potassium iodide. The liberated iodine is then determined by titration
with sodium thiosulphate, as described on page 78. This is a direct
process, although it involves three steps.]
IODIMETRY
The titration of iodine against sodium thiosulphate, with starch as an
indicator, may perhaps be regarded as the most accurate of volumetric
processes. The thiosulphate solution may be used in both acid and
neutral solutions to measure free iodine and the latter may, in turn,
serve as a measure of any substance capable of liberating iodine from
potassium iodide under suitable conditions for titration, as, for
example, in the process outlined in Note 5 on page 74.
The fundamental reaction upon which iodometric processes are based is
the following:
I_{2} + 2 Na_{2}S_{2}O_{3} --> 2 NaI + Na_{2}S_{4}O_{6}.
This reaction between iodine and sodium thiosulphate, resulting in
the formation of the compound Na_{2}S_{4}O_{6}, called sodium
tetrathionate, is quantitatively exact, and differs in that
respect from the action of chlorine or bromine, which oxidize the
thiosulphate, but not quantitatively.
NORMAL SOLUTIONS OF IODINE AND SODIUM THIOSULPHATE
If the formulas of sodium thiosulphate and sodium tetrathionate are
written in a manner to show the atoms of oxygen associated
with sulphur atoms in each, thus, 2(Na_{2}).S_{2}O_{2} and
Na_{2}O.S_{4}O_{5}, it is plain that in the tetrathionate there are
five atoms of oxygen associated with sulphur, instead of the four
in the two molecules of the thiosulphate taken together. Although,
therefore, the iodine contains no oxygen, the two atoms of iodine
have, in effect, brought about the addition of one oxygen atoms to the
sulphur atoms. That is the same thing as saying that 253.84 grams of
iodine (I_{2}) are equivalent to 16 grams of oxygen; hence, since 8
grams of oxygen is the basis of normal solutions, 253.84/2 or 126.97
grams of iodine should be contained in one liter of normal iodine
solution. By a similar course of reasoning the conclusion is reached
that the normal solution of sodium thiosulphate should contain,
per liter, its molecular weight in grams. As the thiosulphate in
crystalline form has the formula Na_{2}S_{2}O_{3}.5H_{2}O, this weight
is 248.12 grams. Tenth-normal or hundredth-normal solutions are
generally used.
PREPARATION OF STANDARD SOLUTIONS
!Approximate Strength, 0.1 N!
PROCEDURE.--Weigh out on the rough balances 13 grams of commercial
iodine. Place it in a mortar with 18 grams of potassium iodide and
triturate with small portions of water until all is dissolved. Dilute
the solution to 1000 cc. and transfer to a liter bottle and mix
thoroughly (Note 1).[1]
[Footnote 1: It will be found more economical to have a considerable
quantity of the solution prepared by a laboratory attendant, and to
have all unused solutions returned to the common stock.]
Weigh out 25 grams of sodium thiosulphate, dissolve it in water which
has been previously boiled and cooled, and dilute to 1000 cc., also
with boiled water. Transfer the solution to a liter bottle and mix
thoroughly (Note 2).
[Note 1: Iodine solutions react with water to form hydriodic acid
under the influence of the sunlight, and even at low room temperatures
the iodine tends to volatilize from solution. They should, therefore,
be protected from light and heat. Iodine solutions are not stable for
long periods under the best of conditions. They cannot be used in
burettes with rubber tips, since they attack the rubber.]
[Note 2: Sodium thiosulphate (Na_{2}S_{2}O_{3}.5H_{2}O) is
rarely wholly pure as sold commercially, but may be purified by
recrystallization. The carbon dioxide absorbed from the air by
distilled water decomposes the salt, with the separation of sulphur.
Boiled water which has been cooled out of contact with the air should
be used in preparing solutions.]
INDICATOR SOLUTION
The starch solution for use as an indicator must be freshly prepared.
A soluble starch is obtainable which serves well, and a solution of
0.5 gram of this starch in 25 cc. of boiling water is sufficient. The
solution should be filtered while hot and is ready for use when cold.
If soluble starch is not at hand, potato starch may be used. Mix about
1 gram with 5 cc. of cold water to a smooth paste, pour 150 cc. of
!boiling! water over it, warm for a moment on the hot plate, and put
it aside to settle. Decant the supernatant liquid through a filter
and use the clear filtrate; 5 cc. of this solution are needed for a
titration.
The solution of potato starch is less stable than the soluble starch.
The solid particles of the starch, if not removed by filtration,
become so colored by the iodine that they are not readily decolorized
by the thiosulphate (Note 1).
[Note 1: The blue color which results when free iodine and starch
are brought together is probably not due to the formation of a true
chemical compound. It is regarded as a "solid solution" of iodine in
starch. Although it is unstable, and easily destroyed by heat, it
serves as an indicator for the presence of free iodine of remarkable
sensitiveness, and makes the iodometric processes the most
satisfactory of any in the field of volumetric analysis.]
COMPARISON OF IODINE AND THIOSULPHATE SOLUTIONS
PROCEDURE.--Place the solutions in burettes (the iodine in a
glass-stoppered burette), observing the usual precautions. Run out 40
cc. of the thiosulphate solution into a beaker, dilute with 150 cc. of
water, add 1 cc. to 2 cc. of the soluble starch solution, and titrate
with the iodine to the appearance of the blue of the iodo-starch.
Repeat until the ratio of the two solutions is established,
remembering all necessary corrections for burettes and for temperature
changes.
STANDARDIZATION OF SOLUTIONS
Commercial iodine is usually not sufficiently pure to permit of its
use as a standard for thiosulphate solutions or the direct preparation
of a standard solution of iodine. It is likely to contain, beside
moisture, some iodine chloride, if chlorine was used to liberate the
iodine when it was prepared. It may be purified by sublimation after
mixing it with a little potassium iodide, which reacts with the iodine
chloride, forming potassium chloride and setting free the iodine. The
sublimed iodine is then dried by placing it in a closed container over
concentrated sulphuric acid. It may then be weighed in a stoppered
weighing-tube and dissolved in a solution of potassium iodide in a
stoppered flask to prevent loss of iodine by volatilization. About 18
grams of the iodide and twelve grams of iodine per liter are required
for an approximately tenth-normal solution.
An iodine solution made from commercial iodine may also be
standardized against arsenious oxide (As_{4}O_{6}). This substance
also usually requires purification by sublimation before use.
The substances usually employed for the standardization of a
thiosulphate solution are potassium bromate and metallic copper. The
former is obtainable in pure condition or may be easily purified by
re-crystallization. Copper wire of high grade is sufficiently pure
to serve as a standard. Both potassium bromate and cupric salts in
solution will liberate iodine from an iodide, which is then titrated
with the thiosulphate solution.
The reactions involved are the following:
(a) KBrO_{3} + 6KI + 3H_{2}SO_{4} --> KBr + 3I_{2} + 3K_{2}SO_{4} + 3H_{2}O,
(b) 3Cu + 8HNO_{3} --> 3Cu(NO_{3})_{2} + 2NO + 4H_{2}O,
2Cu(NO_{3})_{2} + 4KI --> 2CuI + 4KNO_{3} + I_{2}.
Two methods for the direct standardization of the sodium thiosulphate
solution are here described, and one for the direct standardization of
the iodine solution.
!Method A!
PROCEDURE.--Weigh out into 500 cc. beakers two portions of about
0.150-0.175 gram of potassium bromate. Dissolve each of these in 50
cc. of water, and add 10 cc. of a potassium iodide solution containing
3 grams of the salt in that volume (Note 1). Add to the mixture 10 cc.
of dilute sulphuric acid (1 volume of sulphuric acid with 5 volumes of
water), allow the solution to stand for three minutes, and dilute to
150 cc. (Note 2). Run in thiosulphate solution from a burette until
the color of the liberated iodine is nearly destroyed, and then add 1
cc. or 2 cc. of starch solution, titrate to the disappearance of the
iodo-starch blue, and finally add iodine solution until the color
is just restored. Make a blank test for the amount of thiosulphate
solution required to react with the iodine liberated by the iodate
which is generally present in the potassium iodide solution, and
deduct this from the total volume used in the titration.
From the data obtained, calculate the relation of the thiosulphate
solution to a normal solution, and subsequently calculate the similar
value for the iodine solution.
[Note 1:--Potassium iodide usually contains small amounts of potassium
iodate as impurity which, when the iodide is brought into an acid
solution, liberates iodine, just as does the potassium bromate used as
a standard. It is necessary to determine the amount of thiosulphate
which reacts with the iodine thus liberated by making a "blank test"
with the iodide and acid alone. As the iodate is not always uniformly
distributed throughout the iodide, it is better to make up a
sufficient volume of a solution of the iodide for the purposes of the
work in hand, and to make the blank test by using the same volume of
the iodide solution as is added in the standardizing process. The
iodide solution should contain about 3 grams of the salt in 10 cc.]
[Note 2: The color of the iodo-starch is somewhat less satisfactory in
concentrated solutions of the alkali salts, notably the iodides. The
dilution prescribed obviates this difficulty.]
!Method B!
PROCEDURE.--Weigh out two portions of 0.25-0.27 gram of clean copper
wire into 250 cc. Erlenmeyer flasks (Note 1). Add to each 5 cc. of
concentrated nitric acid (sp. gr. 1.42) and 25 cc. of water, cover,
and warm until solution is complete. Add 5 cc. of bromine water and
boil until the excess of bromine is expelled. Cool, and add strong
ammonia (sp. gr. 0.90) drop by drop until a deep blue color indicates
the presence of an excess. Boil the solution until the deep blue is
replaced by a light bluish green, or a brown stain appears on the
sides of the flask (Note 2). Add 10 cc. of strong acetic acid (sp.
gr. 1.04), cool under the water tap, and add a solution of potassium
iodide (Note 3) containing about 3 grams of the salt, and titrate
with thiosulphate solution until the color of the liberated iodine
is nearly destroyed. Then add 1-2 cc. of freshly prepared starch
solution, and add thiosulphate solution, drop by drop, until the blue
color is discharged.
From the data obtained, including the "blank test" of the iodide,
calculate the relation of the thiosulphate solution to the normal.
[Note 1: While copper wire of commerce is not absolutely pure, the
requirements for its use as a conductor of electricity are such that
the impurities constitute only a few hundredths of one per cent and
are negligible for analytical purposes.]
[Note 2: Ammonia neutralizes the free nitric acid. It should be added
in slight excess only, since the excess must be removed by boiling,
which is tedious. If too much ammonia is present when acetic acid is
added, the resulting ammonium acetate is hydrolyzed, and the ammonium
hydroxide reacts with the iodine set free.]
[Note 3: A considerable excess of potassium iodide is necessary for
the prompt liberation of iodine. While a large excess will do no harm,
the cost of this reagent is so great that waste should be avoided.]
!Method C!
PROCEDURE.--Weigh out into 500 cc. beakers two portions of 0.175-0.200
gram each of pure arsenious oxide. Dissolve each of these in 10 cc. of
sodium hydroxide solution, with stirring. Dilute the solutions to 150
cc. and add dilute hydrochloric acid until the solutions contain a few
drops in excess, and finally add to each a concentrated solution of
5 grams of pure sodium bicarbonate (NaHCO_{3}) in water. Cover the
beakers before adding the bicarbonate, to avoid loss. Add the starch
solution and titrate with the iodine to the appearance of the blue of
the iodo-starch, taking care not to pass the end-point by more than a
few drops (Note 1).
From the corrected volume of the iodine solution used to oxidize the
arsenious oxide, calculate its relation to the normal. From the
ratio between the solutions, calculate the similar value for the
thiosulphate solution.
[Note 1: Arsenious oxide dissolves more readily in caustic alkali than
in a bicarbonate solution, but the presence of caustic alkali during
the titration is not admissible. It is therefore destroyed by the
addition of acid, and the solution is then made neutral with the
solution of bicarbonate, part of which reacts with the acid, the
excess remaining in solution.
The reaction during titration is the following:
Na_{3}AsO_{3} + I_{2} + 2NaHCO_{3} --> Na_{3}AsO_{4} + 2NaI + 2CO_{2}
+ H_{2}O
As the reaction between sodium thiosulphate and iodine is not always
free from secondary reactions in the presence of even the weakly
alkaline bicarbonate, it is best to avoid the addition of any
considerable excess of iodine. Should the end-point be passed by a few
drops, the thiosulphate may be used to correct it.]
DETERMINATION OF COPPER IN ORES
Copper ores vary widely in composition from the nearly pure copper
minerals, such as malachite and copper sulphide, to very low grade
materials which contain such impurities as silica, lead, iron, silver,
sulphur, arsenic, and antimony. In nearly all varieties there will be
found a siliceous residue insoluble in acids. The method here given,
which is a modification of that described by A.H. Low (!J. Am. Chem.
Soc.! (1902), 24, 1082), provides for the extraction of the copper
from commonly occurring ores, and for the presence of their common
impurities. For practice analyses it is advisable to select an ore of
a fair degree of purity.
PROCEDURE.-- Weigh out two portions of about 0.5 gram each of the
ore (which should be ground until no grit is detected) into 250 cc.
Erlenmeyer flasks or small beakers. Add 10 cc. of concentrated nitric
acid (sp. gr. 1.42) and heat very gently until the ore is decomposed
and the acid evaporated nearly to dryness (Note 1). Add 5 cc. of
concentrated hydrochloric acid (sp. gr. 1.2) and warm gently. Then
add about 7 cc. of concentrated sulphuric acid (sp. gr. 1.84) and
evaporate over a free flame until the sulphuric acid fumes freely
(Note 2). It has then displaced nitric and hydrochloric acid from
their compounds.
Cool the flask or beaker, add 25 cc. of water, heat the solution
to boiling, and boil for two minutes. Filter to remove insoluble
sulphates, silica and any silver that may have been precipitated as
silver chloride, and receive the filtrate in a small beaker, washing
the precipitate and filter paper with warm water until the filtrate
and washings amount to 75 cc. Bend a strip of aluminium foil (5 cm. x
12 cm.) into triangular form and place it on edge in the beaker. Cover
the beaker and boil the solution (being careful to avoid loss of
liquid by spattering) for ten minutes, but do not evaporate to small
volume.
Wash the cover glass and sides of the beaker. The copper should now be
in the form of a precipitate at the bottom of the beaker or adhering
loosely to the aluminium sheet. Remove the sheet, wash it carefully
with hydrogen sulphide water and place it in a small beaker. Decant
the solution through a filter, wash the precipitated copper twice by
decantation with hydrogen sulphide water, and finally transfer the
copper to the filter paper, where it is again washed thoroughly, being
careful at all times to keep the precipitated copper covered with the
wash water. Remove and discard the filtrate and place an Erlenmeyer
flask under the funnel. Pour 15 cc. of dilute nitric acid (sp. gr.
1.20) over the aluminium foil in the beaker, thus dissolving any
adhering copper. Wash the foil with hot water and remove it. Warm this
nitric acid solution and pour it slowly through the filter paper,
thereby dissolving the copper on the paper, receiving the acid
solution in the Erlenmeyer flask. Before washing the paper, pour 5 cc.
of saturated bromine water (Note 3) through it and finally wash the
paper carefully with hot water and transfer any particles of copper
which may be left on it to the Erlenmeyer flask. Boil to expel the
bromine. Add concentrated ammonia drop by drop until the appearance of
a deep blue coloration indicates an excess. Boil until the deep blue
is displaced by a light bluish green coloration, or until brown stains
form on the sides of the flask. Add 10 cc. of strong acetic acid (Note
4) and cool under the water tap. Add a solution containing about 3
grams of potassium iodide, as in the standardization, and titrate with
thiosulphate solution until the yellow of the liberated iodine is
nearly discharged. Add 1-2 cc. of freshly prepared starch solution and
titrate to the disappearance of the blue color.
From the data obtained, calculate the percentage of copper (Cu) in the
ore.
[Note 1: Nitric acid, because of its oxidizing power, is used as a
solvent for the sulphide ores. As a strong acid it will also dissolve
the copper from carbonate ores. The hydrochloric acid is added to
dissolve oxides of iron and to precipitate silver and lead. The
sulphuric acid displaces the other acids, leaving a solution
containing sulphates only. It also, by its dehydrating action, renders
silica from silicates insoluble.]
[Note 2: Unless proper precautions are taken to insure the correct
concentrations of acid the copper will not precipitate quantitatively
on the aluminium foil; hence care must be taken to follow directions
carefully at this point. Lead and silver have been almost completely
removed as sulphate and chloride respectively, or they too would
be precipitated on the aluminium. Bismuth, though precipitated on
aluminium, has no effect on the analysis. Arsenic and antimony
precipitate on aluminium and would interfere with the titration if
allowed to remain in the lower state of oxidation.]
[Note 3: Bromine is added to oxidize arsenious and antimonious
compounds from the original sample, and to oxidize nitrous acid formed
by the action of nitric acid on copper and copper sulphide.]
[Note 4: This reaction can be carried out in the presence of sulphuric
and hydrochloric acids as well as acetic acid, but in the presence
of these strong acids arsenic and antimonic acids may react with the
hydriodic acid produced with the liberation of free iodine, thereby
reversing the process and introducing an error.]
DETERMINATION OF ANTIMONY IN STIBNITE
Stibnite is native antimony sulphide. Nearly pure samples of this
mineral are easily obtainable and should be used for practice, since
many impurities, notably iron, seriously interfere with the accurate
determination of the antimony by iodometric methods. It is, moreover,
essential that the directions with respect to amounts of reagents
employed and concentration of solutions should be followed closely.
PROCEDURE.--Grind the mineral with great care, and weigh out two
portions of 0.35-0.40 gram into small, dry beakers (100 cc.).
Cover the beakers and pour over the stibnite 5 cc. of concentrated
hydrochloric acid (sp. gr. 1.20) and warm gently on the water bath
(Note 1). When the residue is white, add to each beaker 2 grams of
powdered tartaric acid (Note 2). Warm the solution on the water bath
for ten minutes longer, dilute the solution very cautiously by adding
water in portions of 5 cc., stopping if the solution turns red. It
is possible that no coloration will appear, in which case cautiously
continue the dilution to 125 cc. If a red precipitate or coloration
does appear, warm the solution until it is colorless, and again dilute
cautiously to a total volume of 125 cc. and boil for a minute (Note
3).
If a white precipitate of the oxychloride separates during dilution
(which should not occur if the directions are followed), it is best to
discard the determination and to start anew.
Carefully neutralize most of the acid with ammonium hydroxide solution
(sp. gr. 0.96), but leave it distinctly acid (Note 4). Dissolve 3
grams of sodium bicarbonate in 200 cc. of water in a 500 cc. beaker,
and pour the cold solution of the antimony chloride into this,
avoiding loss by effervescence. Make sure that the solution contains
an excess of the bicarbonate, and then add 1 cc. or 2 cc. of starch
solution and titrate with iodine solution to the appearance of the
blue, avoiding excess (Notes 5 and 6).
From the corrected volume of the iodine solution required to oxidize
the antimony, calculate the percentage of antimony (Sb) in the
stibnite.
[Note 1: Antimony chloride is volatile with steam from its
concentrated solutions; hence these solutions must not be boiled until
they have been diluted.]
[Note 2: Antimony salts, such as the chloride, are readily hydrolyzed,
and compounds such as SbOCl are formed which are often relatively
insoluble; but in the presence of tartaric acid compounds with complex
ions are formed, and these are soluble. An excess of hydrochloric acid
also prevents precipitation of the oxychloride because the H^{+} ions
from the acid lessen the dissociation of the water and thus prevent
any considerable hydrolysis.]
[Note 3: The action of hydrochloric acid upon the sulphide sets free
sulphureted hydrogen, a part of which is held in solution by the acid.
This is usually expelled by the heating upon the water bath; but if it
is not wholly driven out, a point is reached during dilution at which
the antimony sulphide, being no longer held in solution by the acid,
separates. If the dilution is immediately stopped and the solution
warmed, this sulphide is again brought into solution and at the same
time more of the sulphureted hydrogen is expelled. This procedure must
be continued until the sulphureted hydrogen is all removed, since it
reacts with iodine. If no precipitation of the sulphide occurs, it
is an indication that the sulphureted hydrogen was all expelled on
solution of the stibnite.]
[Note 4: Ammonium hydroxide is added to neutralize most of the acid,
thus lessening the amount of sodium bicarbonate to be added. The
ammonia should not neutralize all of the acid.]
[Note 5: The reaction which takes place during titration may be
expressed thus:
Na_{3}SbO_{3} + 2NaHCO_{3} + I_{2} --> Na_{3}SbO_{4} + 2NaI + H_{2}O +
2CO_{2}.]
[Note 6: If the end-point is not permanent, that is, if the blue of
the iodo-starch is discharged after standing a few moments, the cause
may be an insufficient quantity of sodium bicarbonate, leaving the
solution slightly acid, or a very slight precipitation of an antimony
compound which is slowly acted upon by the iodine when the latter is
momentarily present in excess. In either case it is better to discard
the analysis and to repeat the process, using greater care in the
amounts of reagents employed.]
CHLORIMETRY
The processes included under the term !chlorimetry! comprise
those employed to determine chlorine, hypochlorites, bromine, and
hypobromites. The reagent employed is sodium arsenite in the presence
of sodium bicarbonate. The reaction in the case of the hypochlorites
is
NaClO + Na_{3}AsO_{3} --> Na_{3}AsO_{4} + NaCl.
The sodium arsenite may be prepared from pure arsenious oxide,
as described below, and is stable for considerable periods; but
commercial oxide requires resublimation to remove arsenic sulphide,
which may be present in small quantity. To prepare the solution,
dissolve about 5 grams of the powdered oxide, accurately weighed,
in 10 cc. of a concentrated sodium hydroxide solution, dilute the
solution to 300 cc., and make it faintly acid with dilute hydrochloric
acid. Add 30 grams of sodium bicarbonate dissolved in a little water,
and dilute the solution to exactly 1000 cc. in a measuring flask.
Transfer the solution to a dry liter bottle and mix thoroughly.
It is possible to dissolve the arsenious oxide directly in a solution
of sodium bicarbonate, with gentle warming, but solution in sodium
hydroxide takes place much more rapidly, and the excess of the
hydroxide is readily neutralized by hydrochloric acid, with subsequent
addition of the bicarbonate to maintain neutrality during the
titration.
The indicator required for this process is made by dipping strips of
filter paper in a starch solution prepared as described on page 76,
to which 1 gram of potassium iodide has been added. These strips are
allowed to drain and spread upon a watch-glass until dry. When touched
by a drop of the solution the paper turns blue until the hypochlorite
has all been reduced and an excess of the arsenite has been added.
DETERMINATION OF THE AVAILABLE CHLORINE IN BLEACHING POWDER
Bleaching powder consists mainly of a calcium compound which is a
derivative of both hydrochloric and hypochlorous acids. Its formula is
CaClOCl. Its use as a bleaching or disinfecting agent, or as a source
of chlorine, depends upon the amount of hypochlorous acid which it
yields when treated with a stronger acid. It is customary to express
the value of bleaching powder in terms of "available chlorine," by
which is meant the chlorine present as hypochlorite, but not the
chlorine present as chloride.
PROCEDURE.--Weigh out from a stoppered test tube into a porcelain
mortar about 3.5 grams of bleaching powder (Note 1). Triturate the
powder in the mortar with successive portions of water until it is
well ground and wash the contents into a 500 cc. measuring flask
(Note 2). Fill the flask to the mark with water and shake thoroughly.
Measure off 25 cc. of this semi-solution in a measuring flask, or
pipette, observing the precaution that the liquid removed shall
contain approximately its proportion of suspended matter.
Empty the flask or pipette into a beaker and wash it out. Run in the
arsenite solution from a burette until no further reaction takes place
on the starch-iodide paper when touched by a drop of the solution of
bleaching powder. Repeat the titration, using a second 25 cc. portion.
From the volume of solution required to react with the bleaching
powder, calculate the percentage of available chlorine in the latter,
assuming the titration reaction to be that between chlorine and
arsenious oxide:
As_{4}O_{6} + 4Cl_{2} + 4H_{2}O --> 2As_{2}O_{5} + 8HCl
Note that only one twentieth of the original weight of bleaching
powder enters into the reaction.
[Note 1: The powder must be triturated until it is fine, otherwise the
lumps will inclose calcium hypochlorite, which will fail to react with
the arsenious acid. The clear supernatant liquid gives percentages
which are below, and the sediment percentages which are above, the
average. The liquid measured off should, therefore, carry with it its
proper proportion of the sediment, so far as that can be brought about
by shaking the solution just before removal of the aliquot part for
titration.]
[Note 2: Bleaching powder is easily acted upon by the carbonic acid in
the air, which liberates the weak hypochlorous acid. This, of course,
results in a loss of available chlorine. The original material for
analysis should be kept in a closed container and protected form the
air as far as possible. It is difficult to obtain analytical samples
which are accurately representative of a large quantity of the
bleaching powder. The procedure, as outlined, will yield results which
are sufficiently exact for technical purposes.]
III. PRECIPITATION METHODS
DETERMINATION OF SILVER BY THE THIOCYANATE PROCESS
The addition of a solution of potassium or ammonium thiocyanate to one
of silver in nitric acid causes a deposition of silver thiocyanate as
a white, curdy precipitate. If ferric nitrate is also present, the
slightest excess of the thiocyanate over that required to combine with
the silver is indicated by the deep red which is characteristic of the
thiocyanate test for iron.
The reactions involved are:
AgNO_{3} + KSCN --> AgSCN + KNO_{3},
3KSCN + Fe(NO_{3})_{3} --> Fe(SCN)_{3} + 3KNO_{3}.
The ferric thiocyanate differs from the great majority of salts in
that it is but very little dissociated in aqueous solutions, and the
characteristic color appears to be occasioned by the formation of the
un-ionized ferric salt.
The normal solution of potassium thiocyanate should contain an amount
of the salt per liter of solution which would yield sufficient
(CNS)^{-} to combine with one gram of hydrogen to form HCNS, i.e.,
a gram-molecular weight of the salt or 97.17 grams. If the ammonium
thiocyanate is used, the amount is 76.08 grams. To prepare the
solution for this determination, which should be approximately 0.05
N, dissolve about 5 grams of potassium thiocyanate, or 4 grams of
ammonium thiocyanate, in a small amount of water; dilute this solution
to 1000 cc. in a liter bottle and mix as usual.
Prepare 20 cc. of a saturated solution of ferric alum and add 5 cc. of
dilute nitric acid (sp. gr. 1.20). About 5 cc. of this solution should
be used as an indicator.
STANDARDIZATION
PROCEDURE.--Crush a small quantity of silver nitrate crystals in a
mortar (Note 1). Transfer them to a watch-glass and dry them for an
hour at 110°C., protecting them from dust or other organic matter
(Note 2). Weigh out two portions of about 0.5 gram each and dissolve
them in 50 cc. of water. Add 10 cc. of dilute nitric acid which has
been recently boiled to expel the lower oxides of nitrogen, if any,
and then add 5 cc. of the indicator solution. Run in the thiocyanate
solution from a burette, with constant stirring, allowing the
precipitate to settle occasionally to obtain an exact recognition
of the end-point, until a faint red tinge can be detected in the
solution.
From the data obtained, calculate the relation of the thiocyanate
solution to the normal.
[Note 1: The thiocyanate cannot be accurately weighed; its solutions
must, therefore, be standardized against silver nitrate (or pure
silver), either in the form of a standard solution or in small,
weighed portions.]
[Note 2: The crystals of silver nitrate sometimes inclose water which
is expelled on drying. If the nitrate has come into contact with
organic bodies it suffers a reduction and blackens during the heating.
It is plain that a standard solution of silver nitrate (made by
weighing out the crystals) is convenient or necessary if many
titrations of this nature are to be made. In the absence of such a
solution the liability of passing the end-point is lessened by setting
aside a small fraction of the silver solution, to be added near the
close of the titration.]
DETERMINATION OF SILVER IN COIN
PROCEDURE.-- Weigh out two portions of the coin of about 0.5 gram
each. Dissolve them in 15 cc. of dilute nitric acid (sp. gr. 1.2) and
boil until all the nitrous compounds are expelled (Note 1). Cool the
solution, dilute to 50 cc., and add 5 cc. of the indicator solution,
and titrate with the thiocyanate to the appearance of the faint red
coloration (Note 2).
From the corrected volume of the thiocyanate solution required,
calculate the percentage of silver in the coin.
[Note 1: The reaction with silver may be carried out in nitric acid
solutions and in the presence of copper, if the latter does not exceed
70 per cent. Above that percentage it is necessary to add silver in
known quantity to the solution. The liquid must be cold at the time of
titration and entirely free from nitrous compounds, as these sometimes
cause a reddening of the indicator solution. All utensils, distilled
water, the nitric acid and the beakers must be free from chlorides,
as the presence of these will cause precipitation of silver chloride,
thereby introducing an error.]
[Note 2: The solution containing the silver precipitate, as well as
those from the standardization, should be placed in the receptacle for
"silver residues" as a matter of economy.]
PART III
GRAVIMETRIC ANALYSIS
GENERAL DIRECTIONS
Gravimetric analyses involve the following principal steps: first, the
weighing of the sample; second, the solution of the sample; third, the
separation of some substance from solution containing, or bearing a
definite relation to, the constituent to be measured, under conditions
which render this separation as complete as possible; and finally,
the segregation of that substance, commonly by filtration, and the
determination of its weight, or that of some stable product formed
from it on ignition. For example, the gravimetric determination of
aluminium is accomplished by solution of the sample, by precipitation
in the form of hydroxide, collection of the hydroxide upon a filter,
complete removal by washing of all foreign soluble matter, and the
burning of the filter and ignition of the precipitate to aluminium
oxide, in which condition it is weighed.
Among the operations which are common to nearly all gravimetric
analyses are precipitation, washing of precipitates, ignition of
precipitates, and the use of desiccators. In order to avoid burdensome
repetitions in the descriptions of the various gravimetric procedures
which follow, certain general instructions are introduced at this
point. These instructions must, therefore, be considered to be as much
a part of all subsequent procedures as the description of apparatus,
reagents, or manipulations.
The analytical balance, the fundamentally important instrument in
gravimetric analysis, has already been described on pages 11 to 15.
PRECIPITATION
For successful quantitative precipitations those substances are
selected which are least soluble under conditions which can be easily
established, and which separate from solution in such a state that
they can be filtered readily and washed free from admixed material.
In general, the substances selected are the same as those already
familiar to the student of Qualitative Analysis.
When possible, substances are selected which separate in crystalline
form, since such substances are less likely to clog the pores of
filter paper and can be most quickly washed. In order to increase the
size of the crystals, which further promotes filtration and washing,
it is often desirable to allow a precipitate to remain for some time
in contact with the solution from which it has separated. The solution
is often kept warm during this period of "digestion." The small
crystals gradually disappear and the larger crystals increase in size,
probably as the result of the force known as surface tension, which
tends to reduce the surface of a given mass of material to a minimum,
combined with a very slightly greater solubility of small crystals as
compared with the larger ones.
Amorphous substances, such as ferric hydroxide, aluminium hydroxide,
or silicic acid, separate in a gelatinous form and are relatively
difficult to filter and wash. Substances of this class also exhibit
a tendency to form, with pure water, what are known as colloidal
solutions. To prevent this as far as possible, they are washed with
solutions of volatile salts, as will be described in some of the
following procedures.
In all precipitations the reagent should be added slowly, with
constant stirring, and should be hot when circumstances permit.
The slow addition is less likely to occasion contamination of the
precipitate by the inclosure of other substances which may be in the
solution, or of the reagent itself.
FUNNELS AND FILTERS
Filtration in analytical processes is most commonly effected through
paper filters. In special cases these may be advantageously replaced
by an asbestos filter in a perforated porcelain or platinum crucible,
commonly known, from its originator, as a "Gooch filter." The
operation and use of a filter of this type is described on page 103.
Porous crucibles of a material known as alundum may also be employed
to advantage in special cases.
The glass funnels selected for use with paper filters should have an
angle as near 60° as possible, and a narrow stem about six inches in
length. The filters employed should be washed filters, i.e., those
which have been treated with hydrochloric and hydrofluoric acids, and
which on incineration leave a very small and definitely known weight
of ash, generally about .00003 gram. Such filters are readily
obtainable on the market.
The filter should be carefully folded to fit the funnel according to
either of the two well-established methods described in the Appendix.
It should always be placed so that the upper edge of the paper
is about one fourth inch below the top of the funnel. Under no
circumstances should the filter extend above the edge of the funnel,
as it is then utterly impossible to effect complete washing.
To test the efficiency of the filter, fill it with distilled water.
This water should soon fill the stem completely, forming a continuous
column of liquid which, by its hydrostatic pressure, produces a gentle
suction, thus materially promoting the rapidity of filtration. Unless
the filter allows free passage of water under these conditions, it is
likely to give much trouble when a precipitate is placed upon it.
The use of a suction pump to promote filtration is rarely altogether
advantageous in quantitative analysis, if paper filters are employed.
The tendency of the filter to break, unless the point of the filter
paper is supported by a perforated porcelain cone or a small "hardened
filter" of parchment, and the tendency of the precipitates to pass
through the pores of the filter, more than compensate for the possible
gain in time. On the other hand, filtration by suction may be useful
in the case of precipitates which do not require ignition before
weighing, or in the case of precipitates which are to be discarded
without weighing. This is best accomplished with the aid of the
special apparatus called a Gooch filter referred to above.
FILTRATION AND WASHING OF PRECIPITATES
Solutions should be filtered while hot, as far as possible, since
the passage of a liquid through the pores of a filter is retarded by
friction, and this, for water at 100°C., is less than one sixth of the
resistance at 0°C.
When the filtrate is received in a beaker, the stem of the funnel
should touch the side of the receiving vessel to avoid loss by
spattering. Neglect of this precaution is a frequent source of error.
The vessels which contain the initial filtrate should !always! be
replaced by clean ones, properly labeled, before the washing of a
precipitate begins. In many instances a finely divided precipitate
which shows no tendency to pass through the filter at first, while the
solution is relatively dense, appears at once in the washings. Under
such conditions the advantages accruing from the removal of the first
filtrate are obvious, both as regards the diminished volume requiring
refiltration, and also the smaller number of washings subsequently
required.
Much time may often be saved by washing precipitates by decantation,
i.e., by pouring over them, while still in the original vessel,
considerable volumes of wash-water and allowing them to settle. The
supernatant, clear wash-water is then decanted through the filter,
so far as practicable without disturbing the precipitate, and a new
portion of wash-water is added. This procedure can be employed to
special advantage with gelatinous precipitates, which fill up the
pores of the filter paper. As the medium from which the precipitate
is to settle becomes less dense it subsides less readily, and it
ultimately becomes necessary to transfer it to the filter and complete
the washing there.
A precipitate should never completely fill a filter. The wash-water
should be applied at the top of the filter, above the precipitate.
It may be shown mathematically that the washing is most !rapidly!
accomplished by filling the filter well to the top with wash-water
each time, and allowing it to drain completely after each addition;
but that when a precipitate is to be washed with the !least possible
volume! of liquid the latter should be applied in repeated !small!
quantities.
Gelatinous precipitates should not be allowed to dry before complete
removal of foreign matter is effected. They are likely to shrink and
crack, and subsequent additions of wash-water pass through these
channels only.
All filtrates and wash-waters without exception must be properly
tested. !This lies at the foundation of accurate work!, and the
student should clearly understand that it is only by the invariable
application of this rule that assurance of ultimate reliability can
be secured. Every original filtrate must be tested to prove complete
precipitation of the compound to be separated, and the wash-waters
must also be tested to assure complete removal of foreign material. In
testing the latter, the amount first taken should be but a few
drops if the filtrate contains material which is to be subsequently
determined. When, however, the washing of the filter and precipitate
is nearly completed the amount should be increased, and for the final
test not less than 3 cc. should be used.
It is impossible to trust to one's judgment with regard to the washing
of precipitates; the washings from !each precipitate! of a series
simultaneously treated must be tested, since the rate of washing will
often differ materially under apparently similar conditions, !No
exception can ever be made to this rule!.
The habit of placing a clean common filter paper under the receiving
beaker during filtration is one to be commended. On this paper a
record of the number of washings can very well be made as the portions
of wash-water are added.
It is an excellent practice, when possible, to retain filtrates and
precipitates until the completion of an analysis, in order that, in
case of question, they may be examined to discover sources of error.
For the complete removal of precipitates from containing vessels, it
is often necessary to rub the sides of these vessels to loosen the
adhering particles. This can best be done by slipping over the end of
a stirring rod a soft rubber device sometimes called a "policeman."
DESICCATORS
Desiccators should be filled with fused, anhydrous calcium chloride,
over which is placed a clay triangle, or an iron triangle covered with
silica tubes, to support the crucible or other utensils. The cover of
the desiccator should be made air-tight by the use of a thin coating
of vaseline.
Pumice moistened with concentrated sulphuric acid may be used in place
of the calcium chloride, and is essential in special cases; but for
most purposes the calcium chloride, if renewed occasionally and not
allowed to cake together, is practically efficient and does not slop
about when the desiccator is moved.
Desiccators should never remain uncovered for any length of time. The
dehydrating agents rapidly lose their efficiency on exposure to the
air.
CRUCIBLES
It is often necessary in quantitative analysis to employ fluxes to
bring into solution substances which are not dissolved by acids. The
fluxes in most common use are sodium carbonate and sodium or potassium
acid sulphate. In gravimetric analysis it is usually necessary to
ignite the separated substance after filtration and washing, in order
to remove moisture, or to convert it through physical or chemical
changes into some definite and stable form for weighing. Crucibles
to be used in fusion processes must be made of materials which will
withstand the action of the fluxes employed, and crucibles to be used
for ignitions must be made of material which will not undergo any
permanent change during the ignition, since the initial weight of the
crucible must be deducted from the final weight of the crucible and
product to obtain the weight of the ignited substance. The three
materials which satisfy these conditions, in general, are platinum,
porcelain, and silica.
Platinum crucibles have the advantage that they can be employed at
high temperatures, but, on the other hand, these crucibles can never
be used when there is a possibility of the reduction to the metallic
state of metals like lead, copper, silver, or gold, which would alloy
with and ruin the crucible. When platinum crucibles are used with
compounds of arsenic or phosphorus, special precautions are necessary
to prevent damage. This statement applies to both fusions and
ignitions.
Fusions with sodium carbonate can be made only in platinum, since
porcelain or silica crucibles are attacked by this reagent. Acid
sulphate fusions, which require comparatively low temperatures, can
sometimes be made in platinum, although platinum is slightly attacked
by the flux. Porcelain or silica crucibles may be used with acid
fluxes.
Silica crucibles are less likely to crack on heating than porcelain
crucibles on account of their smaller coefficient of expansion.
Ignition of substances not requiring too high a temperature may be
made in porcelain or silica crucibles.
Iron, nickel or silver crucibles are used in special cases.
In general, platinum crucibles should be used whenever such use is
practicable, and this is the custom in private, research or commercial
laboratories. Platinum has, however, become so valuable that it is
liable to theft unless constantly under the protection of the user. As
constant protection is often difficult in instructional laboratories,
it is advisable, in order to avoid serious monetary losses, to use
porcelain or silica crucibles whenever these will give satisfactory
service. When platinum utensils are used the danger of theft should
always be kept in mind.
PREPARATION OF CRUCIBLES FOR USE
All crucibles, of whatever material, must always be cleaned, ignited
and allowed to cool in a desiccator before weighing, since all bodies
exposed to the air condense on their surfaces a layer of moisture
which increases their weight. The amount and weight of this moisture
varies with the humidity of the atmosphere, and the latter may change
from hour to hour. The air in the desiccator (see above) is kept at
a constant and low humidity by the drying agent which it contains.
Bodies which remain in a desiccator for a sufficient time (usually
20-30 minutes) retain, therefore, on their surfaces a constant weight
of moisture which is the same day after day, thus insuring constant
conditions.
Hot objects, such as ignited crucibles, should be allowed to cool in
the air until, when held near the skin, but little heat is noticeable.
If this precaution is not taken, the air within the desiccator is
strongly heated and expands before the desiccator is covered. As the
temperature falls, the air contracts, causing a reduction of air
pressure within the covered vessel. When the cover is removed (which
is often rendered difficult) the inrush of air from the outside may
sweep light particles out of a crucible, thus ruining an entire
analysis.
Constant heating of platinum causes a slight crystallization of the
surface which, if not removed, penetrates into the crucible. Gentle
polishing of the surface destroys the crystalline structure and
prevents further damage. If sea sand is used for this purpose, great
care is necessary to keep it from the desk, since beakers are easily
scratched by it, and subsequently crack on heating.
Platinum crucibles stained in use may often be cleaned by the fusion
in them of potassium or sodium acid sulphate, or by heating with
ammonium chloride. If the former is used, care should be taken not
to heat so strongly as to expel all of the sulphuric acid, since the
normal sulphates sometimes expand so rapidly on cooling as to split
the crucible. The fused material should be poured out, while hot, on
to a !dry! tile or iron surface.
IGNITION OF PRECIPITATES
Most precipitates may, if proper precautions are taken, be ignited
without previous drying. If, however, such precipitates can be dried
without loss of time to the analyst (as, for example, over night), it
is well to submit them to this process. It should, nevertheless, be
remembered that a partially dried precipitate often requires more care
during ignition than a thoroughly moist one.
The details of the ignition of precipitates vary so much with the
character of the precipitate, its moisture content, and temperature to
which it is to be heated, that these details will be given under the
various procedures which follow.
DETERMINATION OF CHLORINE IN SODIUM CHLORIDE
!Method A. With the Use of a Gooch Filter!
PROCEDURE.--Carefully clean a weighing-tube containing the sodium
chloride, handling it as little as possible with the moist fingers,
and weigh it accurately to 0.0001 gram, recording the weight at once
in the notebook (see Appendix). Hold the tube over the top of a beaker
(200-300 cc.), and cautiously remove the stopper, noting carefully
that no particles fall from it, or from the tube, elsewhere than into
the beaker. Pour out a small portion of the chloride, replace the
stopper, and determine by approximate weighing how much has been
removed. Continue this procedure until 0.25-0.30 gram has been taken
from the tube, then weigh accurately and record the weight beneath the
first in the notebook. The difference of the two weights represents
the weight of the chloride taken for analysis. Again weigh a second
portion of 0.25-0.30 gram into a second beaker of the same size as the
first. The beakers should be plainly marked to correspond with the
entries in the notebook. Dissolve each portion of the chloride in 150
cc. of distilled water and add about ten drops of dilute nitric acid
(sp. gr. 1.20) (Note 2). Calculate the volume of silver nitrate
solution required to effect complete precipitation in each case,
and add slowly about 5 cc. in excess of that amount, with constant
stirring. Heat the solutions cautiously to boiling, stirring
occasionally, and continue the heating and stirring until the
precipitates settle promptly, leaving a nearly clear supernatant
liquid (Note 3). This heating should not take place in direct sunlight
(Note 4). The beaker should be covered with a watch-glass, and both
boiling and stirring so regulated as to preclude any possibility of
loss of material. Add to the clear liquid one or two drops of silver
nitrate solution, to make sure that an excess of the reagent is
present. If a precipitate, or cloudiness, appears as the drops fall
into the solution, heat again, and stir until the whole precipitate
has coagulated. The solution is then ready for filtration.
Prepare a Gooch filter as follows: Fold over the top of a Gooch funnel
(Fig. 2) a piece of rubber-band tubing, such as is known as "bill-tie"
tubing, and fit into the mouth of the funnel a perforated porcelain
crucible (Gooch crucible), making sure that when the crucible is
gently forced into the mouth of the funnel an airtight joint results.
(A small 1 or 1-1/4-inch glass funnel may be used, in which case the
rubber tubing is stretched over the top of the funnel and then drawn
up over the side of the crucible until an air-tight joint is secured.)
[ILLUSTRATION: FIG. 2]
Fit the funnel into the stopper of a filter bottle, and connect the
filter bottle with the suction pump. Suspend some finely divided
asbestos, which has been washed with acid, in 20 to 30 cc. of water
(Note 1); allow this to settle, pour off the very fine particles, and
then pour some of the mixture cautiously into the crucible until an
even felt of asbestos, not over 1/32 inch in thickness, is formed. A
gentle suction must be applied while preparing this felt. Wash the
felt thoroughly by passing through it distilled water until all fine
or loose particles are removed, increasing the suction at the last
until no more water can be drawn out of it; place on top of the felt
the small, perforated porcelain disc and hold it in place by pouring a
very thin layer of asbestos over it, washing the whole carefully;
then place the crucible in a small beaker, and place both in a drying
closet at 100-110°C. for thirty to forty minutes. Cool the crucible
in a desiccator, and weigh. Heat again for twenty to thirty minutes,
cool, and again weigh, repeating this until the weight is constant
within 0.0003 gram. The filter is then ready for use.
Place the crucible in the funnel, and apply a gentle suction, !after
which! the solution to be filtered may be poured in without disturbing
the asbestos felt. When pouring liquid onto a Gooch filter hold the
stirring-rod at first well down in the crucible, so that the liquid
does not fall with any force upon the asbestos, and afterward keep the
crucible will filled with the solution.
Pour the liquid above the silver chloride slowly onto the filter,
leaving the precipitate in the beaker as far as possible. Wash the
precipitate twice by decantation with warm water; then transfer it
to the filter with the aid of a stirring-rod with a rubber tip and a
stream from the wash-bottle.
Examine the first portions of the filtrate which pass through the
filter with great care for asbestos fibers, which are most likely to
be lost at this point. Refilter the liquid if any fibers are visible.
Finally, wash the precipitate thoroughly with warm water until free
from soluble silver salts. To test the washings, disconnect the
suction at the flask and remove the funnel or filter tube from the
suction flask. Hold the end of the tube over the mouth of a small test
tube and add from a wash-bottle 2-3 cc. of water. Allow the water to
drip through into the test tube and add a drop of dilute hydrochloric
acid. No precipitate or cloud should form in the wash-water (Note 16).
Dry the filter and contents at 100-110°C. until the weight is constant
within 0.0003 gram, as described for the preparation of the filter.
Deduct the weight of the dry crucible from the final weight, and from
the weight of silver chloride thus obtained calculate the percentage
of chlorine in the sample of sodium c
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