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MANUAL OF CHEMICAL ANALYSIS

BY THE SAME AUTHOR

CHEMICAL LECTURE EXPERIMENTS.

With 230 Diagrams. Crown 8vo.

A TEXT-BOOK OF INORGANIC CHEMIS-
TRY. With 155 Illustrations. Crown 8vo.

QUESTIONS ON THE ABOVE. By G. D.

TlMMONS.

ELEMENTARY PRACTICAL CHEMISTRY.
With 108 Illustrations and 254 Experiments.
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SMALLER CHEMICAL ANALYSIS. Crown
8vo,

LONGMANS, GREEN, AND CO.

LONDON, NEW YORK, BOMBAY, CALCUTTA, AND MADRAS

c

A MANUAL

CHEMICAL ANALYSIS

QUALITATIVE AND QUANTITATIVE

G. S NEWTH, F.I C., F.C.S.

FORMERLY DEMONSTRATOR IN THK ROYAL COLLEGE OF SCIENCE, LONDON
AND ASSISTANT-EXAMINER IN CHEMISTRY, BOARD OF EDUCATION

THIRTEENTH IMPRESSION

*

V

LONGMANS, GREEN AND CO.

39 PATERNOSTER ROW, LONDON

FOURTH AVENUE & 30rH STREET, NEW YORK
BOMBAY, CALCUTTA, AND MADRAS

1921
All rights reserved

PREFACE

A MAN once brought his son to the Royal School of Mines now
the Royal College of Science with the request that he might be
taught to " do copper.'" He did not want his boy to " -waste his
time learning about oxygen and hydrogen, and all that" but he
wished him simply to learn to " do copper."

Although seldom expressed with such refreshing candour, the
desire to do analysis without learning more than the minimum
amount of chemistry is still very prevalent ; and, unfortunately,
chemical analysis is a subject which may be, and frequently is,
taught and practised in such a manner as to degrade it to the level
of a purely mechanical and often quite unintelligible series of rule-
of-thumb operations.

I hope that the student whose aspirations rise no higher than to
learn to do analysis in this fashion, will not find this book suitable
for him. I have done my best to make it as little of a cram-book
as possible, but have endeavoured to teach analytical, chemistry as
well as analysis that is, the theoretical as well as the practical side
of the subject.

With this object in view, I have carefully avoided the use of those
symbolic abbreviated expressions (slang formula, they might be
termed), such as H 2 O (oxalic acid), H 2 T(tartaric acid), HA (acetic
acid), etc., which are becoming so common, and which, so far as
the student is concerned, foster those very evils of cramming which
we as teachers are striving to combat. If such symbols as these
are permitted and recognised, why not H 2 S, HN, H 3 P, for sulphuric,
nitric, and phosphoric acids respectively ? And then, perhaps,
some such hieroglyphic as H 2 S, HN, for sulphurous and nitrous
acids.

For the busy chemist to make use of such shorthand signs in
the privacy of his own note-books is one thing, but to print them in

vi Preface.

a text-book intended for the learner, and thereby put them on the
same footing as such chemical formulae as H 2 SO 4 , H 2 C 2 O 4 , etc., is
quite another thing, and is open to the taunt of chemistry made easy.

The importance to the student of making careful notes of his
analytical work while in progress, cannot well be overrated, as,
perhaps more than anything else, this is calculated to develop in
him those habits of exact observation which are essential qualities
in a scientific man. While insisting on this point more than once,
I have purposely refrained from giving anything of the nature of a
specimen of notes, because notes which are made according to a
stereotyped pattern are practically of no value. Unfortunately, it
is the custom in most laboratories perhaps a necessary custom,
but none the less unfortunate to require from the student some
written account of his analysis by way of proof that he has con-
scientiously done the work. The result of this is that the notes
which he makes are taken primarily with a view to furnish this
required evidence. Points which ought to be observed and noted,
points which he does not at the time understand, are passed over,
while others which are familiar, and which are at once recognised
as useful evidence, are written down. Indeed, as every teacher
has experienced, the midence is often more or less fabricated after
the completion of the analysis.

The consequence of this method is the stereotyped product so
painfully familiar to all who are engaged in teaching or examining,
namely, a certain number of regulation tests, arranged with a
semblance of method in the everlasting three columns, headed
Experiment, Observation, Inference, followed by a more or less
slovenly copy of the time-honoured tables.

The valuelessness of such a written-out record must be patent
to all teachers, and yet we still continue to accept it in lieu of a few
real notes, less quickly and easily examined no doubt, because less
stereotyped, but of infinitely more value to the student himself.

In order to gain as much space as possible for purely analytical
matter, and still to be able to include within the limits of a con-
venient volume, both qualitative and quantitative analysis, I have
carefully excluded all merely descriptive details which have no
direct bearing upon analysis. For example, none of the properties,
either chemical or physical, of such metals as those of the alkalies
or alkaline earths, are made use of in ordinary analytical processes
for the detection or recognition of these elements, therefore any

Preface. vii

descriptive account of their properties in the elemental state is
entirely out of place in a book which is intended only to teach
analytical as distinguished from general chemistry.

In the second portion of the book, devoted to quantitative
analysis, I have confined myself to a comparatively small number
of well-tried typical methods and processes, preferring to describe
and explain in tolerably full detail, a few quantitative determinations
in each of the various sections, such as shall furnish a thoroughly
sound course of practical study, rather than to attempt to cover
necessarily in more sketchy outlines a wider range of subjects.
Whether the choice of processes and examples I have made is as
good or as representative as it might be, will no doubt be a matter
for difference of opinion, and as the book has not been written to
meet any particular syllabus, this is probably a point upon which
hardly any two teachers will exactly agree.

The illustrations throughout have been made from original
photographs of the actual apparatus employed in the various
analytical operations described.

G. S. N.

ROYAL COLLEGE OF SCIENCE, LONDON.
June, 1898.

TABLE OF CONTENTS

BOOK I.
QUALITATIVE ANALYSIS.

PAGE

CHAPTER I. Preliminary exercises Filtration ; Solution ; Evapo-
ration ; Fusion ; Precipitation ; Ignition ; Neutral-
isation I

,, II. Analytical classification 13

,, III. Reactions of the metals of Group V 19

APPENDIX TO CHAPTER III. The rare metals of Group V. . 26

CHAPTER IV. Reactions of the metals of Group IV 30

V. IIlA 36

VI. ,, HlB. .... 5 I

,, VII. The phosphates 63

APPENDIX TO CHAPTER VII. The rare metals of Group III. 70

CHAPTER VIII. Reactions of the metals of Group II., Division I 75

,, IX. ,, ,, ,, II., Division 2 89

APPENDIX TO CHAPTER IX. The rare metals of Group II. . 108

CHAPTER X. Reactions of the metals of Group 1 112

APPENDIX TO CHAPTER X. The rare metals of Group I. . . 118

CHAPTER XI. The non-metals and their acids The halogens . . 120

XII. Sulphur 137

,, XIII. Nitrogen and phosphorus ......... 144

,, XIV. Carbon, silicon, boron 153

,, XV. Systematic detection of the acids 171

,, XVI. Preliminary tests and operations in a systematic

analysis j general analytical tables .... 178

XVII. The results of a qualitative analysis l#6

BOOK IL
QUANTITATIVE ANALYSIS.

PART I. GRAVIMETRIC METHODS.

SECTION I. Preliminary manipulations 189

(I) Weighing.

x Table of Contents.

PAGE

(2) Drying and weighing a filter.

(3) Estimation of water of crystallisation.

(4) Determination of the ash of a filter.

(5) Preparation of pure salts.

SECTION II. Typical gravimetric determinations of metals . 221
Al, Cr, Fe, Ca, Mg, Cu, Ag, Pb, Zn, Mn, Ni, Co, K,

NH 4 , Sn, As, Sb, Cd, Hg.

SECTION III. Typical gravimetric determinations of acids . . 258
H 2 SO 4) HC1, HBr, CO 2 , HNO 3 , H,PO 4> SiO 2 .

SECTION IV. Exercises in gravimetric analysis 265

Silver coin, solder, German silver, bronze, dolomite, zinc

blende, silicate.
SECTION V. Electrolytic methods of analysis 292

PART II. VOLUMETRIC METHODS.

SECTION I. Preliminary manipulations 300

(1) Instruments for measuring liquids Calibration of
vessels.

(2) Standard solutions.

(3) Methods for ascertaining the completion of volumetric
reactions Indicators.

i SECTION II. Volumetric methods based upon saturation

Alkalimetry, acidimetry 316

Typical analyses by means of standard acids and alkalies.
SECTION III. Methods based upon oxidation and reduction . 331
Potassium permanganate Titration with iron, ferrous salt,

oxalic acid.
Typical analyses by means of permanganate.

(1) Iron ores.

(2) Available oxygen in manganese ores.

(3) Estimation of tin.

(4) Estimation of calcium.

Potassium dichromate Titration -w\\h. ferrous salts.
Analyses by means of dichromate.

(1) Estimation of iron in iron ores.

(2) Estimation of chromium in chrome iron ore.
Iodine solution Titration with thiosulphate and arsenious

oxide.
Estimations by means of iodine.

(1) Antimony.

(2) Arsenic.
(2) Tin.

Estimations by means of iodine and thiosulphate

Table of Contents. xi

PAGE

(1) Sulphur dioxide in a sulphite.

(2) Available chlorine in bleaching powder.

(3) Estimation of manganese dioxide.

SECTION IV. Methods based upon precipitation .... 360
(I.) Precipitation with silver nitrate Chlorine, cyanogen.
Indirect estimation of nitric add, calcium, carbon

dioxide, and cadmium.

(II.) Precipitation with sodium chloride Silver.
(III.) Precipitation with ammonium thiocyanate.

(1) Silver.

(2) Indirect estimation of chlorine.

(IV.) Precipitation by means of uranium acetate.

Estimation of phosphoric acid in bone ash.
(V.) Precipitation with sodium sulphide.

Estimation of zinc in zinc ores.

(VI.) Clark's method for estimating hardness of water.
APPENDIX TO SECTION IV. Estimation of copper by potassium

cyanide 374

SECTION V. Gas analysis 377

(1) Estimation by absorption and subsequent titration.

(a) Carbon dioxide in air (Pettenkofer's method).
(l>) Sulphur dioxide in furnace gas.

(2) Estimation by absorption, and measurement of the
residual gas.

Measurement of gases.

Calibration of gas-burettes.

Correction of gaseous volumes.

Collection of samples for analysis.

Absorption in Hempel's gas-pipettes.

Analysis of mixture containing CO 2 , CO, O, N.

(3) Estimation by combustion, and determination of the
products.

Hydrogen by combustion with palladium.

Hydrogen by explosion.

Marsh-gas by explosion.

Analysis of mixtures containing H, CH 4 , N.

The nitrometer.

PART III.

SECTION I. Estimation of carbon, hydrogen, nitrogen,

chlorine, sulphur, and phosphorus in organic compounds . . 414
SECTION II. Miscellaneous physico-chemical determinations . 434
Specific gravity of solids and liquids.
Boiling-point.

xii Table of Contents.

PAGE

Melting-point.

Vapour density.

Molecular weight by freezing-point.

Molecular weight by boiling-point.

APPENDIX

Table of elements and atomic weights 453

Reagents System of standard strengths 454

Table of hardness 459

Expansion of water between o and 25 460

Tension of aqueous vapour from 5 to 25 461

Table of factors for the correction of gaseous volumes . . 462

Table of logarithms , 466

CHEMICAL ANALYSIS

BOOK I.

QUALITATIVE ANALYSIS.

CHAPTER I.
PRELIMINARY EXERCISES.

THE first step that the student must take in approaching the
subject of analytical chemistry, is that of making himself practically
familiar with certain simple operations or manipulations which he
will constantly be required to carry out in the course of his work,
and upon the dexterous and cleanly performance of which much
of his success as an analyst will depend. If he has not had previous
experience in practical chemistry, therefore, he should carefully go
through the following exercises.

i. Filtration. The method by which a liquid is separated
from^any solid substance with which it is mechanically mixed, is
most usually that of filtering the mixture through porous paper,
known as filter-paper.

EXERCISE i. Fold a circular filter-paper into half, and then at
right angles into half again. Open this into a cone having one
thickness of paper on one side and three on the other. This cone
is then placed in a glass funnel of such a size that the glass will
project slightly above the paper. The paper is then moistened
with distilled water, which should not be poured out of the funnel
again, but allowed to run through. After being cautiously pressed
into the glass funnel, the paper should fit close to the glass all
round, leaving no air-spaces. If this is not the case, either another
funnel of the right angle (60 degrees) should be selected, or another
filter-paper folded so that the cone shall be of the same angle as
the funnel. This can be done by making the second fold of the

2 Qualitative Analysis.

paper not quite at right angles to the first. In this way a cone
will be formed having either a more acute or more obtuse angle
than 60 degrees, as the paper is opened out one side or the other.
The funnel is supported by a metal or wooden stand.

Now place some diluted hydrochloric acid in a small beaker,
and stir into it, by means of a glass rod, a quantity of finely
powdered charcoal. When thoroughly mixed, pour upon the filter.
When slowly pouring from a wide vessel like a beaker, there is
risk of some of the liquid being spilt by running down the outside
of the vessel, as. shown in Fig. i. If it be poured quickly, it is

FIG. i. FIG 2.

likely to splash over the funnel. To prevent both of these ecci-
dents, the liquid should be poured down against a glass rod held
lightly against the edge of the beaker, and in such a position that
the liquid does not strike at once against the apex of the paper
cone (Fig. 2).

The filtrate (i.e. the liquid which passes through the filter) may
be received in another beaker, which should be placed close against
the stem of the funnel, so that the liquid shall run down against
the glass. In this way splashing is prevented. The filtrate should
be perfectly clear, the whole of the solid being retained on the
filter. When all the liquid has passed through, the charcoal and
the filter-paper are both still soaked with the hydrochloric acid.
In order to remove this, and so to make the separation of the solid
from the liquid complete, the filter and its contents must be washed

Preliminary Exercises.

with distilled water.* This is done by directing a fine stream of
water from a wash-bottle into the funnel, working downwards from
the upper edges of the paper, and so washing the charcoal down
into the apex of the filter (Fig. 3). Each washing must be allowed
to drain right through before more water is used. This must be
continued until the filtrate is entirely free from acid, which may
be ascertained by allowing one or two drops of it to fall upon a
piece of blue litmus paper.

In practice, the size of the filter should bear a rational relation
to the quantity of solid matter to be separated from a liquid. This
is more especially impor-
tant when the material re-
tained upon the filter has
to be washed. If the
amount of solid is small,
the filter used should be
proportionately small, and
the washing operation will
be more quickly and effec-
tually accomplished than if
an unduly large filter is
employed.

2. Solution. This
term is applied both to the
act of dissolving and to the
product obtained by dis- y
solving.

EXERCISE 2. Place a
little powdered potassium
carbonate in a test-tube,
and add a small quantity
of water. In a few moments
the salt will have entirely
dissolved. The salt has
undergone solution in water. The product is a solution of potas-
sium carbonate. The water is called the solvent. The process of
solution is accelerated by heating the liquid, and it takes place more
quickly the more finely the solid is powdered.

Put a similar quantity of potassium carbonate into another
test-tube, and add a little dilute nitric acid. The salt again under-
goes solution, the acid here being the solvent. But in this case
there is a radical difference. First, a visible difference, in that the

* In the following exercises, and in all analytical operations, distilled
water must always be employed; and when beakers, test-tubes, etc., are
washed up after use, they must be finally rinsed with distilled water.

FIG.

Qualitative A nalysis.

act of solution is accompanied by an effervescence, or rapid evolu-
tion of gas ; and second, an invisible difference ; for the resulting
liquid is not a solution of potassium carbonate, but of potassium
nitrate. In the first case, the process is not accompanied by any
chemical change ; the operation is therefore called simple solution :
the original substance is present in the liquid, and can be obtained
in its former state by evaporating the water. In the second case,
the process is distinguished as chemical solution, because chemical
action took place between the substance dissolved and the solvent,
and the original substance cannot be got back by evaporating the
solvent.

3. Evaporation. The process of changing from the liquid
to the gaseous or vaporous state is known as evaporation. This
operation is greatly accelerated by the application of heat. When
it takes place without the aid of external heat, the process is spoken
of as spontaneous evaporation.

EXERCISE 3. Pour the two solutions obtained in Exercise 2
into separate porcelain evaporating-dishes, and heat them gently
by means of a Bunsen with a " rose " burner (as shown in Fig. 4).
Continue the operation until all the liquid
has evaporated away and a dry residue
is left. This is called evaporating to dry-
ness. As the condition of dryness is ap-
proached, the flame must be turned down
more and more, to prevent the substance
from " sputtering." Try to conduct the
operation so that as little as possible of
the residue is lost in this way.

The two residues may now be exa-
mined by one simple test, which will
prove that the one from the watery solu-
tion is the same as it was before being
dissolved, and that the other is quite dif-
ferent. Add to each a few drops of dilute
nitric acid : the first dissolves with effer-
vescence, as did the original potassium
carbonate ; the other is unacted on by the acid.

Sometimes it is necessary to carry on the operation of evapora-
tion more carefully than can be done by heating the dish in the
manner described. In this case the process is conducted upon
a steam-bath. Water is boiled in a metal vessel (resembling a
saucepan), and the evaporating-dish, supported by a metal ring
which forms the cover, is heated by the steam. The following
exercise is a case in point.

EXERCISE 4. Dissolve some crystals of ammonium nitrate in
a little water ; place half the solution in a dish, and evaporate it

FIG. 4.

Preliminary Exercises. 5

over a rose burner. Evaporate the other portion in a dish upon a
steam-bath. Note the difference in the results in the two cases.

4. Fusion is the term used to denote the process of changing
a substance from the solid to the liquid state by the action of heat.
Thus, when lead is heated it enters into a state of fusion, or, shortly,
it fuses or melts. Fusion must not be confounded with solution.
Chemical action often takes place when one of the reacting sub-
stances is in a condition of fusion, which is incapable of taking
place when they are only in solution. For example

EXERCISE 5. Dissolve a small piece of potassium hydroxide
(caustic potash) in water, and add to the colourless solution a minute
quantity of powdered manganese dioxide. No chemical action
takes place.

Place a similar piece of potassium hydroxide in a dry test-tube,
and heat it : the solid fuses to a colourless liquid. Drop into the
fused mass a few particles of the manganese dioxide. Chemical
action at ouce takes place, resulting in the formation of the deep
green-coloured compound, potassium manganate. (This reaction
is used as a test for manganese compounds.)

EXERCISE 6. Place a small quantity of powdered barium
sulphate in a test-tube, add water, and boil for a minute or two.
If the amount of barium sulphate is quite small, it will be easy to
see that practically none of it dissolves. Allow it to settle, and pour
a few drops of the liquid upon a watch-glass, and set it to evapo-
rate to dryness on a steam-bath.

Treat another similar quantity of the barium sulphate with
dilute hydrochloric acid, and evaporate a few drops of the liquid
in the same way. The result of these two operations will prove
that barium sulphate is insoluble in either "water or hydrochloric acid.

Next dissolve a little sodium carbonate in water, and add to the
clear solution a few particles of barium sulphate ; boil the liquid,
and observe that no change takes place.

Now carefully mix a small quantity of barium sulphate with
about five times as much sodium carbonate ; place the powder in
a platinum crucible, supported on a pipe-clay triangle in the manner
shown in Fig. 5, and heat strongly with a blowpipe. When the
mass has been in complete fusion for a few minutes, allow the
crucible to cool. Then place it on its side in a small beaker with
a little water, and warm gently. The mass in the crucible will soon
become disintegrated, some of it dissolving, while a part remains
undissolved. Filter the liquid as in Exercise i, washing the residue
upon the funnel until the filtrate no longer restores the blue colour
to reddened litmus paper. Now pour a few drops of dilute hydro-
chloric acid upon the residue on the filter, receiving the liquid
which passes through in a fresh beaker or test-tube. Observe that
effervescence at once takes place. But this residue cannot be
sodium carbonate, because that salt, being soluble in water, has

6

Qualitative Analysis.

been all removed ; neither can it be barium sulphate, for that com-
pound has been shown to be insoluble in dilute hydrochloric acid.
By the process of fusion, the sodium carbonate and barium sulphate
underwent a chemical reaction, resulting in the formation of sodium
sulphate and barium carbonate. The former salt, being soluble in
water, was dissolved in that liquid along with the excess of sodium
carbonate. The barium carbonate is insoluble in water, but dis-
solves in dilute hydrochloric acid, forming barium chloride (soluble)

and carbon dioxide, which
escapes as gas. Therefore,
by fusion the insoluble
barium sulphate is converted
into soluble barium carbo-
nate.

5. Precipitation.

When chemical action
takes place between sub-
stances in solution, and one
of the products of the action
is insoluble, the latter sub-
stance is thrown out of solu-
tion, or precipitated. The
substance so thrown down is
termed a precipitate.

EXERCISE 7. Dissolve
a minute quantity of sodium
chloride (common salt) in
water in a test-tube. In
another tube dissolve a small
crystal of silver nitrate, and
mix the two solutions to-
gether. The two compounds
react upon each other, form-
ing sodium nitrate (soluble
in water) and silver chloride

FIG. 5.

(insoluble in water). The insoluble white precipitate is therefore
the silver chloride.

If, in the above example, the two substances are mixed in a
particular proportion, there may have been exactly the amount of
sodium chloride necessary to supply chlorine enough to unite with
the whole of the silver in the silver nitrate used. In this case there
would be nothing left in solution but sodium nitrate, i.e. no excess
of either silver nitrate or of sodium chloride. Ascertain if this
happened to be the case in Exercise 7, by the following experi-
ment :

Preliminary Exercises. 7

EXERCISE 8. Filter the mixture obtained above, and divide
the nitrate into two portions. To one add a single drop of a
solution of sodium chloride, (i) If a precipitate is formed, it
proves that there is some silver nitrate present, and that therefore
an excess of this compound was used in Exercise 7. Continue
adding the sodium chloride solution one drop at a time,* shaking
or stirring the liquid after every addition, so long as it produces
further precipitation.

(2) If no precipitate is thrown down by adding sodium chloride,
add to the second portion of the filtrate a single drop of silver nitrate
solution. If this gives a precipitate, it proves that sodium chloride
is present, and that therefore an excess of this substance had been
employed in Exercise 7. Continue adding the silver solution drop
by drop, with constant stirring, so as to hit off as nearly as possible
the exact point when it just ceases to produce any further pre-
cipitate.

The exact point at which precipitation is complete is not equally
easy to determine in all cases. Some precipitates are heavy,
granular, or crystalline, and settle quickly ; others again are light
or flocculent, and only subside slowly and imperfectly, so that it is
difficult to see whether the addition of more of one of the solutions
does or does not produce any additional precipitate. In such cases
the liquid should be filtered, and the filtrate tested by adding a few
drops more of the precipitant.

Very often several substances present together in one and the
same liquid, form insoluble compounds with another which is
added. These will not be all precipitated simultaneously, but in a
certain order one after the other, the precipitation of one being
more or less complete before that of the next begins. The substance
being added, first selects the compound present for which it has
the greatest chemical affinity, and afterwards that with which it
unites less eagerly. This being so, it will be evident that unless
care be taken to ensure complete precipitation, it might easily
happen that the whole of one of the substances present in the
solution escaped precipitation. It is of the utmost importance,
therefore, in analysis, to be quite sure that precipitation is as com-
plete as possible. On the other hand, the reckless addition of pre-
cipitants is a fault which must be as carefully guarded against, as
it is almost as fruitful a source of trouble as the other.

In most instances, also, it is essential to wash the precipitate
until it is quite free from any of the soluble substances present in

* When solutions are to be added drop by drop, it is best to use a pipette ;
that is, a piece of ordinary glass tube drawn to a point at one end, and about
6 or 8 inches long.

8 Qualitative Analysis.

the liquid, as explained in Exercise i. A precipitate may be re-
moved from the filter either by means of a spatula (preferably
platinum, but, failing this, either glass or porcelain ; iron should
never be used), or by pushing a glass rod through the apex of the
filter, and then washing the precipitate through by means of the
wash-bottle, or by dissolving it off by pouring into the funnel the
liquid to be used for its solution.

EXERCISE 9. Add a solution of sodium carbonate to a solution
of barium chloride, until precipitation is just complete. Barium
carbonate is precipitated, and sodium chloride remains in solution.
Pour the mixture upon three separate filters, and wrfsh the pre-
cipitate on each until quite free from sodium chloride (see Exercises
7 and 8), getting the precipitate well down into the apex.

Take the first filter, and remove a portion of the precipitate with
a spatula. If the quantity in the funnel is small, then carefully draw
the paper cone out of the funnel, spread it open upon a flat sheet of
glass, and scrape off as much of the precipitate as possible with the
spatula, and transfer it to a test-tube. Dissolve it by adding a few
drops of hydrochloric acid.

Through the apex of the second filter push a glass rod, and wash
the precipitate through into a test-tube, using a fine jet of water,
and as little of it as possible. Dissolve this also by adding a few
drops of the same acid.

Upon the third filter pour a small quantity of hydrochloric acid,
collecting the filtrate in a test-tube. Pour the filtrate back over
the precipitate once or twice, until the whole has dissolved.

6. Ignition. Strictly speaking, this word carries with it the
idea of combustion. In common speech it signifies the act of
"setting fire" to an inflammable substance ; and in more scientific
language we speak of the ignition temperature, or the igniting-
Point of a body, meaning thereby the temperature, to which it is
necessary to raise it in order that combustion may be initiated.
Unfortunately, in analytical phraseology the term ignition is used
in a somewhat slipshod way to denote a variety of operations
where substances are simply strongly heated, and where the idea
of combustion is altogether excluded. In this book the words heat
or strongly heat will be used instead of ignite to signify these
operations.

EXERCISE 10. Strongly heating in an open dish. Place a
little solution of ammonium chloride in a small evaporating-dish,
and evaporate to dryness. Then strongly heat the dish with the
dry residue until no more white fumes (consisting of the volatilising
ammonium chloride) are evolved. If the dish has been heated all
over, there should then be nothing left in it. The complete
vaporisation of the salt is more quickly and certainly accomplished

Preliminary Exercises. 9

by using a small platinum capsule or crucible in which to heat the
residue obtained by evaporating the solution to dryness.

EXERCISE u. Strongly heating in a tube closed at one end.
Place a minute quantity of mercuric oxide in a small test-tube
(4 inches x -fa), and apply heat to the compound. Note the change
of colour.; also that it gradually disappears, and that a sublimate
collects on the cool part of the tube, having a white metallic
appearance. Test the evolved oxygen by means of a glowing splint
of wood. By means of a paper " spill " rub the metallic sublimate,
and (if necessary, with a pocket lens) see the globules of mercury.

EXERCISE 12. Heating in the blowpipe flame. Select a piece
of small tubing of lead glass, and heat it in a blowpipe flame,
holding the glass in the extreme tip of the flame until it is red hot.
Then gradually bring it further into the flame, and observe that
when the glass reaches the inner cone of the flame a film begins to
appear upon the red-hot portion. On withdrawing the glass to the
tip of the flame again, this film gradually disappears. Bring the
glass once more into the inner cone of flame, and when the film
has again made its appearance, remove the glass and allow it to
cool. It will then be seen that what appeared like a film when it
was hot, is a black shining metallic-looking deposit in the glass.
This deposit is metallic lead. The lead compound in the glass,
when heated in the inner cone of flame, is reduced to the metallic
state ; and when, after being so reduced, it is heated in the tip of
the flame (i.e. in the outer cone or sheath of the flame), the metal is
again oxidised. The inner flame is therefore called the reducing
flame, and the outer cone is distinguished as the oxidising flame?

EXERCISE 13. Heating on charcoal in the blowpipe flame.
Select a close-grained piece of charcoal, as free as possible from
cracks, and file a flat surface upon it with a broad, flat file.t On
the flat part scoop a small hollow, and place in it a little red-lead
mixed with about an equal quantity of sodium carbonate. Heat
this mixture in the inner blowpipe flame, holding the blowpipe
and the charcoal in the manner shown in Fig. 6, so that the flame
shall play along the surface of the charcoal. Very quickly the lead
oxide will be reduced to the metallic state, and appear in the form
of brilliant silvery globules. When the charcoal is removed, it
will be seen that surrounding the cavity there is a yellowish deposit,
or incrustation. This consists of lead oxide. If the outer tip of
the flame be directed upon this incrustation it will quickly disappear,
and will impart a bluish colour to the end of the flame.

Pick out one or two of the globules of metal, and gently strike
one with a small hammer, or with a pestle, upon some hard surface.

* The memory, of the beginner maybe aided by the alliteration, Outer,
Oxidising. The inner flame is a reducing agent by reason of the fact that
within the cone there is an excess of strongly heated coal-gas ; whereas in the
outer flame there is an excess of heated atmospheric oxygen.

) Specially prepared rectangular blocks of charcoal (6 inches long and i
square inch section) are sold for the purpose. One such block can be used
many times.

10

Qualitative Analysis.

Note whether the metal is hard and brittle, or soft and malleable.
Also further identify the metal as lead by rubbing it upon a piece
of paper, which will be marked by it much as by an ordinary
pencil.

EXERCISE 14. Heat on another piece of charcoal a crystal or
two of zinc sulphate with a little sodium carbonate in the inner
blowpipe flame. No metallic globules are formed in this case,
because zinc is too easily oxidised ; but an incrustation appears on
the charcoal, which is canary-yellow while hot, but turns white on
cooling. Touch the incrustation with a single drop of a solution
of cobalt nitrate, and again heat it, using the outer flame. The

FIG. 6.

incrustation then becomes green. Notice that the incrustation is
not driven off by being thus heated, because zinc oxide is not
volatile.

7. Fusion, with Borax. When borax is strongly heated, it
melts to a clear vitreous mass. In this condition it is capable at a
high temperature of dissolving many metallic compounds, giving in
some cases characteristically coloured glasses.

Preliminary Exercises,

ii

EXERCISE 15. Twist the end of a piece of platinum wire into
a small round loop or eye,* and pick up a little borax upon it by first
heating the wire and then dipping it while hot into the powdered
salt. On heating the borax upon the wire in a blowpipe flame, it
first swells up, and finally fuses, forming a transparent colourless
bead of borax glass. Allow the bead to cool, and touch it with a
glass rod which has been dipped into a solution of any cobalt salt,
so as to bring only a minute quantity of the cobalt compound upon
the bead. Heat the bead once more, and notice that as it melts
the borax loses its transparent appearance. When again allowed
to cool, the bead will appear of an azure blue colour.

If too much of the cobalt salt was employed, the bead may
appear almost black ; in this case a part of it may be shaken off
when it is fluid, and more borax picked up and melted with what
remains of the original bead upon the wire. If too little cobalt is
present, the colour will be correspondingly pale. The colour of
the bead is best examined by holding it against a white object
(such as the bottle of borax itself) in a good light.

Fuse the bead again, holding it first in the outer flame, and
afterwards in the inner flame, and see that in each case when
cold the blue colour remains the same.

EXERCISE 16. Make another borax bead, and touch it with
a small quantity of a solution of manganous sulphate. Heat this
in the outer blowpipe flame. After cooling, examine the colour
carefully. Pale violet, lilac, purple, or amethyst. Heat the bead
again, holding it in the inner flame. Notice that it gradually
loses its opacity ; that as it is heated, something in the fused mass
which seems to give it an appearance of muddiness clears away,
and the molten globule looks clear. When it is in this condition
remove it, and when cold it will be found to have lost its colour
entirely. Manganese compounds there-
fore give a purplish bead in the outer
flame, which becomes colourless upon
being heated in the reducing flame.

8. Neutralisation. When an acid
is carefully mixed with an alkali (the
substances being in solution), a point is

* For greater convenience, as well as economy,
a short piece of wire (about 2 inches) should be
fixed into a glass tube, about the same length,
to serve as a handle. The glass tube is first drawn
out to a point, and the wire inserted into the fine
end. On bringing this into a blowpipe flame,
the glass fuses round the wire and holds it.
Two or three of these should be made, and a
convenient plan is to fit the glass tube into a " IG< 7-

cork, so that when not actually in use the wires

can be kept in small test-tubes containing dilute hydrochloric acid, as in
Fig. 7.

12 Qualitative Analysis.

reached when the mixture no longer possesses the properties of
either the acid or the alkali. The solution is then said to be
neutral. The point of neutrality is ascertained by the use of
certain sensitive colouring matters which have their colour changed
by acids and alkalies. The commonest of these is litmus, the
solution of which in water has a purple colour, capable of being
turned red by acids, and blue by alkalies. The yellow colour of
turmeric is changed to brown by alkalies, but is not altered by
acids, therefore this can only be used to indicate alkalinity, and
will not discriminate between a neutral and an acid liquid.

EXERCISE 17. Add a few drops of litmus solution to a little
dilute hydrochloric acid in a beaker standing upon a piece of white
paper, or a white tile. Add to the red liquid some solution of
sodium hydroxide, adding it cautiously in small quantities, with
constant stirring, until the colour of the litmus is jiist turned blue.
The liquid is now alkaline. By means of a glass rod moistened
with the dilute acid, introduce a minute additional quantity of the
acid, so as to cause the colour of the litmus to become of a purple
tint. The solution is then neutral, and the least trace of either
acid or alkali will at once turn it red or blue, as the case may be.
(Instead of adding litmus solution, papers tinted with litmus may
be dipped into the liquid.)

All substances which redden litmus are not acids, although all
acids will redden litmus. That is to say, there are many things
which have an acid or sour taste which do not belong to that class
of compounds which chemists call acids. An acid may be defined
as a compound containing hydrogen which can be displaced by a
metal, when the latter is presented to it in combination as a
hydroxide. The hydrogen which is thus displaced is not liberated
as free hydrogen, but unites with the hydrogen and oxygen of the
metallic hydroxide to form water, while the metal takes the place
of the hydrogen thus displaced from the acid.

CHAPTER II.
ANALYTICAL CLASSIFICATION.

THE word analysis, in its strict meaning, signifies the breaking
up or separation of a compound substance into its constituent
parts. It is the true antithesis of the word synthesis, which
means the building up of a compound from its constituents.

But the word analysis has come to bear a wider meaning,
and to include all the various processes and operations which
chemists make use of in order to find out what any compound is
composed of, or to enable them to identify the substance, quite
irrespective of whether or not the process involves the breaking up
of the body into its component parts. Thus, a chemist will often
recognise a substance by its particular crystalline form, or from
some other characteristic appearance it may present when examined
under a microscope (microscopic analysis]. Or sometimes he can
detect the presence of certain elements in an unknown substance,
by examining the light which is emitted when the compound is
strongly heated (spectrum analysis'}.

Reactions. Most analytical operations, however, involve
some chemical change. These changes are called reactions.
When the change is effected by strongly heating the substance,
it is described as a dry reaction, or a reaction in the dry way.
This is to distinguish this class of reactions from those which take
place between substances that are dissolved, either in water or
some other liquid, and which are sometimes spoken of as wet
reactions, or reactions in the wet way.

Most analytical reactions are " double decompositions," in which
one of the products of the chemical action is either markedly
different from the others and from the reacting compounds, in its
solubility, or its colour ; or where it is evolved as a gas having
properties by which it may be readily identified. For instance, the
two compounds barium chloride and sodium sulphate are soluble
in water, forming colourless solutions ; if these are mixed together,
"double decomposition" takes place, resulting in the formation of
sodium chloride and barium sulphate, thus

BaCl 2 + Na 2 SO 4 = aNaCl + BaSO 4

14 Qualitative Analysis.

The barium sulphate is practically insoluble in water, and con-
sequently is precipitated. Now, if we know some property belong-
ing to this precipitate of barium sulphate which is so characteristic
of the compound that we could thereby identify it and distinguish
it from all other white precipitates, then this reaction between
barium chloride and sodium sulphate can obviously be used as a
means of testing for the presence of either a soluble barium salt
or a soluble sulphate. For if, on adding a solution of sodium
sulphate to an unknown solution, barium sulphate were precipitated,
the unknown liquid must have contained a soluble barium salt ; or,
on the other hand, if we add barium chloride to an unknown solu-
tion and obtain barium sulphate again, then this unknown solution
must have contained a sulphate.*

Reagents. The materials that are used to bring about
analytical reactions are termed reagents. Thus, in the illustration
given above, the sodium sulphate is the reagent when it is added
to the unknown solution in order to test for barium ; while the
barium chloride is the reagent when it is used to test for a sulphate.
Some reagents are capable of causing reactions of a similar
character with a number of substances ; such are often known as
general reagents. Others, again, are employed' because they pro-
duce a characteristic reaction with some one substance in
particular ; these are distinguished as special reagents.

Reagents are the tools with which the analyst works, and upon
the intelligent and skilful use of them everything depends. In most
laboratories the student finds himself supplied with all the necessary
reagents ready prepared ; but for the help of those who may require
to make them up, brief directions for doing so are given in the
Appendix.

Analytical Classification. Substances are usually divided
into two classes, namely, (i) Metals, and (2) Acid-radicals.
These are also sometimes called positive radicals and negative
radicals respectively. When analysing such a compound as sodium
chloride, NaCl, the sodium and the chlorine are each separately
detected : the sodium is the metal (or positive radical), and the
chlorine is the acid (or negative) radical. But in such a case as
sodium nitrate, NaNO 3 , we do not separately detect the sodium,
the nitrogen, and the oxygen, but the sodium and the negative or
acid radical represented by the formula NO 3 . Or, again, when
such a compound as ammonium sulphate, (NH 4 ) 2 SO 4 , is submitted
to analysis, we do not separately test for the elements nitrogen,

* Sulphuric acid being included, as hydrogen sulphate.

Analytical Classification. 15

hydrogen, sulphur, and oxygen, but for the positive radical NH 4 ,
and the acid-radical SO 4 .

Sometimes the radicals, whether metals or acid radicals, may be
detected by being actually isolated, in which case they are recog-
nised by their known properties in the free state. For example,
from the compound lead chloride, PbClj, it is easy to isolate both
the lead and the chlorine. The metal lead so obtained is readily
identified by its familiar physical properties, while the gas chlorine is
equally easily distinguished by its own well-known characteristics.

In some cases, where a radical is incapable of isolated existence,
it may be detected by the separation of some product of its decom-
position. Thus, in such a compound as ammonium carbonate,
(NH 4 ) 2 CO 3 , neither the positive radical NH 4 , nor the acid-radical
CO 3 can exist in the free or uncombined -state. But we detect the
presence of the former by the evolution of ammonia, NH 3 , and the
latter by the expulsion of carbon dioxide, CO 2 , from the compound.

In the large majority of cases, however, whether the various
radicals are capable of isolated existence or not, they are detected
by causing them to pass into fresh combinations with certain
reagents, whereby new compounds are formed which are readily
recognised by their known properties. Thus, in the case of sodium
chloride above quoted, instead of isolating the chlorine, we can
employ the reagent silver nitrate, AgNO 3 . When this is added to
a solution of sodium chloride, double decomposition takes place,
and silver chloride, AgCl, is formed, which, being insoluble in water,
is precipitated. Silver chloride has properties by which it is easily
recognised, hence by the formation of this compound we can detect
the presence of the chlorine in sodium chloride.

In all such cases as these the interaction is between the

ions into which the compounds dissociate when dissolved in

water. A solution of sodium chloride, for example, contains

+ +

Na and Cl ions ; the silver nitrate contains Ag and NO 3 ions.

When these solutions are mixed, the positive silver ions unite with
the negative chlorine ions to produce the electrically neutral and
insoluble silver chloride, which therefore separates out. The silver
ions, therefore, are the test for chlorine ions, and vice versd
chlorine ions are the reagents for detecting silver ions. Any
chlorine compound which on dissociation furnishes chlorine ions,
will therefore respond to this test with silver ions. There are,
however, many compounds containing chlorine which give no
precipitate of silver chloride on the addition of silver nitrate.
Familiar among these are the chlorates and perchlorates. These

16

Qualitative Analysis.

compounds dissociate on solution, not into simple chlorine ions
but into the complex C1O 3 and C1O 4 ions. Such solutions, there

fore, contain no Cl ions and are therefore incapable of forming AgC

+
with Ag ions.

In analytical classification the term " metal " includes, beside
the metals proper, certain metalloidal elements, such as arsenic
selenium, and others (which, strictly speaking, are not true metals
but which lie on the borderland between the metals and the non
metals), and also the compound positive radical ammonium, NH.
These " metals," then, are divided into a number of groups, base<
on the behaviour of their compounds towards certain reagents.

As, however, more than one scheme or plan of analysis i
possible, due either to a different choice of reagents or to their us
in a different order, so the analytical classification of the metal
which is followed by some chemists varies somewhat from tha
used by others.

In this book the following arrangement will be adopted :

Group I.

Group II.

Group III.

Group IV.

Group V.

Silver

Mercury

Aluminium

Barium

Ammoniun

Mercury v- ^

Lead

Chromium

Strontium

Sodium

Lead*

Bismuth

Iron

XTjnlml

Calcium

Potassium

JhXafnesiui)

Thallium f

oj) j)cr 1 jki/AVA
Cadmium Cobalt

Tungsten

Antimony Manganese

Lithium

Arsenic

Zinc

Rubidium

Tin

Caesium

Gold

Beryllium

Platinum

Zirconium

Thorium

Ruthenium

Cerium

Rhodium

Scandium

Palladium

Yttrium

Osmium

Lanthanum

Iridium

Ytterbium

Tellurium

Titanium

Selenium

Tantalum

Molybdenum

Niobium
Uranium

Indium

Thallium

Vanadium

The reason why certain metals are placed in more than one division \v
appear later.

f The substances printed in small type are usually called rare elements.

Analytical Classification. 17

In the regular course of analysis, the groups are separated in
the order in which they are here numbered. The " silver, lead,
mercury " group is separated first, and the "ammonium, sodium,"
etc., group last. In studying the reactions of the metals, however,
it is more usual to begin with the fifth group, for the reason that
the compounds of these metals are less complex, and the student
is therefore led on gradually from what is comparatively simple
to that which is more difficult. Because of this, some chemists
prefer to number the groups in the reverse order, that is, in the
order in which the preliminary study of them is made, instead of
the order in which they are actually disposed of in the course of
analysis.

The reagents by means of which the elements are separated
nto these groups, and which are known as group reagents, must
ae used in regular order. Each is only capable of separating its
)wn family of metals from those coming- after it in the series, and
not those going before. For example, the group-reagent for
3roup II. is only capable of separating the metals of this family
"rom those of III., IV., and V., but not from those of Group I.
[f, therefore, the metals of Group I. are not first separated, they may
DC precipitated along with the members of the second family by
:he group-reagent for that family.

The various group-reagents (or general reagents) and the
particular compounds of the metals which are precipitated by them
ire indicated in the following table :

GROUP I. General reagent, Hydrochloric acid, HC1,

precipitates

AgCl, Hg' 2 Cl 2 , FbCl 2 (partially soluble).
H 2 WO 4 .)*

3ROUP II. General reagent, Sulphuretted hydrogen, H 2 S,

precipitates in acid solution

/ \ (T?bS, Hg S, Hi S , CuS, CdS \ Insoluble in ammo-
* ; I (Ru 2 S 3 , Rh 2 S 3 , PdS, OsS.) / nium sulphide.

j Sb 2 S 3 , As 2 S 3 , Sn !i S and Sn"S 2 , AuS, PtS 2 ( S a ^ ein

I U I \ /TO ' I , ' H It C 1 \ I dlllIllU.il*

v ' I (Ir 2 S 3 , TeS 2 , Se, MoS,. )

J sulphide.

* The compounds represented by the formulae in small type are those of
the so-called rare elements. They are included in the table in order to give, in
a. bird's-eye view as it were, not only their position in the scheme of classifica-
tion, but also the composition of their compounds, which are precipitated by
the group-reagents.

1 8 Qualitative Analysis.

GROUP III. General reagent, Ammonium sulphide, (NH 4 ) 2 S,

precipitates in presence of ammonium chloride and ammonia

(a) Hydrated I A1 2 (HO), ; , Cr 2 (HO) fi

compounds (Be(HO) 2 , Zr(HO) 4 , Th(HO) 4 , Ce(HO) 3 , Sc(HO) 3 , Y(HO) S ,
( La(HO) 3 , Yb(HO) 3) H 2 TiO 3 . H 3 TaO 4l H 3 NbO 4 .)

( FeS, NiS, CoS, MnS, ZnS.

() Sulphides j (UO 2 S, InS, T1 2 S) ( Vanadium.

converted into soluble ammonium thiovanadate *).

GROUP IV. General reagent, Ammonium carbonate,
(NK 4 ),CO i, precipitates in presence of ammonium chloride
and ammonia

BaCO , SrCO , CaCO .

GROUP V. No general reagent. The group consists of
(NH 4 ), Na, X, 'Mg. (Li, Rb, Cs.)

* Vanadium belongs to the ' ' arsenic and antimony " family in the natural
classification of the elements. The sulphide is, however, not precipitated by
H 2 S, V 2 O S being thereby reduced to V 2 O 4 , which gives a blue colour to the
liquid. NH 4 C1, in presence of ammonia, precipitates white ammonium meta-
vanadate, NH 4 VO 3 , but ammonium sulphide converts this into the soluble
ammonium thiovanadate, which gives a brown colour to the solution. The
true group-reagent, therefore, does not actually precipitate this metal.

CHAPTER III.
REACTIONS OF THE METALS OF GROUP V.

THIS group contains the alkali metals (ammonium being regarded
as a metal), and also the element magnesium, which is more nearly
allied to the metals of the alkaline earths. The members of this
family are not precipitated by any group-reagent, but they are
(with the exception of ammonium) separately tested for in the
solution which is obtained after the metals of Groups I. to IV. have
been removed. By referring to the table on p. 17, it will be seen
that, in the course of separating the various groups, certain am-
monium compounds are employed, therefore it will be obvious that
it is necessary to test for this " metal " in the substance under ex-
amination before adding any ammoniacal compounds.

Ammonium, NH 4 .

DRY REACTIONS. When heated alone in a glass tube, ammo-
nium salts undergo change.

(a) If the acid is readily volatile, the salt dissociates, but the
ammonia and the volatile acid, as they together pass away from
the heated area, immediately reunite, reproducing the original
compound, which then settles or condenses on the cool part of the
tube, forming a sublimate.

[Generally, however, a certain small amount of the dissociated
portions of the compound escapes recombination : e.g. heat a small
quantity of ammonium chloride in a dry test-tube ; notice that
white fumes are produced, which sublime up the tube. Now hold
a moistened red litmus paper in the mouth of the tube, and it will
be turned blue, showing that a portion of the ammonia escapes
from the tube before it meets the hydrochloric acid from which
it has been dissociated. For a moment discontinue heating, and
presently the blued paper will be reddened, for the molecules of
hydrochloric acid which have lost their partners (the escaped
ammonia) now make their way up the tube and act on the litmus
paper.]

2O Qualitative Analysis.

() If the acid is non-volatile, or volatile only at a high tempera-
ture, then the ammonium salt is decomposed, ammonia being
evolved, while the acid remains.

{E.g. Heat a little ammonium sulphate or phosphate in a test-
tube ; ammonia is rapidly evolved, and may be detected by its
characteristic smell.]

(c] The ammonium salts of certain oxyacids which readily part
with oxygen (such as ammonium nitrate, nitrite, chromate) are also
decomposed by heat, the ammonia being oxidised to nitrogen or
oxides of nitrogen.

[E.g. Heat a few crystals of ammonium nitrate in a test-tube.
Examine the gas with a taper and a glowing splint of wood.]

NH 4 NO 3 = 2H 2 O + N 2 O
NH 4 NO 2 = 2H 2 O + N 2
(NH 4 ) 2 Cr 2 O 7 = Cr 2 O s + 4H 2 O + ^T 2

Ammonium is separated from the other members of the group
by evaporating the solution to dryness, and strongly heating the
residue until the ammonia is completely expelled,which may generally
be regarded as accomplished when fumes are no longer given off.

WET REACTIONS. Ammonium salts are all soluble in water,
therefore it is only in concentrated solutions that any precipitations
with reagents can be formed. Use ammonium chloride.

Caustic alkalies (NaHO or KHO) and oxides or hydroxides
of metals of the alkaline earths (e.g. CaO, Ba(HO) 2 ), when heated
with an ammonium salt, cause the evolution of ammonia gas, NH S .

(NH 4 ) 2 SO 4 + 2NaHO = Na 2 SO 4 + 2H 2 O + 2NH 3
2NH 4 C1 + CaO = CaO 2 + H 2 O + 2NH 3

In practice, sodium hydroxide solution is added either to the solid
salt or to its solution in water, and the mixture gently warmed.
The evolved ammonia may be recognized (i) by its characteristic
odour if present in sufficient quantities ; (2) by its power of restoring
the blue colour to moist reddened litmus paper, or of turning tur-
meric paper brown ; (3) by the formation of white fumes when a
glass rod moistened with strong hydrochloric acid is held in the
mouth of the test-tube.

[In special cases, as in the examination of natural waters, minute
traces of ammonia are detected by the use of Nessler's solution (a
solution of potassium mercuric iodide in potash), which gives either
a brown precipitate or a coloration, according to the amount of
ammonia present, 2(HgI 2 ,2KI) + 3KHO + NH 3 = NHg 2 "I,H 2 O
+ 7KI + 2H 2 0.]

Group V. 21

Hydrogen platinum chloride* (chloroplatinic acid) H 2 PtCl 6 ,
precipitates from concentrated solutions a yellow crystalline com-
pound, ammonium chloroplatinate (or ammonium platinic chloride)
(NH 4 ) 2 PtCl 6 , soluble in 170 parts of water at 10 ; insoluble in
alcohol and ether. This compound is distinguished from the
similar potassium salt in that, when strongly heated, it leaves a
residue of spongy platinum only.

Tartaric acid, H 2 (C 4 H 4 O 6 ), or hydrogen sodium tartrate,
HNa(C 4 H 4 O 6 ), produces in concentrated solutions a white pre-
cipitate of hydrogen ammonium tartrate, H(NH 4 )(C 4 H 4 O 6 ). Soluble
in water, in mineral acids, and in alkalies ; insoluble in alcohol.
This compound is distinguished from the corresponding potassium
salt by the fact that when strongly heated, the carbonaceous residue
is without any alkaline reaction.

[When tartaric acid is used, the acid previously in combination
with the ammonia is liberated by the double decomposition, thus

NH 4 C1 + H 2 (C 4 H 4 6 ) = H(NH 4 )(C 4 H 4 6 ) + HC1

And as the precipitate is soluble in mineral acids, the delicacy of the
reaction is increased by employing hydrogen sodium tartrate as the
reagent, in which case no free acid is formed in the reaction, thus

NH 4 C1 + HNa(C 4 H 4 0) = H(NH 4 )(C 4 H 4 O 6 ) + NaCl]

Sodium, Na.

DRY REACTION. Sodium compounds, when heated upon a
platinum wire in a Bunsen flame, undergo volatilisation, and
impart to the flame a brilliant golden yellow colour. This flame-
reaction is the most characteristic and delicate test for this metal, f

WET REACTIONS. All sodium salts are soluble ; sodium platino-
chloride is soluble in water, in alcohol, and in ether. Hydrogen
sodium tartrate also is freely soluble in water. Sodium pyroanti-
monate,$ however, is less soluble in water than the corresponding
potassium salt, and is therefore precipitated by the addition of a

* By long habit this reagent is called platinum chloride : it may be
regarded as platinum chloride plus two molecules of HC1 ; PtCl 4 , 2HC1. In

+ _

solution it yields H and PtCl 6 ions.

f When the light emitted by heating a sodium salt in the Bunsen flame is
examined by the spectroscope (see p. 27), it is found to be monochromatic, i.e.
to consist of one colour only, namely, pure yellow light. Many common
coloured materials, such as indigo, have the power of absorbing yellow rays,
hence if the sodium flame be viewed through a thin stratum of such a coloured
solution, the yellow light is entirely intercepted. For the use that is made of
this property, see potassium (p 22).

J Formerly misnamed sodium metantimonate.

22 Qualitative Analysis.

strong solution of potassium pyroantimonate to a strong solution of
a sodium salt, such as sodium chloride, thus

H 2 K 2 Sb 2 O 7 + 2NaCl = H 2 Na 2 Sb 2 O 7 + 2KC1
Potassium, K.

DRY REACTION. When potassium compounds are heated upon
a platinum wire in a Bunsen flame, they impart to the flame a pale
violet or lilac colour. This delicate tint, however, is completely
masked by the intense yellow colour which the presence of even
minute quantities of sodium compounds impart to the flame.

Introduce a fragment of potassium nitrate into the Bunsen
flame upon a loop of clean platinum wire ; * notice the lilac colour
imparted to the flame. Now look at the flame through a potassio-
scope^ and observe that it appears a brilliant crimson-red colour.
Upon another wire introduce a particle of sodium chloride into the
flame, and notice that when this is examined through the potassio-
scope, the intense golden yellow light is absolutely cut off, and is
invisible. Now touch the wire containing the nitre with a fragment
of sodium chloride, and again bring it into the flame. The yellow
of the sodium completely overpowers and masks the violet of the
potassium when viewed direct, but if looked at through the potassio-
scope, the red colour due to the potassium shines up as brilliantly as
before, while the yellow is completely intercepted J (see also p. 33).

WET REACTIONS. Most potassium salts are soluble in water.
Use a solution of potassium chloride.

* By merely touching the wire with the fingers, it contracts sufficient sodium
compounds to give the yellow flame. To clean it, it should be dipped in hydro-
chloric acid and heated until it ceases to
impart any colour to the flame.

f The potassioscope consists merely
of a small, flat, glass cell, containing a
dilute solution of one of the aniline blue
dyes, known as "soluble blue X.L."
The advantage of this over ordinary blue
glass or the older indigo prism lies in the
fact that no other metal but potassium
(except the extremely rare element rubi-
dium) gives a flame which appears red
when viewed through the potassioscope,
whereas lithium, barium, strontium, and
calcium all give flames which appear red
through indigo or blue glass.

J When studying flame reactions, it is
often of the greatest convenience to use a
stand on which to support the platinum
wire, so that the hands may be free ; a
simple stand is readily constructed as
shown in Fig. 8. A piece of glass tube
or glass rod is inserted in a large cork

p IG g (rubber, being heavier, makes a steadier

foot), and a piece of galvanized iron wire
is twisted two or three times round the rod with one end projecting at right

Group V. 23

Hydrogen platinum chloride (chloroplatinic acid\ H 2 PtCl G ,
produces, with concentrated solutions of potassium salts, a yellow
crystalline precipitate of potassium chloro-platinate (or potassium
platinic chloride), K 2 PtCl 6 , soluble in no parts of water at 10
(therefore more soluble than the corresponding ammonium com-
pound). Soluble in alkalies (therefore the solutions used should
be acid) ; nearly insoluble in alcohol ; quite insoluble in a mixture
of alcohol and ether (therefore the precipitation of this compound
is promoted by the addition of alcohol).

Hydrogen sodium tartrate, HNa(C 4 H 4 O 6 ), gives, with solu-
tions of potassium salts, a white crystalline precipitate of hydrogen
potassium tartrate, HK(C 4 H 4 O 6 ) ; soluble in much water, and also
in acids and alkalies (therefore the solution should be both concen-
trated and neutral). The precipitate is insoluble in alcohol.

Hydrofluosilicic acid (or silico-fluoric acid), H 2 SiF 6 ,
throws down a white precipitate of gelatinous appearance, consisting
of potassium silicofluoride, K 2 SiF 6 , sparingly soluble in water.

Magnesium, Mg.

In the " natural classification " of the elements, magnesium is
associated with the rnetals of the alkaline earths (Be, Ba, Sr, Ca)
on the one hand, and with zinc and cadmium on the other. Its
position along with the alkalies in Group V. of the analytical
classification, is simply because it differs from the other members of
its own natural family in that the presence of ammonium chloride
prevents the precipitation of magnesium hydroxide by ammonia
in Group III., and also of magnesium carbonate by the group
reagent of Group IV.

DRY REACTION. When magnesium salts are strongly heated
in the outer blowpipe, a white infusible residue of the oxide is left.
If, after cooling, the residue be moistened with a drop or two of
cobalt nitrate solution and again strongly heated in the outer blow-
pipe flame, the mass acquires a pink colour. This reaction is
reliable only in the absence of other metallic oxides.

WET REACTIONS. Of the common salts of magnesium, the
sulphate, chromate, nitrate, and halogen salts are soluble in water.
One prominent characteristic of magnesium compounds is the
readiness with which they form " double " salts, many of which are
soluble in water. Use magnesium sulphate.

angles to the upright. The little glass tube into which the platinum wire i
fused, is then slipped over the projecting iron wire. This arrangement acimiis
of the wire being raised or lowered as desired, while at the same time it readily
remains in any position.

24 Qualitative Analysis.

Alkaline hydroxides (NH 4 HO, KHO, NaHO, Ca(HO) 2 , or
Ba(HO) 2 ) precipitate from solutions of magnesium sulphate or
chloride, white magnesium hydroxide, Mg(HO) 2 . Almost insoluble
in water ; soluble in acids, soluble in ammonium chloride

Mg(HO) 2 + 4NH 4 C1= 2NH 4 HO + MgCl 2 ,2NH 4 Cl (soluble
double salt)

Owing to the solubility of magnesium hydroxide in ammonium
chloride, Only half the magnesium is precipitated from magnesium
chloride by means of ammonia, thus

2MgCl 2 + 2NH 4 HO = Mg(HO) 2 + MgCl 2 ,2NH 4 Cl

If ammonium chloride is previously present in sufficient quantity
the alkaline hydroxide gives no precipitate.

Stated in terms of the ionic theory, the solubility of magnesium
hydroxide in ammonium chloride is due to two causes ; first the

slight solubility of the compound in water, with a corresponding

+

slight dissociation into Mg and HO ions ; and second the com-
paratively small extent to which ammonium hydroxide is ionised.
The addition of ammonium chloride throws into the solution a large

+
excess of NH 4 ions, one effect of which is to reduce the ionisation

of ammonium hydroxide to such an extent that there are no HO

ions available for union with Mg ions ; hence no precipitate of

Mg(HO) 2 occurs. On the other hand, if ammonium chloride be

+
added after the precipitation of Mg(HO) 2 , then the NH 4 ions thus

introduced unite with the HO ions provided by the slightly ionised
magnesium hydroxide, yielding NH 4 HO (practically unionised).

To restore the equilibrium disturbed by this removal of HO ions,
more of the Mg(HO) 2 undergoes ionisation, and this process
continues until the whole of the Mg(HO) 2 has passed into solution.
The above equation then becomes

Mg, 2HO + 4NH 4 , 4C1 = 2NH 4 HO + [Mg, Cl, Cl, 2N + H 4 , 2C1]

ionised double salt.

Alkaline carbonates (K 2 CO 3 , Na 2 CO 3 , (NH 4 ) 2 CO 3 ) pro-
duce in solutions of magnesium salts, in the absence of ammonium
salts, precipitates of basic carbonates of magnesium, the composition
of which varies with conditions of temperature and concentration.
The precipitate with (NH 4 ) 2 CO 3 only separates out after a short

Detection of tJie Metals of Group V. 25

time. In the presence of ammonium chloride these reagents give no
precipitate.

Hydrogen disoditun phosphate, HNa 2 PO 4 , precipitates
hydrogen magnesium phosphate, HMgPO 4 , and tri-magnesium
phosphate, Mg 3 (PO 4 ) 2 . In the presence of ammonium chloride,
however, the double ammonium magnesium phosphate is thrown
down as a white crystalline precipitate, NH 4 MgPO 4 . It is
appreciably soluble in water, but insoluble in ammonia; hence
ammonia must be previously added.

In very dilute solutions the precipitation only takes place on
long standing. It is accelerated by stirring with a glass rod, the
deposition first appearing where the rod has rubbed the glass
vessel. The precipitate is soluble in acids, even acetic acid, but
reprecipitated by ammonia.

SYSTEMATIC PLAN OF ANALYSIS FOR THE METALS OF
GROUP V.

After having carefully gone through the various reactions for
the metals of Group V., the student should proceed to the examina-
tion of a few solutions containing mixtures of two or more salts of
these metals.

The various special tests by which the individual members of
this group are recognised are not, for the most part, interfered with
by the presence of the rest of the group. Therefore a complete
separation of all the metals is not necessary. Thus, the test for
ammonium (evolution of ammonia by heating with sodium
hydroxide) can be made in the presence of Mg, K, and Na com-
pounds ; and obviously must be made in a separate portion of the
solution under examination from that in which sodium is to be
tested for, as it involves the addition of a sodium compound.

The test for magnesium, likewise (precipitation of NH 4 MgPO 4 ),
may be made in the presence of all the other members ; and
clearly must be made also in a separate portion of the solution, as
it involves the addition of both ammonium and sodium compounds.
Similarly, the flame tests for potassium and sodium are not inter-
fered with by the presence of ammonium or magnesium. The
flame reaction for potassium, however, unless examined by the aid
of the potassioscope, must always be corroborated by the forma-
tion of potassium chloroplatinate But before this test can be
applied, ammonium salts must first be removed.

Solutions may be examined for the metals of Group V.,
NH 4 , Na, K, Mg, by the following system :

26 Qualitative Analysis.

DETECTION OF THE METALS OF GROUP V.

Operation I. To a portion of the solution add NaHO, and heat
in a test-tube. The evolution of ammonia (detected by its odour,
and its action on test-papers) proves the presence of NH 4 .

Operation 2. To a second portion add NH 4 C1, NH 4 HO, and
HNa 2 PO 4 . A white crystalline precipitate of NH 4 MgPO 4 proves
the presence of Mg.

Operation 3. Evaporate another (and larger) portion to dryness
in a porcelain dish. If ammonium salts are present * (already
ascertained in Operation i), they must be removed. For this pur-
pose, scrape the residue out of the dish and strongly heat it on the
lid of a platinum crucible (or a piece of platinum foil), until, on
momentarily withdrawing it from the flame, fumes are no longer
visible.

Dissolve the residue in a small quantity of water, and add one
drop of HC1. Dip a clean platinum loop into the solution, and heat
it in a Bunsen flame. An intense f yellow coloration proves the
presence of Na. A lilac colour indicates the presence of K.

In either case examine the flame through the potassioscope ; a
crimson flame indicates K.

Add to the solution of the residue a few drops of H 2 PtCl 6 , and
stir with a glass rod. A yellow precipitate of K 2 PtCl fi confirms K.

APPENDIX TO CHAPTER III

LITHIUM, RUBIDIUM, AND CAESIUM.

These three elements are usually placed in the category of rare
metals. It must be remembered, however, that there are degrees of
rarity ; and while the compounds of rubidium and caesium are
certainly among the very rare substances with which the chemist
comes into contact, those of lithium, on the other hand, are very
widely distributed and are much more frequently met with.|

* In a complete analysis, ammonium salts are always present here, as they
will have been introduced in the process of separating the other groups.
Under these circumstances, therefore, the operation of removing ammonium
compounds is always necessary. The substance under analysis is tested for
ammonium before the ammoniacal reagents are introduced.

t More or less of a yellow flame is usually obtained, owing to the presence
of traces of sodium compounds as impurities in the reagents previously used in
separating the groups in the course of a complete analysis.

J Perhaps a rough idea of the relative rarity of the compounds of these
metals might be gained by a comparison of their cost. Lithium salts can be
obtained for about 121. per pound, while rubidium and caesium salts cosi
about 5^. per drachm, i.e. at the rate of ^64 per pound.

The Rare Metals of Group V 27

Lithium, Li.

DRY REACTION. Lithium salts impart to the flame a brilliant
carmine-red colour.

WET REACTIONS. All the common salts are readily soluble in
water, except the carbonate, phosphate, and oxide, which are
soluble with difficulty. The chloride and nitrate are soluble in a
mixture of alcohol and ether (distinction from No. and A", the
chlorides and nitrates of which are not soluble).

Na 2 CO 3 and K 2 C0 3 precipitate Li 2 CO 3 from cold moderately
concentrated solutions (i part dissolves in 100 parts of water).

HNa 2 PO 4 gives a white precipitate, on boiling, of Li a PO 4 . The
precipitation is complete in the presence of NaHO.

PtCl 4 gives no precipitate (distinction from NH 4 , K, Rb, Cs).

HNa(C 4 H 4 O 6 ), hydrogen sodium tartrate, gives no precipitate.

Rubidium, Rb, and Caesium, Cs.

The compounds of these metals present the very closest
resemblance to those of potassium, and there are scarcely any
chemical reactions by which they can be distinguished. The
separation of these metals is based on the different degrees of
solubility of their chloro-platinates. In this respect, as well as in
their other properties,* rubidium stands intermediate between
potassium and caesium.

When heated in a Bunsen flame, rubidium and caesium salts
impart to it a lilac colour, which to the unaided eye is absolutely
indistinguishable from that produced by potassium compounds ;
and when compounds of these metals, as well as those of lithium,
are mixed with comparatively minute quantities of sodium salts, the
colours they give to the flame are completely overpowered and
masked by the yellow of the sodium. By means of the spectroscope,
however, not only are the apparently identical colours given by
potassium, rubidium, and caesium proved to consist of light of
different quality or composition, but the presence of any or all of
them is easily and certainly detected even when admixed with
sodium salts.

The spectroscope is an instrument by means of which the light
emitted by strongly heated substances can be examined after it has
been made to pass through a glass prism. Its use depends upon
the fact that different coloured lights possess different degrees of
refrangibility ; that is to say, different coloured rays of light are
bent out of their straight course, by passage through a prism, at
different angles. Ordinary white light is composed of rays of all
degrees of refrangibility, hence, when such light passes through a

* Atomic weights, melting-points, etc. Also in the optical properties of
their crystallized salts (Tutton, J. C. S., May, 1896).

28 Qualitative Analysis.

prism, the various coloured rays are separated, and spread out in the
order of their refrangibility, the least refrangible red at one extreme,
to the deep violet at the other. This familiar " rainbow " coloured
band of light is called the continuous spectrum.

In the spectroscope the light is passed through a narrow slit at
one end of a small telescope, and an image of the slit is received
upon a glass prism. This bends the light out of its straight course,
and spreads it out into the various colours of which it is composed,
If white light be admitted, then the continuous spectrum is seen,

BLUE.. GREEN.

T3I*

(V.:

I !IIH!~ZI>

IT

FIG. 9.

which is an infinite number of images of the slit arranged side by
side ; if such a ;;w#<?chromatic light as that given by heating sodium
salts in the flame be used, then one image of the slit is seen in that
part of the spectrum which corresponds to the particular degree of
refrangibility of the light. In this case we say that the spectrum
of sodium consists of one line (that is, one image of the slit) in the
yellow, or one yellow line* The light which passes out of the prism
is usually examined by a second telescope, which can be revolved
round the prism so that it can sweep the whole length of the
spectrum.

When the red colour imparted to the Bunsen flame by a

* When this sodium line is examined by a higher dispersive power, it is
found to consist of a group of lines.

The Rare Metals of Group V. 29

lithium salt is examined by the prism, it is seen to consist of light
of two degrees of refrangibility, therefore two images of the slit are
seen, one in the bright orange (a little on the red side of the sodium
line), and the other an extremely brilliant image in the red. The
spectrum of lithium, therefore, is one bright orange line and one
brilliant red line (see Fig. 9).

Similarly, when the lilac flame of potassium is examined in this
way, the light is found to be composed of three colours ; therefore
three lines are seen one extremely bright line far down in the red ;
a second, less brilliant, about halfway between the first and the
yellow sodium (or D) line ; while the third is far away in the deep
violet, and not easy to see.*

If a mixture of salts of sodium, potassium, and lithium be heated
in a flame, and the light examined by the spectroscope,! then all
six lines are seen ; sodium, one yellow ; potassium, two red and one
violet; lithium, one red and one orange.

In such a simple mixture as this there is no difficulty in
identifying the various lines of the different elements, but in more
complex spectra it is only possible to do so by recording the exact
position they each occupy on some fixed scale.

When the spectra of rubidium and caesium are in this way com-
pared with that of potassium, a striking difference is at once
observed. In the first place, they are obviously much more complex.
Rubidium shows two strongly marked red lines (hence the name
of the element) very close together, and near to the red line of
potassium. Then a number of lines in the orange-red and in the
green, and lastly two characteristic and brilliant violet lines.

Caesium gives red lines (not so far down as those of rubidium)
orange and green lines, and two most characteristic brilliant
lines in the bright blue part of the spectrum (ccesium signifies
sky -blue).

In the chart (Fig. 9), the most prominent lines J in the spectra
of the five metals of the alkalies are compared together. The
letters at the top refer to lines in the solar spectrum mapped by
Fraunhofer, which correspond to those given by known metals.
Thus, the D line in the solar spectrum corresponds to the yellow
line of sodium.

* The sodium D line will be also visible, not only on account of the presence
of traces of sodium compounds in the other salts, but owing also to the
presence of particles of salt which are always floating about in the air, and
which tinge the Bunsen flame.

t Where the equipment of the laboratory will allow, every student should
have the opportunity of examining the spectra of the metals of the alkalies and
alkaline earths.

% Speaking generally, it may be said that at higher temperatures than can be
obtained in the Bunsen flame, lines will appear in the spectrum of a particular
metal which are not seen at the lower temperature.

CHAPTER IV.
REACTIONS OF THE METALS OF GROUP IV.

THIS group consists of the three metals, barium, strontium,
and calcium, known as the metals of the alkaline earths.

The group-reagent, ammonium carbonate, precipitates the
carbonates of the metals in an ammoniacal solution, even in the
presence of ammonium chloride. Hence the group-reagent separates
these three metals from magnesium.

Barium, Ba.

DRY REACTION. Barium compounds, heated on platinum wire
in the Bunsen flame, impart a pale apple-green colour to the flame,
which becomes more distinct if the substance on the wire is
moistened with strong hydrochloric acid. The test is not very
reliable.

Barium sulphate, BaSO 4 (also SrSO 4 and CaSO 4 ), when heated
on charcoal or with carbon, is reduced to the sulphide.

WET REACTIONS. Of the common salts of barium, the chloride,
bromide, iodide, nitrate, chlorate, acetate, and sulphide are soluble
in water.

Ammonium carbonate, (NH 4 ),CO 3 , group-reagent (also
Na 2 CO 3 and K 2 CO 3 ) precipitates barium carbonate, BaCO 3 , as a
white amorphous powder. Insoluble in water ; readily dissolved,
with evolution of carbon dioxide, by dilute acids ; slightly soluble in
NH 4 C1.

It is also dissolved by a solution of carbon dioxide in water,
forming hydrogen barium carbonate, H 2 Ba(CO 3 )2.*

* On this account complete precipitation of barium carbonate cannot be
accomplished by hydrogen ammonium carbonate, HNH 4 CO 3 , for carbon
dioxide is evolved, which dissolves a portion of the precipitate

2 HNH 4 CO 3 + BaCl 2 = BaCO 3 -f 2NH 4 C1 + H 2 O f CO 2

And since the solution of ammonium carbonate used as the group-reagent
consists chiefly of HNH 4 CO 3 (because the normal salt undergoes decomposition
into the hydrogen salt and free ammonia when in aqueous solution), it is always
necessary to first add NH 4 HO before applying the group-reagent.

Group IV. 31

H 2 SO 4 , or any soluble sulphate, produces a white granular
precipitate of BaSO 4 , practically insoluble in water, insoluble also
in acids and alkalies. (Boiling concentrated H 2 SO 4 slowly dis-
solves it, forming hydrogen barium sulphate, H 2 Ba(SO 4 ) 2 .) In-
soluble in solutions of (NH 4 ) 2 SO 4 . BaSO 4 , being practically insoluble
in water, is precipitated by a saturated solution of SrSO 4 , although
such a solution contains only I part of salt in 7000 parts of water.

Potassium chromate, K 2 CrO 4 , produces a primrose-yellow
precipitate of barium chromate, BaCrO 4 , practically insoluble in
water (distinction from SrCrO 4 , which is MODERATELY soluble, and
CaCrO 4 , which is VERY soluble in water). It is insoluble in acetic
acid (distinction from SrCrO 4 ) ; soluble in HNO 3 and in HCL
By boiling with K 2 CO 3 , barium chromate is converted into the
carbonate

BaCrO 4 + K 2 CO 3 = BaCO 3 + KjCrO,

Hydrofluo silicic acid, H 2 SiF 6 , gives a white crystalline pre-
cipitate of barium silicofluoride, BaSiF 6 , slightly soluble in water,
but insoluble on the addition of alcohol (distinction from SrSiF 6
and CaSiF 6 , which are readily soluble in water).

Strontium, Sr.

DRY REACTION. When heated in the Bunsen flame, volatile
strontium salts, such as SrCl 2 , Sr(NO 3 ) 2 , impart a rich crimson
colour to the flame ; other salts require to be moistened upon the
wire with strong HCL Strontium sulphate must be converted into
the sulphide, which is then dissolved in HC1 and the chloride
tested in the flame.*

WET REACTIONS. The same common salts of strontium as of
barium are soluble in water. The chromate and sulphate are
somewhajt soluble.

(1TH 4 ) 2 CO 3 (Na 2 CO 3 and K 2 CO 3 ) precipitates SrCO 3 , exactly
similar to the barium compound in its reactions.

H 2 SO 4 or soluble sulphates precipitate SrSO 4 . The precipitate
is slightly soluble in water (i : 7000), considerably soluble in HC1
or HNO 3 , but almost insoluble in a solution of (NH 4 ) 2 SO 4 (dis-
tinction from CaSO 4 ). SrSO 4 is precipitated by a solution of

* Small quantities of SrSO 4 upon a filter may be converted into SrS by first
drying the filter, then folding it into a small roll, and twisting a platinum wire
round it so as to hold it in a flame. The paper is burnt to an ash, and the
latter is then heated upon the wire in an ordinary smoky gas-flame, when the
reduction of the sulphate takes place. The ash is then touched with a glass
rod dipped in strong HC1, and on bringing it into a Bunsen flame the crimson
flame is seen.

32 Qualitative Analysis.

CaSO 4 * (a saturated solution of which only contains i part in
4.30 parts of water) ; the precipitation does not take place at once
in cold solutions, but appears quickly on heating.

K 2 CrO 4 and H_SiF (see barium reactions with these reagents)

Calcium, Ca.

D RY REACTIONS. Calcium compounds, when heated in a Bunsen
flame, impart to it a reddish colour, especially if previously moistened
with hydrochloric acid. The presence of strontium masks the red
colour given by calcium compounds. When calcium carbonate is
strongly heated it loses carbon dioxide, and is converted into
calcium oxide

CaCO 3 = CO 2 + CaO

This process is carried on in the lime-kilns during the operation of
" burning lime." Limestone (calcium carbonate] is thus converted
into lime (calcium oxide). SrCO 3 and BaCO 3 undergo similar
decomposition ; SrCO 3 less readily than CaCO s , while BaCO 3
requires prolonged heating to a white heat to effect the change.

WET REACTIONS. The same common salts of calcium are
soluble in water as of strontium and barium ; and, generally
speaking, calcium salts are more readily soluble. Thus, the chloride
and nitrate are extremely deliquescent ; the sulphate dissolves to
the extent of I : 430 (compare Sr and Ba). The solubility of the
oxalates, however, is in the reverse order, calcium oxalate being
the most insoluble, and barium oxalate the most soluble. A
peculiarity of calcium salts is that many of them appear to be less
soluble in hot than in cold water.

(NH 4 ) 2 CO 3 (also Na 2 CO 3 and K 2 CO 3 ) precipitates CaCO 3 ,
similar to the barium and strontium compounds in its reactions.

H.SO, or soluble sulphates,! added to a strong solution of a
calcium salt,J give an immediate precipitate of calcium sulphate.

* SrSO 4 is itself soluble in water to the extent of i : 7000, but obviously the
addition of a solution of CaSO 4 or H 2 SO 4 to this solution does not precipitate
the SrSO 4 . Neither is the strontium precipitated from this solution by any
reagent which forms with it a compound which is more soluble in water than
is the sulphate, e.g. oxalic acid or oxalates.

f The soluble sulphate must be more soluble than calcium sulphate ; SrSO 4
therefore, being less soluble, would not give a precipitate with a solution of a
calcium salt, although, being more soluble than BaSO 4 , it throws down this
compound from solutions of barium salts. For the same reason, a solution of
CaSO 4 will throw down SrSO 4 and BaSO 4 from solutions of strontium and
barium salts.

J Only in the case of calcium salts which arc more soluble than calcium
sulphate.

Groiip IV. 33

From more dilute solutions the precipitate only separates after
some time, or, if still more dilute, not at all. The precipitate is
insoluble in alcohol, therefore the addition of this liquid in
considerable bulk favours the precipitation.

Calcium sulphate is readily soluble in a concentrated solution
of ammonium sulphate, especially when hot (distinction from
SrSO 4 and BaSO 4 ). Boiling with potassium carbonate easily
converts it into calcium carbonate.

Oxalic acid, H 2 C 2 O 4 , or soluble oxalates, gives a white
crystalline precipitate -of calcium oxalate. The precipitate is
soluble in mineral acids ; hence, when oxalic acid is used as the
precipitant, precipitation is not complete, owing to the liberation
of the mineral acid of the calcium salts, thus

CaCl.; + H 2 C 2 O 4 = CaC 2 O 4 + 2HC1

Calcium oxalate is insoluble in NH 4 HO, therefore, if the solution
be first rendered alkaline, the whole of the oxalate is thrown down.
In practice it is usual to employ ammonium oxalate, (NH 4 ) 2 C 2 O 4 .
In cold dilute solutions precipitation is not immediate. The pre-
cipitate is insoluble in water and in acetic acid. Strontium and
barium oxalates are soluble, to a small extent, both in water and
in acetic acid.

SPECTRA OF BARIUM, STRONTIUM, AND CALCIUM.

The spectra given by these metals, when their salts are heated
in the Bunsen flame, are extremely characteristic. For calcium,
the chloride may be used, with occasional moistening with hydro-
chloric acid. For barium and strontium, the nitrates, or, better
still, the chlorates give the best result, although in the absence of
these salts the chlorides can be employed.

From the charts of their spectra given in Fig. 10, it will be seen
that they each give lines in the red portion of the spectrum. Owing
to this fact, if salts of these metals are heated in a Bunsen flame,
and the flame examined through an indigo firism, the flame will
appear red'm each case, because indigo does not intercept or absorb
the red portion of the spectrum. The potassioscope, on the other
hand, has the property of cutting off the rsd rays emitted by these
metals. It absorbs all the red part of the spectrum down very
nearly to the prominent red line (K.o) of potassium (compare the
chart on p. 28), beyond the lithium and caesium red lines. Hence,
if flames which are coloured by the alkaline earths are examined
through this instrument, no red will be visible.

34

Qualitative Analysts.

The importance of this lies in the following fact : When Groups
IV. and V. are separated by precipitating the carbonates of barium,
strontium, and calcium, minute traces of these metals pass through
into Group V., as their carbonates are not absolutely insoluble
Hence, when the final residue is obtained (in which the tests for

FIG. 10.

potassium are made), it will contain these traces of these alkaline
earths. If this residue be tested in the flame, and the flame
examined by the potassioscope, the presence of these members of
Group IV. will not vitiate the reaction ; but with the indigo prism
the red light they emit would be mistaken for, and returned as that
given by potassium.*

SEPARATION OF GROUPS IV. AND V

The solution is first rendered alkaline by the addition of NH 4 HO.
NH 4 C1 is then added (to prevent the precipitation of magnesium
carbonate), after which ammonium carbonate is added until the
carbonates of the metals of Group IV. are completely thrown down.
The mixture may be gently warmed. [It must not be boiled, or
the precipitated carbonates will react with the NH 4 C1, forming
soluble chlorides, while NH 3 and CO 2 will escape with the steam ;
thus, BaCO 3 + 2NH 4 C1 = BaCl. 2 + CO 2 + aNH 3 + H 2 O.] The
mixture is filtered. The filtrate is examined for the metals of
Group V. by the method given on p. 25, while the precipitate is
treated in the following way :

* The only element which could be mistaken for potassium, when using the
potassioscope, is the exceedingly rare metal rubidium, whose spectrum contains
red lines still lower down than those of potassium.

Group IV. 35
SEPARATION OF THE METALS OF GROUP IV.

te precipitate, consisting of PaCO,, SrCO,, and CaCO,, is washed,
and then dissolved in a small quantity of warm dilute acetic acid,
HCC HO) To the solution thus obtained potassium chromate,
;, is added ; the mixture gently warmed and filtered.

The precipitate con-
sists of yellow barium
chromate, BaCrO 4 .

Boil the precipitate with
K a CO a , which converts
BaCrO, into BaCO 3 .
Filter, and wash the
precipitate free from
the K 2 CrO 4 . Dissolve
the precipitate in HC1,
and confirm by SrSO4
or H 2 SiF 6 .

The filtrate t contains strontium and calcium
acetates. Add a strong solution of
(NH 4 ) 2 SO 4 , and boil for a short time to
ensure the solution of CaSO 4 . Filter.

-The precipitate con-
sists of SrSO 4 .

Wash the precipitate
thoroughly, and con-
firm by the flame
reaction.

The solution contains
calcium sulphate.
Add ammonium oxa-
late, (NH 4 ) 2 C S O 4 (or
(NH 4 )HO and oxa-
lic acid, H 2 C 2 O 4 ).
A white precipitate
of CaC 2 O 4 confirms
calcium.

Other methods of separating the metals of Group IV., based on
the insolubility of certain of their common salts in alcohol, may
be employed. Thus, the nitrates of barium and of strontium are
insoluble in alcohol, while that of calcium is soluble in this liquid
(or mixtures of alcohol and ether).

The precipitated carbonates are dissolved in dilute HNO 3 , and
the solution evaporated to dryness. The residue, consisting of the
mixed nitrates, Is then treated with a mixture of alcohol and ether,
in about equal volumes, which dissolves the Ca(NO 3 ) 2 , leaving
Ba(NO 3 ) a and Sr(NO 3 ) 2 .

Again, barium chloride is insoluble in the same mixture, whereas
the chlorides of strontium and calcium are both soluble. Therefore,
if the mixed carbonates are dissolved in dilute HC1, and the solu-
tion evaporated to dryness, the SrCl 2 and CaCl 2 may be dissolved
out, leaving a residue of BaCl 2 .

* Instead of adding the chromate at once to the whole of the solution, a
small separate portion may be used, and the presence or absence of the Ba
thus ascertained. If present, then the whole of the solution is treated as above,
and the precipitation of barium chromate may then be taken as confirmatory.
If absent, then a second small portion of the solution is taken, and tested for Sr
by adding CaSO 4 . If Sr is present, then the separation of Sr and Ca is carried
out as shown above ; but if CaSO 4 gives no precipitate even on warming,
thus proving the absence of Sr, the remainder of the solution is at once tested
for Ca by the oxalate reaction.

t If barium has been separated as chromate, the filtrate will be coloured
yellow by the excess of K 2 CrO 4 used, and a white precipitate may appear
yellow. The process is more thorough if the carbonates of Sr and Ca are
again thrown down by adding ammonium carbonate. The precipitate is then
washed free from chromate, dissolved in a few drops of HC1, and the sulphates
separated as directed above.

CHAPTER V.
THE METALS OF GROUP III.

THIS group comprises the metals aluminium, chromium, iron,
manganese, zinc, nickel, and cobalt, besides a number of
the rare metals (see Classification, p. 16).

Although all these metals are precipitated by the group-reagent,
namely, ammonium sulphide, in the presence of ammonia, the
compounds that are thrown down are not all of similar composition,
as was the case in Group IV.

The addition of ammonia to a solution containing the metals
of this group, results in the precipitation of the hydroxides (or
hydrated oxides) of Al, Cr, and Fe, corresponding to the oxides
R 2 O 3 .* Thus, taking aluminium in potash alum as an example

K 2 SO 4 ,A1 2 (SO 4 ) 3 + 6NH 4 HO = K 2 SO 4 + 3(NH 4 ) 2 SO 4 + A1 2 (HO) 6

In the presence of ammonium chloride, the hydroxides of the
metals Mn,t Zn, Ni, and Co (which have the general formula
R(HO) 2 ) are not precipitated by ammonia, for the same reasons as
apply in the case of magnesium. (See p. 24.)

If ammonium sulphide be added to the mixture after the
ammonia, the already precipitated hydroxides of aluminium and
chromium are unchanged, but the ferric hydroxide is converted
into ferrous sulphide, with the elimination of sulphur ; thus

Fe 2 (HO) 6 + 3H 2 S = 2FeS + 6H 2 O + S
and the remaining metals are also thrown down as sulphides.

* The separation of the metals of Group III. is complicated by the fact that
if Mn, Zn, Ni, and Co are present in the form of phosphates, the addition of
ammonia causes the partial precipitation of these phosphates. And, further,
if the members of Group IV. and Mg are present as phosphates, they also are
thrown down by ammonia, and therefore appear in Group III. The special
treatment of phosphates will be discussed later.

f In the case of manganese, the separation is not complete under all
conditions (see reactions of manganese).

Group III. Division A. 37

The result, therefore, of adding NH 4 C1, NH 4 HO, and (NH 4 ) 2 S is
as follows :

Precipitated by]
NH 4 HOinthe

Precipitated on the
addition of (NH 4 ) 2 S

, >u- a (,HU),

J Fe 2 (HO) 6 converted into FeS

MnS
ZnS
NiS
CoS

The metals of Group III V therefore, may be subdivided into
two families, based upon their behaviour towards ammonia : Divi-
sion A, consisting of the three metals, aluminium, chromium, and
iron ; and Division B, of manganese, zinc, nickel, and cobalt.

REACTIONS OF THE METALS OF GROUP III. DIVISION A.
Aluminium. Al.

DRY REACTION. When aluminium compounds are strongly
heated on charcoal in the outer flame, aluminium oxide is formed,
and if this be moistened with a solution of cobalt nitrate, and again
strongly heated, either upon the charcoal or upon a loop of plati-
num wire, the mass assumes a rich blue colour, due to the formation
of cobalt aluminate.

This test is, however, greatly masked if other metallic oxides
which are coloured are present at the same time. It may be
employed as a confirmatory test when aluminium is separated from
iron and chromium in the course of analysis.

WET REACTIONS. Of the common salts of aluminium, the
chloride, A1 2 C1 6 , and sulphate, A1 2 (SO 4 ) 3 , are soluble in water.
The important salts, however, are the double sulphates of
aluminium with ammonium or potassium, known as ammo-
nium alutn, (NH 4 ) 2 SO 4 ,A1 2 (SO 4 ) 3 ,24H 2 O, and potassium alum,
K 2 SO 4 ,A1 2 (SO 4 ) 3 ,24H 2 O, respectively.! A solution of either of
these alums may be used for the following reactions.

* In reality the compounds precipitated are the hydrated sulphides, ex-
pressed by the general formula R(HS)(HO), or RS,H 2 O. To avoid unnecessary
complication in reactions, the molecule of H 2 O may be left out of consideration.

f The alums constitute a large class of double sulphates, having the general
formula M^SO^R^SO^s^I^O, where M is a monovalent element or group,
such as potassium or ammonium, and R is either aluminium, iron, chromium,
or manganese. The commonest of all the salts is potassium aluminium alum,
K 2 SO 4 ,Al2(SO 4 )3,24H 2 O ; this is the salt, therefore, that is distinguished by
the single word ALUM. From their formulae it may be seen at a glance that

3 8 Qualitative Analysis.

NH 4 HO throws down a white translucent precipitate of the
hydrated oxide, or hydroxide,* A1 2 O 3 , 3H 2 O, or A1 2 (HO) 6 . Soluble
in a large excess of the reagent, but on gently boiling, the hydroxide
in entirely precipitated. [Prolonged boiling, however, causes partial
dissociation of the ammonium salt in solution into ammonia, which
escapes, and free acid, which then begins to dissolve the preci-
pitate.] In the presence of ammonium chloride, the precipitation
of A1 2 (HO) 6 by ammonia is complete. The precipitate is readily
soluble in mineral acids, and in acetic acid.

KHO or NaHO produces the same precipitate, readily soluble
in an excess of the reagent, forming potassium or sodium aluminate
(Al 2 O 3 ,3Na 2 O or Na 6 Al 2 O 6 ). The ready formation of these
aluminates is due to the fact that aluminium hydroxide can exhibit
feeble acidic properties ; that is to say, it is ionised to some extent

into H and A1O 3 ions. With strong bases, therefore, we have the
reaction

H 6 A1 2 O 6 + 6NaHO = 6H 2 O + Na 6 Al 2 O 6 or

In the case of the slightly ionised base ammonium hydroxide it is
probable that ammonium aluminate is first formed, and then
undergoes hydrolysis, this latter change being aided by the gentle
warming which determines the complete precipitation of the
hydroxide. These aluminates are decomposed by acids, even by
such feeble acids as carbonic acid or hydrosulphuric acid
(sulphuretted hydrogen), with re-precipitation of the aluminium
hydroxide. Thus, with carbonic acid

In the case of stronger acids, such as HC1, the same action
takes place, but any excess of the acid beyond that required to

these salts are composed of a molecule of each of the two sulphates, together
with twenty-four molecules of water. With compounds of this description
there, is unfortunately, a tendency in certain quarters to add the formulas of
the two salts together, and then to divide the numerals by their greatest com-
mon measure ; thus, K 2 SO 4 ,A1 2 (SO 4 )3, 24^0 = K 2 A1 2 (SO 4 ) 4 ,24H 2 O, which,
divided by 2= KAl)SO 4 ) 2 ,i2H 2 O. Presumably this plan is adopted with
a view to simplification, but as it obscures the origin and the nature of the
compounds, and as there is not the smallest evidence that such formulas are
more exact representations of the molecular constitution of the compounds,
their use is greatly to be deprecated.

* Three hydrated oxides of aluminium are known, obtainable by precipi-
tation under different circumstances, namely, A1 2 O 3 ,3H 2 O, A1 2 O3,2H 2 O,
and A1 2 O 3 ,H 2 O; these may also be formulated A1 2 (HO) S , A1 2 O(HO) 4 , and
A1 2 O 2 (HO) 2 respectively. The first of these, A1 2 (HO) 6 is sometimes written
A1(HO) 3 . (See previous note. )

Group III. Division A. 39

combine with the sodium of the aluminate at once re-dissolves the
A1 2 (HO) 6 . When, therefore, this acid is used, a slight excess is
added, and the aluminium hydroxide is re-precipitated by means
of ammonia.

Sodium and potassium aluminates are also decomposed by am-
monium chloride, with the precipitation of aluminium hydroxide ; *
the precipitation is complete on boiling. The compound thrown
down under these circumstances consists mainly of the di-hydrated
oxide, A1 2 O 3 ,2H 2 O ; thus

A1 2 O 3 , Na 2 O + 6NH 4 C1 = 6NaCl + 6NH 3 + H 2 O + ALO 3 , 2H 2 O

BaCO suspended in water, precipitates A1 2 (HO) 6 , carbon di-
oxide being evolved. The precipitation is complete even in the
cold.f If alum or aluminium sulphate is used, the precipitate is
mixed with insoluble barium sulphate

A1 2 (SO 4 ) 3 + 3BaCO 3 + 3H 2 O = A1 2 (HO) 6 + 3BaSO 4 + 3CO 2 %

K 2 CO 3 and Na^CO. precipitate an uncertain mixture of the
hydroxide and basic carbonates.

(NH ( ) ,S precipitates aluminium hydroxide, with evolution of
sulphuretted hydrogen (compare Fe).
A1 2 (S0 4 ) 3 + 3(NH 4 ) 2 S + 6H 2 = A1 2 (HO) 6 + 3 (NH 4 ) 2 SO 4 + 3 H 2 S I

[Aluminium forms no sulphide in the wet way. A1 2 S 3 (obtained
by the union of Al and S) is decomposed instantly by water,
forming the trioxide, and evolving H 2 S.]
Chromium, Cr.

DRY REACTIONS. Chromium compounds impart to a borax
bead a grass-green colour, when heated either in the outer or inner
blowpipe flame.

* The determining cause of this action of ammonium chloride is doubtless
the instability of ammonium aluminate, and the readiness with which it under-
goes hydrolysis. The ionised sodium aluminate and ammonium chloride may
be regarded as first undergoing "double decomposition " forming ammonium
aluminate which is immediately hydrolysed, thus

6Na , 2AK) 3 + 6NH 4 , 6C1 = 6Na , 6cT+ 6N H 4 , 2A1O 3
6NH 4 , 2A1O 3 = 6NH 3 + H 2 O + H 4 A1 2 O 5 (or A1 2 O 3 , 2H 2 O)

f In the presence of certain organic acids, as oxalic, tartaric, or citric acids,
aluminium hydroxide is only more or less imperfectly precipitated by the
above-mentioned reagents, owing to the formation of soluble double salts
of the organic acid with aluminium and the alkali metal ; such, for example,
as the double tartrate of aluminium and sodium, Na 2 (C 4 H 4 O 6 ),Al 2 (C4H 4 O 6 ) 3 .
This applies also in the case of the corresponding chromium and iron com-
pounds.

J In these equations simple aluminium sulphate is given instead of alum,
in order not to unnecessarily load the equation with materials taking no part
in the reaction.

40 Qualitative Analysis.

When fused in a platinum capsule with five or six times their
weight of a mixture consisting of I part of KNO 3 and 2 parts of
dry Na 2 CO 3 or K 2 CO 3 (or I part of KC1O 3 with 6 parts of Na 2 CO 3 ),
chromium compounds are converted into alkaline chromates, which
appear as a yellow mass, soluble in -water to a yellow solution.
In the case of chromic oxide, for instance, Cr 2 O 3 , the reaction is the
following:

Cr 2 O 3 + 2K 2 CO 3 + KC1O 3 = 2K 2 CrO 4 + KC1 + 2CO 2

The chief natural source of chromium is the mineral chrome iron
ore, Cr 2 O 3 ,FeO. When this is fused with either of the above
mixtures, the same reaction takes place as regards the chromium,
while the iron is changed to Fe 2 O 3 ; thus

6Cr 2 O 3 , FeO + I2K 2 CO 3 + 7KC1O S = i2K 2 CrO 4 + 3Fe 2 O 3 +

yKCl + i2CO 2

[Sodium peroxide, Na 2 O 2 , may be substituted as the oxidizing
material, in which case the fusion should he carried out in a silver
capsule.]

WET REACTIONS. The two best-known classes of chromium
salts are derived from the two oxides, namely

Chromium sesquioxide (or chromic oxide), Cr 2 O 3
Chromium trioxide (chromic anhydride), CrO 3

Chromic oxide, Cr 2 O 3 , is basic, uniting with acids to form the
chromic salts, such as chromic hydroxide, Cr 2 (HO) 6 or Cr 2 O 3 ,3H 2 O ;
chromic chloride, CrCl 3 ; chromic sulphate, Cr 2 (SO 4 ) 3 ; double potas-
sium and chromium sulphate (chrome alum}, K 2 SO 4 ,Cr 2 (SO 4 ) 3 ,
24H 2 O. Of these salts, the hydroxide alone is insoluble in water.

Chromium trioxide, CrO 3 , is the anhydride of the hypo-
thetical chromic acid,* H 2 CrO 4 , which gives rise to salts known as
chromates, analogous in constitution to the sulphates.

Chromates of the metals of Groups IV. and V. are all soluble in
water, except BaCrO 4 . The other chromates are insoluble.

a. Chromic Salts. These salts are mostly of a purplish
or violet-grey colour when solid, giving either a purple or green
solution when dissolved, the colour depending upon the conditions
of solution. Thus chrome alum dissolved in cold water gives a

* At first it may confuse students to find that chromzV anhydride, and
chromzc acid with its salts, should not be in the class of chromic compounds.
It must be remembered that the classification is not based upon the nomen-
clature of the substances. "Chromic" compounds are those containing
chromium as the " base," or the positive radical ; while in those compounds
derived from CrO 3 the element is in the "acidic " or negative group. They
may, therefore, be conveniently distinguished as " chromic acid " compounds.

Group III. Division A. 41

purple solution, which on boiling turns green,* and on long standing
again becomes purple.

NH,HO produces a bluish or greenish-grey precipitate of
chromic hydroxide, Cr 2 (HO) 6 t, partially dissolved by excess of
ammonia in the cold, giving a lilac-coloured liquid, but completely
precipitated on gently boiling. Cr 2 (HO) 6 is readily soluble in acids.!

KHO and NaHO precipitate Cr 2 (HO) 6 , readily soluble in
excess, giving a deep green solution. Reprecipitated by neutraliza-
tion with HC1, and by boiling with NH<C1, as in the case of Al.

BaCO 3 precipitates a mixture of the hydroxide and basic
carbonate. Complete precipitation only after some hours.

KjCO, and Na 2 CO 3 give a similar precipitate, the composition
of which varies with the conditions of precipitation.

(NH 4 ),,S precipitates Cr 2 (HO) 8 . Precipitation complete. [Cr,
like Al, is incapable of forming a sulphide in the wet way.]

Oxidation of Chromic Compounds. By means of suitable
oxidising agents, chromic compounds are readily converted into
compounds of chromic acid, the mechanism of the change in all
cases being the oxidation of the sesquioxide into the trioxide ; thus

Cr 2 O 3 + 30 = 2CrO 3

One method, namely, by fusion with oxidising agents, has been
explained under Dry reactions. The Cr 2 O 3 in that instance is
oxidised into the potassium salt of chromic acid. The oxidation
may be accomplished in the wet way by the following reactions :

(1) Boiling chromic hydroxide with potassium hydroxide and
manganese dioxide

Cr 2 (HO) 6 + 4KHO + 6MnO 2 = 2K 2 CrO 4 + 3Mn 2 O 3 + 5H 2 O

(2) By substituting lead peroxide for manganese dioxide
Cr 2 (HO) 8 + 4KHO + 3PbO 2 = 2K 2 CrO 4 + 3?bO + 5H 2 O

In this case secondary reactions take place, for PbO is soluble
in KHO, and the solution so formed then reacts upon the K 2 CrO 4 ,
producing PbCrO 4 , which also dissolves in KHO.

* The green colour is said to be due to the formation of a basic salt, by the
action of water upon the normal compound.

f The composition of the precipitate depends on the conductions under
which it is formed. There are several hydrated chromic oxides ; compare
also Al.

J For the influence exerted by the presence of organic acids, see footnote,

P- 39-

The soluble compounds produced are similar to those given by
aluminium ; the potassium salt has the composition Cr 2 O 3 ,K 2 O or K 2 Cr 2 O 4 .
They are known as chromites. Chrome iron ore is ferrous chromite, Cr 2 O 3> FeO.
Although analogous to the allumina/ifj-, they must not be called chromafes, as
this name is reserved for the salts of chromic acid.

4-2 Qualitative Analysis.

(3) By the action of hypochlorites (or hypobromites) in the
presence of caustic alkalies, either employed as such, or formed in
the solution by the use of chlorine or bromine in the presence of
the caustic alkali

Cr 2 (HO) 6 + 4KHO + 3KC1O = 3KC1 + 2K 2 CrO 4 + 5H 2 O

(4) By the action of sodium peroxide. If a small quantity of
Na 2 O 2 be added to chromium hydroxide suspended in water, and
the mixture gently warmed, the chromium compound is immediately
converted into the yellow sodium chromate ; thus

Cr 2 (HO) 6 + 3Na 2 O 2 = 2Na 2 CrO 4 + 2NaHO + 2H 2 O

/3. Chromic Acid and Chromates. The acid, H 2 CrO 4 , has
aever been isolated. The anhydride, CrO 3 , is readily obtained by
adding strong H 2 SO 4 to a cold strong solution of potassium di-
chromate, when the oxide is deposited in the form of red silky needles.
It forms two classes of salts, viz. the normal chromates, of which
K 2 CrO 4 is a type ; and the dichromates,* of which K 2 Cr 2 O 7 is a
familiar example. The chromates are mostly yellow or red in
colour, and those which are soluble in water (see p. 40) impart a
yellow or orange colour to the liquid. The most important of the
insoluble chromates made use of in analysis, and which are all
precipitated by the addition of potassium chromate to solutions of
the metallic salts, are the following :

Barium chromate, BaCrO 4 (see Ba reactions, p. 31).

Lead chromate, PbCrO 4 (see Pb reactions, p. 81). PbCrO 4
melts without decomposition, and solidifies on cooling to a brown
crystalline mass. At higher temperatures it gives off oxygen

2PbCrO 4 = Cr 2 O 3 + 2PbO + 30

PbCrO 4 (known as chrome yellow), when digested with NaHO, or
with K 2 CrO 4 , is converted into a red basic lead chromate (known
as chrome red)

2PbCrO 4 + 2NaHO = Na2CrO 4 + H 2 O + Pb 2 CrO 6 (or PbCrO 4 ,PbO)
2PbCr0 4 + K 2 Cr0 4 = PbCrO 4 ,PbO + K 2 CrO 4 ,CrO 3

* The constitution of the dichromates (sometimes wrongly called fo'chro-
mates) may be expressed thus, K 2 CrO 4) CrO 3 . They are strictly analogous
to the pyrosulphates, K 2 S 2 O 7 , or K 2 SO 4 ,SO 3 , and on this account should
consistently be named pyrochromates. By the action of strong acids, the
normal potassium chromate is converted into the dichromate ; thus, 2K 2 CrO 4
+ H 2 Sp 4 = K 2 SO 4 + H 2 O + K 2 CrO 4 ,CrO 3 . And the dichromate is re-con-
verted into the normal salt by the action of potash

K 4 CrO 4 ,CrO 3 + 2KHO = 2K 2 CrO 4 + H 2 O

Group III. Division A. 43

Silver chromate, Ag 2 CrO 4 . A dark chocolate-red precipitate,
soluble in ammonia and nitric acid.

Mercurous chromate (basic), Hg 2 CrO 4 ,Hg 2 O. A brick-
red precipitate, which, when dried, and heated in a tube, gives a
mercury sublimate, evolves oxygen, and leaves a residue of Cr 2 O 3 .

Oxidation of Chromic Acid. Although CrO 3 is such a
highly oxygenated compound, it appears to be capable of still
further oxidation by hydrogen peroxide, giving rise to a compound
which is believed by some to be perchromic acid, HCrO 4 , or
2CrO 3 ,H 2 O 2 , and by others to be a compound of CrO 3 and H 2 O 2
in undetermined proportions. The interest of the compound lies
in the fact that it has an intense azure-blue colour, and its formation
affords an extremely delicate test for either chromic acid or
hydrogen peroxide. A few drops of H 2 O 2 (or a few particles of
Na 2 O 2 ) are added to half a test-tube of water, and the mixture
acidified with one or two drops of HC1. A single drop of potassium
dichromate solution added to this produces an intense blue colour.
[The compound is very unstable in aqueous solution, but less so in
ether ; therefore, in testing for very minute quantities, ether should
be added before the dichromate ; and on shaking the mixture,
the ethereal layer which rises to the surface will be coloured blue.]

Reduction of Chromic Acid. CrO 3 is a powerful oxidising
agent, giving up oxygen to oxidisable substances, and being itself
reduced to Cr 2 O 3 ; that is, to the condition of a " chromic " com-
pound. Thus, by sulphur dioxide it is reduced to chromium
sulphate

2CrO 3 + 3SO 2 = Cr 2 (SO 4 ) 3

The same action takes place in an acidified solution of potassium
dichromate

K 2 Cr 2 O 7 + H 2 SO 4 + 3H 2 SO 3 = Cr 2 (SO 4 ) s + K 2 SO 4 + 4H 2 O

Similarly, chromic acid and chromates are. reduced by HC1, oxidis-
ing the hydrogen of the acid, and liberating chlorine, after the
manner of peroxides ; thus

CrO 3 + 6HC1 = CrCl 3 + 3H 2 O + 3d
K 2 Cr 2 O 7 + I4.HC1 = 2CrCl 3 + 2KC1 + ?H 2 O + 3C1 2

On account of this reaction, a mixture of potassium dichromate
and hydrochloric acid is capable of " oxidising " FeCl 2 into FeCl s ;
SnCl 2 into SnCl 4 ; As 2 O 3 into As 2 O 5 . In all cases of oxidation by
chromic acid, the reduction of the chromic acid compound to the
state of a " chromic " compound is evidenced by the change of

44 Qualitative Analysis.

colour from the yellow or orange of the former, to the green colour
of the latter. This reduction and change of colour is at once seen
by passing sulphuretted hydrogen through acidified potassium
dichromate

K 2 Cr 2 O 7 + sH 2 S + 8HC1 = 2CrCl 3 + 2KC1 + 7H,O + 38

Many organic substances also reduce chromic acid, such as
oxalic acid, and alcohol. Thus, one molecule of oxalic acid, C 2 H 2 O 4 ,
requires one atom of O to convert it into CO 2 and H 2 O

C 2 H 2 O 4 + O = 2CO 2 + H 2 O

Potassium dichromate, in being reduced, has three available atoms
of oxygen to give up ; thus, K 2 Cr 2 O 7 - Cr 2 O 3 ,K 2 O,3O (in the
presence of dilute acids the Cr 2 O 3 and, K 2 form salts). Therefore
one molecule of K 2 Cr 2 O 7 can oxidise three molecules of oxalic acid,
resulting in the evolution of six molecules of CO 2 ; thus

K 2 Cr 7 + 4H 2 S0 4 + 3C 2 H 2 O 4 = K 2 SO 4 + Cr 2 (SO 4 ) 3 + 7H 2 O

+ 6C0 2

If alcohol be added to a mixture of potassium dichromate and
sulphuric acid, the alcohol (C 2 H 6 0) is oxidised first to aldehyde,
C 2 H 4 O, and then to acetic acid (C 2 H 4 O 2 ), and the colour of the
mixture changes from orange-red to green.

Iron, Fe.

DRY REACTIONS. Iron compounds impart to a borax bead
heated in the outer flame, a colour which appears chocolate when
hot, and yellow when cold. After heating in the reducing flame,
the colour changes to a bottle-green (the green colour of common
bottle glass is caused by the presence of iron). When heated on
charcoal with Na-jCOg in the inner blowpipe flame, iron com-
pounds become reduced, and a dark grey magnetic mass is obtained.
If this be washed with water in a small mortar, and the end of
a magnet applied, it will be attracted after the manner of iron
filings.

WET REACTIONS. The salts of iron are derived from the two
oxides FeO and Fe 2 O 3 .* They are both basic oxides, and give
rise to two classes of salts, namely, ferrous and ferric respectively.
Ferrous salts readily take up oxygen, and become converted into

* The oxide known as magnetic oxide of iron, or ferroso-ferric oxide,
Fe 3 O 4 or Fe 2 O 3 ,FeO, yields a mixture of ferric and ferrous salts.

. Group III. Division A. 45

ferric compounds ; while the latter, under the influence of suitable
reducing agents, easily pass back again to the ferrous condition.

(a) Ferric Compounds. The common ferric salts that are
soluble in water are the chloride, FeCl 3 ; nitrate, Fe 2 (NO 3 ) 6 , and
sulphate, Fe 2 (SO 4 ) 3 . These all give yellowish-brown solutions.

NH 4 HO, KHO, and NaHO throw down a brown voluminous
precipitate of ferric hydroxide,* Fe 2 (HO) 6 , insoluble in excess,
or in NH 4 Cl.t

K 2 CO 3 , Na 2 CO 3 , and BaCO 3 give the same precipitate, CO 2
being liberated

2FeCl 3 + 3Na 2 CO 3 + 3H 2 O = Fe 2 (HO) 6 + 6NaCl + 3CO 2
The precipitate is soluble in a concentrated solution of K 2 CO 3 ,
giving a deep reddish solution of unknown composition. On the
addition of water the hydroxide is reprecipitated. (With BaCO 3
basic carbonates are also precipitated.)

(NHJoS produces a black precipitate of ferrous sulphide.
The action may be considered as taking place in two stages :
(i) the reduction of the iron to the ferrous- state, and (2) the
formation of the ferrous sulphide ; thus, using a dissected formula
for ammonium sulphide

(i) :NH 3 NH 3 : H 2 S + 2FeCl 3 = 2FeCl 2 + 2HC1 + S
'(2) (NH 4 ) 2 S + FeCl 2 = FeS + 2NH 4 C1

In the first equation the hydrochloric acid formed unites with
the ammonia, producing 2NH 4 C1.

Sulphuretted hydrogen, H 2 S, brings about the first stage
in the above action, reducing the iron from the ferric to the ferrous
state with precipitation of sulphur, but in the presence of the free
acid which is developed by the action, ferrous sulphide cannot be
formed. [Ferric sulphide cannot be produced in the wet way.']

Potassium ferrocyanide, K 4 Fe(CN) 6 , or K 4 FeCy 6 ,J pro-
duces with ferric salts a dark blue precipitate {Prussian blue}

3 K 4 (FeCy 6 ) + 4 FeCl 3 = I2KC1 + Fe 4 (FeCy 6 ) 3 '

* Several hydrated ferric oxides are known, e.g. Fe 2 O3,3H 2 O ; Fe 2 O3,2H 2 O;
Fe 2 O3,H 2 O. The composition of the precipitate produced by alkalies depends
upon the conditions of precipitation,

f See footnote on p. 39 as to the influence of organic compounds.

J ''Cy" is a recognised and convenient symbol for the radical (CN) ;
cyanogen. The use of this reagent as a test for iron is unique, as being the
only case in which the reagent is itself a compound containing the very metal
it is employed to detect. The ferrocyanides and the ferricyanides, however,
although compounds of iron, do not yield on solution either ferrous or ferric

ions, but the complex anion Fe(CN) 6 . These give no reaction with the

4-6 Qualitative Analysis.

This test is extremely delicate, but where the amount of iron
is very small, a blue or greenish coloration only is produced.
" Prussian blue " is insoluble in hydrochloric acid, but readily
dissolves in oxalic acid. It is decomposed by NaHO or KHO,
with precipitation of ferric hydroxide

Fe 4 (FeCy 6 ) 3 + i 2 KHO = 2Fe 2 (HO) 6 + 3 K 4 (FeCy 6 )

Potassium ferricyanide, K 3 (FeCy 6 ), gives no precipitate with
ferric salts.

Potassium thiocyanate, K(CN)S, produces with ferric
salts a rich wine-red coloration, owing to the formation of ferric
thiocyanate, Fe(CNS) 3 , which is soluble in water. The colour of
this compound is very intense, hence the reaction may be employed
to detect very small quantities of iron.*

Reduction of Ferric to Ferrous Compounds. FerrzV
compounds are readily reduced to the ferrous state ; they are
therefore oxidising agents of some importance. The action of
(NH 4 ) 2 S and of H 2 S has been already mentioned. Nascent
hydrogen reduces tHem in the same way ; therefore, when metallic
iron is dissolved in HC1 or H 2 SO 4 , the salts produced are ferrous
chloride and sulphate respectively. Nitric acid, on the other
hand, converts the iron into the " ferric " state.

A ferric salt already in solution is reduced by nascent hydrogen,
generated by introducing zinc into the acidified liquid.

In passing from FeCl 3 to FeCl 2 , one atom of chlorine is
available for oxidising purposes, and is capable of bringing about
such actions as the following

The " oxidation " of stannous chloride, SnCl 2 , to stannic
chloride, SnCl 4 .

The oxidation of sulphurous acid or thiosulphuric acid into
sulphuric acid ; thus

2FeCl 3 + H 2 SO 3 + H 2 O = H 2 SO 4 + 2HC1 + 2FeCl 2
2FeCl s + Na 2 SSO 3 + H 2 O = Na^O, + 2HC1 + 2FeCl2 + S

reagents employed for detecting either ferrous or ferric ions, and therefore
before the iron in such compounds will give any of the ordinary reactions, its
union with the cyanogen radical must be first destroyed (see Cyanides, p. 163).
* This is a "reversible" reaction, and therefore, when equilibrium is
established, there will be present in the liquid both ferric chloride and potas-
sium thiocyanate, thus, 3KCNS + FeClj J 3KC1 + Fe(CNS) 3 . That this is
so may be proved by the following experiment : Add to a little moderately
dilute ferric chloride a small quantity of potassium thiocyanate ; then dilute
the liquid with water so that the intensity of the red colour is greatly reduced,
and divide it into two portions. To one add more ferric chloride, and to the
other add more potassium thiocyanate. In each case the liquid becomes a
deeper red colour.

Group III. Division A. ' -47

(b] Perrons Compounds. Ferrous salts are usually pale green
when crystallised, and white when anhydrous.* Of the common salts
"the chloride and sulphate afe soluble. The latter readily forms
. double salts with the sulphates of the alkalies (such as ferrous ammo-
nium sulphate, FeSo 4 ,(NH 4 ) 2 SO 4 ,6H 2 O), which are also soluble in
water, and are less readily oxidised on exposure to the air than ferrous
sulphate, which they otherwise closely resemble in appearance.

NH,HO. KHO, and NaHO produce a precipitate of ferrous
hydroxide, Fe(HO) 2 , which is at first a dirty white colour, but which
rapidly turns first pale greenish-grey, then a dirty grey, and finally
brown, owing to its oxidation by atmospheric oxygen. The presence
of ammonium salts renders the precipitation incomplete. The
precipitate is not soluble in excess of the reagents ; boiling with
KHO turns it black, converting it into Fe 3 O 4 .

K 2 CO 3 and Na 2 CO 3 give a white precipitate of ferrous carbonate,
FeCO 3 , which on exposure to the air quickly absorbs oxygen.

(NH,),S precipitates black ferrous sulphide, FeS. Readily
soluble in acids, with evolution of sulphuretted hydrogen ; insoluble
in alkalies. The precipitate in the moist state is oxidised on exposure
to the air into ferrous and basic ferric sulphate.

X 4 (FeCy 6 ) precipitates potassium ferrous ferrocyanide,
FeK 2 (FeCy 6 ), thus

K 4 (FeC ya ) + FeCl 2 = 2KC1 + FeK 2 (FeCy 6 )
When the solutions are mixed in test-tubes in the ordinary way,
the precipitate has a greenish-blue colour ; but when the reaction
is made in an atmosphere free from oxygen, and the solutions are
previously boiled so as to entirely expel all dissolved oxygen, the
precipitate is perfectly white. It rapidly absorbs oxygen and
becomes blue, and is also easily oxidised to " Prussian " blue by
nitric acid or chlorine ; thus

4FeK 2 (FeCy 6 ) + 2 C1 2 = Fe 4 (FeCy 6 ) 3 + K 4 FeCy 6 + 4KC1

Potassium ferricyanide, K 3 (FeCy 6 ), gives, with ferrous
salts, a precipitate of ferrous ferricyanide, Fe 3 (FeCy 6 ) 2 (known as
TurnbulVs blue'}, which is indistinguishable by its appearance from
Prussian blue

2 K 3 (FeCy 6 ) + 3 FeCl 2 = Fe 3 (FeCy 6 ) 2 + 6KC1

The precipitate is insoluble in hydrochloric acid, but is decom-
posed by caustic alkalies, with the precipitation of ferrous hydroxide ;
thus

Fe 3 (FeCy fl ) 2 + 6KHO = 2K 3 (FeCy 6 )

48 ' Qualitative Analysis,

Oxidation of Ferrous to Ferric Compounds. The ferric
salts being the more stable, the ferrous compounds undergo oxida-
tion even more readily than the ferric' salts become reduced. Mere "
exposure to the air in many cases causes the change. In analysis
the oxidation is usually accomplished either by chlorine (or bromine)
or by nitric acid.

The chlorine may be employed in the form of its aqueous
solution (chlorine water), or more conveniently by generating the
gas in contact with the ferrous compound by means of hydrochloric
acid and potassium chlorate. The solution of the ferrous salt is
acidified with concentrated HC1, and heated. A few particles of
potassium chlorate are then dropped into the mixture, and the
heating continued for a short time.

A mixture of HC1 and KC1O 3 evolves both chlorine and chlorine
peroxide ; thus

4KC1O 3 + I2HC1 = 6H 2 O + 4KC1 + 3C1O 2 + gC\

Both the free chlorine and the chlorine of the chlorine peroxide
are available for oxidising the ferrous compound ; hence the equation
may be simplified as follows :

KC10 3 + 6HC1 + 6FeCl 2 = 6FeCl s + KC1 + sH 2 O

When the oxidation is accomplished with nitric acid, the strong
acid is added, a few drops at a time, to the hot acidulated solution
of the ferrous salt. The solution becomes dark in colour, and nitric
oxide is disengaged ; thus

6FeS0 4 + 3H 2 S0 4 + 2HNO 3 = 3Fe 2 (S0 4 ) 3 + 4H. 2 O + 2NO
3FeCl 2 + 3HC1 + HNO 3 = 3FeCl 3 + 2H 2 O + NO

Unless the solution of the ferrous salt is acidified, a portion ot
the iron is converted into Fe 2 O 3 , which is taken up, in the case
of the sulphate, by the ferric sulphate, forming insoluble basic ferric
sulphates, Fe 2 (SO 4 ) 3 , rFe 2 O s .

SEPARATION OF THE METALS OF GROUP II I A.

The separation of the metals of this subdivision from the other
metals of Group III., and also from those of Groups IV. and V., is
based upon the fact that their hydrated sesquioxides are precipi-
tated by ammonia in the presence of ammonium chloride.*

* The separation of Group IIlA. from Group IIlB. by means of NH 4 HO is
not sharp and complete in all cases (see Manganese reactions).

Separation of the Metals of Group III A 49

The separation of the three metals of this group from each
other is based upon

1. The oxidation of chromic oxide to chromic acid ; and

2. The solubility of aluminium hydroxide in caustic alkalies.
To the solution add NH 4 C1 in considerable quantity ; heat the

mixture to boiling, and add NH 4 HO carefully until precipitation is
complete. Bring the liquid once more " to the boil," when, if suffi-
cient ammonia has been added, the steam will smell of it. Filter
the mixture while hot.*

The precipitate consists of A1 2 (HO) 6 , Cr 2 (HO) 8 , and Fe 2 (HO) 4 . Wash
the precipitate, and transfer it (or a portion of it) to a test-tube with
a small quantity of water. Add to the mixture a little sodium per-
oxide, and boil for a moment, until the temporary effervescence ceases.
The chromium is oxidised to chromate, and the A1 2 (HO) 6 dissolves
in the NaHO, which is formed by the action of the sodium peroxide
upon the water. Filter.

The filtrate contains sodium chromate,Na 2 CrO 4 ,
and sodium aluminate, Al 2 O 3 ,3Na 2 O. The
former shows itself by the yellow colour.
Divide into two portions

(1) Acidify with acetic acid, and confirm
chromium by special reactions, e.g. lead
acetate.

(2) Acidify with dilute nitric acid, and add
NH 4 HO. A white precipitate of A1 2 (HO; 6
confirms aluminium.

The residue consists of
Fe 2 (HO) 6 . Dissolve
in a little hot dilute
HC1, and confirm
iron by special re-
actions, e.g. K 4 FeCy s
or KCNS.f

The following alternative methods of separation may also be
used.

(a) The precipitated hydroxides are washed and dried. The
residue is then mixed with at least six times its weight of fusion
mixture^, and -fused in a platinum capsule. In this way the
chromium is converted into alkaline chromate ; a variable pro-
portion of the aluminium into aluminates.

* In the regular course of a complete analysis, the filtrate obtained here
will contain the metals of Group IIlB., IV., and V.

f At this stage in the process, the iron will be in the "ferric" condition.
To ascertain whether it was originally present as a "ferrous" or "ferric"
compound, separate tests must be made in the solution before it has been
subjected to the action of either reducing or oxidising agents.

I Fusion mixture is a mixture of Na 2 CO 3 and K 2 CO 3 in equivalent propor-
tions (or about 10 parts Na 2 CO 3 to 13 of K 2 CO 3 ). It is used in preference to
Na 2 COa alone, because it has the property of melting more easily than either
carbonate separately.

5O Qualitative Analysis.

The fused mass is then dissolved in water, and filtered. The
nitrate is tested for aluminium and chromium, while the residue is-
dissolved and tested for iron, as in the foregoing scheme.

(b) The precipitated hydroxides are dissolved in a little warm
dilute HC1, and pure NaHO * added in quantity considerably
more than sufficient to produce precipitation. The mixture is then
boiled for a few minutes, and filtered.

The filtrate contains sodium aluminate, AUO 3 ,3Na. 2 O. Add
dilute HC1 until just acid, and reprecipitate A1 2 (HO) 6 with
ammonia.

The precipitate contains Cr 2 (HO) 6 and Fe 2 (HO) 6 . This is
dried, and fused with fusion mixture. The fused mass is dissolved
in water and filtered. The solution contains sodium chromate,
while the Fe 2 O 3 remains on the filter. These are confirmed as
in the above methods.

* The commercial caustic soda usually employed in the laboratory always
contains more or less sodium aluminate. The student should test a sample of
the reagent by neutralising it with HC1, and then adding NH 4 HO. In the
method of separation given above, this difficulty is avoided, as the sodium
peroxide of commerce is usually free from this impurity.

CHAPTER VI.
REACTIONS OF THE METALS OF GROUP III. DIVISION B,

Manganese, Mn.

DRY REACTIONS. Manganese compounds, when heated in a
borax bead in the oxidising flame, impart to the bead a violet or
lilac colour. When heated in the reducing flame, the bead again
becomes colourless.

A more characteristic reaction is based upon the oxidation of
manganese to manganic acid. When a manganese compound is
fused with KHO, or with Na 2 CO s and a little KNO 3 or KC1O 3
upon a platinum capsule, the manganese undergoes oxidation, and
a deep green-coloured mass is obtained, consisting of manganates.
of the alkali metals

MnO 2 + Na 2 CO 3 + O (from KC1O S or KNO 3 ) = Na 2 MnO 4 + CO 2 .
MnO 2 + 2KHO + O = K 2 MnO 4 + H 2 O

The green mass (especially when obtained by fusion with KHO)
dissolves in a small quantity of cold water to a deep green solution.
When this is either acidified, or warmed, or even largely diluted
with water, its colour changes from green to pink, owing to the
conversion of the manganate into permanganate ; thus