UC-NRLF
SB 27fl Db7
ALLER
AL ANALYSIS
NEWTtt
Edmund O'Heill
SMALLER CHEMICAL ANALYSIS
BY THE SAME AUTHOR
CHEMICAL LECTURE EXPERI-
MENTS. With 230 Illustrations. Crown
8vo, 6s.
A TEXT-BOOK OF INORGANIC
CHEMISTRY. With 155 Illustrations. Crown
8vo, 6s. 6d.
A MANUAL OF CHEMICAL ANALY-
SIS, QUANTITATIVE AND QUALITA-
TIVE. With 100 Illustrations. Crown 8vo,
6s. 6d.
ELEMENTARY PRACTICAL CHE-
MISTRY. With 108 Illustrations and 254
Experiments. Crown 8vo, 2f. 6d.
LONGMANS, GREEN, AND CO.
LONDON, NEW YORK, AND BOMBAY
SMALLER
CHEMICAL ANALYSIS
BY
G. S. NEWTH, F.I.C, F.C.S.
it
DEMONSTRATOR IN THE ROYAL COLLEGE OF SCIENCE, LONDON, ASSISTANT-
EXAMINER IN CHEMISTRY, BOARD OF EDUCATION
LONGMANS, GREEN, AND CO.
39 PATERNOSTER ROW, LONDON
NEW YORK AND BOMBAY
1906
All rights reserved
Nt
JN MEMORJAM
PREFACE
THIS little book is practically an abridged edition of the
qualitative section of my Manual of Chemical Analysis, and is
designed for the use of students who are taking a less advanced
stage than those for whom the larger book is intended.
I have included in this book, however, a chapter on such
simple volumetric processes as it is customary to introduce into
moderately elementary courses of Practical Chemistry.
G. S. N.
ROYAL COLLEGE OF SCIENCE, LONDON,
October, 1906.
889771
TABLE OF CONTENTS
CHAPTER PAGE
I. PRELIMINARY EXERCISES I
II. ANALYTICAL CLASSIFICATION 16
III. REACTIONS OF THE METALS OF GROUP V 21
IV. REACTIONS OF THE METALS OF GROUP IV .... 28
V. REACTIONS OF THE METALS OF GROUP IIlA .... 32
VI. REACTIONS OF THE METALS OF GROUP Ills .... 43
VII. REACTIONS OF THE METALS OF GROUP II .... 54
VIII. REACTIONS OF THE METALS OF GROUP I 77
IX. THE NEGATIVE OR ACID RADICALS 83
X. PRELIMINARY EXAMINATION FOR METALLIC RADICALS . 107
XI. PRELIMINARY EXAMINATION FOR ACID RADICALS . .114
XII. SIMPLE VOLUMETRIC DETERMINATIONS 119
Involving the Use of Standard Acids and Alkalies, Potas-
sium Permanganate, Potassium Bichromate, Iodine,
Sodium Thiosulphate, and Silver Nitrate.
INDEX 145
SMALLER
CHEMICAL 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 practi-
cally familiar with certain simple op^ja/K^s^q^^ijajiipulations
which he will constantly be required 'to carry' out* in tne course
of his work, and upon the dexterous and cleanly performance
of which much of his success 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 (6o*degrees) should be selected, or another
filter-paper folded so that the cone shall be of the same angle as
the funnel. The funnel is supported by a metal or wooden stand.
Now place some diluted hydrochloric acid in a small beaker,
B
2 Smaller Chemical Analysis
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
likely to splash over the funnel. To prevent both of these acci-
dents, the liquid should be poured down against a glass rod held
FIG. i.
FJG. 2.
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
Preliminary Exercises
filter and its contents must be washed with distilled water. 1
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 espe-
cially important when the
material retained 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
FIG. 3.
quickly and effectually accomplished than if an unduly large
filter is employed. 4
2. Solution. This term is applied both to the act of dis-
solving and to the product obtained by dissolving.
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 potassium carbo-
nate. The water is called the solvent. The process of solution is
1 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.
4 Smaller Chemical Analysis
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
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 spontaneo^ls evaporation.
EXERCISE 3. Pour the two solu-
tions obtained in Exercise 2 into sepa-
rate 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
dryness. As the condition of dryness is
FIG approached, 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 examined by one simple test,
which will prove that the one from the watery solution is the same
as it was before being dissolved, and that the other is quite different.
Add to each a few drops of dilute nitric acid : the first dissolves
Preliminary Exercises 5
with effervescence, as did the original potassium carbonate ; the
other is unacted on by the acid.
Sometimes it is necessary to carry on the operation of
evaporation 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 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 substances 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 potasJi) 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 once 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 evaporate
to dryness on a steam-bath.
Treat another similar quantity of the barium sulphate with
dilute hydrochloric acid, and evaporate a few drops in the same
Smaller Chemical Analysis
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 or nickel cru-
cible, 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 re-
mains 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 hydrochloric
acid upon the residue on
FlG . 5> 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 been all removed ; neither can it
be barium sulphate, for that compound 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 dissolves in dilute
Preliminary Exercises 7
hydrochloric acid, forming barium chloride (soluble) and carbon
dioxide, which escapes as gas. Therefore, by fusion the insoluble
bariiim sulphate is converted into soluble barium carbonate.
5. Precipitation. When chemical action takes place
between substances in solution, and one of the products of the
action is insoluble, the latter substance 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 together.
The two compounds react upon each other, forming sodium nitrate
(soluble in water) and silver chloride (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 experiment :
EXERCISE 8. Filter the mixture obtained above, and divide the
filtrate 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, 1 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
1 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 Smaller Chemical Analysis
nearly as possible the exact point when it just ceases to produce any
further precipitate.
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 select the com-
pound 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 precipi-
tation. It is of the utmost importance, therefore, in analysis,
to be quite sure that precipitation is as complete as possible.
On the other hand, the reckless addition of precipitants 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 the liquid, as explained in Exercise i. A precipitate may
be removed from the filter either by means of a spatula (pre-
ferably 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. For
example
EXERCISE 9. Add a solution of sodium carbonate to a solution
of barium chloride, until precipitation is just complete. Barium
Preliminary Exercises 9
carbonate is precipitated, and sodium chloride remains in solution.
Pour the mixture upon three separate niters, and wash 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 care-
fully 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
by using a small platinum capsule or crucible in which to heat the
residue obtained by evaporating the solution to dryness.
EXERCISE n. Strongly heating in a tube closed at one end.
Place a minute quantity of mercuric oxide in a small test-tube
io Smaller Chemical Analysis
(4 inches x j^), 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, hold-
ing 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. 1
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. 2 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
1 The memory of the beginner may be aided by the alliteration, Center,
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 j whereas in
the outer flame there is an excess of heated atmospheric oxygen.
2 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.
Preliminary Exercises
II
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.
FIG. 6.
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 found 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
incrustation then becomes green. Notice that the incrustation is
12
Smaller Chemical Analysis
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.
EXERCISE 15. Twist the end of a piece of platinum wire into a
small round loop or eye, 1 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
z'nnerf[a.me, and see that in each case when
cold the blue colour remains the same.
1 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 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.
FIG. 7.
Preliminary Exercises 13
EXERCISE 16. Make another borax bead, and touch it with a
small quantity of solution of manganous sulphate. Heat this in
the outer blowpipe flame. After cooling, examine the colour care-
fully : 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. Man-
ganese compounds therefore 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 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 bhte 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 just 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.)
9. Oxidation and Reduction. Processes which convert
lower oxides (or compounds derived therefrom) into higher
oxides of the same element (or compounds derived from them)
are processes of oxidation. Thus, when either sulphur^j acid
14 Smaller Chemical Analysis
or a sulph//<? is converted into sulphur/V acid or a sulphate, the
processes is one of oxidation
H 2 SO 3 + O = H 2 SO 4
Similarly, when ferrous sulphate is changed to ferric sulphate,
we say that the iron in the ferrous salt has been oxidised
2 FeS0 4 + H 2 S0 4 + O = H 2 O + Fe 2 (SO 4 ) 3
We say exactly the same, also, when ferrous chloride passes
into ferric chloride, for although there is no oxygen in these
compounds, the change is brought about by means of oxygen
2FeCl 2 + 2HC1 + O = H 2 O + 2FeCl 3
Agents which bring about oxidation are termed oxidising agents,
and those most frequently employed in qualitative analysis are
nitric acid, sodium peroxide, and chlorine.
EXERCISE 18. Make a solution of ferrous chloride by dissolving
a few fragments of iron wire in strong hydrochloric acid in a test-
tube, gently boiling the acid. To this solution add a few drops of
nitric acid. Notice that the nearly colourless solution of ferrous
chloride changes to the yellow ferric chloride, two molecules of
nitric acid supplying three atoms of oxygen
2HNO 3 = H 2 O + 2NO + 30
EXERCISE 19. Dissolve a small crystal of chrome alum in cold
water (this compound is a salt derived from chromium sesquioxide,
Cr 2 O 3 ). To this purple solution add a little sodium peroxide, and
gently warm. The solution becomes bright yellow, owing to the
formation of sodium chromate, a compound derived from the higher
oxide, CrO 3 .
Reduc i.he reverse process, namely, the withdrawal of
oxygen, or ivalent negative constituent, from a compound.
Exercise 13 i example of reduction by heating on charcoal.
The commoi i educing agents employed in analysis for the
reduction of substances in solution are hydrogen, sulphurous
acid, and stannous chloride. Sulphuretted hydrogen, also, is
a powerful reducing agent, although it is not often employed
expressly for this purpose.
When metals such as iron, tin, or copper are dissolved
in hydrochloric acid, the nascent hydrogen by its reducing
Preliminary Exercises 15
influence prevents the formation of the higher chlorides of the
metals, and the compounds obtained are the ferrous, stannous,
and cuprous chlorides respectively. The reducing action of
nascent hydrogen may even go beyond the removal of oxygen
or its equivalent; it sometimes itself unites with the element
undergoing reduction. Nitric acid, for instance, may not only
be reduced to the lower oxides of nitrogen, but can be further
reduced to ammonia.
EXERCISE 20. Place a little granulated zinc in a small beaker,
cover it with water, and add a few drops of sulphuric acid so as to
produce a slow evolution of hydrogen ; now add a little very dilute
nitric acid, and notice that the evolution of hydrogen becomes
much slower with a little care it can be made to cease altogether.
The nascent hydrogen is attacking the nitric acid and reducing it
to ammonia, which unites with the sulphuric acid present, forming
ammonium sulphate. After about a quarter of an hour take a little
of the liquid in a test-tube, add sodium hydroxide, and warm ;
ammonia will be evolved, which may be detected by its smell or
its action on red litmus paper.
A similar action takes place when nascent hydrogen acts
upon arsenious oxide. This compound is reduced first to
elemental arsenic, and then still further reduced to arsine,
AsH 3 .
EXERCISE 21. To a solution of mercuric chloride, add a little
stannous chloride. A white precipitate is first obtained consisting
of mercun?#.r chloride. If the mixture be now gently warmed the
white precipitate turns grey, owing to its further reduction to
metallic mercury.
(1) 2HgCl 2 + SnCl 2 - Hg 2 Cl 2 + SnCl 4
(2) Hg 2 Cl 2 + SnCl 2 = 2Hg + SnCl 4
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.
For example, when an electric current is passed through a
solution of common salt, the compound is separated into its
two component elements, sodium and chlorine. The operation
is therefore "analytical." When sodium and chlorine are
brought together under suitable conditions, the two elements
unite, and reproduce the compound sodium chloride, the pro-
cess in this case being a synthetical one.
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 miscroscope (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.
16
Analytical Classification 17
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 com-
pounds, 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
BaQ 2 + Na 2 SO 4 = 2NaCl + BaSO 4
The barium sulphate is practically insoluble in water, and con-
sequently is precipitated. Now, if we know some property
belonging 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 solution and obtain
barium sulphate again, then this unknown solution must have
contained a sulphate. 1
Reagents. The materials that are used to bring about
analytical reactions are called reagents. Thus in the illustra-
tion 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
1 Sulphuric acid being included, as hydrogen sulphate.
C
1 8 Smaller Chemical Analysis
often known as general reagents. Others, again, are employed
because they produce 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.
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, 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 recognised by their known properties in the free state.
For example, from the compound lead chloride, PbClu, 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 cannot be isolated, 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 be isolated, 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
Analytical Classification 19
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 isolat-
ing the chlorine, we can employ the reagent silver nitrate,
AgNO,j. 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 versa 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 per-
chlorates. These compounds dissociate on solution, not into
simple chlorine ions, but into the complex C1O 3 and C1O 4 ions.
Such solutions, therefore, contain no Cl ions, and are therefore
incapable of forming AgCl with Ag ions.
In analytical classification the metallic or positive radicals
are classified into groups, based upon their behaviour towards
certain chosen reagents, known as group-reagents or general-
reagents, which are used in a certain order.
Group I. {hydrochloric acid group). Metals whose chlorides
are precipitated by dilute hydrochloric acid.
Group II. (sulphuretted hydrogen group). Metals whose
sulphides are precipitated from acid solutions by sulphuretted
hydrogen.
20
Smaller Chemical Analysis
Group H!A. (ammonia group). Metals whose hydroxides
are precipitated by ammonia in presence of ammonium
chloride.
Group Ills, (ammonium sulphide group). Metals whose
sulphides are precipitated by ammonium sulphide in presence
of ammonia.
Group IV. (ammonium carbonate group). Metals whose
carbonates are precipitated by ammonium carbonate in pre-
sence of ammonium chloride.
Group V. (no group-reagent). Metals which are not pre-
cipitated by any of the above reagents.
Group I.
Group II.
Group I II A.
Group Ills.
Group IV.
Group V.
Lead 1
Lead 1
Aluminium
Manganese
Barium
Magnesium
Silver
Mercury(ic)
Chromium
Zinc
Strontium
Sodium
Mercury
Copper
Iron
Nickel
Calcium
Potassium
(ous)
Bismuth
Cobalt
Ammonium
Cadmium
Arsenic
Antimony
Tin
The order in which the group-agents are used is an indis-
pensable condition of this classification. They can only be
employed in the order in which the groups are here numbered,
for the reason that each is only capable of separating its own
particular group from those groups which follow, and not from
those which go before. For example, if the metals of Group I.
are not first removed by precipitation with hydrochloric acid
before the group-reagent for Group II. is added, they would all
be precipitated as sulphides along with the metals of Group II.,
and thus no separation would be effected. There are con-
ditions under which the group-reagent for Group IIlA. fails to
separate this group from those which follow. When these con-
ditions are present, a special procedure has to be adopted,
which will be explained in its proper place later on.
1 The reason why lead is placed in both Groups I. and II. is explained
under the reactions for that metal.
CHAPTER III
REACTIONS OF THE METALS OF GROUP V
THIS group contains the alkali metals (ammonium being re-
garded 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 classification table, it will be seen that, in
the course of separating the various groups, certain ammonium
compounds are employed, therefore it will be obvious that it
is necessary to test for this " metal " in the substance under
examination before adding any ammoniacal compounds.
Ammonium, NH 4
DRY REACTIONS. When heated alone in a glass tube,
ammonium salts undergo change.
(a) If the acid is readily volatile (e.g. hydrochloric acid),
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.
(b) If the acid is non-volatile, or volatile only at a high
temperature (e.g. sulphuric or phosphoric), then the ammonium
salt is decomposed, ammonia being evolved, while the acid
remains.
(c) The ammonium salts of certain oxyacids which readily
part with oxygen (such as ammonium nitrate, nitrite, chromate)
21
22 Smaller Chemical Analysis
are also decomposed by heat, the ammonia being oxidised to
nitrogen or oxides of nitrogen.
NH 4 NO 3 = 2H 2 O + N 2 O
NH 4 NO 2 = 2H 2 O + N 2
(NH 4 ) 2 Cr 2 7 = Cr 2 3 + 4H 2 O + N 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.
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 3 .
(NH 4 ) 2 S0 4 + 2 NaHO = Na 2 SO 4 + 2H 2 O + 2NH 3
2NH 4 C1 + CaO = CaCl a + 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 recognised (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 turmeric paper brown ; (3) by the forma-
tion of white fumes when a glass rod moistened with strong
hydrochloric acid is held in the mouth of the test-tube.
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.
WET REACTIONS. All sodium salts are soluble ; sodium
Reactions of the Metals of Group V 23
platino-chloride is soluble in water, in alcohol, and in ether.
Hydrogen sodium tartrate also is freely soluble in water.
Sodium pyroantimonate, however, is less soluble in water than
the corresponding potassium salt, and is therefore precipitated
by the addition of a strong solution of potassium pyroanti-
monate 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 im-
parted 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
1 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 hydrochloric acid and heated until it ceases to impart any colour
to the flame.
2 The potassioscope consists merely of a small flat glass cell, contain-
ing 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 rubidium) gives a flame which appears red when
viewed through the potassioscope, whereas lithium, barium, stronium, and
calcium all give flames which appear red through indigo or blue glass.
24 Smaller Chemical Analysis
potassioscope, the red colour due to the potassium shines up as
brilliantly as before, while the yellow is completely intercepted. 1
WET REACTIONS. Most potassium salts are soluble in
water. Use a solution of potassium chloride.
Platinum chloride, 2 PtCl 4 , produces, with concentrated
solutions of potassium salts, a yellow crystalline precipitate
of potassium chloro-platinate (or potassium platinic chloride),
K 2 PtQ 6 , soluble in no parts of water at 10 (therefore more
soluble than the corresponding ammonium compound). 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
solutions 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 concentrated and neutral). The precipitate
is insoluble in alcohol.
1 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 con-
structed as shown in Fig. 8. A piece
of glass tube or glass rod is inserted in
a large cork (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 angles to the upright. The little
glass tube into which the platinum wire
is fused is then slipped over the projecting
iron wire. This arrangement admits of
the wire being raised or lowered as desired,
while at the same time it readily remains
in any position.
2 In reality the reagent used is chloroplatinic acid, H 2 PtCl 6 , which by
long habit is called platinum chloride : it is platinum chloride PtCl 4 plus
two molecules of hydrochloric acid, PtCl 4 , 2HC1.
FlG - 8 -
Reactions of the Metals of Group V 25
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
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 blowpipe 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.
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 = 2 NH 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 magne-
sium chloride by means of ammonia, thus
2 MgCl 2 + 2 NH 4 HO = Mg(HO) 2 + MgCl 2 , 2 NH 4 Cl
If ammonium chloride is previously present in sufficient quantity
to combine with the whole of the magnesium hydroxide, the
alkaline hydroxides give no precipitate.
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 am-
monium salts, precipitates of basic carbonates of magnesium,
the composition of which varies with conditions of temperature
26 Smaller Chemical Analysis
and concentration. The precipitate with (NH 4 ) 2 CO 3 only
separates out after a short time. In the presence of ammonium
chloride these reagents give no precipitate.
Hydrogen disodium 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.
A solution is examined for metals of Group V. by Table
V. on the page following.
Reactions of the Metals of Group V
^H
^>
3 bJD.S
Ji
s
[o _P
S J5
S P-i
P ^
rfl
i i
I s^ 53 .
4)
Fn v
w
If
- 1
^fc
tfl 13
^ P
8 S
^S
<L> CJ
^J
If
1 s
.S.S
So
^^
iJ
-2 I
1 1
il
CS O
in o
I.N
g ^
a a rt
a a g
CJ 03 S
fll
CHAPTER IV
REACTIONS OF THE METALS OF GROUP IV
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 except as a confirmatory test.
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 ) 2 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.
H 2 S0 4 , or any soluble sulphate, produces a white granular
precipitate of BaSO 4 , practically insoluble in water, insoluble
also in acids and alkalies (except by prolonged boiling with
strong acids and concentrated sodium carbonate, when it is
slowly dissolved). Insoluble 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 con-
tains only i part salt in 7000 parts of water.
Potassium chromate, K 2 CrO 4 , produces a primrose-yellow
precipitate of barium chromate, BaCrO 4 . It is insoluble in
acetic acid (distinction from SrCrO 4 ) ; soluble in HNO 3 and
in HC1.
28
Reactions of the Metals of Group IV 29
Hydrofluosilicic acid, H 2 SiF 6 , gives a white crystalline
precipitate of barium silicofluoride, BaSiF 6 , slightly soluble in
water, but insoluble on the addition of alcohol.
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 HC1.
WET REACTIONS. The same common salts of strontium
as of barium are soluble in water. The chromate and sulphate
are somewhat soluble.
(NH 4 ) 2 C0 3 (Na 2 CO 3 and K 2 CO 3 ) precipitates SrCO 3 , exactly
similar to the barium compound in its reactions.
H 2 S0 4 or soluble sulphates precipitate SrSO 4 . The pre-
cipitate is slightly soluble in water but almost insoluble in
a solution of (NH 4 ) 2 SO 4 (distinction from CaSO 4 ). SrSO 4 is
precipitated by a solution of CaSO 4 ; the precipitation does
not take place at once in cold solutions, but appears quickly
on heating.
Calcium, Ca
DRY 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.
WET REACTIONS. The same common salts of calcium
are soluble in water as of strontium and barium.
(NH 4 ) 2 C0 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 2 S0 4 added to a strong solution of a calcium salt give an
immediate precipitate of calcium sulphate. 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
3O Smaller Chemical Analysis
solution of ammonium sulphate, especially when hot (distinction
from SrSO 4 and BaSO 4 ). Boiling with potassium carbonate
easily converts it into calcium carbonate.
Ammonium oxalate, (NH 4 ) 2 C 2 O4, gives a white crystalline
precipitate of calcium oxalate. The precipitate is soluble in
mineral acids, but insoluble in acetic acid and in ammonia.
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 com-
pletely 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 -f 2NH 3 + H 2 O.] The mixture is filtered.
The filtrate is examined for the metals of Group V. by the
table given on p. 27, while the precipitate is examined by
Table IV. on the opposite page.
Reactions of the Metals of Group V 31
e
Jr!
a d
1/3 IO
Jtf
g
2 5
I a
*^ o
-
U o
5"-^ 'a
d
SI
ri
S o
acetates
ure the
rt o
s
I!
<D O
n
<D Z
ISO
3
^ ^J ,
^1.^0
5 | 2 ft
~ 73 , ^
8 S W W S
PP
*
fills
d ' ^ o 1>
* " "
fi o S
CHAPTER V
REACTIONS OP THE METALS OF GROUP III 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 platinum 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 ammonium
alum, (NH 4 ) 2 SO 4 ,A1 2 (SO 4 ) 3 ,24H 2 O, and potassium ahim,
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.
NH 4 HO throws down a white translucent precipitate of
aluminum hydroxide, or A1 2 (HO) 6 . Soluble in a large excess
of the reagent, but on gently boiling, the hydroxide is entirely
precipitated. In the presence of ammonium chloride, the pre-
cipitation by ammonia is complete. The precipitate is readily
soluble in mineral acids, and in acetic acid.
KHO or NaHO produces the same precipitate, which readily
combines with an excess of the reagent, forming a soluble
32
Reactions of the Metals of Group I II A 33
alurninate of potassium or sodium (A] 2 O 3J 3Na 2 O or Na 6 Al 2 O 6 ).
These aluminates are decomposed by acids
Al 2 O 3 ,3Na 2 O + 6HC1 = A1 2 (HO) 6 -f 6NaCl
but any excess of the acid beyond that required to combine with
the sodium of the alurninate at once re-dissolves the A1 2 (HO) 6 .
BaCO.j suspended in water, precipitates A1 2 (HO) 6 , carbon
dioxide being evolved. The precipitation is complete even in
the cold. 1 If alum or aluminium sulphate is used, the precipi-
tate is mixed with insoluble barium sulphate
A1 2 (S0 4 ) 3 -f 3 BaC0 3 4- 3 H 2 O = A^HO), + 3BaSO 4 4 3CO 2
(NH 4 ) 2 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
[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, Or
DRY REACTIONS. Chromium compounds impart to a
borax bead a grass-green colour, when heated either in the
outer or inner blowpipe flame.
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 a 8 4 2K 2 C0 3 + KC10 3 = 2 K 2 CrO 4 + KC1 + 2CO
1 In the presence of certain organic acids, as oxalic, tartaric, ot
citric acids, aluminium hydroxide is only more or less imperfectly pre-
cipitated 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 (C 4 H 4 O 6 ) 3 . This applies also in the case of the corre-
sponding chromium and iron compounds.
D
34 Smaller Chemical Analysis
WET REACTIONS. Two classes of chromium compounds
will be considered, namely the chromic compounds, derived
from chromium sesquioxide, Cr 2 O 3 ; and the chromates^ derived
from chromium trioxide (or chromic anhydride), CrO 3 .
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.
NH 4 HO produces a bluish or greenish-grey precipitate of
chromic hydroxide, Cr 2 (HO) 6 , partially dissolved by excess of
ammonia in the cold, giving a lilac-coloured liquid, but com-
pletely 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
neutralisation with HC1, and by boiling with NH 4 C1, as in the
case of Al.
BaC0 3 precipitates a mixture of the hydroxide and basic
carbonate. Complete precipitation only after some hours.
K 2 C0 3 and Na 2 C0 3 gave a similar precipitate, the composi-
tion of which varies with the conditions of precipitation.
(NH 4 ) 2 S precipitates Cr 2 (HO) 6 . 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 3 + 30 = 2Cr0 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 oxida-
tion may be accomplished in the wet way
(i) 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 + sKCIO = aKCl + 2K 2 CrO 4 -f sH a O
Reactions of the Metals of Group III A 35
(2) 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
never been isolated. The anhydride, CrO 3 , is readily obtained
by adding strong H 2 SO 4 to a cold strong solution of potassium
dichromate, 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 salts are mostly
yellow or red in colour. Both the chromates and dichromates
of the alkalies are soluble in water. The former (yellow) are
converted into the latter (red) by the addition of the requisite
amount of sulphuric acid
2 K 2 CrO 4 + H 2 SO 4 = K 2 Cr 2 O 7 + K 2 SO 4 + H 2 O
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. 28).
Lead chromate, PbCrO 4 (see Pb reactions, p. 58). PbCrO 4
melts without decomposition, and solidifies on cooling to a
brown crystalline mass. At higher temperatures it gives off
oxgen
2PbCrO 4 = Cr 2 O 3 + 2PbO + 30
Silver chromate, Ag 2 CrO 4 . A dark chocolate-red precipi-
tate, 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 3 .
Reduction of Chromic Acid. CrO 3 is a powerful oxidising
agent, giving up oxygen to oxidisable substances, and being
36 Smaller Chemical Analysis
itself reduced to Cr 2 O 3 ; that is, to the condition of a ' chromic "
compound. Thus, by sulphur dioxide it is reduced to chromium
sulphate
2 Cr0 3 -f 3 S0 2 = Cr 2 (S0 4 ) 3
The same action takes place in an acidified solution of potassium
dichromate
K 2 Cr 2 7 + H 2 S0 4 + 3H 2 S0 3 = Cr 2 (SO 4 ) 3 + K 2 SO 4 + 4H 2 O
Similarly, chromic acid and chromates are reduced by HC1,
oxidising the hydrogen of the acid, and liberating chlorine, after
the manner of peroxides ; thus
Cr0 3 + 6HC1 = CrCl 3 + 3 H 2 O + 3 C1
K 2 Cr 2 7 + I4HC1 = 2CrQ 3 + 2KC1 + 7H 2 O + 3C1 2
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 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 7 + 3 H 2 S + 8HC1 - 2 CrCl 3 + 2 KC1
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 2 CO 3 in the inner blowpipe flame,
iron compounds become reduced, and a dark grey magnetic
mass is obtained. If this be washed with water in a 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^. 1 They are both basic oxides, and
give rise to two classes of salts, namely, ferrous and ferric
1 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.
I
Reactions of the Metals of Group II I A 37
respectively. Ferrous salts readily take up oxygen, and become
converted into 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 volumin-
ous precipitate of ferric hydroxide, Fe 2 (HO) 6 , insoluble in
excess, or in NHjCl. 1
K 2 C0 3 , Na 2 CO 3 , and BaCO 3 give the same precipitate, CO 2
being liberated
2 FeCl 3 + 3Na 2 CO 3 + 3H 2 O = Fe 2 (HO) 6 +6NaCl + sCO 2
(NH 4 ) 2 S produces a black precipitate of ferrous sulphide,
the iron being reduced from the ferric state
2 FeCl 3 + s(NH 4 ) 2 S = 2 FeS + 6NH 4 C1 + S
Sulphuretted hydrogen, H 2 S, reduces 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.]
2 FeCl 3 + H 2 S = 2FeCl 2 + 2HC1 + S
Potassium ferrocyanide, K 4 Fe(CN) 6 , or K 4 FeCy 6 , pro-
duces with ferric salts a dark blue precipitate (Prussian blue]
3 K 4 (FeCy 6 ) + 4FeCl 3 = I2KC1 + Fe 4 (FeCy 6 ) 3
This test is extremely delicate, but where the amount of
iron is very small, a blue or greenish coloration only is pro-
duced. "Prussian blue " is insoluble in hydrochloric acid, but
readily dissolves in oxalic acid. It t is decomposed by NaHO
or KHO, with precipitation of ferric hydroxide
Fe 4 (FeCy 6 ) 3 + I2KHO = 2 Fe 2 (HO) 6 + 3K 4 (FeCy 6 )
Potassium ferricyanide, K 3 (FeCy 6 ), gives no precipitate
with ferric salts.
1 See footnote on p. 33, as to the influence of organic compounds.
38 Smaller Chemical Analysis
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. Ferr/V com-
pounds are readily reduced to the ferwits state ; they are there-
fore 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, $nG 2 , to stannic
chloride, SnCl 4 .
The oxidation of sulphurous acid or thiosulphuric acid into
sulphuric acid ; thus
2 FeCl 3 -f H 2 S0 3 + H 2 = H 2 SO 4 + 2 HC1 + 2FeCl 2
2 FeCl 3 -h N^SSO, + H 2 O = Na 2 SO 4 + 2HC1 + 2FeCl 2 + S
(b) Ferrous Compounds. Ferrous salts are usually pale
green when crystallised, and white when anhydrous. Of the
common salts the chloride and sulphate are soluble.
NH 4 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 pre-
cipitation 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 :j and Na 2 CO, give a white precipitate of ferrous
Reactions of the Metals of Group III A 39
carbonate, FeCO 3 , which on exposure to the air quickly absorbs
oxygen.
(NH 4 ) 2 S precipitates black ferrous sulphide, FeS. Readily
soluble in acids, with evolution of sulphuretted hydrogen.
K/FeCy 6 ) precipitates potassium ferrous ferrocyanide,
FeK 2 (FeCy 6 ), thus
K 4 (FeCye) + 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 dis-
solved 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
4 FeK 2 (FeCy 6 ) + 2C1 2 = Fe 4 (FeCy 6 ) 3 + K 4 FeCy^ 4- 4KC1
Potassium ferricyanide, K 3 (FeCy 6 ), giveV with ferrous
salts, a precipitate of ferrous ferricyanide, Fe 3 (FeCy 6 )2 (known
as Turnbulfs blue), which is indistinguishable by its appearance
from Prussian blue
2 K,(FeCy 6 ) + 3 FeCl 2 = Fe 3 (FeCy 6 ) 2 + 6KC1
The precipitate is insoluble in hydrochloric acid, but is
decomposed by caustic alkalies, with the precipitation of ferrous
hydroxide ; thus
Fe 3 (FeCy 6 ) 2 + 6KHO = 2K 3 (FeCy 6 ) + 3Fe(HO) 2
Oxidation of Ferrous to Ferric Compounds. The ferric
salts being the more stable, the ferrous compounds undergo
oxidation 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 chlorine water ^
or more conveniently by gently warming the ferrous compound
with hydrochloric acid and adding a few crystals of potassium
chlorate.
When the oxidation is accomplished with nitric acid, the
4O Smaller Chemical Analysis
strong acid is added, a few drops at a time, to the hot acidu-
lated solution of the ferrous salt. The solution becomes dark
in colour, and nitric oxide is disengaged ; thus
6FeS0 4 + 3H 2 S0 4 + 2 HNO,, = 3Fe 2 (SO 4 ) 3 + 4 H 2 O + 2NO
3 FeCl 2 + 3 HC1 + HN0 3 = 3FeCl 3 + 2H 2 O + NO
Unless the solution of the ferrous salt is acidified, a portion
of 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 , ^Fe 2 O 3 .
SEPARATION OF GROUP IIlA FROM GROUPS IIlB, IV., V
To the solution add NH 4 C1 in considerable quantity ; heat
the mixture to boiling, and add NH 4 HO carefully until precipi-
tation is complete. 1 Bring the liquid once more " to the boil,"
when, if sufficient ammonia has been added, the steam will
smell of it. Filter the mixture while hot. The precipitate
contains the metals of Group IIlA. in the form of hydroxides,
and is examined according to the table on the opposite page.
The filtrate contains the metals of Groups IIlB., IV., and V.
1 The NH 4 C1 prevents the precipitation of hydroxides of the metals of
Group IIlB. and of magnesium j it must be therefore added plentifully.
Excess of ammonia, however, must be avoided, otherwise manganous
hydroxide will be precipitated in spite of the ammonium chloride. (See
footnote, p. 44.)
Reactions of the Metals of Group I II A 41
IP
S S.S
PH Cj !/5
Pi
*'$$
* 8 3
S ^^
I a g
<u -2 o
S . Jq
x .S3 "
'a -
*"
O
On
*
5 fc
U
c
O D
o -H
"
is
a ,2
.1 1
^
1 8
a o
2-S
8 S
a
sodi
ic
t
e nitric acid
s alumini
wo portion
y with acet
cetate
into
Acid
lead
Acidify wth dil
of A1 2 (HO) 6 confi
H . .2,
'S
O *5
r^ m
A
I fi
2
'^ ^ I
III
8 I
T~
g-S
&
=
P IH
Jj <u
II
42 Smaller Chemical Analysis
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
fiision mixtiire? and fused in a platinum capsule. In this way
the chromium is converted into alkaline chromate ; a variable
proportion of the aluminium into aluminates.
The fused mass is then dissolved in water, and filtered.
The filtrate 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 2 added in quantity con-
siderably more than sufficient to produce precipitation. The
mixture is then boiled for a few minutes, and filtered.
The filtrate contains sodium aluminate, AL 2 O 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 dis-
solved 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.
1 Fusion mixture is a mixture of Na 2 CO 3 and K 2 CO 3 in equivalent
proportions (or about 10 parts Na 2 CO 3 to 13 of K 2 CO 3 ). It is used in
preference to Na 2 CO 3 alone, because it has the property of melting more
easily than either carbonate separately.
2 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 in Table IIlA., this difficulty
is avoided, as the sodium peroxide of commerce is usually free from this
impurity.
CHAPTER VI
REACTIONS OF THE METALS OF GROUP
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 to fuse the manganese
compound with KHO, or with Na 2 CO 3 and a little KNO 3 or
KC1O 3 upon a piece of platinum foil ; the manganese undergoes
oxidation, and a deep green-coloured mass is obtained, con-
sisting of manganates of the alkali metals
MnO 2 +Na 2 CO 3 +O (from KC1O 3 or KNO 3 ) = Na 2 MnO 4 +CO 2
WET REACTIONS. Of the common manganous l salts, the
chloride, MnCl 2 ,4H 2 O, and the sulphate, MnSO4,5H 2 O, are
soluble in water. In the crystallised state they have all a pink
colour. Manganous salts which are soluble in water do not
undergo atmospheric oxidation.
NH 4 HO, KHO, and NaHO produce a white precipitate of
manganous hydroxide, Mn(HO) 2 . Insoluble in excess of the
reagent, the precipitate quickly absorbs oxygen, and is con-
verted into hydrated manganic oxide, Mn 2 O 3 ,H 2 O, having a
brown colour.
Freshly precipitated Mn(HO) 2 , while still white, is soluble
in NH 4 C1; therefore in the presence of NH 4 C1, manganous
hydroxide is not precipitated by NH 4 HO, and only incom-
pletely by KHO or NaHO.
The ammoniacal solution of the double chloride, however, is
capable of absorbing oxygen just as the precipitated Mn(HO 2 )
1 The mangaiwV salts are extremely unstable in solution, and do not
exist under the ordinary conditions of analysis.
43
44 Smaller Chemical Analysis
does, and the liquid quickly becomes muddy, owing to the
precipitation from it of the brown hydrated manganic oxide. 1
(NH 4 ) 2 S precipitates manganous sulphide, MnS, as a pale
pinkish-white compound, easily soluble in dilute acids (dis-
tinction from Ni and Co), soluble also in acetic acid (distinction
from Zn). Precipitation with (NH 4 ) 2 S is only complete in the
presence of NH 4 C1. Owing to the ready solubility of MnS in
acids, H 2 S is incapable of precipitating manganous sulphide
from neutral solutions ; for, by double decomposition, the acid
of the manganous salt would be set free, and would imme-
diately redissolve the sulphide.
Manganese Compounds as Oxidising Agents. When
manganese dioxide is acted upon by acids, a manganous salt
is formed, and available oxygen is eliminated, which either
appears as free oxygen gas, or as the product of the oxidation
of the acid ; thus
MnO 2 + H 2 SO 4 = MnSO 4 + H 2 O + O
MnO 2 -f 4HC1 = MnCl 2 + 2H 2 O + Cl,
The manganates and permanganates are still more powerful
oxidizing agents, the former containing 2 and the latter z\
atoms of available oxygen in the molecule
K 2 MnO 4 = KAMnOA
2 KMnO 4 - K 2 O, 2 MnO,5O
Or, when acted upon by hydrochloric acid the equivalent
quantity of chlorine is evolved
KMn0 4 + 8HC1 - KC1 + MnCl. 2 + 4 H 2 O + 5C1
Potassium permanganate in acid solution is therefore capable
of oxidising a great variety of oxidisable compounds, giving
up its available oxygen, and becoming itself reduced to a
1 For this reason the complete separation of manganese from the metals
of Group IIlA. is difficult to accomplish. If care be taken to ensure the
presence of sufficient NH 4 C1, and if the precipitation with NH 4 HO be
made as quickly as possible with the least possible excess of ammonia in a
hot solution, and the excess of ammonia at once boiled off, and the liquid
filtered immediately, the risk of precipitating the manganese may be re-
duced to a minimum. The student will do well to practise the separation
of manganese from the metals of Group IIlA. by using solutions containing
known small proportions of a manganous salt, mixed with large quantities
of iron or chromium or aluminium.
Reactions of the Metals of Group IIlB 45
manganous salt, the characteristic purplish colour of the per-
manganate of course being destroyed. The reaction with
ferrous sulphate may be taken as a typical example. Two
molecules of this salt in presence of sulphuric acid require
one atom of oxygen for their oxidisation
2FeSO 4 + H 2 SO 4 -f O = H 2 O + Fe 2 (SO 4 ) 3
Hence one molecule of KMnO 4 is capable of oxidising five
molecules of FeSO 4 ; or, in the presence of HC1, of oxidising
five molecules of FeCl 2 into FeCl 3 .
Zinc, Zn
DRY REACTIONS. Zinc compounds give no characteristic
borax bead.
When heated on charcoal with sodium carbonate in the
reducing flame, zinc compounds are reduced, but the metal
is too volatile to appear in the form of globules. As it is
reduced, it volatilises; and the vapour burns as it passes
through the outer flame, which thereby becomes tinged a
bluish-white colour. The zinc oxide which is thus produced,
deposits as an incrustation upon the charcoal, which is canary-
yellow while hot, becoming white on cooling. Zinc oxide is
not volatile, and therefore the incrustation does not disappear
when the oxidising flame is made to play upon it. If the zinc
oxide be moistened with a drop of cobalt nitrate, and again
heated in the oxidising flame, it assumes a green colour.
WET REACTIONS. Of the common salts the chloride,
sulphate, and nitrate are soluble in water.
NaHO, or KHO, throws down a white precipitate of
Zn(HO) 2 , readily soluble in excess of the reagent, forming
double salts (sometimes called zincates)^ such as ZnNa 2 O 2 .
Moderately strong solutions of these zincates may be boiled
without undergoing any change, but from dilute solutions
Zn(HO) 2 is reprecipitated. NH 4 HO precipitates the same
compound, also soluble in excess of ammonia. Zn(HO) 2 is
soluble in NH 4 C1, owing to its readiness to form soluble double
salts with the alkaline chlorides, having the general formula
ZnCl 2 ,2RCl
Zn(HO) 2 + 4NH 4 C1 = ZnCl 2 ,2NH 4 Cl + 2NH 4 HO
46 Smaller Chemical Analysis
Hence in the presence of much ammonium chloride, NH 4 HO
gives no precipitate, and the precipitation with KHO or NaHO
is incomplete.
K 2 C0 3 , Na 2 CO 3 , or (NH 4 ) 2 CO 3 produces a white precipitate
of basic carbonate; the precipitate is soluble in excess of
(NH 4 ) 2 CO 3 . The presence of much NH 4 C1 prevents the
precipitation.
(NH 4 ) 2 S throws down a white precipitate of zinc sulphide,
ZnS. In the presence of NH 4 C1 the precipitation is complete
even from dilute solutions. ZnS is soluble in dilute mineral
acids, hence H 2 S is incapable of completely precipitating this
sulphide from neutral solutions of the zinc salts of such acids
ZnCl 2 + H 2 S> ZnS + 2HC1
ZnS is insoluble in acetic acid (contrast MnS), therefore
from the acetate, or other zinc salts in presence of an alkaline
acetate, ZnS is completely precipitated by H 2 S ; thus
Zn(C 2 H 3 O 2 ) 2 + H 2 S = ZnS + 2H(C 2 H 3 O 2 )
ZnCl 2 + 2Na(C 2 H 3 O 2 ) + H 2 S = ZnS + 2NaCl + 2H(C 2 H 3 O 2 )
Nickel, Ni
DRY REACTIONS. Nickel compounds impart a dark red-
brown colour to the borax bead when heated in the oxidising
flame, the colour becoming brownish-yellow on cooling. In
the reducing flame the borax bead becomes opaque and grey.
In a bead of microcosmic salt, the red-brown colour persists in
both flames.
The presence of other colour-producing oxides renders this
test uncertain, while even traces of cobalt entirely mask it.
Heated on charcoal with Na 2 CO 3 , metallic nickel is obtained
as a grey feebly magnetic mass.
WET REACTIONS. Only one of the oxides of nickel is
basic, namely NiO, hence only one series of salts exists. The
sesquioxide, Ni 2 O 3 , behaves like a peroxide ; thus
Ni 2 O 3 + 2H 2 SO 4 = 2 NiSO 4 + 2H 2 O + O
In the crystalline or hydrated condition the nickel salts have a
Reactions of the Metals of Group II IB 47
green colour, and dissolve to green solutions. The anhydrous
salts are pale yellow. Of the common salts, the chloride,
nitrate, and sulphate are soluble in water.
KHO or NaHO gives a pale bluish-green precipitate of
nickelous hydroxide, Ni(HO) 2 , insoluble in excess of either
reagent ; soluble in ammonium salts. Ni(HO) 2 is not oxidised
on exposure to air, but it is converted into black hydrated
sesquioxide, Ni 2 O 3 ,3H 2 O, by hypochlorites, or by the action of
chlorine or bromine in the presence of caustic alkalies; thus
2Ni(HO) a + NaCIO + H 2 O = NaCl + Ni 2 O 3 ,3H 2 O
2Ni(HO) a 4- 2NaHO + C1 2 = 2NaCl + Ni 2 O 3 ,3H 2 O
NH 4 HO gives a greenish precipitate in neutral solutions of
a basic compound, readily soluble in excess of ammonia to a
greenish-blue solution. In the presence of ammonium salts,
no precipitate is produced by ammonia. The nickel in this
solution is not oxidised by hypochlorites, but it is completely
precipitated, as Ni(HO) 2 , on the addition of KHO.
K 2 C0 3 , Na 2 CO 3 , or (NH 4 )CO 3 produces a pale-green pre-
cipitate of basic carbonate, ^NiCO 3 ,jNi(HO) 2 , soluble in excess
of ammonia to a bluish solution.
(NH 4 ) 2 S, or H 2 S in presence of ammonia, produces a black
precipitate of NiS, soluble to a slight extent in excess ; more
readily soluble if ammonia or polysulphides of ammonia are
present, yielding a brown solution. From this solution the
dissolved NiS is reprecipitated slowly on boiling, more quickly
after acidifying with acetic acid, or the addition of ammonium
acetate.
NiS is only difficultly soluble in strong HC1, and almost
insoluble in the dilute acid; also in acetic acid. Readily
soluble in aqua, regia, or in HC1 and a crystal of KC1O 3 , yield-
ing NiCl 2 ; soluble also in HNO 3 .
H 2 S only produces complete precipitation of NiS from a warm
solution of the acetate, or from other nickel salts in the presence of
an alkaline acetate. In the case of neutral solutions of nickel salts
with mineral acids, the precipitation is only partial, while in acid
solutions it does not take place at all.
KCy gives a pale-green precipitate of nickelous cyanide,
48 Smaller Chemical Analysis
NiCy 2 , readily soluble in excess of the reagent forming the
double cyanide, K 2 NiCy 4 . From this solution (NH 4 ) 2 S fails
to precipitate the sulphide. Dilute mineral acids cause the
reprecipitation of nickelous cyanide
K 2 NiCy 4 + 2HC1 = 2 KC1 + 2HCy + NiCy 2
Boiling with hydrochloric acid decomposes the metallic cyanide
altogether
K 2 NiCy 4 + 4HC1 = 2KC1 + 4HCy + NiCl 2
Oxidising agents, such as hypochlorites, chlorine, or bromine
precipitate the black hydrated sesquioxide. (Distinction from
cobalt.}
2 K 2 NiCy 4 + NaCIO + 5H 2 O = Ni 2 O 3 ,3H 2 O + NaCl
+ 4 KCy
Cobalt, Co
DRY REACTIONS. Cobalt compounds impart to a borax
bead a rich blue colour, when heated either in the oxidising
or reducing flame. The test is characteristic and delicate.
WET REACTIONS. Cobalt forms a number of oxides, two
only of which are basic. The cobaltous salts are derived from
CoO, while the feebly basic sesquioxide Co 2 O 3 forms the very
unstable cobaltic salts.
Of the common cobaltous salts, the sulphate, nitrate, and
haloid salts are soluble in water. In the hydrated condition
they have a pink colour, dissolving to pink solutions. In the
anhydrous state they are blue. A strong aqueous solution
(pink) will,' however, turn blue when boiled, and return to its
original pink colour when again cooled.
KHO or NaHO gives a greenish-blue precipitate of a basic
salt. On boiling with excess of the alkali, the precipitate is
converted into the pink hydroxide, Co(HO) 2J which, however
is coloured more or less brown by the oxidation of a portion,
of it (by atmospheric oxygen) into the hydrated cobaltic oxide,
Co 2 3 ,3H 2 0, or Co 2 (HO) 6 .
Co(HO) 2 is oxidised by hypochlorites, in the same manner
as the corresponding nickel compound.
Separation of Group fill? 49
NH 4 HO causes partial precipitation of a basic salt, which
dissolves easily in excess of ammonia or in ammoniacal salts.
The solution, which has a brownish colour, absorbs oxygen,
and becomes darker in colour.
K,C0 3 , Na. 2 CO 3 , or (NH 4 ) 2 CO 3 precipitates a lilac-coloured
basic carbonate, .#CoCO 3 , yCo(HO) 2 , readily soluble in excess
of (NH 4 ) 2 CO 3 to a reddish solution.
HKC0 3 gives a pinkish precipitate of normal cobalt car-
bonate, CoCO 3 . The precipitate so obtained is soluble in
hydrogen peroxide, yielding a deep green solution.
(NH 4 ) 2 S gives a black precipitate of cobaltous sulphide,
CoS. The precipitation is complete in the presence of NH 4 C1
(contrast NiS). CoS is soluble in HNO 3 , in "aqua regia,"
and in HC1 with the addition of a crystal of KC1O 3 ; difficultly
soluble in strong HC1 ; practically insoluble in dilute HC1.
H. 2 S precipitates CoS under the same conditions as apply
in the case of NiS.
KCy gives a reddish precipitate of CoCy 2 soluble in excess
of the reagent, forming potassium cobaltocyanide, K 4 CoCy 6
(corresponding to potassium ferrocyanide) (contrast nickel).
The addition of oxidising agents hypochlorites, chlorine, or
bromine to this solution produces no precipitation of cobalt
sesquioxide, but oxidises the cobaltocyanide into cobalticyanide.
(Distinction from nickel.)
2K 4 CoCy 6 + NaCIO + H 2 O = NaCl + 2KHO + 2 K 3 CoCy 6
SEPARATION OF GROUP IIlB. FROM GROUPS IV. AND V
Add NH 4 C1, NH 4 HO, 2 and pass H 2 S through the solution
(or add (NH 4 ) 2 S, drop by drop, avoiding excess ; see Nickel)
until precipitation is complete. Warm gently, and filter.
1 Solutions of these double cyanides do not give precipitates with the
reagents usually employed for the detection of the metals, because on dis-
sociation they do not yield ions of the metal, but complex ions of the
metal and cyanogen. For example, such cyanides as K 2 NiCy 4 yield
2K + NiCy 4 , while compounds such as the ferrocyanides and cobaltocyanides
yield the ions FeCy e and CoCy 6 respectively.
2 In the ordinary course of analysis, these reagents have already been
added for the separation of Group II IA.
E
50
Smaller Chemical Analysis
S
|i
-V * a
5 **" H
nl
ci ^ VM
30.
3 **
Jt *'G "&
8J!
"o* u.!5
U ? g
T3 (U 2.2
g-s^S
^s.S 'g U
^* O 'o
CJ <U
H
lve
i|*1
S' 3 "
"^ S ^ >r 3 oj c '2
B.3tta trt eS 5
IS^alfr 8 -
- 'sf-S'.sS
ll^lll
1.90*11
&^'|
^ ^ v> csr^ o"5
0-3 ^;^ S^
.y I, a c
- H
.5^-5 o!-; o
i- o ^-rffi-B
CS
ate
r
i
u
y
e
lu
aHO
cu
S
pitate (Ni
Blue colo
quantit
eutralise
ade sol
, add N
ntil the
n
y m
d
u
2n"l11'~
o -
Separation of Phosphoric Acid 5 1
SEPARATION OF PHOSPHORIC ACID.
As stated on p. 20, there are conditions under which the
group reagent for Group IIlA. fails to separate the metals of
this group from those which follow. The presence of phos-
phates is such a condition. The phosphates of all the metals
of Groups III. and IV., as well as of magnesium, are soluble
in hydrochloric acid, and are reprecipitated on the addition
of ammonia. Hence the precipitate obtained by ammonia, in
the process of separating Group IIlA., may contain any or all
of these phosphates, in addition to the hydroxides of iron,
chromium, and aluminium. Before proceeding to the separa-
tion of Groups IIlA. and IIlB., therefore, it is essential first
to ascertain whether or not any phosphates are present ; and
second, if they are found present, to adopt measures to remove
the phosphoric acid.
Test for Phosphoric Acid. To a small quantity of the solu-
tion add a large excess (four or five times the volume) of a
nitric-acid solution of ammonium molybdate (NH 4 ) 2 MoO 4 , and
gently warm the mixture, when a bright yellow precipitate of
ammonium phospho-molybdate separates out. 1
The removal of Phosphoric Acid is based upon the fact that
ferric phosphate (also the phosphates of Al and Cr) is insoluble
in acetic acid. Thus, if acetic acid or sodium acetate (in
practice, a mixture of the two is employed) be added to a
hydrochloric acid solution of ferric phosphate, and the mixture
boiled, the ferric phosphate is precipitated; also, if ferric
chloride is added to an acid solution of any of the phos-
phates soluble in that acid (i.e. the phosphates of metals of
Groups IIlB., IV., and magnesium), by double decomposition
ferric phosphate is thrown down, and chlorides of the other
metals are left in solution. For example
Ca 3 (PO 4 ). 2 4- 2FeCl 3 (in presence of acetic acid) = Fe 2 (PO 4 ) 2
The process is carried out according to the scheme given in
Table I He. on the following page.
1 Arsenic acid gives a similar yellow precipitate of ammonium arseno-
molybdate ; therefore the test for phosphoric acid must not be applied until
arsenic has been removed in Group II.
TABLE
THE SEPARATION OF
The precipitate produced by NH 4 HO in presence of NH 4 C1 may consist
of Groups III. and IV., and of Mg. Dissolve in a little dilute HC1,
and acetic acid * (reagent).
The Precipitate may consist of phos-
phates of Fe, Al, Cr (along with
basic acetate of iron).
This precipitate may be treated with
Na 2 O 2 exactly as the hydroxides
(Table IIlA., p. 41) A1 2 (PO 4 ) 2
dissolves, the Cr 2 (PO 4 ) 2 is oxidised
to Na 2 CrO 4 . On filtering, the
iron remains on the filter. Alu-
minium is detected by neutralising
and reprecipitating with ammonia ;
the chromium by barium chloride
in acetic acid solution.
The Filtrate. To a small portion
If a precipitate is produced,
drop, until the whole of the
the precipitation is seen by the
mixture, 3 and filter.
If the preliminary test with ferric
contains no more phosphoric
The Precipitate consists of ferric
phosphate and basic acetate.
Throw away,
1 If the addition of sodium acetate gives no precipitate, it follows
analysis (see previous page).
If, after the addition of the acetate reagent, and before the mixture
allowing the precipitate to settle somewhat), this means that ferric acetatt
solution) ; and if this is formed, it follows that there is present in the
present ; hence all the phosphoric acid will have passed into the precipitat<
excess of ferric chloride upon the sodium acetate. Thus
2 Boiling the mixture at this point ensures the conversion of the solubL
solution should be colourless, or entirely free from the red colour.
* The boiling here is for the same reason as that in the foregoing note
precipitated ferric phosphate is soluble in ferric chloride and in ferri
IIIc
PHOSPHORIC ACID
of hydroxides of Al, Fe, Cr, as well as phosphates of any of the metals
avoiding excess. Nearly neutralise with Na 2 CO 2 , and add sodium acetate
Boil the mixture, 2 and filter.
of the liquid, which should be colourless, add a drop of ferric chloride,
then ferric chloride is added to the main portion of the filtrate, drop by
phosphoric acid is thrown down as ferric phosphate. The completion of
liquid becoming red, by the formation of ferric acetate. Boil the
chloride gave no precipitate, but only a red coloration, then the solution
acid, and is at once treated as in the next step.
The Solution. Add NH 4 C1, heat to boiling, and add NH 4 HO.
Filter.
The Precipitate may
consist of A1 2 (HO) 6
and Cr 2 (HO) 6 .
magnesium in the usual way.
Examine in usual way. i
Table IIlA. Tables IIlB., IV., and V.
The Filtrate.
Examine for Groups IIlB. and IV., and for
that no iron, aluminium, or chromium are present in the substance under
has been boiled, the liquid itself appears red (which is easily seen by
is being produced (which is soluble in sodium acetate, forming a red
mixture more than enough iron to unite with all the phosphoric acid
along with the iron. This ferric acetate is formed by the action of the
FeCl 3 + 3Na(C 2 H 3 O 2 ) = Fe(C 2 H 3 O 2 ) 3 + sNaCl
ferric acetate into the insoluble basic acetate, so that when filtered the
It is important to avoid any unnecessary excess of FeCl 3 , because the
acetate.
CHAPTER VII
REACTIONS OF THE METALS OF GROUP II
THIS group is conveniently divided into two sections
Subdivision i. Mercury, lead, bismuth, cadmium, copper.
Subdivision 2. Arsenic, antimony, tin.
SUBDIVISION i
Mercury, Hg
DRY REACTIONS. When heated alone in a tube, many
mercury compounds (those with the halogens, for example)
volatilise unchanged, giving sublimates of the same compound.
The iodide (red) when heated forms a sublimate, consisting
chiefly of the yellow allotropic form of HgI 2 , which when cold
changes to red if scratched or rubbed. Some mercury com-
pounds, e.g. the oxide, yield a sublimate of metallic mercury.
If a mercury salt be mixed with several times its weight of
sodium carbonate (both being as dry as possible), and the
mixture be strongly heated in a dry narrow test-tube, a sub-
limate of metallic mercury will be obtained. The sublimed
mercury will present the appearance of a bright metallic mirror,
but if examined by means of a lens, or if rubbed with a glass
rod, distinct globules of liquid metal will be visible.
WET REACTIONS. Mercury forms two classes of salts,
namely, the mercurzV and the mercuric salts. The former
compounds contain the divalent atom or radical Hg", while
the mercurous salts contain the divalent double atom or
radical (Hg a )".
The metal in its mercuric compounds belongs to Group II.,
54
Group //. Subdivision l 55
while in its mercurvus salts it falls in Group I. For conveni-
+ +
ence the reactions of both the Hg ion and the (Hg 2 ) ion will
be studied in this place.
(a) Mercuric Compounds. Of the common salts, the nitrate,
sulphate, chloride, and bromide (but not the iodide] are soluble
in water, but the solubility is not very great.
KHO or NaHO gives with mercuric compounds a yellow
precipitate 1 of mercuric oxide, HgO
HgCl 2 + 2KHO = HgO + H 2 O + 2KC1
The precipitate is insoluble in excess of the reagent.
NH 4 HO produces a white precipitate of an ammoniacal
mercuric compound, where two atoms of hydrogen from the
ammonium radical are replaced by the divalent atom Hg;
thus
HgCl 2 + 2 NH 4 HO = NH 2 HgCl + NH 4 C1 + 2 H 2 O
H 2 S produces a black precipitate of HgS. At first the
precipitate is white, changing rapidly through yellow and
brown to black. The white compound consists of HgCl 2 ,
2 HgS (or Hg(NO 3 ) 2 , 2 HgS when the nitrate is used). These
colour changes are characteristic. The precipitation is only
complete after some time, and when the solution is consider-
ably dilute. The compound is insoluble in HC1, and in
HNO 3 even when boiling. (The prolonged action of boiling
HNO (J partially converts it into the white compound Hg(NO 3 ) 2 ,
2 HgS.) Mercuric sulphide dissolves in aqua regia, forming
mercuric chloride. In the presence of caustic alkalies it dis-
solves in sodium or potassium sulphide (not in ammonium
sulphide), forming the double sulphides, HgS,Na 2 S and
HgS,K 2 S.
KI precipitates HgI 2 as a rich scarlet compound, soluble
1 On the first addition of the reagent, the precipitate appears a brownish
colour (probably due to the momentary formation of the hydroxide, which
is incapable of existing), but almost immediately it becomes yellow. Why
the oxide obtained by precipitation should be yellow, while that prepared
in the dry way is brick-red, is not known. Compare also the sulphide.
56 Smaller Chemical Analysis
in excess of either solution. When first precipitated it appears
yellow, but quickly turns salmon-red, and then scarlet.
Reduction of Mercuric Compounds. By reducing agents
mercuric compounds may be converted into mercurous salts,
or the reduction may go a stage further and result in the
precipitation of mercury. Thus, on the addition of stannous
chloride, SnCl 2 , a white precipitate of mercurous chloride is
produced
2HgCl a + SnCl 2 = Hg 2 Cl 2 + SnCl 4
On gently warming with an excess of stannous chloride, the
precipitated mercurous chloride changes to a grey deposit of
mercury in a condition of fine powder
Hg 2 Cl 2 4- SnCla = SnCl 4 + 2Hg
Many metals are capable of displacing mercury from its
solutions, the mercury being deposited upon the metal. Thus,
if a strip of clean copper be immersed in a neutral or slightly
acid solution of a mercury salt, it becomes coated with a
greyish deposit, from which the mercury can be readily volati-
lised and obtained as a metallic sublimate by heating the
copper in a dry test-tube.
(b) Mercurous Compounds. Of the common salts mer-
curous nitrate is the only one which is readily soluble, and
this only so long as the water is acid with nitric acid. The
addition of much water results in the precipitation of a basic
nitrate. Mercurous sulphate is soluble with difficulty.
KHO or NaHO throws down a black precipitate of (Hg 2 )O.
Mercurous oxide is very unstable. When gently warmed, or
even upon exposure to light, it is converted into HgO and Hg.
NH 4 HO precipitates an ammoniacal mercurous compound,
which is black.
Hg 2 (N0 3 ) 2 + 2NH 4 HO = NH 2 (Hg 2 )N0 3 + NH 4 NO 3
H 2 S produces a black precipitate, which is a mixture of
HgS and Hg. This precipitate, therefore, behaves, on treat-
ment with nitric acid, in the same way as that obtained from
a mercuric solution, giving the white insoluble compound
Hg(N0 3 ) 2 ,2HgS.
Group II. Subdivision i 57
HC1 and soluble chlorides precipitate white mercurous
chloride, Hg 2 Cl 2 . Insoluble in dilute acids ; soluble in boiling
HNO 3 , being converted into HgCl 2 and Hg, and the mercury
then dissolves to mercuric nitrate, with evolution of oxides of
nitrogen.
Ammonia converts it into black mercurous ammonium
chloride, NH 2 (Hg 2 )Cl. (This constitutes one of the most
characteristic reactions for mercurous compounds.)
Mercurous salts are reduced to metallic mercury by the
reducing agents which reduce the mercuric compounds ; thus,
with stannous chloride a grey precipitate of mercury is at once
produced
Hg a (N0 8 ) a + SnCl 2 + 2HC1 = SnCl 4 = 2 HNO 3 + 2 Hg
Lead, Pb
DRY REACTIONS. Lead compounds are very readily re-
duced when heated upon charcoal before the blowpipe flame,
either alone or mixed with sodium carbonate or potassium
cyanide. Globules of metallic lead are thus obtained, and at
the same time a yellowish incrustation is formed, consisting of
the oxide PbO (litharge). When cold, one of the globules can
be removed and the properties of the metal examined. Lead
may be recognised by its malleability and softness, the latter
property enabling it to leave a black mark when rubbed upon
paper.
WET REACTIONS. The only salts of lead which are met
with in analysis are derived from plumbic oxide, PbO, in
which the metal is divalent. Of the common salts, the nitrate
and acetate are readily soluble in water ; the chloride, bromide,
and iodide are sparingly soluble.
KHO, NaHO, or NH 4 HO gives a white precipitate of
lead hydroxide, Pb(HO) 2 (usually admixed with a basic com-
pound), .soluble in excess of KHO or NaHO, but not in
NH 4 HO.
KjCOg, NaaCO 3 , or (NH 4 ) 2 CO 3 gives a precipitate of basic
carbonate of lead.
HJ3 gives a black precipitate of lead sulphide, PbS. In
58 Smaller Chemical Analysis
the presence of much hydrochloric acid, the precipitate first
formed consists of a brown compound having the composition
PbCl 2 ,2PbS, which by the further action of H 2 S is converted
into the black PbS.
PbS is insoluble in cold dilute acids, in alkalies, or in the
sulphides of the alkalies. It is readily dissolved by hot dilute
HNO 3 , giving lead nitrate and free sulphur. As the strength
of the acid is increased, this sulphur begins to be oxidised into
sulphuric acid, which causes the precipitation of lead sulphate.
Strong nitric acid converts lead sulphide entirely into the
sulphate.
H 2 S0 4 and soluble sulphates give a white precipitate of
lead sulphate, PbSO 4 . Very slightly soluble in water; less
soluble in the presence of either dilute sulphuric acid or
alcohol; hence, in very dilute solutions, precipitation is ac-
celerated by the addition of alcohol. PbSO 4 dissolves by long
boiling with strong HC1, yielding PbCl 2 . It dissolves more
readily in strong ammoniacal solutions of ammonium acetate
or tartrate, as well as in hot KHO or NaHO. From these
it is again precipitated on addition of H 2 SO 4 .
HC1 and soluble chlorides give a white precipitate of lead
chloride, PbCl 2 . The precipitate is slightly soluble in cold
water, moderately freely in boiling water ; from which solution
it separates on cooling in long white needle-shaped crystals.
The presence of free HC1 diminishes its solubility in cold
water. Owing to this partial solubility of the chloride, lead
is not completely separated by the group-reagent for Group I.,
and therefore is met with also among the metals of Group II.
Kl gives a yellow precipitate of PbI 2 . Soluble, but to a
less extent than the chloride, in boiling water to a colourless
solution.
KjCrC^ precipitates yellow lead chromate, PbCrO 4 , insoluble
in acetic acid. Soluble in dilute HNO 3 and in caustic alkalies
(see reactions for chromium).
Bismuth, Bi
DRY REACTIONS. Bismuth compounds are easily reduced
when heated with Na 2 CO 3 upon charcoal. The metal, however,
Group IL Subdivision i 59
rapidly oxidises when strongly heated, hence the charcoal
becomes covered with an incrustation of the pale yellow oxide,
Bi 2 O 3 , the colour of which (as is the case with most coloured
oxides) appears darker (orange-yellow) while hot. Globules
of the metal, if detached from the charcoal, may be at once
distinguished from lead or silver by their brittleness. Bismuth
dissolves easily in HNO 3 , Sut is scarcely attacked by HC1, or
by dilute H. 2 SO 4 .
WET REACTIONS. Of the common salts of bismuth none
are soluble in water in the ordinary sense, but the nitrate and
chloride are readily soluble in water acidified with the re-
spective acids. Water alone, converts these salts into basic
compounds, which are soluble in acid ; the action, therefore, is
reversible
Bi(N0 3 ) 3 + 2H 2 O$; Bi(HO) 2 N0 3 + 2 HNO 3
In the case of bismuth chloride, the oxychloride is thrown
down
BiCl 3 + H 2 OBiOCl + 2HC1
This compound is not so easily dissolved by HC1 as the basic
nitrate is by HNO 3 , therefore the whole of the bismuth will be
precipitated if the solution is dilute.
The precipitate is not dissolved by tartaric acid. (Distinc-
tion from antimony. )
KHO or NaHO precipitates the white hydroxide Bi(HO) 3 ,
or Bi 2 O 3 ,3H. 2 O. Insoluble in excess of the precipitants. From
boiling solutions, or on heating to boiling, the basic compound
is formed, BiO(HO) or Bi 2 O 3 ,H 2 O.
K 2 C0 3 , Na,,CO 3 , or (NH 4 ) 2 CO ; 5 throws down a white basic
carbonate, (BiO). 2 CO ;j . Insoluble in excess of the reagents.
H. 2 S or (NH 4 ) 2 S precipitates bismuthous sulphide, Bi. 2 S 3 , as
a dark brown, almost black, compound. Soluble in HNO 3 ;
insoluble in alkaline sulphides.
K 2 Cr 2 7 precipitates basic bismuth dichromate, (BiO) 2 Cr 2 O 7 .
Insoluble in KHO. (Distinction from lead.}
60 Smaller Chemical Analysis
Cadmium, Cd
DRY REACTIONS. Cadmium compounds, heated on char-
coal with sodium carbonate, are easily reduced; but, owing
to the ready volatility of the metal, the latter is converted into
the oxide, which is deposited as a brown incrustation upon
the charcoal.
WET REACTIONS. Of the common salts, the nitrate,
sulphate, chloride (bromide, iodide, and acetate) are soluble
in water.
KHO, NaHO, or NH 4 HO precipitates the white hydroxide,
Cd(HO) 2 . Insoluble in excess of KHO or NaHO, but soluble
in NH 4 HO.
K 2 C0 3 , Na 2 CO 3 , or (NH 4 ) 2 CO 3 gives a white precipitate of
CdCO 3 . The presence of NH 4 HO prevents the precipitation.
H 2 S or (NH 4 ) 2 S precipitates cadmium sulphide, CdS, dis-
tinguished from the sulphides of all the other metals of this
division by its pure yellow colour. It is more easily soluble
in acids than the other sulphides of the group, and therefore,
to ensure complete precipitation, the solution must not be too
strongly acid.
CdS is insoluble in potassium cyanide, and is precipitated
by H 2 S from a solution of CdCy 2 in excess of KCy (see Method
of Separation from Copper). CdS is insoluble in alkaline
sulphides, which distinguishes it from arsenious sulphide, which
is the only other yellow sulphide (see p. 55).
Copper, Cu
DRY REACTIONS. Copper compounds are reduced to
metallic copper when strongly heated upon charcoal along
with sodium carbonate in the reducing flame. Reddish scales,
or even globules, of metal will be found. Heated in a borax
bead, copper salts impart a colour which is green while the
bead is hot, but bluish when cold. Heated on a platinum
wire they give a green flame, which appears bright blue if the
substance is moistened with strong HC1 and reintroduced into
the flame.
WET REACTIONS. Copper forms two series of salts, cuprous
Group II. Subdivision i 61
and cupric, derived from the two oxides Cu 2 O and CuO. The
former readily pass by oxidation into cupric compounds.
(a) Cupric Salts. Of the common cupric salts, the sulphate,
nitrate, chloride (bromide and acetate) are readily soluble in
water. In the crystallised or hydrated condition they are blue
or green, but in the anhydrous state either white or pale
yellow.
KHO or NaHO produces a pale blue precipitate of cupric
hydroxide Cu(HO) 2 . Insoluble in excess. On boiling the
mixture, the hydroxide is converted into black cupric oxide.
NH 4 HO or (NH 4 ) 2 CO 3 precipitates a light blue basic com-
pound, readily soluble in excess to a deep blue solution
(characteristic of copper compounds).
K 2 C0 3 or Na 2 CO 3 gives a greenish precipitate of the basic
carbonate, Cu(CO 3 ),Cu(HO) 2 . Insoluble in excess of the
reagent.
KCy gives with both cupric and cuprous salts a white pre-
cipitate of cuprous cyanide, Cu 2 Cy 2 ; soluble in excess, forming
cuprous potassium cyanide, Cu 2 Cy 2 ,6KCy or K 6 Cu 2 Cy 8 .
This double cyanide is also formed when excess of potas-
sium cyanide is added to blue ammoniacal copper solution,
the blue colour disappearing in consequence.
Sulphuretted hydrogen fails to precipitate copper sulphide
from the solution of this double cyanide (separation from
cadmium).
H 2 S or (NH 4 ) 2 S produces a nearly black precipitate of
cupric sulphide, CuS, which, when exposed to the air in a
moist condition, absorbs oxygen and is converted into the
sulphate. The precipitate is slightly soluble in ammonium
sulphide. It readily dissolves in potassium cyanide, forming
cuprous potassium cyanide, Cu 2 Cy 2 ,6KCy (compare cadmium).
(b) Cuprous Salts. The common cuprous salts are all
insoluble in water. For the reactions, a solution of cuprous
chloride in hydrochloric acid may be used.
KHO or NaHO gives a yellow precipitate of cuprous
hydroxide, Cu 2 (HO) 2 or Cu 2 O,H 2 O. If the mixture be heated,
the precipitate is converted into the red cuprous oxide. The
reduction of a cupric salt with the precipitation, the red cuprous
62 Smaller Cftemical Analysis
oxide, is brought about by many organic substances. Thus,
if KHO be added to a solution of CuSO 4 in the presence of
grape sugar, the Cu(HO) 2 first precipitated dissolves in excess
of KHO to a blue solution. If this be gently warmed, a bright
red precipitate of cuprous oxide Cu 2 O is obtained. 1
NH 4 HO gives no precipitate, but forms a soluble compound
having the composition Cu 2 Cl 2 ,2NH 3 . The solution is colour-
less, but rapidly absorbs oxygen from the air, first becoming
brown, and finally depositing a greenish precipitate of cupric
oxychloride, CuCL 2 ,3CuO,4H 2 O.
SEPARATION OF THE METALS OF GROUP II., SUBDIVISION i
Acidify the solution with a small quantity of dilute HC1. 2
and pass sulphuretted hydrogen through the liquid until
precipitation is complete. The precipitate is then examined
according to Table HA. on the next page.
1 This reaction is utilised as a test for sugar. The reaction is more
delicate when an alkaline solution of cupric tartrate (Fehling's solution) is
employed.
2 In the ordinary course of analysis HC1 will already have been added
for the separation of Group I., and most of the lead will have been pre-
cipitated. When the exercise is confined to Group II., if the addition of
HC1 gives a white precipitate, it should be filtered off and examined
separately for lead.
Group II. Subdivision
is I
'"*' o c
T3 i; C
tfl ^3
rt .2
<u *
^4 fl >i
^ S.S
gS^H
s .
"Ill
1^2
j X-
^lli
^ S o
liK
llll
.
P* rt cj
c3 .
Si *
12
cO
3w
!
S-o
s a
a g
s5 S
tJJ **
ffi fl
o o
s
'3 -^
t
S-f-a
o ^ _
^S^u
U
^ -S
S 4.--
_c **
T3 .'bJDvMO
:^iK
a |5Sg
T3 2 ts^
fe fc*.
rt ^ ^
*-. O i ,Q
fllii
&HG S ">'&
ifofl
?l!|i
-| ^I'o
S$I11
2 - 9> s
a ^ J ^
g
64 Smaller Chemical Analysis
SUBDIVISION 2
Arsenic, As
DRY REACTIONS. Compounds of arsenic, when heated
upon charcoal with Na 2 CO 3 and KCy, are reduced; but the
metal, being extremely volatile and readily combustible, is for
the most part burnt to arsenious oxide, As 4 O 6 , which passes
off as a white fume. At the same time some of the vapour of
the element itself is carried away with the fumes of the oxide,
and is readily recognised by its characteristic garlic-like odour.
The reduction may be made by heating the arsenic com-
pound in a glass tube with KCy, or a mixture of Na 2 CO 3 and
KCy. The reaction is conveniently studied by using arsenious
oxide. A small fragment (about the size of a pin's head) is
placed in a narrow test-tube 2 and covered by adding a mixture
of N^COg and KCy (equal parts), the materials being as dry
as possible. The total quantity of material in the tube should
not occupy more space than is shown in Fig. 9. On the
FIG. 9.
application of a gentle heat, the first effect will be the expul-
sion of moisture from the imperfectly dried materials, which
1 For such experiments small test-tubes 4 X ^ g inches answer admirably.
Bulb tubes are neither necessary nor desirable.
Group II. Subdivision 2 65
condenses upon the sides of the tube. This may be driven
up the tube by gently warming it, and finally removed by
introducing a " spill " of blotting-paper. When no more mois-
ture collects, the mixture may be steadily heated in the tip of
a small Bunsen flame. The arsenic sublimes upon the tube as
a metallic mirror. Sufficient of the vapour escapes condensation
to enable the strong garlic odour to be detected.
The reaction which takes place may be expressed by the
equation
As 4 O 6 + 6KCy = As 4 + 6KOCy
WET REACTIONS. All the salts of arsenic are such as
contain this element in the acidic or negative portion of the
molecule ; such, for example, as the arsenites and arsenates of
various metals. No oxysalts of arsenic are known in which
the element plays the part of a base.
(a) Arsenious compounds, derived from arsenious oxide,
As 4 O 6 . The arsenites of sodium, potassium, and ammonium
alone are soluble in water. For the following reactions, potas-
sium arsenite, K 3 AsO 3 , or a solution of As 4 O 6 in dilute HC1,
may be employed.
H. 2 S or (NH 4 ) 2 S precipitates from slightly acid solutions
yellow arsenious sulphide, As 2 S 3 . Soluble in excess of (NH 4 ) 2 S,
forming ammonium thio-arsenite, (NH 4 ) 3 AsS 3 .
It dissolves also in caustic alkalies, ammonia, and am-
monium carbonate, yielding a mixture of arsenite and thio-
arsenite, e.g.
As 2 S 3 + 6KHO = K 3 AsO 3 + K 3 AsS 3
On addition of an acid to such solutions, arsenious sulphide
is precipitated.
When As 2 S 3 is dissolved in yellow ammonium sulphide,
ammonium thio-arsen^/^ is formed, and from this solution
acids precipitate As 2 S 5 .
Arsenious sulphide is practically insoluble in HC1 (contrast
S6%S 3 ), but readily dissolves in HNO 3 or in HC1 with the
addition of a crystal of KC1O 3 .
CuS0 4 produces, in a solution of potassium arsenite, a
F
66 Smaller Chemical Analysis
green precipitate of hydrogen cupric arsenite, HCuAsO 3
(Scheetts green) ; soluble in ammonia and caustic alkalies.
If the solution be boiled, the arsenite is oxidised to arsenate
and cuprous oxide precipitated.
AgN0 3 gives a pale-yellow precipitate of silver arsenite,
Ag 3 AsO 3 , soluble both in NH 4 HO and in HNO 3 . When the
ammoniacal solution is boiled for some time, metallic silver
is precipitated, and the arsenite is oxidised to arsenate.
Precipitation by Copper (Reinsch's test). If a strip of
clean copper foil be introduced into a solution of arsenious
oxide in HC1, or an arsenite acidified with the same acid, and
the mixture be warmed, metallic arsenic is deposited upon the
copper, at the same time uniting with it, forming copper
arsenide, Cu 5 As 2 . If the copper be then dried, and gently
heated in a dry test-tube, the arsenic will be volatilised, and
at the same time oxidised, giving, therefore, a white crystalline
sublimate of As 4 O 6 (contrast antimony).
(b) Arsen/c compounds, derived from arsenic pentoxide,
As 2 O 5 . Arsenates of sodium, potassium, and ammonium are
soluble in water.
H 2 S. From acidified solutions of an arsenate, H 2 S gives
a yellow precipitate after a short time, which is either As 2 S 5 or
a mixture of As 2 S 3 and S, depending upon conditions.
If the solution is strongly acid, and the gas is passed
rapidly, the precipitate which slowly comes down is the penta-
sulphide. On the other hand, if the solution is less strongly
acid, and the H 2 S is passed slowly, the arsenic acid is first
reduced to arsenious acid, with deposition of sulphur, and the
arsenious acid as it forms is converted into arsenious sulphide
thus
(1) H 3 AsO 4 + H 2 S = H 2 O + S + H 3 AsO 3
(2) 2 H 3 AsO 3 + 3H 2 S = As 2 S 3 + 6H 2 O
This reducing action of H 2 S is very slow, therefore complete
precipitation requires considerable time. Warming the liquid
hastens the action. The addition of a more powerful reducing
agent, such as sulphurous acid, produces the effect at once.
As 2 S 5 dissolves in alkaline sulphides and in caustic alkalies,
Group II. Subdivision 2 67
forming thio-arsenates, from which the sulphide is repreci-
pitated by acids
2 K 3 AsS 4 + 6HC1 = 6KC1 + As 2 S 5 + 3 H 2 S
CuS0 4 gives a pale bluish precipitate of hydrogen cupric
arsenate, HCuAsO 4 . Soluble, like the corresponding arsenite,
in ammonia; but the copper is not reduced on heating the
solution, for the reason that arsenates are incapable of further
oxidation in other words, they do not act as reducing agents
(contrast ar smites).
AgN0 3 produces a chocolate-coloured precipitate of silver
arsenate, Ag 3 AsO 4 , which, like the arsenite, dissolves in NH 4 HO
and in HNO 3 . The ammoniacal solution is not reduced on
boiling, for the same reason that the copper salt is not reduced
(contrast ar smites).
MgS0 4 , in presence of NH 4 C1 and NH 4 HO, gives a white
crystalline precipitate of ammonium magnesium arsenate,
(NH 4 )MgAsO 4 ; practically insoluble in water. (This reaction
serves to distinguish an arsenate from an arsenite.)
Marsh's Test
In the presence of nascent hydrogen, both arsenic and
arsenious compounds are reduced, and arsine (or arsenuretted
hydrogen), AsH 3 , is evolved. Thus, if a solution of arsenious
or arsenic oxide be subjected to electrolysis, or if such solutions
are introduced into a mixture from which hydrogen is being
generated (e.g. zinc or magnesium with dilute acid), this com-
pound is produced. The properties of arsenuretted hydrogen
which are made use of in analysis are the following :
(1) The deposition of metallic arsenic from the flame of
the burning gas when a cold object is depressed upon the
flame.
(2) The decomposition of the compound on passing through
a heated tube, with deposition of an arsenical mirror.
(3) The action of the gas upon a solution of silver nitrate,
resulting in the precipitation of metallic silver.
The reaction is made in a small hydrogen generating
68
Smaller Chemical Analysis
apparatus, in which hydrogen is slowly generated from zinc
and dilute sulphuric acid, both materials being free from
arsenic. The issuing gas is passed through a piece of com-
bustion tube which has been drawn out so as to produce one
or two constricted places
in its length, as shown
^| in Fig. 10. As soon as
dp the air is all expelled
from the apparatus, the
issuing hydrogen is in-
flamed. 1
A small quantity of
the arsenical solution is
now introduced through
the thistle-tube. The
first effect of this is to
suddenly cause a greatly
increased rate of evolu-
tion of hydrogen. 2 The
colour of the hydrogen
flame will be seen to
change, and to assume a lilac tint, and at the same time white
fumes of As 4 O 6 escape from the tip of the flame. If now a
porcelain dish be depressed upon the flame, a rich brown-black
metallic-looking stain will be deposited. The deposit being
volatile, and the flame very hot, the stain will again disappear
if the flame be allowed to impinge for more than a moment
or two on the same spot.
If the drawn-out tube be heated near one of the constrictions,
the arsenuretted hydrogen will be decomposed as it passes the
FIG. 10.
1 A small test-tube should be filled by upward displacement, and tested
by a flame before igniting the gas at the exit-tube of the apparatus. As an
additional precaution, it is well to throw a duster lightly over the bottle
before applying a light, so that, should an explosion happen, the broken
glass will be prevented from flying about.
2 On this account, it is necessary that the generation of hydrogen before
adding the arsenic solution should be quite slow ; and also that the quantity
of the arsenic solution added at a time should be small.
Group II. Subdivision 2 69
hot spot, and an arsenic mirror will be deposited in the tube.
Minute traces of arsenic can be detected in this way.
It will be noticed that the deposition takes place entirely
on that part of the tube which is on the side of the flame
farthest from the generating vessel (antimony is deposited from
its hydride on both sides of the heated spot).
Since antimony also forms a gaseous compound with
hydrogen which gives similar stains, it is necessary to employ
further confirmatory tests.
i. The arsenic stains are readily dissolved by a solution of
a hypochlorite. If, therefore, a solution of bleaching powder
be poured over such stains they immediately disappear
SCa(OCl), + 6H 2 O + As 4 = 5CaCl 2 + 4H 3 AsO 4
Antimony stains do not dissolve in hypochlorite solutions.
2. When a stream of H 2 S is passed through the tube
containing the deposit of either arsenic or antimony, slightly
warmed, in each case the sulphide is formed. Yellow arsenious
sulphide volatilises along from the warm region and condenses
on the cold distant part of the tube; antimonious sulphide,
reddish or nearly black, remains unmoved, being non-volatile.
(If present together, they can in this way be separated.)
If a stream of gaseous HC1 be now passed through the
tube, antimonious sulphide is converted into antimonious
chloride, which passes on with the HC1, and may be led
into water and again precipitated as the red sulphide with
H. 2 S. The yellow arsenious sulphide remains in the tube,
being unattacked by HC1.
3. Arsenuretted hydrogen can also be distinguished from
the antimony compound, by the difference in the behaviour
of the two gases towards silver nitrate. When passed into the
silver solution, each gas produces a black precipitate. In
the case of arsenic this consists of metallic silver, while with
antimony it consists of antimonide of silver ; thus
6AgN0 3 + sH 2 + AsH 3 = 3 Ag 2 + 6HNO 3 + H 3 AsO 3
3AgN0 3 + SbH 3 = SbAg 3 + sHNO 3
On filtering, the arsenic is found in the filtrate, while the antimony
70 Smaller Chemical Analysis
would be in the precipitate. If the filtrate is neutralised by
the cautious addition of ammonia, a yellow precipitate of
silver arsenite is produced by interaction with the excess of
silver nitrate present.
The antimony in the black silver antimonide may be
detected by boiling with a solution of tartaric acid. The liquid
thus obtained is acidulated with HC1, and sulphuretted hydrogen
passed into it, which gives a precipitate of red antimonious
sulphide.
Fleitmann's Test
When an arsenite, or a solution of arsenious oxide, is
warmed in a test-tube with a solution of sodium hydroxide and
metallic zinc, arsenuretted hydrogen is evolved, which can be
detected by means of a piece of filter-paper moistened with
silver nitrate held to the mouth of the tube. A black stain of
precipitated silver is produced. Antimoniuretted hydrogen is
not produced from antimony compounds under similar con-
ditions.
Antimony, Sb
DRY REACTIONS. Antimony compounds may be reduced
to metallic antimony by heating them with Na 2 CO 3 and KCy
upon charcoal. Globules of the metal are thus obtained, which
burn in the blowpipe flame, producing white fumes of anti-
monious oxide, Sb 4 O 6 . The charcoal at the same time receives
a white incrustation. The bead of metal will be found to be
very brittle, and, when broken, to exhibit a highly crystalline
appearance. Antimony is unacted upon by dilute HC1 or
H 2 SO 4 . Nitric acid oxidises it into antimonic acid, or anti-
monious oxide, depending upon conditions of concentration.
WET REACTIONS. Antimony forms two series of com-
pounds, " antimonious " and " antimonic," which may be
regarded as being derived respectively from the two oxides,
antimonious oxide, Sb 4 O 6 , and antimony pentoxide, Sb. 2 O 5 .
(a) Antimonious Compounds. Antimonious oxide is feebly
basic, forming salts in which the metal constitutes a part of the
positive radical. Of these salts the tartrate, (SbO).XC 4 H 4 O 6 ), and
Group II. Subdivision 2 71
the double potassium tartrate (tartar emetic), (SbO)K(C 4 H 4 O 6 ),
are the most familiar. For the following reactions l an acid
(HC1) solution of antimonious chloride may be employed :
KHO, NaHO, NH 4 HO, as well as alkaline carbonates, pre-
cipitate antimonious oxide, Sb 4 O 6 .
The precipitate redissolves in excess of either potassium or
sodium hydroxide, forming the respective metantimonites
Sb 4 O 6 + 4NaHO = 4NaSbO 2 + 2H 2 O
H 2 0, added in considerable quantity to the acid solution of
SbCl 3 , gives a white precipitate of an oxychloride, SbOCl. The
precipitate dissolves in tartaric acid forming antimonyl tartrate,
(SbO) 2 (C 4 H 4 O 6 ) (distinction from bismuth).
H 2 S and (NH 4 ) 2 S give a red or orange-red precipitate of
antimonious sulphide, Sb 2 S 3 , soluble in excess of ammonium
sulphide forming ammonium thio-antimonite. It dissolves also
in caustic alkaline, and from all these solutions Sb 2 S 3 is repre-
cipitated on addition of HC1.
When dissolved in yellow ammonium sulphide it forms
ammonium thio-antimontf/<?, and from this solution acids pre-
cipitate Sb 2 S 5 .
Antimonious sulphide is not dissolved by ammonia or by
ammonium carbonate (contrast arsenic).
Antimonious sulphide is decomposed by hot hydrochloric
acid, with evolution of sulphuretted hydrogen (contrast arsenic)
SbA + 6HC1 = 2SbCl 3 + sH 2 S
(b) Antimonic Compounds. These are derived from the
pentoxide, Sb 2 O 5 , the pyro-antimonates (e.g. K 4 Sb 2 O 7 ) and met-
antimonates (e.g. KSbO 3 ) being the most familiar salts.
Potassium pyro-antimonate is readily soluble in water, while
the sodium salt is difficultly soluble, hence the potassium com-
pound is used as a reagent for sodium (p. 23).
H.J3 and (NH 4 ) 2 S produce with acidified solutions of anti-
monic compounds an orange-red precipitate consisting of
Sb 2 S 5 , Sb 2 S 3 , and S in varying proportions.
1 Except that with AgNO a , in which obviously the presence of chlorine
would interfere.
72 Smaller Chemical Analysis
The precipitate dissolves in alkaline sulphides and caustic
alkalies, and from these solutions Sb 2 S 5 is reprecipitated by
HC1.
Precipitation of Metallic Antimony. If one or two drops
of an antimony solution acidified with HC1 are placed upon a
piece of clean platinum foil, and a small fragment of zinc im-
mersed in the liquid, a black stain is produced upon the platinum
by the deposition of metallic antimony. HC1 has no action
upon the deposit (distinction from tin).
Antimoniuretted hydrogen is evolved when antimony
compounds are acted upon by nascent hydrogen. The com-
pound undergoes reactions similar to those of the correspond-
ing arsenic compound. The methods for distinguishing between
them are described under arsenic.
Tin, Sn
DRY REACTIONS. Compounds of tin are reduced to the
metallic state by being heated on charcoal with Na 2 CO 3 and
KCy in the reducing flame. A portion of the metal is oxidised
by the flame, and produces a white incrustation of SnO 2 upon
the charcoal ; this, on being moistened with cobalt nitrate, and
reheated, assumes a greenish appearance.
The beads of reduced metal are malleable (therefore easily
distinguished from Bi or Sb), but are not soft enough to mark
paper in the manner of lead.
The metal dissolves in hot strong HC1, forming stannous
chloride, SnCL Strong nitric acid converts it into white in-
soluble metastannic acid, (H 2 SnO 3 ) 5 .
WET REACTIONS. Tin forms two classes of compounds,
distinguished as " stannous " and " stannic," derived respectively
from stannous oxide, SnO, and stannic oxide, SnO a .
(a) Stannous Compounds. Of these the chloride, sulphate,
and nitrate are soluble in water.
HKO, NaHO, NH 4 HO, as well as alkaline carbonates, give
with stannous chloride a white precipitate of hydrated stannous
oxide (basic hydroxide) j thus
4KHO = 2SnO,H 2 O + 4KC1 + H 2 O
Group II. Subdivision 2 73
The precipitate is soluble in excess of KHO or NaHO (but not
in the other precipitants), forming alkaline stannites ; thus
2SnO,H 2 O + 4KHO = 2KaSnO a + 4H 2 O
H 2 S or (NH 4 ) 2 S gives, with dilute solutions of stannous
chloride, a deep brown precipitate of stannous sulphide;
soluble in caustic alkalies, but reprecipitated on acidification.
Soluble also in yellow ammonium sulphide forming the thio-
Stannous sulphide is insoluble in colourless ammonium
sulphide and in ammonium carbonate. Boiling HC1 converts
it into SnCl-2 ; while aqua regia oxidises it to stannic chloride,
SnCl 4 .
Oxidation of Stannous Compounds. These substances
readily pass, by oxidation, into stannic compounds ; they there-
fore act the part of powerful reducing agents in a number of
reactions, of which the following are important :
(1) Mercuric chloride, see Reactions for Mercury, p. 56.
(2) Ferric salts are reduced to the " ferrous " state thus,
ferric sulphate, in the presence of hydrochloric acid, gives
ferrous sulphate and stannic chloride
Fe 2 (SO 4 ) 3 -f SnCl 2 + 2HC1 - 2FeSO 4 + H 2 SO 4 + SnCl 4
(3) Gold chloride, in the presence of acid, is reduced to
metallic gold
2 AuCl 3 + 3SnCl 2 = 3SnCl 4 + 2Au
In dilute neutral solutions, a reddish or purple coloration
is produced, known as purple of Cassius.
(b] Stannic Compounds. Stannic oxysalts, such as the
nitrate or sulphate, are unstable, being converted by water into
insoluble metastannic acid ; and as stannic oxide, SnO 2 , is in-
soluble in HC1, a stannic solution for the following reactions is
best obtained by oxidising an HC1 solution of stannous chloride
either by adding bromine water to it, or by warming with a
crystal of KC1O 3 .
HKO, NaHO, NH 4 HO, as well as alkaline carbonates, give
74 Smaller Chemical Analysis
a white precipitate of hydrated stannic oxide, or stannic acid,
SnO 2 ,H 2 O, or H 2 SnO 3 ; thus
SnCl 4 + 4NaHO = H 2 SnO 3 + 4NaCl + H 2 O
Soluble in HNO 3 and in HC1. Soluble in KHO and NaHO,
with formation of the respective stannates, K 2 SnO 3 and
NaaSnOsj thus
H 2 SnO 3 + 2NaHO = Na 2 SnO 3 + 2H 2 O
H 2 S or (NH 4 ) 2 S precipitates yellow stannic sulphide, SnS 2
(with H 2 S the precipitate appears nearly white at first, and is
only complete in dilute solutions). The precipitate is soluble
in caustic alkalies, in ammonium sulphide, and sulphides of the
alkalies, forming thio-stannates ; thus
3 SnS 2 + 6KHO = 2KnSs + K 2 SnO 3 + 3H 2 O
SnS 2 + (NH 4 ) 2 S = (NH 4 ) 2 SnS 3
From these solutions the yellow stannic sulphide is repre-
cipitated on the addition of HC1
(NH 4 ) 2 SnS 3 + 2HC1 = 2NH 4 C1 + SnS 2 + H 2 S
Stannic sulphide is insoluble in ammonium carbonate, but
dissolves in hot strong HC1.
Precipitation of Metallic Tin. When zinc is immersed in
an acid solution of stannous or stannic chloride, the tin is pre-
cipitated as a grey-black deposit upon the zinc ; or, if the whole
of the zinc becomes dissolved, the tin is left as a scaly powder.
It dissolves in warm HC1 (contrast antimony].
SEPARATION OF GROUP II. INTO SUBDIVISIONS i AND 2
The solution, if neutral or alkaline, is acidified with HC1, 1
and sulphuretted hydrogen passed through until the precipita-
tion of all the metals of Group II. is complete. The precipitate
is washed and is then transferred to a small beaker, and gently
1 If the solution under examination is alkaline, it may contain thio-salts
of As, Sb, or Sn ; the addition of HC1 will result in the precipitation of
the sulphides of these metals. If it is neutral, basic salts of antimony might
be precipitated at first, but redissolve on warming with a slight excess of the
acid.
Separation of Grotip II 75
warmed with yellow ammonium sulphide for a few minutes.
The liquid is then filtered. The residue consists of the undis-
solved sulphides of the metals of subdivision i, the filtrate con-
tains the thio-salts of the metals of subdivision 2.
1 Ammonium sulphide dissolves CuS to a slight extent (see Reactions,
p. 61), hence, if this element is present, a small quantity of it may find its
way into the solution along with As, Sb, and Sn.
7 6
Smaller Chemical Analysis
S o
W o
o oj Q
fjiw
I
.2 2 *e -S
S ^ 2.2
1 a a
d 3 .2
*3 . d - 1 -
rt ^2 '13 T3
,co 4S S
S ^
U
22 cs is
o *>
1^1
-Ma
I! Hi
S, I an
i ^ ^
"S c
:.S"S
500,
4-> O
^ G 2
/*\ r- ^
E
So
a
ii
II"
T3 c/3
1,0
ion.
ragm
ent,
der o
*<<3~i
U "^
add
is p
rem
g
12
b.2 ^
S-g
* 8 J
1-1
_ 'o o
U ^,co
ffi w g
fl P
cj
CO fl HH
O
CJ
2 -
'S S
CHAPTER VIII
REACTIONS OF THE METALS OF GROUP 2
Silver, Ag
DRY REACTIONS. Compounds of silver, when heated on
charcoal with sodium carbonate in the reducing flame, yield
metallic silver ; which, being non-oxidisable, is not accompanied
by any oxide incrustation upon the charcoal. The metal, how-
ever, is slightly volatile in the blowpipe flame, and sometimes
a faint red-brown incrustation is thus obtained.
The reduced metal may be removed to a watch-glass, dis-
solved in nitric acid, and precipitated as chloride.
WET REACTIONS. Of the common salts of silver, the nitrate
is readily soluble, the acetate and sulphate sparingly soluble, in
water.
HC1, and soluble chlorides, give a white curdy precipitate
of silver chloride, AgCl, which, on being warmed or stirred,
becomes granulated in appearance, and very quickly settles.
On exposure to light, the white compound assumes a slate
colour or drab tint, which gradually deepens to a violet, and
finally appears brown or black.
Silver chloride is quite insoluble in water and in nitric acid.
It is soluble to a slight extent in strong HC1, but reprecipitated
completely on dilution. It readily dissolves in ammonia, form-
ing the compound 2AgCl,3NH 3 , but reprecipitated on the
addition of nitric acid.
Silver chloride is soluble also in KCy, being first converted
into siver cyanide, which dissolves in excess of KCy, forming
the double cyanide KCy,AgCy. It also dissolves in sodium
77
78 Smaller Chemical Analysis
thiosulphate, with the formation of a double thiosulphate ;
thus
AgCl + Na 2 S 2 O 3 = NaCl + NaAgS 2 O 3
[Reactions with bromides, iodides, and cyanides are de-
scribed under the respective acids.]
KHO, NaHO, or NH 4 HO gives a greyish-black precipitate
of silver oxide, Ag 2 O. Insoluble in excess of the caustic
alkalies, but readily soluble in ammonia.
H 2 S or (NH 4 ) 2 S produces a black precipitate of silver sul-
phide, Ag 2 S. Insoluble in dilute acids, except boiling dilute
nitric acid, which converts it into nitrate.
Silver sulphide is insoluble in ammonia, ammonium sulphide,
or potassium sulphide.
Reduction of Silver Salts. Silver compounds are readily
reduced to the metallic state ; for example, if silver chloride is
placed in a little dilute sulphuric acid, and strip of zinc intro-
duced, the nascent hydrogen converts the white chloride into
grey metallic silver. The reduction is complete when a particle
of the grey solid dissolves completely in nitric acid.
Lead, Mercury
The reactions of these metals have already been considered
in connection with the metals of Group II., Subdivision i,
PP- 54, 57-
Separation of the Metals of Group I 79
PQ w
<1 W
H H
K
o
II
H
^ g
'
S ^K
td 1
{i-i tuO
1 !
II
X
* s 4
C o
;i
1 r|
TJ "
^ .a w ^
*"
!f
r
1
. -
1 -a
o w
"S 4)
1,8
I
>
a&
n . ^j
oJ
1
i s *
8| I
8o
Smaller Chemical Analysis
GENERAL TABLE FOR THE SEPARATION
To the solution of the substance under analysis (see pp. 112, et seq.) add a few
ployed in effecting the solution of the substance, this first step in the general
reagent gradually until the precipitation is complete. Warm gently
The Precipitate
may consist of
AgCl
H g2 Cl 2
PbCl 2
Group I.
Examine by
Table I.
(P- 79).
The Filtrate is gently warmed, and a stream of sulphuretted
cipitation is complete (watch the precipitation carefully,
A small portion of the mixture should be filtered, and the
etted hydrogen passed into it. Should any further pre-
filtered, and treated to more gas.
The Precipitate may consist
of
(i)PbS;HgS;Bi 2 S 3 ;CuS;
CdS.
(2) Sb 2 S 3 ; As 2 S 3 ; SnS ;
SnS 2 .
Wash thoroughly, and then
transfer it to a small beaker
and warm gently with
yellow ammonium sul-
phide. See Note 3.
The Besidue
may contain
the sulphides
of Division I.
Examine by
Table IlA.
(P- 63).
The Filtrate
may contain
the thio salts
of As, Sb,
Sn.
Examine by
Table IlB.
(p. 76).
The Filtrate. Boil until sul-
and boil again for a few
the sulphuretted hydrogen),
the extent to which it has
liminary tests, the evapora-
to destroy the latter, and
Test a small portion of the
To the main portion of the
NH 4 HO until precipitation
The Precipitate.
I. In the absence of phos-
phoric acid, may consist
of
A1 2 (HO) 6 ; Cr 2 (HO) 6 ;
Fe 2 (HO) 6 .
Examine by Table IIlA.
(P. 41).
II. In the presence of phos-
phoric acid, besides the
above hydroxides, the pre-
cipitate may contain the
phosphates of any or all of
the metals of Groups III.
and IV., and of magnesium.
Examine by Table IIIc.
(P- 52).
Separation of Metals into Groups
81
OF METALS INTO GROUPS
drops of dilute HC1. [Obviously in cases when HC1 has already been em-
separation is omitted.] If any precipitate is produced, continue adding the
(see footnote, p. 74), then thoroughly cool again, and filter.
hydrogen slowly bubbled through the liquid, with frequent stirring, until pre-
and note any colour changes). See Note 2, following page.
nitrate diluted with two or three times its volume of water, and more sulphur-
cipitation result, the main portion must be similarly diluted, without being
phuretted hydrogen is entirely expelled. Add two or three drops of HNO 3 ,
moments (to oxidise any iron or chromium which may have been reduced by
and then evaporate the liquid to about half its volume, or less, according to
become diluted. If silica or organic compounds have been detected by pre-
tion must be carried down to dryness, and the residue gently heated in order
render the silica insoluble. The residue is then extracted with HC1 and water,
solution for phosphoric acid (see p. 51).
solution add a considerable quantity of NH 4 C1, and heat to boiling. Add
is complete (see p. 40), boil for a moment, and filter while hot.
The Filtrate. Pass sulphuretted hydrogen (or add ammonium sulphide) until
precipitation is complete. Gently warm the liquid (see Ni, p. 47), and filter.
The Precipitate
may consist of
MnS ; ZnS ; NiS
CoS.
Examine by
Table IIlB.
(P- So).
The Filtrate. Boil to expel H 2 S. If ammonium sulphide
has been employed, add a little HC1 before boiling. If
necessary, concentrate the solution by evaporation. Add
NH 4 HO until alkaline, and (NH 4 ) 2 CO 3 until precipitation
is complete. Warm the liquid, but do not boil (see p. 30).
The Precipitate
may consist of
BaCO 3 ;SrCO 3 ;
CaC0 3 .
Examine by
Table IV.
(p. 3D-
The Filtrate.
Examine for Mg, K, Na, by Table V.
(P- 27).
NOTE i. In testing for Mg at this point, it
will obviously be unnecessary to add more
NH 4 C1 and NH 4 HO.
NOTE 2. Any failure to effect complete group
separations will usually result in the precipita-
tion of some metallic phosphate at this stage,
other than magnesium phosphate
G
82 Smaller Chemical Analysis
NOTES ON THE GENERAL TABLE OF SEPARATION
1. Take great care to ensure complete precipitation in every group
separation, otherwise the object of the separation is defeated. For the
same reason every group precipitate should be thoroughly washed to free it
from adhering solution which contains groups that follow. Such washings
need not be mixed with the first filtrates, otherwise the liquid becomes too
much diluted.
2. To ensure complete precipitation here it is necessary that the solution
be not too strongly acid (hence unnecessary excess of HC1 in precipitating
Group I . must be avoided). See Cd, p. 60 ; also As, p. 66.
3. Copper sulphide is slightly soluble in ammonium sulphide (see p. 61).
An alternative reagent which may be employed is caustic soda ; this does
not dissolve CuS, but it dissolves HgS more freely, hence cannot be used
unless mercury is known to be absent ascertained by the preliminary
tests. When both copper and mercury are present ammonium sulphide
should be used.
CHAPTER IX
THE NEGATIVE OR ACID RADICALS
THESE are the negative ions (anions) which are produced when
the acids (hydrogen salts) or salts (metallic salts) undergo dis-
sociation when dissolved in water. Thus, the negative radical
in hydrochloric acid or in metallic chlorides is the chloride ion,
Cl; the tests for chlorides, therefore, are tests for this ion.
Similarly, sulphates and nitrates dissociate into their positive
ions, and the negative ions SO 4 and NO 3 respectively ; the tests
for sulphuric and nitric acids are thus, in reality, tests for these
negative ions, although, in ordinary language, we often speak
of them as tests for the various adds from which these ions are
derived. The negative radicals are classified into groups on
the basis of their behaviour towards certain chosen reagents,
but these reagents are not employed as group-reagents to separate
one group of acid radicals from another, but are merely used
in order to discover, by a single operation, the absence or
otherwise of an entire group, whereby the necessity for applying
a number of separate tests may be obviated.
The acids which will be included in this section are the
following :
Hydrochloric acid, Chloric acid.
Hydrobromic acid.
Hydriodic acid.
Hydrofluoric acid.
Sulphuretted hydrogen, Sulphuric acid, Sulphurous acid.
Nitric acid, Nitrous acid.
Phosphoric acid.
Carbonic acid.
Silicic acid.
Boric acid.
83
84 Smaller Chemical Analysis
Certain acids, such as Arsenious, Arsenic, Chromic, Per-
manganic, have already been discussed under their respective
metals.
Hydrochloric Acid and Chlorides
Hydrogen chloride is a colourless gas having a sharp
choking smell. It fumes in contact with moist air, is strongly
acid, but has no bleaching properties. It is extremely soluble
in water, the solution constituting the ordinary reagent, hydro-
chloric acid.
The chlorides are all soluble in water, except those of the
metals of Group I. (PbCl 2 being soluble in hot water), and
certain others which are decomposed by water, such as the
chlorides of antimony, bismuth, and tin.
Silver nitrate, AgNO 3 , gives, in solution of chlorides or
hydrochloric acid, a white precipitate of silver chloride. In-
soluble in nitric acid. Readily soluble in ammonia, even dilute
(for further properties, see Silver reactions, p. 78).
AgCl is distinguished from either AgBr or Agl by the fact
that chlorine water is without action upon it (see Bromides and
Iodides). It may also be distinguished in the following way :
If the washed precipitate of AgCl be mixed with a little very
dilute sulphuric acid, and a strip of zinc placed in the mixture,
the silver chloride turns grey, owing to its reduction to metallic
silver, while zinc chloride passes into solution. This, on treat-
ment with manganese dioxide and sulphuric acid, will yield
chlorine.
Fusion with sodium carbonate converts AgCl into metallic
silver and sodium chloride. On treatment with water, chlorine
can be liberated from the solution, as in the foregoing.
Liberation of Chlorine from Chlorides. The chloride is
mixed with MnO 2 and H 2 SO 4 , and the mixture gently warmed
in a test-tube. The chlorine which escapes may be detected
by its characteristic smell and by its bleaching properties
(litmus paper, or, better, paper coloured red by an alkaline
solution of carmine, may be used). Small quantities may be
detected by fitting the test-tube with a cork and delivery tube,
and passing the evolved gas into water in a second tube. The
Hydrochloric Acid and Chlorides 85
presence of free chlorine in the water may be detected by
adding a few drops of Kl solution and then starch paste. A
blue coloration results from the liberated iodine (set free by
the chlorine) uniting with the starch.
Liberation of Hydrogen Chloride. When chlorides (except
those of tin, lead, mercury, and silver) are gently heated with
strong H 2 SO 4 , hydrogen chloride is evolved.
The presence of hydrochloric acid in a solution containing
a soluble chloride may be detected by gently warming the
liquid with MnO 2 (without sulphuric acid), when chlorine is
evolved, which may be detected as described above
4 HC1 + Mn0 2 = MnCl 2 + 2H 2 O + C1 2
Formation of Chromyl Chloride. When a mixture of a
chloride and potassium dichromate is gently warmed with
strong sulphuric acid, a red-brown vapour is disengaged (resem-
bling bromine in colour, but very different in smell) consisting
of chromyl chloride, CrO 2 Cl 2
4 KC1 + K 2 Cr 2 7 + 3H 2 S0 4 = 3 K 2 SO 4 + 3H 2 O + 2CrO 2 Cl 2
If the reaction be made in a test-tube fitted with a delivery
tube, and the vapour of the chromyl chloride be passed into a
second test-tube containing an alkaline hydroxide, a chromate
of the alkali is formed
CrOaCla + 4NH 4 HO = (NH 4 ) 2 CrO 4 + 2 NH 4 C1 + 2H 2 O
The presence of the chromate is indicated by the yellow
colour which the liquid assumes, which may be confirmed by
acidifying with acetic acid and adding lead acetate. The
presence of the chromate is proof of the presence of a chloride
in the first test-tube.
By means of this test it is possible to detect a chloride in
the presence of either a bromide or iodide, 1 as neither bromine
nor iodine form similar chromyl compounds.
1 The former tests (p. 84), which enable one to distinguish between a
chloride, bromide, and iodide, will not be confounded with a test such as
86 Smaller Chemical Analysis
Hydrobromic Acid and Bromides
Gaseous hydrobromic acid closely resembles hydrochloric
acid. The properties of the gas are not used in analysis.
All bromides are soluble in water, except mercurous bromide
and silver bromide; lead bromide dissolves in boiling water
less easily than the chloride.
Silver nitrate, AgNO 3 , precipitates from solutions of bro-
mides or hydrobromic acid, pale-yellow silver bromide, AgBr
(the colour is indistinguishable from white by gaslight). It is
insoluble in nitric acid, and difficultly soluble in ammonia
(scarcely soluble in dilute ammonia. Contrast AgCl).
AgBr may be distinguished from AgCl by shaking up a little
of the washed precipitate with a few drops of carbon disulphide
and chlorine water.
Silver bromide is decomposed by metallic zinc in the
presence of dilute sulphuric acid, in the same manner as the
chloride. Zinc bromide goes into solution, from which the
bromine can be separated by either of the methods given
below.
Prolonged boiling with a strong solution of sodium car-
bonate (or, better, heating the dry substances strongly in a
glass tube) decomposes silver bromide. On filtering (after
extraction with water in the case of the dry reaction), the
aqueous solution containing sodium bromide may be tested as
above.
Liberation of Bromine from Bromides. When gently
warmed with MnO 2 and H 2 SO 4 , bromides evolve bromine,
which escapes as a brown-red vapour having an irritating smell,
condensing on a cold surface to dark brown-red drops of liquid.
In contact with starch, it gives a yellow colour ; if, therefore,
the reaction is made in a small beaker which is covered with a
piece of moistened filter-paper upon which a starch flour is
dusted, this yellow colour is produced.
Bromine is also liberated by means of chlorine. By adding
the above, which permits of the detection of one class of salts in the presence
Bothers.
Hydriodic Acid and Iodides 87
chlorine water to a solution of a bromide, the liquid becomes
brownish ; and if a little carbon disulphide is added and the
mixture shaken, the bromine is taken up by the disulphide,
which settles down as a brown layer at the bottom.
By this reaction a bromide can be detected in the presence of a
chloride.
When bromides (except Hg 2 Br 2 and AgBr) are acted upon
with strong sulphuric acid, bromine is liberated along with
hydrobromic acid and sulphur dioxide
KBr + H 2 SO 4 = HBr -f HKSO 4
2 HBr + H 2 S0 4 = Br 2 + SO 2 + 2H 2 O
The detection of hydrobromic acid in solution in presence
of a dissolved bromide is accomplished by gently warming the
liquid with MnO 2 . Bromine is liberated from the acid (not
from the bromide), and may be detected by the starch test.
Hydriodic Acid and Iodides
Gaseous hydriodic acid closely resembles HBr and HC1.
The properties of the gaseous compound are not utilised in
analysis.
All iodides are soluble in water except those of silver,
mercury, copper (gold, platinum, and palladium). Those of
bismuth and lead are sparingly soluble.
Silver nitrate, AgNO 3 , precipitates from solutions of iodides
or hydriodic acid a pale-yellow precipitate of silver iodide,
Agl, insoluble in HNO 3 , and more difficult of solution in
NH 4 HO than AgBr. Agl may be distinguished from either
AgBr or AgCl by shaking up the precipitate with a little CS 2
and chlorine water.
Copper sulphate, CuSO 4 , gives a dirty white precipitate of
cuprous iodide, Cu 2 I 2 , coloured by free iodine. In presence of
suitable reducing agents, such as sulphurous acid, the whole
of the iodine is precipitated as white Cu 2 I 2 . Bromides and
chlorides give no precipitate with CuSO 4 , hence by this
88 Smaller Chemical Analysis
reaction an iodide may be separated from a mixture of halogen
salts.
Liberation of Iodine from Iodides. Iodine is more easily
set free from combination than either bromine or chlorine, and
the methods which are applicable for the liberation of these
apply also in the case of iodine. Thus, manganese dioxide
and dilute sulphuric acid decompose iodides in a manner
precisely similar to that explained on p. 84 for chlorides.
Strong acids, as nitric and sulphuric, also expel iodine from
iodides, with evolution of oxide of nitrogen, or sulphur dioxide ;
2 KI + 2 H 2 S0 4 = K 2 S0 4 + 2 H 2 + S0 2 + I 2
The comparative ease with which iodine is liberated from
combination, affords the basis of most of the tests by which
this element is detected. The following are the reactions
most used in analysis :
1. Chlorine water, when added to a solution of an iodide,
expels the iodine. The test may be applied as described
under bromine, the carbon disulphide in this case being
coloured violet.
The presence of the liberated iodine may also be recognised
by means of starch. A small quantity of starch paste is mixed
with the solution of the iodide, and one or two drops of chlorine
water added, when the deep indigo-blue compound of iodine
with starch is produced. On the addition of an excess of
chlorine, the colour is destroyed. Boiling the liquid also
destroys the compound, hence, when small quantities of iodine
are being tested, it is necessary to avoid using the starch while
hot.
2. Nitrous Acid. When a solution of an iodide is acidified
with dilute sulphuric acid, and a few drops of a solution of
sodium nitrite added, the nitrous acid generated (by the action
of the acid upon the nitrite) decomposes the iodide, with the
liberation of iodine. The action is in reality between the
nitrous acid and hydriodic acid ; thus
HNO 2 + HI = H 2 O + NO + I
Neither chlorine nor bromine is liberated by nitrous acid.
Chloric Acid and Chlorates 89
Detection of Bromides and Iodides*n Solution together. When
carbon disulphide is added to a solution of an iodide and bromide
in a test-tube, and chlorine water added in small quantities at a
time, with agitation, the iodine will be liberated first. If this be
done carefully, it is not difficult to see when the further addition of
a drop of chlorine water produces no further precipitation of iodine.
At this point the carbon disulphide is coloured deep violet with dis-
solved iodine. A portion of the aqueous liquid is then withdrawn
by means of a small pipette and transferred to another test-tube.
A fresh quantity of carbon disulphide is now added to this, and a
few drops of chlorine water. If the whole of the iodine had been
liberated in the first tube, the bromine now begins to be expelled,
and the carbon disulphide becomes brown. If a small quantity of
iodine were still left, the first drop of chlorine water causes its
liberation, and, on shaking, the disulphide will show a pale-violet
colour. A few more drops of chlorine water, however, will destroy
this, and afterwards liberate the bromine.
Detection of Iodides, Bromides, and Chlorides in Solution together.
The solution containing the three salts, to which a little carbon
disulphide has been added, is acidified with two or three drops of
dilute sulphuric acid, and a dilute solution of sodium nitrite added
drop by drop until the whole of the iodine has been liberated. On
shaking the mixture this will be dissolved by the carbon disulphide,
giving the violet solution. The aqueous liquid is then withdrawn
with a pipette and divided into two portions. The first is neutralised
by the cautious addition of ammonia drop by drop. It is then
shaken with chlorine water and carbon disulphide. The bromine
is thereby liberated, and imparts its brownish colour to the
disulphide. The second portion is evaporated down, mixed with
potassium dichromate and sulphuric acid, and the chromyl-
chloride test made as described on p. 85.
Chloric Acid and Chlorates
Chloric acid is unstable except in dilute solutions.
The chlorates are all soluble in water, therefore no reagents
give precipitates by double decomposition. They are all
decomposed by heat, evolving oxygen (in some cases mixed
with chlorine), and leaving either a metallic chloride or
oxide.
When heated with oxidisable substances (e.g. charcoal),
deflagration of the mixture results.
go Smaller jCltemical Analysis
4
Hydrochloric acid decomposes chlorates, with the evolution
of chlorine and chlorine peroxide
4KC1O 3 -f I2HC1 = 4KC1 + 6H 2 O + 9C1 + 3C1O 2
(The use of this mixture as an oxidising agent has frequently
been referred to.)
Sulphuric acid decomposes chlorates, with the evolution
of chlorine peroxide (a deep yellow unpleasant-smelling gas),
which on very slight elevation of temperature, explodes with
violence
3KC1O 3 + 2H 2 SO 4 = KC1O 4 + 2HKSO 4 -f H 2 O + 2C1O 2
On adding a few drops of strong sulphuric acid to a small
crystal of potassium chlorate, the mixture immediately becomes
yellow, and on very gently warming explodes with a sharp
detonation. This is characteristic of chlorates.
Separation of a Chloride and Chlorate. Add a solution
of silver sulphate; this precipitates silver chloride, which is
removed by filtration. Sodium carbonate is then added to
remove the excess of silver (and any metals other than alkalies),
and the solution is evaporated to dryness and heated until the
chlorate is converted into chloride. The presence of the
chloride in the residue is ascertained by silver nitrate.
Hydrofluoric Acid and Fluorides
The fluorides of the alkali metals, and of silver, mercury
(iron, aluminium, tin), are soluble in water. Those of the
alkaline earths, and of lead (copper, zinc, manganese), are
insoluble.
Calcium chloride gives with soluble fluorides a transparent
gelatinous precipitate of calcium fluoride, CaF 2 , partially soluble
in hydrochloric acid.
Barium chloride throws down a white precipitate of
barium fluoride, BaF 2 , partially soluble in HC1.
Silver nitrate gives no precipitate, as silver fluoride is
soluble in water (distinction between a fluoride and the other
halides).
Liberation of Hydrogen Fluoride. Fluorides are
Hydrofluoric Acid ana Fluorides 91
decomposed by strong sulphuric acid, with evolution of gaseous
hydrogen fluoride
CaF 2 + H 2 SO 4 = CaSO 4 + 2HF
The gas is a colourless, fuming, and highly corrosive com-
pound ; its presence may be detected in the following ways :
(a) Etching Glass. The powdered fluoride is mixed with
strong sulphuric acid in a small dish or tray, made of lead (or
a platinum capsule). It is covered with a small piece of sheet
glass which has been coated on one side with wax, and some
marks or words scratched upon the wax. In a few minutes
the exposed parts of the glass will have become eaten into or
dissolved away by the acid gas ; so that, on removing the wax
with a little hot water, the marks or letters will be. found to be
etched into the glass.
This effect is due to the action of the acid upon silica (and
silicates) forming gaseous silicon fluoride
SiO 2 + 4HF = SiF 4 + 2 H 2 O
(b) The Decomposition of Silicon Fluoride by Water.
The powdered fluoride is mixed with sand, and gently warmed
in a test-tube with a little strong sulphuric acid; a glass rod
with a drop of water upon the end is lowered into the mouth
of the tube. The gas, on coming in contact with the water,
is decomposed, and a white deposit of silicic acid is formed
upon the rod
3 SiF 4 + 3 H 2 = 2H 2 SiF 6 -f H 2 Si0 3
(c) The Formation of Boron Fluoride. When a fluoride
(finely powdered) is mixed with powdered borax, and the
mixture moistened with strong sulphuric acid, gaseous boron
fluoride, BF 3 , is evolved. The action takes place between the
hydrofluoric acid and boric acid, which are disengaged by the
action of the sulphuric acid upon the respective compounds
B(HO) 3 4- 3HF = 3 H 2 + BF 3
If the mixture be introduced into the edge of a Bunsen
flame upon a loop of platinum wire, the flame is tinged a grass-
green colour by the escaping boron fluoride.
92 Smaller Chemical Analysis
Sulphuretted Hydrogen 1 and Sulphides
Sulphuretted hydrogen is a colourless gas, easily distinguished
from all other gases by its unmistakable odour. It is soluble
in water, and imparts its own smell to the liquid. The solution,
however, is unstable, undergoing oxidation and depositing
sulphur. The gas burns with a flame resembling that of
burning sulphur, and yields water and sulphur dioxide.
Liberation of sulphuretted hydrogen takes place when
certain sulphides (see below) are acted upon by acids. The
gas may be recognised (i) by its odour; (2) by its action
upon solutions of metallic salts, e.g. lead acetate. The reaction
is made in a test-tube, and a piece of paper moistened with
lead acetate is held over the mouth of the tube. The sul-
phuretted hydrogen causes a black stain of lead sulphide.
Sulphides of the alkalies and alkaline earths are soluble in
water ; all other metallic sulphides are insoluble (see Analytical
Classification of the Metals).
Soluble sulphides are decomposed by dilute acids (HC1 or
H 2 SO 4 ), with liberation of sulphuretted hydrogen ; in the case
of polysulphides, sulphur is also precipitated
K 2 S + 2HC1 + Aq = 2 KC1 + H 2 S + Aq
CaS 5 + 2HC1 = CaCl 2 + H 2 S + 4 S
Insoluble Sulphides. The behaviour of these towards
acids has already been considered in detail, in studying the
separation of the metals. It may be briefly summarised as
follows :
(a) Sulphides decomposed by dilute acids (HC1 or H 2 SO 4 ),
with liberation of sulphuretted hydrogen : namely, ZnS, MnS,
FeS.
(b) Sulphides unacted upon by dilute acid, but decomposed
by hot strong hydrochloric acid with more or less difficulty :
Sb 2 S 3 , PbS, SnS, NiS, CoS.
(c) Sulphides unacted upon by strong hydrochloric acid,
1 Sometimes called hydrosulphuric acid; the solution of the gas in
water has a feeble acid reaction.
Sulphuric Acid and Sulphates 93
but decomposed by aqua regia, or by a mixture of hydrochloric
acid and potassium chlorate : HgS, As 2 S 3 .
The sulphides of class (b), when treated with hydrochloric
acid in the presence of zinc, or, better, of reduced iron, readily
evolve sulphuretted hydrogen.
Oxidising agents, e.g. nitric acid, convert many of the
sulphides into oxides or sulphates, sulphur being first separated
and afterwards oxidised into sulphuric acid.
When a sulphide is added (in small quantities at a time)
to a fused mixture of sodium carbonate and potassium nitrate
in a platinum crucible, the sulphide is immediately oxidised.
After the mass has cooled, and been extracted with water, the
aqueous liquid may be tested for a sulphate. The sulphides
of classes (a) and ($), when fused upon a piece of platinum foil
(or, better, silver) with sodium hydroxide, are decomposed,
with the formation of sodium sulphide. If a fragment of the
fused mass, after cooling, be placed upon a silver coin and
moistened with a drop of water, or upon a piece of paper
which has been moistened with a solution of lead acetate,
in either case a black stain will be produced ; silver sulphide
on the coin, and lead sulphide upon the paper.
Most sulphides, when heated in a glass tube open at both
ends, and held in a slightly inclined position in order to cause
an air-current to pass through the tube, are decomposed, and
evolve sulphur dioxide.
Sulphuric Acid and Sulphates
Sulphuric acid is an oily, highly corrosive acid liquid. It
combines with water with evolution of heat, and is able to
abstract the elements of water from many organic compounds.
Thus paper, straw, etc., are blackened or charred by the strong
acid. This property is made use of in testing for the free
acid in the presence of soluble sulphates : a piece of paper
is moistened here and there with the solution, and then care-
fully dried, when it becomes charred where it had been wetted.
Or the solution may be mixed with a little white sugar, and
evaporated down in a porcelain dish upon a steam-bath, when
a charred residue will be left.
94 Smaller Chemical Analysis
Most sulphates are soluble in water. Barium, strontium,
calcium, and lead sulphates are insoluble, or nearly so.
Soluble Sulphates. Barium chloride, BaCl 2 , gives with
sulphuric acid or soluble sulphates, a white precipitate of
barium sulphate (see Barium reactions), insoluble in hydro-
chloric acid. The solutions should be dilute, as barium
chloride, being insoluble in strong hydrochloric acid, may
otherwise be thrown out of solution; the addition of water
dissolves it.
Insoluble sulphates may be decomposed by fusion with
sodium carbonate, sodium sulphate being formed. The residue
is extracted with water, and the aqueous solution tested with
barium chloride after being acidified.
When a sulphate is fused with sodium carbonate (which
must be free from sulphates as impurities) upon charcoal
in the reducing flame, a sulphide of the alkali metal is
obtained. If this be placed upon a piece of paper moistened
with acetate of lead, and touched with a drop of dilute hydro-
chloric acid, sulphuretted hydrogen is liberated, and the lead
paper stained black.
[This test is only conclusive evidence of a sulphate when
other sulphur compounds are proved to be absent.]
Sulphurous Acids and Sulphites
Sulphurous acid, H 2 SO 3 , is only known in solution, being
produced when sulphur dioxide is passed into water, or when
this gas is liberated from combination (as from sulphites) in
the presence of water ; thus
(1) In dilute solution
Na 2 SO 3 + 2HC1 + Aq = 2NaCl + H 2 SO 3 + Aq
(2) In stronger solution
Na 2 SO 3 + 2HC1 = 2NaCl + H 2 O + SO 2
The anhydride, SO 2 , is recognised by its characteristic suffo-
cating odour (familiar as " the smell of burning sulphur ").
Reducing Action of Sulphurous Acid. Sulphurous acid
easily takes up oxygen, and passes into sulphuric acid, and
Sulphurous Acids and Sulphites 95
some of its most important reactions are those in which it
thus acts as a reducing agent. Thus, potassium permanganate
is reduced with formation of manganous sulphate
2KMnO 4 + sH 2 SO 3 = KjSO, + 2MnSO 4 + 2H 2 SO 4 + 3H,O
This reaction affords a delicate test for sulphur dioxide.
The gas is cautiously decanted (being much heavier than air)
into a test-tube containing water slightly tinted with a minute
quantity of potassium permanganate. On shaking the gas and
water, the pink colour will be destroyed.
Oxidising Action of Sulphurous Acid. Sulphurous acid is
also capable of undergoing reduction, acting therefore towards
more powerful reducing agents in the capacity of an oxidising
substance. Thus, stannous chloride in presence of hydro-
chloric acid, is oxidised into stannic chloride, the sulphurous
acid being reduced to sulphuretted hydrogen. This latter then
reacts upon the stannic chloride, with precipitation of stannic
sulphide
(1) 3 SnCl 2 + 6HC1 + H 2 S0 3 = 3SnCl 4 + 3 H 2 O + H 2 S
(2) SnCl 4 + 2H 2 S = SnSa + 4HC1
Nascent hydrogen, obtained by the action of hydrochloric
acid upon zinc, also reduces sulphurous acid to sulphuretted
hydrogen
H 2 S0 3 + 3 H 2 = 3 H 2 + H 2 S
The test may be made by adding a minute trace of sulphurous
acid (or a solution of a sulphite) to a mixture of zinc and
hydrochloric acid in a test-tube, and applying acetate of lead
paper to the mouth of the tube.
Sulphites. The only sulphites soluble in water are those
of the alkali metals. They are all decomposed by dilute acids,
with evolution of sulphur dioxide (see above). Oxidising
agents convert them into sulphates.
When heated by themselves, most sulphites are converted
into sulphides and sulphates
4 K 2 SO 3 = K 2 S + 3 K 2 S0 4
96 Smaller Chemical Analysis
Those of the alkaline earths leave an oxide, and evolve sulphur
dioxide
BaSO 3 = BaO -f- SO 2
Barium chloride gives a white precipitate of barium sulphite,
BaSO 3 , soluble in dilute HC1 (distinction from BaSO 4 ).
Lead acetate precipitates white lead sulphite, PbSO 3 .
The salt undergoes no change when boiled (contrast lead
thiosulphate).
Silver nitrate gives a white precipitate of silver sulphite,
which, on boiling, is converted into black metallic silver
Ag 2 SO 3 + H 2 O = H 2 SO 4 + Ag 2
Separation of a Sulphate and Sulphite. The solution
(dilute) is acidulated with hydrochloric acid, and barium
chloride added. The precipitated sulphate (insoluble in acid)
is removed by nitration. To the solution, which now contains
barium chloride and sulphurous acid, an oxidising agent, such
as chlorine water, is added, when a precipitate of barium
sulphate is again thrown down
BaCl 2 + H 2 SO 3 + H 2 O + C1 2 = BaSO 4 + 4HC1
Separation of a Sulphide, Sulphate, and Sulphite. The
sulphide is first separated as an insoluble metallic sulphide
by shaking up the solution with a little lead carbonate (or
cadmium carbonate). The precipitated sulphide is then
removed by nitration. Very small traces of sulphuretted
hydrogen will produce a distinct coloration in the white
carbonate of lead.
If cadmium carbonate is used the precipitate should be
treated with acetic acid, which dissolves the excess of white
CdCO 3 , leaving the yellow CdS.
The sulphate and sulphite in the filtrate are then separated
as described above.
Nitric Acid and Nitrates
Nitric acid is a fuming corrosive liquid. It readily dis-
solves most metals, converting them into nitrates or oxides,
Nitric Acid and Nitrates 97
with evolution of oxides of nitrogen, and in some cases with
the formation of ammonia.
Nitric acid also oxidises many of the non-metals; thus
sulphur, phosphorus, and iodine, are converted respectively
into sulphuric, phosphoric, and iodic acids. It is capable also
of oxidising indigo, which thereby loses its blue colour, being
bleached.
Nitrates are all soluble in water ; their recognition, there-
fore, is based upon the oxidising reactions of which they, or
the nitric acid which they yield, are capable.
Reduction by Ferrous Salts. When ferrous sulphate is
brought into contact with a mixture of a nitrate and strong
sulphuric acid, the solution assumes a deep brown colour.
Three chemical changes go to make up the reaction : (i) the
liberation of nitric acid by the action of sulphuric acid upon the
nitrate ; (2) the reduction of the nitric acid by the ferrous salt,
with elimination of nitric oxide ; and (3) the absorption of the
nitric oxide so formed by a further portion of ferrous salt,
forming an unstable brown compound having the composition
NO,2FeSO 4 . The test is extremely delicate, and is carried out
in the following manner : The solution of the nitrate is mixed
with about its own 'volume of strong sulphuric acid in a test-
tube, and the mixture cooled. To this a little ferrous sulphate
solution is cautiously added, the tube being held in an inclined
position, so that the ferrous sulphate shall float upon the denser
liquid already in the tube. Where the two liquids meet, the
brown colour will be developed. By a gentle movement of
the tube, so as to cause a slight admixture of the liquids at
the point where they meet, the brown ring will be still more
apparent. The coloured compound is decomposed by heat,
with evolution of nitric oxide, hence the necessity for making
the test with cold solutions.
Reduction by Sulphurous Acid. When copper (or mer-
cury) is heated with sulphuric acid in the presence of a nitrate,
nitric oxide is evolved, which, in contact with the air, gives red
vapours of nitrogen peroxide
3Cu + 4H,SO 4 + 2KNO 3 = sCuSO 4 + K,SO 4 -f 4H 2 O + 2NO
98 Smaller Chemical Analysis
The sulphur dioxide (developed by the action of the acid upon
the copper) is oxidised by the nitric acid (simultaneously
generated by the action of the acid upon the nitrate) to sul-
phuric acid ; thus
(1) Cu + 2H 2 SO 4 = CuSO 4 + 2H 2 O + SO 2
(2) 2 HN0 3 + sSO a + 2H 2 = 3 H 2 S0 4 + 2NO
The nitrate is mixed with a little strong sulphuric acid, and
a few fragments of copper foil or turnings are introduced. On
boiling the mixture, red fumes of nitrogen peroxide, NO 2 , will
appear in the tube, which will be more easily seen by looking
down through the mouth of the tube.
Decomposition by Heat. Nitrates all undergo decomposi-
tion when strongly heated. Nitrates of alkali metals and
alkaline earths, when gently heated, are reduced to nitrites,
with evolution of oxygen.
Ammonium nitrate passes into water and nitrous oxide.
Other nitrates, e.g. lead nitrate, leave an oxide of the metal, and
give off oxygen and nitrogen peroxide.
When heated with oxidisable substances (carbon, sulphur,
etc.) the decomposition is propagated with explosive violence.
Thus, when nitrates are heated before the blowpipe on char-
coal, deflagration of the charcoal takes place.
Nitrates and chlorates, when present together, are ex-
amined by being first converted by heat into nitrites and
chlorides. If present as salts of metals other than the alkalies,
sodium carbonate is added, and the dry mixture heated until
the evolution of oxygen is at an end. The residue is extracted
with water, and the solution examined for nitrites and chlorides.
If chlorides are originally present as well as nitrates and
chlorates, they must be first removed by precipitation with
silver sulphate, as explained on p. 90.
Nitrous Acid and Nitrites
The acid is not known in the pure state. Even when
liberated in dilute solutions, it speedily breaks up into nitric acid,
nitric oxide, and water. Hence when nitrites are decomposed
Nitrotis Acid and Nitrites 99
by acids, nitric oxide is evolved, which, in contact with atmo-
spheric oxygen, passes into the brown gas NO 2 ; thus
6NaNO a + 3H 2 S0 4 = sNa 2 SO 4 + 2 HNO 3 + 2H 2 O + 4 NO
Nitrites are all soluble in water, but the silver salt is suffi-
ciently difficult of solution to be precipitated, on the addition
of silver nitrate, to a (not too dilute) solution of a nitrite.
All nitrites are easily decomposed by dilute acids in the
cold, with evolution of nitric oxide, as shown in the above
equation. If the action takes place in the presence of a ferrous
salt, the same brown-coloured compound is produced as in the
case of a nitrate. [Nitrites therefore give a " brown ring,"
when dilute sulphuric, or even acetic acid is used (distinction
from nitrates)^
Oxidation Reactions. Nitrous acid and nitrites part with
oxygen, and are converted into nitric oxide. Thus, in contact
with potassium iodide, the latter is oxidised with liberation of
iodine
2 KI + H 2 O + O = 2KHO + I a
Similarly, sulphuretted hydrogen is oxidised by a nitrite in
presence of an acid, with precipitation of sulphur
H 2 S + 2 HNO 2 = 2H 2 O + 2NO + S
Cobaltous nitrite is oxidised to cobaltic nitrite. When a strong
solution of cob<ous chloride mixed with acetic acid is added to
a solution of potassium nitrite, a yellow precipitate is obtained,
consisting of potassium cobaltzV nitrite, 3KNO 2 , Co(NO 2 ) 3 .
Reduction Reactions. Nitrous acid, by absorption of
oxygen, passes into nitric acid; it therefore is capable of
reducing other compounds, such as chromates, permanganates,
mercurous (but not mercuric) salts.
Detection of Nitrites and Nitrates in the same Solution.
Owing to the ready decomposition of nitrites by dilute
acids, they are easily detected in presence of nitrates, either
by the liberation of iodine, the oxidation of ferrous salts, or
reduction of potassium permanganate. To find a nitrate when
nitrites are present is less simple. The dilute solution of the
ioo Smaller Chemical Analysis
mixed nitrate and nitrite is acidified with three or four drops
of dilute sulphuric acid, and a little ferrous sulphate solution
(or a small crystal of the salt) is added. The solution at once
becomes dark brown (owing to the absorption, by the ferrous
salt, of the nitric oxide liberated from the nitrite). It is then
heated (but not allowed to boil), with frequent shaking, when
nitric oxide is expelled, and the liquid gradually becomes
colourless. The mixture is cooled, and one drop more dilute
acid added, and a little more ferrous sulphate. (If all the
nitrite present had been decomposed, this addition gives no
further coloration.) This solution is now poured carefully on
to a small quantity of strong sulphuric acid in a test-tube, so as
to float upon the acid, and where the liquids meet a " brown
ring " will be formed, due to the nitrate present.
Phosphoric Acid and Phosphates
Three phosphoric acids (each with its series of phosphates)
are known, namely, orthophosphoric acid, H 3 PO 4 ; pyrophos-
phoric acid, H 4 P 2 O 7 ; and metaphosphoric acid, HPO 3 .
Orthophosphoric acid. The only orthophosphates which
are soluble in water are those of the alkali metals.
Silver nitrate gives a yellow precipitate with soluble
phosphates, of silver phosphate, Ag 3 PO 4 , which distinguishes
ortho from pyro and meta compounds.
Ammonium molybdate gives a yellow precipitate, consist-
ing of ammonium phospho-molybdate (see p. 51). Pyro and
meta phosphates also give the test, because on warming in con-
tact with nitric acid they are transformed into ortho salts.
Magnesium sulphate, in presence of NH 4 C1 and NH 4 HO,
gives a white precipitate of ammonium magnesium phosphate,
NH 4 MgPO 4 . (Magnesium reaction, p. 26.)
Action of Heat. Orthophosphates containing either one
or two acidic hydrogen atoms (as HNa 2 PO 4 , or H 2 NaPO 4 ), or
the volatile radical NH 4 , yield when heated either pyro or meta
phosphates. Normal phosphates containing only nonvolatile
positive radicals are not decomposed.
Pyrophosphoric Acid and Pyrophosphates are produced when
orthophosphoric acid or certain orthophosphates are heated.
Carbonic Acid and Carbonates 101
Boiling with acids retransforms pyrophosphates into orthophos-
phates.
Only the pyrophosphates of the alkalies are soluble in water.
Silver nitrate gives a white pr^c'fjaiiafte, of silvf^r 'pyjrbphosphate,
Ag 4 P 2 7 .
Magnesium sulphate precipitate^ w}i<te^ na^gn^si'Jiiii /^yrophos-
phate, Mg 2 P 2 O 7 , soluble irr jexcebs^ of ^ magnesikrA ' sulphate, and
not reprecipitated in the cold by ammonia (distinction from ortho-
phosphate^).
Ammonium molybdate gives no precipitate until the " pyro "
acid has been changed to " ortho " by the action of the nitric acid
present.
Metaphosphoric Acid and Metaphosphates. The acid is formed
when the " ortho " or " pyro " acids are strongly heated, whereby
water is expelled. It is known as glacial phosphoric acid
H 3 P0 4 = H 2 + HP0 3
The reaction is reversible ; for when metaphosphoric acid is dis-
solved in water, it passes back to the ortho acid slowly in the
cold, quickly when boiled.
Silver nitrate gives a white precipitate of silver metaphosphate,
AgP0 3 .
Magnesium sulphate, in presence of ammonium chloride, gives
no precipitate (distinction Jrom "pyro" and " ortho " acids}.
Albumen (white of egg) is coagulated when shaken up with
metaphosphoric acid (or metaphosphates acidified with acetic acid)
(distinction from "pyro " and " ortho " acids ^ which are without
action upon albumen}.
Carbonic Acid and Carbonates
Carbonic Acid, H 2 CO 3 , is an unstable compound only
capable of existence in dilute aqueous solution. It is formed
when carbon dioxide is dissolved in water, and has a feeble
acid reaction. It is capable of dissolving the normal carbonates
of the alkaline earths and of magnesium, forming the so-called
" acid " carbonates or " bicarbonates." All normal carbonates
are insoluble in water, except those of the alkalies. The " acid "
carbonates are all soluble, but on boiling their solutions, they
are converted into normal salts, with evolution of carbon
dioxide.
IO2 Smaller Chemical Analysis
All carbonates are decomposed by dilute hydrochloric acid
(and by nearly all acids) with effervescence, due to the rapid
escape of carbpn dioxide^ The gas is identified by its action
upon lime-wa'tcr (or baryta- V/^ter).
The test is made, by adding a few drops of acid to the
carbonate in a test-tube, anj decanting the evolved (heavy)
gas into a second test-tube containing a little lime-water,
Ca(HO) 2 . On shaking the lime-water with the gas, the liquid
becomes milky, owing to the precipitation of calcium carbonate. 1
When strongly heated, the normal carbonates of the alkali
metals (not ammonium) remain unchanged. Those of the
alkaline earths are converted at a high temperature into oxides,
with evolution of carbon dioxide (illustrated in the process of
lime-burning). All other carbonates are more readily decom-
posed by heat.
Silicic Acid and Silicates
Silicic acid, H 2 SiO 3 , is obtained, when soluble silicates are
decomposed by acids, as a white gelatinous substance, slightly
soluble in water, and still a little more soluble in acids
Na 2 SiO 3 + 2HC1 = H 2 SiO 3 + 2NaCl
Silicic acid is also precipitated from a solution of an alkali
silicate by the addition of ammonium carbonate (ammonia does
not cause any precipitate)
Na 2 Si0 3 + (NH 4 ) 2 C0 3 ' = Na^COg + 2 NH 3 + H 2 SiO 3
Silicic acid is a very feeble acid, and when heated to
130 C. it parts with a molecule of water, and is converted into
1 The only other gas which gives a white precipitate with lime-water is
sulphur dioxide. This is easily distinguished from carbon dioxide by its
smell. If the two gases are present together, they may be passed through
a little dilute potassium permanganate solution. The sulphur dioxide is
absorbed (being oxidised by the permanganate into sulphuric acid), and the
carbon dioxide passes on, and can be detected by means of lime-water. If
the passage of the gases through the permanganate be continued for a few
minutes, the colour of the solution becomes entirely destroyed ; and the
liquid may then be tested for sulphuric acid in the usual way.
Silicic Acid and Silicates 103
silicon dioxide (silica), SiO 2 , a compound which is insoluble in
water and in acids. 1
Silica, SiO 2 , occurs in a more or less pure state in nature,
in the form of quartz, flint, agate, sand, etc. When prepared
artificially by heating the hydrated compound, it is a white
amorphous powder. It is extremely stable, and capable of
standing a very high temperature. It is insoluble in water and all
acids except hydrofluoric acid. It dissolves in caustic alkalies ;
and, when in the amorphous state, in boiling carbonates of the
alkalies, yielding in all cases the soluble alkali silicates.
The only silicates which are soluble in water are the alkali
silicates (known as water glass, or soluble glass).
The presence of silica is detected by means of its behaviour
when heated with microcosmic salt. A clear bead of this salt
is made upon a loop of platinum wire, and a few small particles
of the powdered silicate are heated in it. The metallic oxides
are dissolved by the fused microcosmic salt, but not the silicon
dioxide, particles of which (the skeletonic remains of the
mineral) remain floating about in the molten bead.
Silica may also be detected by the formation of silicon
fluoride when the compound is acted upon by hydrofluoric acid.
The test is made as described under hydrofluoric acid (p. 9i). 2
In this case, however, the drop of water which is suspended in
the gas in order to detect the silicon fluoride, should be held
on a loop of platinum wire, as the action of the hydrofluoric
acid upon the glass rod might be mistaken for a deposition of
silica.
Insoluble silicates may be classified for analytical purposes
into (i) those which are decomposed by acids (other than
hydrofluoric acid) ; and (2) those which are unattacked, which
comprises by far the larger class.
(i) Silicates which are decomposed by acid may have their
silica removed by treatment with hydrochloric acid. The
mineral, in as finely powdered a condition as possible,, is
1 Silicic acid is sometimes spoken of as soluble silica, and silicon dioxide
as insoluble silica.
2 Except that in this case the test must not be made in glass vessels
(which are themselves silicates) but in a lead or platinum capsule.
IO4 Smaller Chemical Analysis
digested with strong hydrochloric acid at a gentle heat, when
gelatinous silicic acid separates, and the powdered mineral
gradually dissolves. The mixture is next evaporated to dry-
ness (preferably on a steambath), and then gently heated over
a flame for a short time in order to ensure the entire conversion
of the silicic acid into silica. The residue is treated with a small
quantity of strong hydrochloric acid, water is added, and the
solution containing the metals present as chlorides is separated
by nitration from the insoluble residue of silica.
(2) Silicates which are unattacked by acids are decom-
posed by fusion with alkali carbonates. The finely powdered
silicate is mixed with several times its weight oi fusion mixture,
and the mixture heated in a platinum crucible. The fusion is
continued until effervescence ceases, the temperature being
raised towards the end of the operation. The residue, after
cooling, is extracted with water, which dissolves the alkali
silicate (and excess of carbonate), leaving the metallic oxide (or
carbonate). Hydrochloric acid is then gradually added, which
causes the precipitation of silicic acid, and at the same time
dissolves the metallic oxides. The mixture is then evaporated,
and treated in the manner described above.
Fusion with alkali carbonates is obviously inadmissible in
the case of natural silicates which are suspected of containing
the alkali metals, and which are insoluble in hydrochloric acid.
In this case, one of the following plans may be employed :
(a) Heating with Ammonium Chloride and Calcium Car-
bonate. A small quantity of the powdered silicate is mixed
with about an equal weight of ammonium chloride and about
eight times its weight of calcium carbonate (precipitated), and
the mixture gently heated in a platinum crucible until no more
fumes of ammonium chloride are given off. It is then strongly
heated with the blowpipe for about fifteen minutes, after which
the mass is treated with water. The alkalies, in the form of
chlorides, together with a small quantity of calcium chloride,
pass into solution, and are separated by filtration. The calcium
is removed by precipitation with ammonium carbonate, and the
filtrate evaporated and examined for alkalies.
(b) Decomposing- the Silicate with Hydrofluoric Acid.
Boric Acid and B orates 105
This may be accomplished by treating the powdered mineral
with aqueous hydrofluoric acid in a platinum crucible, gently
evaporating the liquid (in a draught cupboard) to dryness, add-
ing fresh acid and evaporating again, continuing the operation
until the residue is entirely soluble in hydrochloric acid.
Boric Acid and Borates
Boric acid, H 3 BO 3 , is a white crystalline solid, sparingly
soluble in cold, but more readily soluble in hot, water. It is
deposited from its solutions in the form of pearly white scales.
It is also soluble in alcohol, and when either the aqueous or
alcoholic solution is boiled, the acid vaporises along with the
solvent.
When heated, boric acid (ortho), H 3 BO 3 , loses water, pass-
ing first into metaboric acid, H 2 B 2 O 4 , and finally into pyroboric
acid, H. 2 B 4 O 7 .
The borates of the alkalies are readily soluble in water ;
most other borates are insoluble. A few (e.g. magnesium
borate) are difficultly soluble. The most familiar salts are
those of pyroboric acid, e.g. ordinary borax, Na 2 B 4 O 7 .
The insoluble borates obtained by precipitation are not
characteristic, and are not used in analysis for the detection of
borates.
All borates are decomposed by mineral acids, with libera-
tion of orthoboric acid ; thus
Na 2 B 4 O 7 + 2HC1 + 5H 2 O = 2NaCl + 4H 3 BO 3
Reaction with Turmeric. Boric acid produces upon tur-
meric paper a characteristic red-brown stain. This coloration
is distinguished from that produced by alkalies (which it closely
resembles in appearance) by the fact that when touched with
an alkali the brown colour is changed to a greenish-black, but
is restored to its original tint by dilute acids (HC1 or H 2 SO 4 ).
The borate is moistened with hydrochloric or sulphuric acid in
order to liberate the boric acid, and a drop or two of the liquid
is poured upon the turmeric paper.
Flame Reactions. Volatile boron compounds impart a
characteristic green colour to the flame ; thus, when an alcoholic
106 Smaller Chemical Analysis
solution of boric acid is boiled, and the alcohol vapour inflamed,
the green colour due to the volatilised boric acid is apparent.
The test is made in the following manner :
The borate (borax) is moistened with a little strong sulphuric
acid in a test-tube or small flask, and alcohol is added. The
test-tube is closed with a cork carrying a short straight glass
tube. The contents of the tube are then heated, and as the
alcohol boils off it is inflamed at the exit tube, when the green
colour of the flame is observed.
A small quantity of the powdered borate (or boric acid) is
mixed with about its own weight of powdered calcium fluoride,
and the mixture moistened with one or two drops of strong
sulphuric acid. A little of the paste is introduced into a Bunsen
flame upon a loop of platinum wire. By the action of the acid
upon the fluoride, hydrofluoric acid is formed ; and this in the
presence of the borate gives boron fluoride, BF 3 , which causes
a green coloration of the flame. The presence of copper salts
masks the reaction.
Permanganic Acid and Permanganates
The acid is not met with in analysis. The permanganates
are all soluble in water, therefore no precipitates by double
decomposition are produced with these salts. Their chief pro-
perties are their oxidising powers, which have been already
described under the reactions for manganese (p. 44).
Arsenates and Chromates have already been treated under
Arsenic (p. 64) and Chromium (p. 35).
CHAPTER X
PRELIMINARY EXAMINATION FOR METALLIC
RADICALS
A. When the Substance under Examination is a Liquid.
Before proceeding to the group separation, the liquid should be
carefully tested with litmus paper, in order to ascertain whether
it is neutral, acid, or alkaline.
(1) If neutral, the liquid might be simply water. To
ascertain whether or not it contains anything in solution, a few
drops should be placed upon a watch-glass and carefully
evaporated to dryness. If there is no residue left in the watch-
glass, it contained no salts in solution, and was therefore simply
water.
(2) If acid) the solution may contain either a free acid, or
salts possessing an acid reaction. [Either certain normal salts,
e.g. copper sulphate, alum, etc., or certain " acid " salts, as
hydrogen sodium sulphate.]
(3) If alkaline, the liquid may contain either free alkali, or
salts having an alkaline reaction.
B. When the substance is a solid, it should be critically
examined, in order, if possible, to gain any information from
its general physical properties which may help to identify it.
If crystalline, the colour, shape, etc., of the crystals should be
noted. If powdered, it may be examined with a pocket-lens, in
order to discover whether or not it is homogeneous ; i.e.
whether it is a single compound, or a mixture of more than one.
The substance should then be subjected to the following
general tests x :
1 Every smallest detail of these preliminary tests should be carefully
and systematically noted down, whether the interpretation of the observa-
tion is obvious to the student or not.
107
io8 Smaller Chemical Analysis
I. The Flame Reaction. A small quantity of the substance
is introduced into the edge of the Bunsen flame (first in the
cooler region near the base of the flame, and afterwards in the
hotter parts near the top of the interior cone) upon a loop of
clean platinum wire. The flame is coloured
Intense yellow \ by compounds of sodium.
Violet, appearing red through the potassioscope, potassiiim.
Crimson, compounds of strontium.
Orange red calcium.
Pale green barium.
Green ,, copper, boric acid.
Obviously the presence of one of these substances may
modify or mask the colour reaction due to another.
II. The Borax Bead. A minute particle of the substance
is heated in a borax bead in the blowpipe flame, both inner
and outer flame. The bead is coloured
Blue (both outer and inner flame) indicates cobalt.
Brown (outer), grey (inner), compounds of nickel.
Violet (outer), colourless (inner), compounds of manganese
(also mixture of Co and Ni).
Brown, cooling to yellow (outer), green (inner), iron.
Green (outer and inner), compounds of chromium.
Green, cooling to bluish (outer), red (outer), copper.
Here, as with the flame test, one substance modifies or
masks the reaction due to another.
[Should this test lead to the suspicion that either manga-
nese or chromium is present, it should be followed up by the
fusion of a portion of the substance with sodium carbonate
and nitre upon platinum foil. See Manganese and Chromium
reactions.]
III. Blowpipe Reactions upon Charcoal. (a) When heated
alone. A considerable amount of information may be obtained
by heating a little of the substance by itself upon charcoal.
(1) If it melts and is absorbed into the charcoal, it points to
the substance consisting of salts of the alkalies. [Chlorates
and nitrates cause vivid combustion of the charcoal.]
(2) If a white infusible residue is obtained, the substance
may consist of oxides (or salts which yield oxides when
Preliminary Examination for Metallic Radicals 109
heated) of the alkaline earth metals, alumina, zinc (ZnO yellow
while hot, white when cold), or of silica. [If the residue,
when placed upon a piece of tumeric or litmus paper and
moistened with water, shows an alkaline reaction, it will con-
tain one of the alkaline earth metals. Silica may be specially
tested for by the bead of microcosmic salt.]
(3) If a coloured residue is left, it points to one of the metals
already indicated by the borax-bead test.
(4) If reduction takes place, resulting in the formation of
fumes and an incrustation upon the charcoal, without indica-
tion of a metallic bead, it points to the presence of compounds
of very volatile metals, as arsenic (white incrustation accom-
panied by garlic odour), cadmium (red-brown incrustation),
zinc (incrustation yellow when hot, and forming very close to
the substance). Such a volatile compound as ammonium
chloride gives white fumes and an incrustation (the latter being
formed at a considerable distance from the heated spot). In
this case the blowpipe flame appears of a yellow-ochre colour
as it impinges upon the ammonium salt.
(b) When heated with Reducing Agents. When the sub-
stance is mixed with sodium carbonate and potassium cyanide,
and heated upon charcoal in the reducing flame, compounds of
copper and silver are reduced to the metallic state without
giving any incrustation upon the charcoal ; while compounds of
antimony, bismuth, tin, and lead give metallic beads accom-
panied by incrustations (for the characteristics of the incrusta-
tions and the metals, see Special relations of each).
IV. The Action of Heat. The behaviour of a substance
when heated alone in a dry narrow tube closed at one end (a
small test-tube), will generally afford important information
respecting its composition.
A. If the substance simply melts, and solidifies on cooling, without
giving off any gases or vapours, a large number of compounds
are obviously at once excluded. In this case the substance
may consist of salts of the alkalies or the alkaline earth metals ;
a few white salts of the heavy metals, e.g. silver chloride ; or a
few coloured salts, e.g. lead chromate.
B. If water is given off, collecting in drops upon the upper
no Smaller Chemical Analysis
part of the tube, it may be due (i) to hygroscopic ] moisture (in
this case the amount will probably be small), (2) to the decom-
position of metallic hydroxides, or (3) to water of crystallisa-
tion. [Substances containing much water of crystallisation,
when heated, often melt twice first in the water of crystallisa-
tion, and as this is expelled they resolidify, but on the applica-
tion of a stronger heat they once more undergo fusion.]
The water should be carefully tested by introducing a small
strip of litmus paper.
Alkalinity would indicate ammonium compounds (e.g.
(NH 4 ) 2 MgPO 4 , HNa(NH 4 )PO 4 , etc.).
Acidity might be due to the decomposition of certain acid
salts, either alone or by interaction with other salts (e.g. hydro-
gen potassium sulphate, when heated, is converted into the
normal salt and sulphuric acid, and if present along with a salt
containing water of crystallisation, the water would become
acid ; or if mixed with sodium chloride or nitrate, hydrochloric
acid or nitric acid would be evolved by double decomposition).
C. If the substance changes colour
(1) From white to yellow, indicates oxide of zinc, tin,
bismuth ;
(2) From yellow to brown (fusing at a red heat), oxide of
lead.
Most coloured oxides become much darker when heated
(e.g. HgO, Fe 2 3 ).
D. If gases or vapours are evolved, they must be identified by
special tests.
(1) Oxygen and nitrous oxide (re-ignition of glowing splint
of wood). The former from chlorates, nitrates, or peroxides ;
the latter from ammonium nitrate, or salts which by double
decomposition give ammonium nitrate.
(2) Nitrogen (extinguishes flame, and gives no reaction with
lime-water), from ammonium nitrite or chromate, or from
mixtures which yield these salts by double decomposition.
(3) Chlorine, bromine, and iodine (recognised by colour,
1 That is moisture due to the substance being " damp " mechanically
adhering moisture. Almost all powders attract a little moisture in this
way from the atmosphere.
Preliminary Examination for Metallic Radicals 1 1 1
smell, etc.), evolved from certain halogen compounds when
heated alone, or in admixture with acid salts and peroxides
(e.g. a mixture containing such compounds as NaCl, HKSO 4 ,
and MnO. 2 , when heated, evolves chlorine).
(4) Carbon dioxide (action on lime-water), from most car-
bonates other than the normal salts of the alkalies.
(5) Sulphur dioxide (characteristic odour), from certain
sulphites and sulphates.
(6) Ammonia (odour, and action on test-paper), from
ammonium salts.
(7) Sulphuretted hydrogen (odour, and action on paper
moistened with lead acetate), from the decomposition of
hydrated metallic sulphides.
(8) Nitrogen peroxide, from the decomposition of nitrates
of heavy metals.
E. If a sublimate is produced, it indicates such volatile solids
as the following :
(1) Giving a white sublimate ammonium haloid salts,
arsenious oxide, mercuric chloride, mercurous chloride
(yellowish white hot, white on cooling).
(2) Giving a coloured sublimate mercuric iodide (red and
yellow), arsenious sulphide (yellow).
(3) Giving a black metallic-looking sublimate iodine
(violet vapours), mercuric sulphide (red streaks if rubbed with a
glass rod), metallic mercury (runs into liquid globules when
rubbed with a glass rod), metallic arsenic (garlic smell). [If
this test leaves it doubtful whether arsenical or mercurial com-
pounds are present, a little of the substance should be mixed
with sodium carbonate and heated in a dry tube. If mercurial,
the sublimate consists of minute globules.]
Ammoniacal compounds may be at once tested for by
heating a portion of the substance with a little sodium hydroxide,
when ammonia is evolved.
F. If no change takes place, it will be evident from the fore-
going that the number of substances which can possibly be
present is extremely limited ; and if in addition the compound
is white, the range of probable substances is still further
narrowed down to the oxides of a few of the metals, such as
the alkaline earths, alumina, silica, etc.
112 Smaller Chemical Analysis
To obtain a Solution of the Solid Substance. (i) In
Water. A small quantity of the substance, in a finely powdered
condition, is treated with water in a large test-tube, and the
mixture boiled for a few moments.
If the substance does not appear to dissolve, it is either
wholly or partly insoluble in water. To ascertain whether any
of it has dissolved, the mixture should be allowed to settle,
and a few drops of the clear liquid evaporated to dryness upon
a watch-glass. If a residue is obtained on evaporation (thus
showing that partial solution in water has taken place), the
aqueous liquid is decanted off, and the insoluble portion again
boiled with water and filtered.
(2) In Acids. The portion insoluble in water is then
treated with dilute hydrochloric acid and boiled. If it does
not entirely dissolve, the dilute acid is decanted off into another
test-tube, and strong hydrochloric acid substituted, the mixture
being submitted to prolonged boiling, if necessary. If the
substance wholly dissolves, the two acid solutions may be mixed
together.
If, on adding a few drops of the aqueous extract to a small
portion of the acid solution, no precipitation takes place, the
main portions of the two solutions may be mixed together.
On the other hand, if the two extracts contain compounds
which will interact with the formation of an insoluble precipi-
tate, they must be examined separately.
If the substance is not entirely dissolved by strong hydro-
chloric acid, it may consist of one of the few compounds which
are only dissolved by aqua regia. 1 The residue is therefore
treated with a small quantity of mixed nitric and hydrochloric
acids, and boiled. The solution should be evaporated down
in a porcelain dish until only a small bulk remains, in order to
expel as much acid as possible, and then diluted with water ;
this solution may then be mixed with the hydrochloric acid
extract.
(3) Treatment of Substances insoluble in Water or in
1 Such as sulphides of nickel, cobalt, mercury. The presence of these
should have been indicated by the preliminary tests. The use of aqua regia
should only be resorted to when absolutely necessary.
Preliminary Examination for Metallic Radicals 113
Acids. The number of compounds insoluble in water and
acids is very limited. Of commonly occurring substances, the
following may be present :
(a) Sulphates of barium, calcium, strontium, and lead (the
three last named being sufficiently soluble to be detected in
the aqueous extract).
(b) Silica, and many natural silicates.
(c) Fluor spar, and other natural fluorides.
(d) Silver chloride.
(e) Oxides of aluminium and chromium (which have been
strongly heated). Native stannic oxide (tinstone).
(f) A few arsenates and phosphates.
These insoluble substances are converted into soluble
compounds by fusion with alkali carbonates. The insoluble
residue (in the absence of lead and silver) is dried, and mixed
with about four times its weight of fusion mixture and a little
potassium nitrate, and fused in a platinum crucible 1 until all
effervescence is at an end. The crucible is then allowed to
cool.
The fused mass is then boiled with water until nothing
further dissolves, and the solution filtered.
The aqiieo2is solution is examined for such acid radicals as
in the nature of the case could be present namely, sulphuric,
silicic, hydrofluoric, chromic, arsenic, phosphoric, hydrochloric. 2
The residue, after being thoroughly washed with hot water,
until the wash-water is no longer alkaline, is dissolved in
hydrochloric acid, and the solution examined for such metallic
radicals as can be present.
1 If lead or silver is present (ascertained during the preliminary exami-
nation), a platinum crucible must not be used.
2 The fusion-mixture must obviously be free from sulphates, chlorides,
etc., or tests for these acids are of no value.
CHAPTER XI
PRELIMINARY EXAMINATION FOR ACID RADICALS
(i) By the Action of Dilute Acids. Dilute HC1 or H 2 SO 4
is added to the substance, and the mixture gently warmed.
Effervescence may take place owing to the escape of
CO 2 from carbonates. Apply lime-water test.
SO 2 sulphites. Recognised by smell. If along with
CO 2 , apply permanganate test.
NO nitrites. Shows itself as brown vapour, NO 2 .
H 2 S sulphides. Recognised by smell, and acetate of
lead paper.
Cl (when the acid used is HC1, and the liquid boiled)
from chlorates , chromates, nitrates ; peroxides.
(2) By the Action of Strong Sulphuric Acid. On gently
warming the solid with a little strong H 2 SO 4 in a test-tube, the
following gases or vapours may be evolved : *
HC1 from chlorides.
HF fluorides (in contact with the glass test-tube,
SiF 4 will be evolved).
HNO 3 (with more or less brown fumes) from nitrates.
Cl from a chloride^ in presence of peroxides or chromates.
C1O 2 chlorate (a deep-yellow gas, the mixture deto-
nates).
Br bromide (brown vapour, with generally sufficient
HBr to fume in moist air).
CrO 2 Cl 2 from a mixed chloride and chromate. (Brown
vapour resembling Br in colour.)
I from an iodide (violet vapour, accompanied by fumes
due to HI).
1 Care must be taken not to heat the mixture so strongly as to volatilise
the sulphuric acid, the acid fumes from which might be mistaken for other
114
Systematic Detection of the Acids 115
CLASSIFICATION AND SYSTEMATIC DETECTION OF ACID
RADICALS
The acids are classified into three main groups based upon
the solubility of their barium and silver salts.
Group I. Acids whose barium salts are precipitated from
neutral solutions by barium chloride.
(a) Whose barium salts are insoluble in dilute hydrochloric
acid
Sulphuric acid, H 2 SO 4
(b) Whose barium salts are soluble in hydrochloric acid
Carbonic acid, H 2 CO 3 Phosphoric acid, H 3 PO 4
Sulphurous H 2 SO 3 Boric H 3 BO 3
Silicic H 2 SiO 3 Hydrofluoric HF
Chromic H 2 CrO 4
Group II. Acids whose silver salts are precipitated by
silver nitrate from solutions acidified with nitric acid
Hydrochloric acid, HCl Hydriodic acid, HI
Hydrobromic HBr Sulphuretted hydrogen, H 2 S
Group III. Acids whose barium salts are soluble in water,
and whose silver salts are not precipitated in a nitric acid
solution, namely
Nitric acid, HNO 3 Chloric acid, HC1O 3
Nitrous HNO 2
As already mentioned (p. 83), these general reagents,
barium chloride and silver nitrate, are not employed to effect
the separation of groups of acids, but rather as indicators of
the presence or absence of entire groups.
In most cases, each acid is individually tested for in
separate portions of specially prepared solutions (see below).
The examination for acids should be made after the metals
have been detected, for two reasons : firstly, because during
the course of the systematic examination for the metals, the
presence or absence of quite a number of acids will be inci-
dentally ascertained; and secondly, because a knowledge of
n6 Smaller Chemical Analysis
what metals are present will be a guide to the student in
deciding what acids may and what cannot possibly be present.
As stated above, several acids will be detected during
the examination for metals; thus on acidifying with hydro-
chloric acid and gently warming (for separation of Group I.),
the presence of carbonates, sulphites, sulphides, or nitrites will
be indicated. (See preliminary tests.)
On treatment with sulphuretted hydrogen, in the separation
of the metals of Group II., indications will have been obtained
of the presence of chromates and iodates : the former by the
change of colour from orange to green, with simultaneous pre-
cipitation of sulphur ; the latter by the elimination of iodine,
which gives a dark brown colour to the solution, the colour
gradually disappearing as the excess of sulphuretted hydrogen
converts the iodine into hydriodic acid.
On the preparation of the solution for separating the metals
of Group III., the presence or absence of phosphates and
silicates will have been ascertained.
Besides these, indications of the presence of several acids
will have been obtained during the preliminary examination
for metallic radicals.
Preparation of the Solution for the Detection of Acids.
Before testing for the acids, it is advisable (in many cases
it is necessary) that they should be present in the solution as
salts of the alkalies (or alkaline earths); that is to say, the
metals with which the various acids are united in the substance
under analysis should be exchanged, by double decomposition,
for an alkali metal ; for the reason that the presence of these
other metals would in many cases mask the reactions by which
the acids are to be detected. To accomplish this, the solution
is boiled, 1 and sodium carbonate added in quantity slightly in
excess of that required to effect complete precipitation. The
mixture is then filtered, and the acids, now present as their
sodium salts, are detected in the solution. 2
1 In cases where the substance under analysis is insoluble, the product
obtained by fusion with sodium carbonate is extracted with water, and the
solution so obtained is employed for the detection of the acids (see p. 113)-
2 This method for separating the metals from their acids is not,
ever, of universal application. Thus, in the case of many of the phosphates
Systematic Detection of the Acids 117
GENERAL TESTS. (i) A small portion of this solution is
carefully neutralised by first adding dilute nitric acid drop by
drop (the mixture being heated to expel carbon dioxide), until
the liquid is jus f add, and then adding a drop or two of dilute
ammonia.
This neutral solution is then tested by the addition of
barium chloride. If no precipitate is obtained, the acids of
Group I. are absent. If a precipitate is formed which re-
dissolves on the addition of hydrochloric acid, sulphuric acid
is excluded.
(2) A second small portion of the solution is acidified with
nitric acid, and tested with silver nitrate. A negative result
proves the absence of the acids of Group II.
If these general tests show that acids of both Groups I.
and II. are present, special tests must then be applied in
separate portions of the solution, for such acids as, from
information already gained, are considered likely to be in the
substance under analysis, and which have not been definitely
discovered during the course of the examination. The tests
may be made in accordance with the following outline scheme,
in which most of those acids which must certainly have been
detected in the earlier stages are not again mentioned.
A. In portions of the solution acidulated with hydrochloric
acid.
Sulphuric Acid. Barium chloride precipitates white barium
sulphate : not dissolved by the addition of strong hydrochloric
acid, and boiling the liquid.
Silicic Acid. Ammonium carbonate (but not ammonia)
precipitates silicic acid.
Arsenic Acid. If arsenic has been found among the metals
and its state of oxidation (i.e. whether present as an arsenite
or arsenate) has not been determined, the following test may
be applied : Ammonium chloride, ammonia, and magnesium
held in solution by acids, the action of the sodium carbonate is to cause
the precipitation of the phosphates themselves ; in such instances, therefore,
the acid is not found in the solution. Since, however, phosphoric acid will
have been discovered during the examination for the metals, this fact is of
little consequence.
1 1 8 Smaller Chemical A nalysis
sulphate are added a white precipitate may be due to either
a phosphate or arsenate. If phosphoric acid has been proved
to be absent (by the molybdate reaction), it must be the
arsenate. In either case it should be filtered, and after being
washed free from ammonium chloride, it is dissolved in a little
nitric acid, and silver nitrate added (or a drop or two of silver
nitrate may be poured upon the washed precipitate in the
funnel). A brown precipitate indicates an arsenate.
B. In portions acidulated with nitric acid.
Hydrochloric. Silver nitrate gives a white precipitate.
Hydrobromic and Hydriodic Acids. Silver nitrate gives
a yellowish precipitate.
(For the methods of discriminating between these, see
p. 89.)
C. In portions acidulated with acetic acid.
Hydrofluoric Acid. Calcium sulphate or chloride precipi-
tates calcium fluoride : only slightly soluble in hydrochloric
acid.
(Chromic, phosphoric, and arsenic acids may also be looked
for in portions of this solution. These acids, however, will
have been detected at an earlier stage of the analysis.)
CHAPTER XII
SIMPLE VOLUMETRIC DETERMINATIONS
WHEN sulphuric acid is added to sodium hydroxide we know
that the acid gradually neutralises the alkali, and that a point
is reached when the solution shows neither acid nor alkaline
properties it is, in fact, netitral. We indicate this action by
the equation
H 2 S0 4 + 2NaHO = Na 2 SO 4 + 2 H 2 O
We know also that this means that 2 + 32 -f- 64 = 98 grams of
sulphuric acid are capable of neutralising 2 (23 + i 4- 16)
= 80 grams of sodium hydroxide. If therefore we have a solu-
tion of sulphuric acid, the strength of which is such that i litre
contains 98 grams of H 2 SO 4 ; and suppose, further, that it took
exactly 100 c.c. of this acid to neutralise a given solution of
caustic soda of unknown strength, we should then have deter-
mined the actual weight of sodium hydroxide present in the
latter solution to be 8'o grams. Or, again, if the alkali were in
the solid state, a weighed quantity of it might be dissolved in
water so as to make a known volume of solution, and the
volume of acid necessary to exactly neutralise either the whole
or a measured portion of it would enable us to find the per-
centage of alkali in the solid. By this operation of measuring the
volume Q! acid, we ftxu&determine the weight of caustic soda present.
The reagents of known strength (the acid in the above
illustration) which are used in volumetric analysis are called
standard solutions.
Weighing. It will be evident that Weighing must be at the
root of all volumetric processes, for not only must the materials
to be analysed be accurately weighed, but the standard solutions
themselves can only be prepared (in most cases) by exact
weighings of the materials' contained in them.
In order to obtain the weight of a substance with sufficient
119
I2O Smaller Chemical Analysis
exactness for these volumetric purposes, a moderately delicate
balance must be used, and the operations conducted with some
care. The object to be weighed is placed upon the left scale-
pan, and a weight, which by a guess is judged to be rather
greater than that of the object, is placed upon the opposite
scale by means of the forceps 1 the balance being at rest. 2
The beam is then liberated from its supports by means of the
lever. Suppose the 20-gram weight had been taken, and it is
found to be too little, the beam is brought to rest, and a 10-
gram weight added. If this is too much, it is returned to its
place in the weight-box, and the 5-gram substituted. If this is
too little, the 2 -gram is added if still too little, the i-gram is
added; and should this be too much, the weight then lies
between 27 and 28 grams. The same systematic process is
continued with the subdivisions of the gram, until the object is
so counterpoised that, as the beam gently swings, the pointer
oscillates along the scale to practically the same distance in
both directions. The result of a weighing operation should be
recorded before removing the weighed object from the scale. The
value of the weights should first be read off from the empty
spaces in the weight-box? and the result then checked as the
weights are returned in order from the highest to the lowest to
their places in the box.
With the exception of pieces of metal, alloys, etc., sub-
stances to be weighed must never be placed directly on the
balance, but must be contained in a suitable vessel. When
taking a weighed quantity of a substance for analysis it is usual
to obtain its weight by difference, in one of two ways : (i) some
suitable vessel, such as a watch-glass or porcelain crucible, is
weighed being perfectly dry and clean and then a sufficient
quantity of the substance is placed in it, and the whole re-
1 None of the weights, or the moving parts of the balance, must be
handled with the fingers.
2 The balance must always be brought to rest before placing anything
upon, or removing anything from, the pans.
3 This implies two things, (i) that the box of weights is complete, and
(2) that the student tidily returns each weight to its proper place which has
been taken out but is not in actual use upon the balance when the weighing
is completed.
Simple Volumetric Determinations 12 1
weighed ; or (2) the substance contained in a light, thin glass
bottle (a weighing bottle} is weighed, and a sufficient quantity of
it is then carefully tipped out into a beaker or flask, as the case
may be, and the bottle and remaining contents re-weighed. In
either case the difference between the two weighings is the
weight of the substance employed.
Standard Solutions are usually made of such a strength that
the quantities of the positive or negative constituents, or of
what may be called the active constituents of the compounds in
the solutions shall bear the same relation to each other as the
numbers which express their chemical equivalents ; that is to
say, equal volumes of the different solutions will contain equiva-
lent proportions of the effective constituent of the substance in
solution. For example, standard solutions of NaCl and AgNO 3
will be of such strengths that whatever volume of the first con-
tains 35-5 grams of chlorine, the same volume of the other
shall contain 108 grams of silver ; or, again, standard solutions
of HC1, NaHO, and Na 2 CO 3 will be of such strengths that what-
ever volume of the first contains i gram of hydrogen, the same
volume of the others shall each contain 23 grams of sodium.
Normal Standard Solutions are solutions of such a strength
that the particular volume which contains i gram of hydrogen,
23 grams of sodium, 35*5 grams of chlorine, etc., is i litre.
A normal solution of HC1, to contain i gram of hydrogen in
the litre, must obviously contain 1 + 35*5 = 3^'5 grams of
HC1 per litre. Similarly, normal NaHO, to contain 23 grams
of sodium to the litre, must contain 23 + 1 + 16 = 40 grams
NaHO per litre ; while a solution of Na 2 CO 3 , to contain 23
grams of sodium to the litre, must obviously contain
= 53 grams of the salt in the same volume, and normal H 2 SO 4
must contain 2 "*" ^ 2 _4 = 49 grams of that acid.
In some cases it is only a portion of some particular radical
or element present in the solution which takes an active part in
the particular chemical reactions for which the solution is used.
For example, potassium dichromate, K 2 Cr 2 O7, is employed as
an oxidising agent, but only three out of the seven oxygen
122 Smaller Chemical Analysis
atoms in the salts are available for this purpose ; the compound
may be regarded as breaking down into K 2 O, Cr 2 O 3 , 30. A
normal solution of this salt, therefore, when it is to be employed
for oxidation reactions, must contain 8 grams of available
oxygen (i.e. the weight of oxygen equivalent to i gram of
hydrogen) ; it must therefore contain (K 2 O) ^ + (Cr 2 O 3 )-^|
+ (30) } y =49*13 grams of the dichromate in i litre.
If a standard solution of potassium dichromate were
required for precipitating an insoluble chromate by double
decompositon such as lead chromate, or barium chromate, then
the normal solution would have quite a different strength. In
this case the reaction would be merely the replacement of
potassium by another metallic radical, and the normal solu-
tion would be required to contain 39 grams of K to the litre ;
therefore (K^-f + (Cr 2 )^ + (O 7 )~ = 147 '4 grams of
potassium dichromate would be the weight of salt in i litre.
Standard solutions of one-half, one-tenth, and one-hundredth
of the strength of a normal solution are called respectively semi-
normal, deci-normal, and centi-normal solutions. In preparing
standard solutions the weighed quantity of salt is not added to
and dissolved in the measured volume of water, but it is first
dissolved in a moderate quantity of water, and the solution then
made up to the exact volume required by adding water. This
operation is carried out in a flask capable of holding either i
litre or J litre when filled to a graduation mark upon the neck.
Pipettes are practically glass tubes with a bulb or enlargement
upon them, and drawn to a point at one end. Fig. 1 1 shows
two forms of pipette. Like the flasks, they have one graduation
mark upon the stem, and are made of various capacities such as
5, 10, 15, 20, 50, and 100 c.c. A pipette is filled by sucking the
liquid up until it is somewhat higher than the graduation mark
(avoiding sucking it up into the mouth), and then quickly cover-
ing the upper end with the finger. Then, by slightly releasing
the pressure of the finger, the liquid is allowed to drip slowly
out until the level is exactly opposite the graduation mark. When
Simple Volumetric Determinations
123
the contents of the pipette are allowed to flow out, the drop or
two which remain in the pointed end may be blown out.
Burettes. The burette is a long straight glass tube, one end
of which is drawn down and terminated by a glass stop-cock,
or connected to a jet by
means of a caoutchouc tube
which can be closed by means
FIG.
FIG. 12.
of a pinch-cock. Fig. 12 shows the two forms of apparatus.
The burette is graduated almost throughout the length, the
graduations being usually tenths of a cubic centimetre. The
size most commonly used has a capacity of 50 c.c. For
ordinary use the common retort-stand and clamp shown to the
left in the figure make a very convenient stand for holding
burettes. If desired, several can be supported upon the same
retort-stand. The instrument is filled by means' of a small
funnel placed in the top (which should be removed afterwards,
lest any adhering drops fall into the burette), until the liquid is
considerably above the topmost graduation. The tap or pinch-
cock is then momentarily opened, in order that the liquid may
124
Smaller Chemical Analysis
sweep out before it the air which is contained in the tap and
jet. This will not be successfully accomplished if the tap or
pinch-cock is only gradually opened, as the liquid then slips
down the walls of the narrowed portions and leaves air-bubbles
in the tubes, which it is absolutely necessary to remove before
the instrument can be used with exactness.
To read a burette. Owing to the action of capillarity, the sur-
face of a liquid contained in a glass tube is not plane, but curved.
This curved surface is called the meniscus. Fig. 13 shows
the meniscus in the 'case of water contained in a burette.
In reading the burette, it is usual to take the graduation
FIG. 13.
18-
FIG. 14-
which coincides with the lowest point of the meniscus. Thus,
in Fig. 14 the correct reading would be 16-5. If, as sometimes
happens in certain lights, the line of the curve is not very
clearly visible, it may usually be made quite distinct by holding
behind the tube a piece of white paper, inclining it slightly
upwards, as in Fig. 13. In order to make a correct reading, it
is necessary that the eye of the observer should be in the same
horizontal plane as the surface of the liquid. The reason for
this will be evident from the diagram (Fig. 14), where an error
of one graduation would arise by reading from either of the
positions A or A' instead of from B.
Simple Volumetric Determinations
125
FIG. 15.
A simple device to ensure that the readings of a burette
shall be consistently made from a correct point of observation
is shown in Fig. 15. It con-
sists merely of a narrow
strip of card folded in the
middle, with the two free
ends pinned together with
a paper-fastener. This is
slipped over the burette,
and while it allows of being
easily slid up and down the
tube, it will also hold itself
in any position in which it
may be put. When taking
a reading, this little clip is
placed so that its upper edge
is just a little below the level
of the liquid : then, when
the eye is in such a position
that the back and front edges of the clip just coincide, it will
also be practically in the same horizontal plane as the bottom
of the meniscus.
Indicators. It is essential to a volumetric process that the
exact point when the chemical action is complete should be
readily discerned. This is often accomplished by the presence
of some substance which will change its colour as soon as
the reagent is in the slightest excess. Such substances are
called indicators. The most familiar example is litmus, which
is reddened by acid and rendered blue by alkalies. Sometimes
the reagent itself is the indicator ; potassium permanganate, for
instance ; so long as this reagent is being utilised in the chemical
action of oxidation, so long will the violet colour of each added
drop be destroyed, but the first drop after the completion of
the reaction imparts its violet tint to the liquid. When acids
or alkalies are being determined, either litmus or methyl orange
will answer the purpose of an indicator. When carbon dioxide
is eliminated in the neutralisation, as when alkali carbonates
are being used, litmus cannot be used tmless the liquid be boiled,
126 Smaller Chemical Analysis
since otherwise the CO 2 dissolving in the liquid would exert an
add action on the litmus. Under these circumstances methyl
orange is preferable.
Methyl orange gives an orange-coloured solution (yellow in
presence of alkali) which is changed to a pink-red by acids.
Only one or two drops of a very dilute solution should be used.
Most volumetric processes belong to one of the three
following classes :
1. Those that are based upon the neutralisation of adds and
alkalies.
2. Those based upon processes of oxidation or reduction.
3. Methods by precipitation.
In the following section two or three typical examples of
each of these classes will be given.
I. Volumetric Exercises based on Neutralisation of Acids
and Alkalies.
The first requisite is a standard solution of some acid or
alkali which can readily be prepared with a reasonable degree
of accuracy, to serve as a foundation or basis upon which to
make up any other acids or alkalies that may be required, for
obviously we must start with a solution whose strength is exactly
known before it can be used as a measure of the acidity or alka-
linity of some other solution. To attempt to start with either
sulphuric or hydrochloric acid would involve serious complica-
tions, for although we may exactly weigh out 49 grams of strong
sulphuric acid in order to dilute it to the strength required for a
normal solution, we do not know what is actually the strength of
the strong sulphuric we are weighing. It certainly is not H 2 SO 4
and nothing else, but is H 2 SO 4 + an unknown amount of water.
Caustic soda, again, rapidly absorbs moisture from the air
how much of this it has already absorbed while in its bottle
is an unknown quantity, and, moreover, it quickly absorbs
more while it is actually being weighed out. Therefore, with
however much care and expedition we weigh out 40 grams of
caustic soda and make it up to a litre of solution, we are still
ignorant of the exact weight of the alkali the solution contains.
These difficulties are avoided by starting with sodium carbonate.
Neutralisation of Acids and Alkalies 127
Normal Sodium Carbonate
(53 grams of Na 2 CO 3 /^ litre)
In order to obtain this salt in a state of purity, it is prepared
by heating the purest sodium bicarbonate to a dull red heat
until no further loss of carbon dioxide and water takes place. 1
Theoretically, 84 grams of bicarbonate should yield 53 of the
normal salt; a slight excess of this proportion, therefore,
should be employed. To prepare half a litre of the standard
solution, about 43 grams of the pure sodium bicarbonate are
heated in a weighed platinum dish to a low red heat for about
10 or 15 minutes. The salt must not be allowed to fuse. It
is then cooled in a desiccator and weighed. To ensure that
the decomposition has been completed, the dish is again heated
for another 10 minutes, and, after cooling in the desiccator,
weighed again. The weight of the salt (after deducting the
weight of the dish) will be a little over 26*5 grams. By means
of a clean spatula or pen-knife, a small quantity is removed,
so as to bring the weight to exactly 26-5 grams, the operation
being performed without undue exposure of the dry salt. The
contents of the dish are then washed out into a beaker, the
dish being thoroughly rinsed with warm water, and the salt
completely dissolved by stirring the mixture with a glass rod.
The solution is then carefully poured into a half-litre flask, the
beaker being several times rinsed with water. The flask is
then carefully filled to the graduation mark, after which the
stopper is inserted, and the contents thoroughly mixed by
shaking.
Instead of bringing the weight of the sodium carbonate to
the exact quantity required for half a litre of solution, the
whole of the salt in the dish may be employed, and the exact
volume of water which will be required in order to make the
solution of normal strength is calculated. Thus, suppose the
weight of sodium carbonate in the dish after heating is 26749
grams, instead of 26-5 ; then
26-5 : 26749 : : 500 c.c. : 5047 c.c.
1 Of the so-called " pure " salts of commerce, the bicarbonate is of a
higher degree of purity than the normal salt.
128 Smaller Chemical Analysis
Hence, after the solution has been made up to the graduation
mark in the manner described above, 47 c.c. of water are
added from a burette, and the solution then finally shaken up
to ensure thorough mixing.
This solution, if its preparation has been carefully carried
out, should now be of exact "normal" strength, and, as all
analytical determinations in this class are made on the basis
of this assumption, it is evident that great care should be
taken to ensure accuracy.
i c.c. of this solution then contains 0*053 gram Na. 2 CO 3 ;
it is therefore capable of neutralising, or is equivalent to 0*049
gram H 2 SO 4 , and 0*0365 gram HC1.
By means of this standard sodium carbonate, it is now
easy to prepare either standard sulphuric or hydrochloric acid.
Normal Sulphuric Acid
(49 grams 0/H 2 SO 4 /*r litre)
The specific gravity of ordinary oil of vitriol being about
1*8, 49 grams will be rather less than 30 c.c. Hence, if this
volume be measured out and diluted up to a litre, a solution
will be obtained which will have a rough approximation to the
required strength. If this solution is then titrated with the
standard sodium carbonate, its actual strength can be ascer-
tained; and by calculation, the volume of water which must
be added in order to bring it to the exact normal strength is
determined.
Dilution of the Acid. Thirty cubic centimetres of pure
sulphuric acid are gradually poured into about 150 c.c. of
water in an ordinary flask. The mixture is then cooled by
holding the flask under the water-tap, and allowing a stream
of water to run over it, at the same time shaking the liquid
round within the vessel. When quite cold, the solution is
transferred to a litre flask, and the volume made up to the
1000 c.c. mark by the addition of cold water.
Titration of the Acid. A burette is first filled with the
dilute acid, 1 care being taken to remove all air-bubbles from
1 Whenever burettes or pipettes are employed for measuring standard
solutions, they must either be dry before use, or, if moist, they must be
Neutralisation of Acids and Alkalies 129
the tap, as explained on p. 123. Twenty-five cubic centi-
metres of the normal sodium carbonate are then transferred
by means of a pipette * to a small beaker, and a single drop
of the methyl orange indicator is added. The beaker is then
placed upon a white glazed tile beneath the burette, and the
acid gradually run into the alkaline liquid, the solution being
gently rotated in order to ensure thorough mixing after each
addition of acid. At first 2 or 3 c.c. of the acid may be
added at a time ; but as the point of neutrality is approached,
smaller and smaller quantities are added at once, until they
are reduced to a few drops only, and at last the addition of
a single drop produces a permanent pink colour in the liquid.
Correction of the Acid. The exact volume of the acid
which has been used in order to neutralise the 25 c.c. of normal
sodium carbonate is noted, and from it the volume of water
which must be added in order to make the acid exactly normal
is calculated. Thus, suppose instead of 25 c.c. of acid, 23-8
c.c. were used in neutralising 25 c.c. of the normal alkali ; then
25 : 23*8 : : 1000 : 952
That is to say, 952 c.c. of the acid contain as much sulphuric
acid as should be contained in 1000 c.c., if it were exactly
normal. If, therefore, 952 c.c. of this acid be measured out
into a litre-flask, and the volume be then made up to the litre
by the addition of water, a correct normal acid will be obtained.
The solution thus obtained should be once more titrated
with the sodium carbonate.
When the strength of the acid is very nearly, but not exactly ',
normal, instead of attempting to bring it to the precise normal
strength, it may be used as it is, and a correction introduced
every time by means of a factor. For example, suppose every
25 c.c. of normal sodium carbonate required 24*6 c.c. of the
acid instead of 25 c.c. in order to neutralise it; then
24 6 : 25 : : i : roi6
That is to say, every cubic centimetre of this acid is equal to
first rinsed out with a small quantity of the solution ; otherwise that portion
which is measured out for use would be slightly diluted by the water
adhering to the walls of the instrument.
1 See above note.
K
130 Smaller Chemical Analysis
1*016 c.c. of an exactly normal acid; therefore, if the number
of cubic centimetres of this acid used in any titration be
multiplied by the factor 1-016, the result will be the number
of cubic centimetres which would have been required if the
acid had been strictly normal.
Or again, if the acid should be a little weaker than the
exact normal, a similar correction can be made. Thus, suppose
25 c.c. of normal sodium carbonate required 25*2 c.c. of acid
instead of 25 c.c. in order to reach the neutral point, then
25*2 : 25 : : i : 0*992
That is to say, each cubic centimetre of the acid is in reality
only equivalent to 0-992 c.c. of normal acid; therefore in this
case 0*992 is the factor by which the number of cubic centi-
metres of this acid used would have to be multiplied in order
to convert them into cubic centimetres of normal acid.
Titration of the Acid, using Litmus as Indicator. In
the absence of methyl orange, the titration of the acid by
means of sodium carbonate may be carried out with litmus
as the indicator. In this case, however, it is necessary to
boil the solution in order to expel the carbon dioxide. The
acid is added gradually, until the colour of the litmus changes
from blue to purple-red. The solution is then boiled, and as
the carbon dioxide is expelled the colour returns to the original
blue shade. After boiling for a few minutes, acid is admitted
in small quantities, and the liquid boiled up after each addition,
until at last the addition of a single drop gives a permanent
bright-red colour, which does not change on boiling.
Having now prepared and standardised the normal sulphuric
acid, its exact strength or value should be plainly indicated on
a label. Thus, in the case of the first of the two examples
above, the factor i'oi6 should be written on the label. It is
also useful to indicate the actual value of the acid per cubic
centimetre in terms of alkali, which is readily calculated. For
example, 24*6 c.c. of the acid in question contains as much
H 2 SO 4 as would be contained in 25 c.c. were the solution
strictly normal.
One litre of this slightly stronger acid, therefore, instead
Neutralisation of Acids and Alkalies 131
of containing 49 grams H 2 SO 4 , would contain 49*8 grams, or
i c.c. = 0*0498 gram H 2 SO 4 .
Then to find its value in terms of Na. 2 CO 3 , we have the
proportion
0-049 : 0-0498 : : 0*053 : 0-0538
Similarly for NaHO
0*049 : 0*0498 * : 0*040 : 0*0406
The label for this acid should therefore carry these figures
i c.c. = 0*0538 gram Na 2 CO 3
= 0*0406 NaHO
= 0-0315 Na 2 O
Examples of Analyses by means of Standard Acids and
Alkalies
i. Determination of the Total Alkali in a sample of
soda-ash
From 5 to 10 grams of the soda-ash are weighed out into
a flask, and dissolved in cold water, and the solution made up
to 500 c.c. in a half-litre flask, and thoroughly shaken. Then,
by means of a pipette, 50 c.c. of this solution are transferred
to a small flask and a drop of methyl orange added. Standard
acid is run in from a burette (see note on p. 128) until the pink
colour is visible. 1
At least two determinations should be made after the first
rough trial experiment, and if they fairly agree the mean may
be taken as the basis for calculating the result.
1 When neutralising a solution of an entirely unknown degree of
alkalinity, one must either add the reagent a few drops at a time from
the beginning often a slow process, which may occupy much time or, if
it is added in larger quantities one must run the risk of overstepping the
mark. In practice it is best to adopt the second plan for the first experi-
ment, so as to ascertain roughly the volume of acid required. Then a
second measure of 50 c.c. is taken, when the acid may be fairly quickly
run in until within I or 2 c.c. of the required volume, after which it should
be added drop by drop.
132 Smaller Chemical Analysis
EXAMPLE. Weight of soda-ash taken = 875 grams.
Dissolved and made up to 500 c.c.
(1) 50 c.c. taken, vol. of acid reqd. between 13 c.c. and 14 c.c. (rough).
(2) > ,, I3' 2 c - c -
\3/ ?> > *3 4 -
Mean = 13'3
Strength of standard acid, i c.c. = 0*0315 Na 2 O
Therefore weight of Na 2 O in 50 c.c.) m'3 x 0*0315 = 0*41895
of the alkali solution >\ gram
And since 50 c.c. = one-tenth the)
total, therefore the weight of Na 2 O> = 4*1895 grams
in the original weight of soda-ash J
Hence the percentage of Na 2 O (on _ 4*1895 x 100 _
total alkali) }~ 875
A volumetric determination of an acid by means of standard
sodium carbonate is carried out exactly as in the operation
of standardising the normal acid solution, except that a
measured volume of the acid of unknown strength is transferred
by means of a pipette to a small flask, a drop of methyl orange
added, and the standard alkali added from a burette until the
pink colour of the indicator is just destroyed.
Other Standard Acids and Alkalies. It will be obvious
that the standard sodium carbonate and sulphuric acid can
be used for preparing other standard acids and alkalies, such
as caustic soda or hydrochloric acid. In the former case
about 23 or 24 grams of the purest available caustic soda are
dissolved in water and made up to 500 c.c. This gives a
solution stronger than the normal, but whose strength is prac-
tically an unknown quantity. To ascertain its exact strength
it is titrated with the standard acid, using 25 c.c. of the alkali
measured with a pipette, and one drop of methyl orange as
indicator. From this determination the volume of water
necessary to add in order to dilute the alkali to the normal
is calculated, as explained in the case of sulphuric acid.
II. Exercises based on Processes of Oxidation and Reduction.
Three oxidising agents will be considered, namely, potas-
sium permanganate, potassium dichromate, and iodine.
Processes of Oxidation and Reduction
133
A. Decinormal Potassium Permanganate (3'i6 grams of
KMnO 4 per litre).
To prepare this solution 3*2 grams (as near as possible)
of the ordinary " pure " salt are weighed out and dissolved
in water in a litre flask. When entirely dissolved the flask
is filled to the graduation mark, and the contents well shaken.
Titration of Potassium Permanganate by means of
Ferrous Sulphate. An exactly deci-normal permanganate
solution will contain 0*0008 gram available oxygen per i c.c.
Therefore i c.c. is capable of oxidising 0-0056 gram of ferrous
iron to the ferric state ; it is then equivalent to this weight of
iron. The aim of this titration is to determine with the greatest
possible care the exact strength of the permanganate. For this
reason the ferrous sulphate is prepared by dissolving 0*5 gram
of the purest soft iron wire in
sulphuric acid, with exclusion
of air, the wire being clean
and free from rust. 1
About 80 c.c. of dilute sul-
phuric acid (i part acid to 5
parts water) are placed in a
2 50 -c.c. flask fitted with a
rubber cork and bent glass
tube. The air in the flask is
then expelled by removing
the cork and introducing two
or three crystals of pure
sodium carbonate, the flask
being in a vertical position.
As soon as the carbonate has
dissolved, the weighed quan-
tity of iron is dropped in.
The cork is then inserted,
and the flask supported in the
manner shown in Fig. 1 6, with
the tube dipping into a solution of sodium carbonate in a small
1 The fine iron binding wire used for flowers is the best for the purpose ;
it contains 99'6 per cent of iron.
FIG. 16.
134 Smaller Chemical Analysis
beaker. The flask being in this inclined position, the fine spray
thrown up during the solution of the iron strikes against the sides
of the flask and falls back into the liquid. The flask is gently
heated by means of a small flame until the iron is wholly
dissolved, and only a few minute particles of carbon remain.
The lamp is then withdrawn, and the flask allowed to cool.
As it does so, the solution in the beaker is gradually drawn
up the tube, but the first drops which enter the flask at once
cause an effervescence of carbon dioxide which drives the
liquid down again, and at the same time fills the flask with
carbon dioxide. When it has partially cooled in this way,
the cork is removed, and air-free distilled water (prepared
by boiling the water, and again quickly cooling it) is added
until the solution is within about 20 or 30 c.c. of the gradua-
tion mark. The flask is then closed with a rubber stopper,
and the contents made quite cold by holding the vessel in
a stream of cold water. The solution is then made up to
250 c.c. by the further addition of cold air-free water.
Fifty cubic centimetres of this solution are transferred by
means of a pipette to a small flask, and diluted by the addition
of about half the volume of air-free distilled water. The flask
is placed upon a white tile, and the deci-normal permanganate
solution added from a burette 1 until the colour of the reagent
ceases to be destroyed, and a faint pink tint is imparted to
the solution.
Four separate experiments should be made, taking 50 c.c.
of the iron solution each time, in order to gain practice in
judging when the first appearance of the permanent pink
colour takes place. After the experience thus gained, in
subsequent duplicate titrations the volume of the reagent
used should agree to o'i of a cubic centimetre. From the
results obtained, the exact strength of the permanganate is
calculated; thus
0*5 gram of iron wire was dissolved in 250 c.c. of liquid.
Fifty cubic centimetres of the solution therefore contain o'i
1 For permanganate solutions a burette with a glass tap must be used,
as this liquid acts upon rubber, thereby, of course, becoming altered in
strength.
Processes of Oxidation and Reduction 135
gram of iron. But since the iron wire contained 99*6 per
cent, of Fe, the actual weight of iron present in 50 c.c. of the
solution was not 0*1 gram, but 0*0996 gram.
The mean of four titrations gave 17*69 c.c. as the volume
of permanganate required.
Then 17*69 c.c. : i c.c. : : 0*0996 gram : 0*00563 gram
Therefore the solution is very slightly stronger than the
exact deci-normal, since i c.c. should be equivalent to 0*0056
gram of iron ; it should therefore carry on its label its equivalent
value, thus
i c.c. = 0*00563 Fe
Estimation of Iron in a Ferrous Salt. As an exercise
in the use of standard permanganate, an estimation of iron
may be made in a ferrous salt, the exact composition of which
is not known to the student. It might be ferrous sulphate,
or one of the double sulphates of ferrous iron and the alkalies.
Three or four grams of the salt are weighed out, dissolved in
water, and made up to 250 c.c. ; 50 c.c. of this are then trans-
ferred to a small flask, and 10 to 15 c.c. of dilute sulphuric acid
added. This is then titrated with the deci-normal permanganate.
Suppose the following data obtained :
Weight of salt taken = 3*5 grams
Dissolved and made up to 250 c.c. ; 50 c.c. employed for
each experiment.
Deci-normal permanganate
used (mean of four experi- =17*8 c.c.
ments)
Value of permanganate, i c.c. = 0*00563 Fe
Therefore weight of Fe in)
c f= 0*00563 X 17*8 = 0*1002 gram
Hence weight of iron in 3-5
grams of salt (i.e. in 250
0*1002 X 5 = 0*501 gram
c.c. of the solution)
Therefore percentage of I 0*501 X 100 _
iron in compound I 3*5
B. Deci-normal Potassium Dichromate^'^^ grams per litre)
(see p. 122).
136 Smaller Chemical Analysis
4^913 grams of the pure, dry, powdered salt are exactly
weighed out, dissolved in water in a litre flask, and the volume
made up to the graduation mark. Being deci-normal, one litre
will contain one-tenth of an equivalent of available oxygen, i.e.
0-8 gram ; hence i c.c. = 0*0008 gram available oxygen, and
is equivalent to 0*0056 gram Fe.
This solution may be used in a burette with a rubber tube
and pinchcock.
Titration of Potassium Bichromate by means of Ferrous
Sulphate. The ferrous sulphate is prepared by dissolving
pure iron in dilute sulphuric acid, with exclusion of air,
precisely as described for permanganate, p. 133. An aliquot
part of the solution say 50 c.c. is withdrawn by means of
a pipette, and transferred to a small flask, and the dichromate
solution gradually added from a burette.
In this process the end of the reaction is ascertained by
means of a freshly made and dilute solution of potassium
ferricyanide, used as an outside indicator. A number of drops
of the ferricyanide are placed about upon a white plate or tile,
and from time to time, during the addition of the dichromate,
a drop of the mixture is withdrawn upon a glass rod and
brought into contact with one of the drops of the indicator.
At first a strong blue coloration is produced, but as the amount
of ferrous salt is gradully diminished by the addition of the
dichromate, the blue becomes less and less intense, until at
last a drop of the liquid so tested fails to give any coloration.
At this point the whole of the ferrous salt has been oxidised,
and the reaction is therefore complete. 1
The mean of two or three titrations is taken, and from it
the exact strength of the dichromate calculated,' as in the case
of permanganate, and the true value of the solution in terms
of iron is indicated on the label, e.g. i c.c. = 0*00559 Fe,
which would mean that the solution was just a little below the
true normal strength.
In a number of instances, potassium dichromate may be
1 It will be evident that it is absolutely essential to the success of this
operation that the ferricyanide should be perfectly free from ferrocyanide,
otherwise the oxidised iron will itself give rise to a blue coloration.
Processes of Oxidation and Reduction 137
substituted for permanganate in volumetric analysis. This
is the case, for example, with all estimations that are based
upon the oxidation of ferrous to ferric salts.
Estimation of Iron in Iron Ores. About 2 grams of
finely powdered and dry red haematite are weighed out into
a flask and boiled with a small quantity of strong hydrochloric
acid, diluted with about half its own volume of water, until
the whole of the iron has been extracted, and the residue is
free from dark-coloured particles.
The next step consists in reducing the iron, which is at
present either partially or wholly in the ferric state. This may
be done by first diluting the liquid somewhat and introducing
into it a few fragments of pure zinc (i.e. free from iron). A
cork with a leading tube is then inserted. Hydrogen is evolved
by the solution of the zinc in the acid liquid, and the iron
existing in the ferric state is thereby reduced to the ferrous
condition. The action is allowed to continue until the zinc
is entirely dissolved, the process being aided towards the end
by the application of heat. In order to test whether the
reduction of the iron is complete, a drop of the liquid is with-
drawn upon the end of a fine glass rod, and brought into
contact with a drop of a solution of ammonium thiocyanate
upon a white tile. Any remaining ferric salt will be revealed
by the formation of the red colour ; in which case the process
must be continued by the addition of more zinc, and, if
necessary, of more acid also. When the reduction is complete,
the liquid is quickly cooled, and made up to 250 c.c. ; 25 c.c.
of this solution are then withdrawn with a pipette, transferred
to a small flask, and titrated with the deci-normal dichromate.
The presence of the zinc salt in the liquid interferes somewhat
with the delicacy of the indicator, 1 so that several titrations
should be made in order to practise the use of the indicator
under these conditions.
Made in this way, the result of the analysis is the estimation
1 For this reason the reduction is often brought about by the use of
stannous chloride, or by adding an alkali sulphite when iron is to be
determined by dichromate. The presence of zinc salts is immaterial when
ferrous permanganate is employed for the titration.
K 2
138 Smaller Chemical Analysis
of the total iron in the ore. Very often iron ores contain a
portion of the iron in the ferrous, and the remainder in the
ferric, state. These may be separately determined by first
dissolving a weighed quantity of the ore with exclusion of air,
in the arrangement shown on p. 133. In this solution the
ferrous iron is determined. Then a second weighed quantity is
dissolved and subjected to the reducing agent as above de-
scribed, and in this the total iron is estimated. The difference
between these gives the ferric iron originally present.
This analysis can be carried out equally well with perman-
ganate, but in this case an unnecessary excess of hydrochloric
acid must be avoided, and the liquid well diluted before titra-
tion, for the reason that permanganate reacts upon hydrochloric
acid with evolution of chlorine unless the acid is quite weak.
C. Deci-normal Iodine Solution (127 grams of iodine per
litre).
The solution is prepared by weighing out as exactly as
possible 127 grams of pure iodine, and adding to it in a litre
flask a solution of 20 grams of potassium iodide in 200 c.c. of
water. As soon as the iodine is entirely dissolved, the liquid
is diluted up to i litre. If strictly deci-normal, i c.c. would
contain 0-0127 gram iodine, and would be equivalent to
o '003 5 5 gram chlorine and 0-0008 gram oxygen.
In presence of water and certain oxidisable substances
(i.e. reducing agents), iodine unites with the hydrogen of the
water, and the oxygen so eliminated oxidises the reducing
agent. Thus arsenzV&r are converted into arsenal, and thio-
sulphates into tetrathionates. As a deci-normal solution of
sodium thiosulphate is required for many of the processes in
which iodine is used, this may be used in order to standardise
the iodine solution prepared as above.
Deci-normal Sodium Thiosulphate (24-8 grams of
Na 2 S 2 O 3 ,sH 2 O per litre).
24*8 grams of the purest salt are weighed out into a litre
flask, and dissolved in a moderate quantity of water. When
wholly dissolved, the volume is made up to the graduation
mark. If strictly normal, i c.c. will be completely oxidised by
i c.c. of the deci-normal iodine solution.
Processes of Oxidation and Reduction 139
Titration of Deci-normal Iodine with Sodium Thiosul-
phate. Twenty-five cubic centimetres of the thiosulphate
solution are transferred by means of a pipette to a small
beaker, diluted with a little water, and two or three drops of
clear starch solution added, to serve as an indicator. The
iodine is then run in from a stoppered burette until a single
drop gives a permanent blue colour.
This titration can equally well be performed in the reverse
order, and as in many analyses with standard iodine the solu-
tions are so used, the operation should here be practised in
both ways. Transfer 25 c.c. of the iodine solution to a small
beaker, and, before adding the starch indicator^ run in the thio-
sulphate solution from a burette until the brown colour of the
iodine has become paled to straw colour, then add the indicator,
which, of course, produces a blue coloration. The thiosulphate
is now added drop by drop until the colour is just discharged.
The result obtained must be the same whichever way the
operation is conducted.
If both solutions are exactly deci-normal, 1 i c.c. of iodine
= i c.c. thiosulphate. If there is only a slight discrepancy,
then the factor (see p. 129) should be indicated upon the
label of each solution ; for example, suppose 25 c.c. of the
iodine required 24-9 c.c. thiosulphate, then i c.c. of the latter
equals 25 -*T 24*9 = 1*004 c.c. iodine, and i c.c. of the iodine
= 0*996 c.c. thiosulphate. On the other hand, should the
titration show a wide disparagement between the two solutions,
then one of them (the iodine) must be tested by another
method (arsenious oxide).
Titration of Deci-normal Iodine with Arsenious Oxide.
The weight relation is expressed by the equation
As 2 O :j + 2l a + 2H 2 O - 4HI + As 2 O 5
That is to say, 127 parts of iodine will oxidise 49*5 parts of
arsenious oxide.
1 It is, of course, possible, although very improbable, that the two
solutions might be of exactly equivalent strength, although not exactly
deci-normal ; that is to say, an error to exactly the same extent might have
been made in preparing each.
140 Smaller Chemical Analysis
4*95 grams of resublimed arsenious oxide are weighed out
into a litre flask, and about 500 c.c. of water added. Then
30 grams of pure sodium bicarbonate are added, and the
mixture slightly warmed and continually shaken, until the oxide
is completely dissolved. The liquid is then cooled and diluted
up to i litre : i c.c. of this deci-normal solution should be
exactly equivalent to i c.c. of the iodine solution; 25 c.c. are
transferred to a small beaker, a few drops of the starch added,
and the iodine solution run in until the blue coloration is
permanent. From the result the exact strength of the iodine is
calculated.
The estimation of arsenic in an unknown solution of an
arsenious compound, or in an arsenite, is carried out by means
of deci-normal iodine solution exactly as this titration.
Estimation of Sulphur Dioxide or a Sulphite. As an illus-
tration of the use of standard iodine and thiosulphate in con-
junction, the following exercise may be carried out.
A dilute solution of sodium sulphite of unknown strength is
taken, and 50 c.c. of it are transferred to a small beaker by
means of a pipette. A measured volume of deci-normal iodine
is added, well in excess of what is required to oxidise the
sulphite, which is seen by the mixture having a brown colour.
The excess of iodine present is then ascertained by titration
with deci-normal thiosulphate added from a burette, the starch,
as before, being added when the brown colour changes to
yellowish. The excess of iodine thus determined, deducted
from the total originally taken, gives the iodine which has been
used to oxidise the sulphite, according to the equation
Na 2 SO 3 + I 2 + H 2 O = 2HI + Na 2 SO 4
That is to say, 127 parts of iodine are capable of oxidising 32
parts of SO 2 into SO 3 ; i c.c. of the iodine solution, therefore,
is equivalent to 0*0032 gram SO 2 .
Thus, suppose 50 c.c: deci-normal iodine to have been
originally used, and that the titration required 15 c.c. of the
thiosulphate, then 50 15 =35 c.c. of iodine solution were
used in oxidising the SO 2 present in 50 c.c. of the unknown
solution.
Processes based on Precipitation 141
Then, since i c.c. iodine = 0-0032 gram SO 2
0*0032 x 35 = 0*112 gram SO 2 in 50 c.c., i.e. 2*24 grams per
litre
Estimation of Chlorine. The element chlorine decomposes
potassium iodide, liberating its equivalent of iodine; i.e. 127
parts of iodine are liberated by 35-5 parts of chlorine. If the
iodine thus liberated is estimated by titration with a standard
solution of sodium thiosulphate, indirectly the weight of
chlorine which caused its liberation will be determined. In
this way the strength of a solution of chlorine water could be
determined; or the amount of available chlorine in bleaching
powder (i.e. the chlorine which is evolved when the compound
is acted upon by a dilute acid).
Estimation of Available Oxygen. Many substances contain-
ing oxygen (e.g. certain peroxides, chromates, permanganates),
when heated with hydrochloric acid, oxidise the acid with the
evolution of an amount of chlorine equivalent to the oxygen so
used. Thus MnO 2 contains one atom of available oxygen ;
and by the action of hydrochloric acid upon this oxide, this
one atom causes the evolution of 2 atoms of chlorine; or
8 parts of O result in the evolution of 3 5 '5 parts of Cl. If
this chlorine is made to act upon potassium iodide, it in its
turn liberates an equivalent of iodine, which can be estimated
by means of the standard thiosulphate. The iodine is there-
fore the indirect measure of the oxygen, and from this
obviously the amount of the oxygen compound can be
calculated. 1
III. Processes based on Precipitation.
Deci-normal Silver Nitrate (17*00 grams per litre). To
prepare a quarter-litre, 4*25 grams of pure silver nitrate are
weighed out into a 25o-c.c. flask and dissolved in water, and
the solution diluted up to the graduation mark ; i c.c. of this
solution, if strictly deci-normal, will contain o'oioS gram Ag,
and is equivalent to 0*00355 gram of Cl.
1 For details of such processes as are here merely hinted, the student
must consult larger manuals of analysis.
142 Smaller Chemical Analysis
Estimation of Chlorine in a Soluble Chloride. As an
exercise upon this process, a dilute solution of common salt of
unknown strength may be analysed.
Twenty-five cubic centimetres of the solution are transferred
to a small beaker by means of a pipette, and three or four drops
of a solution of potassium chromate (the normal salt) are added.
The deci-normal silver solution is then gradually run in until a
permanent reddish tinge is visible. The red colour is due to
the formation of silver chromate, which does not begin to form
until the whole of the chloride present has been precipitated as
silver chloride. The titration should be repeated once or twice
to gain practice in the use of this indicator.
Estimation of Silver. This process is exactly the reverse
of the former, and is carried out by means of a deci-normal
solution of sodium chloride. This solution is prepared by dis-
solving 5*85 grams of pure sodium chloride in water, and dilut-
ing the solution up to i litre. In conducting the titration the
indicator must not be added to the silver solution to be
estimated, but to the sodium chloride. A measured volume,
25 c.c., of the latter is transferred to a beaker, coloured with the
indicator, and the silver solution delivered from a burette. One
cubic centimetre deci-normal NaCl contains 0*00355 gram Cl,
and is therefore equivalent to 0*0108 gram Ag.
Deci-normal Ammonium Thiocyanate (7 -6 grams (NH 4 )CNS
per litre). When ammonium thiocyanate is added to silver
nitrate, a white precipitate of AgCNS is formed ; and if a drop
of a solution of a ferric salt (not the chloride) be added to the
silver solution, the development of the familiar blood-red colour
will indicate the completion of the precipitation.
About 8 grams of the thiocyanate are weighed out and dis-
solved to make i litre of solution; 25 c.c. of deci-normal silver
nitrate are transferred to a small flask, and 3 or 4 c.c. of
ferric sulphate solution added (previously made by dissolving a
crystal of ferrous sulphate in a little water in a test-tube, adding
about half its volume !of strong nitric acid and boiling for a few
minutes, then diluting with about twice the volume of water).
The ammonium thiocyanate is then run in from a burette. As
each drop enters, a red colour momentarily appears, but
Processes based on Precipitations 143
disappears on gently shaking the flask. The precipitation is
complete when a single drop causes a permanent red tint.
From the volume used the real strength of the thiocyanate is
ascertained, and the amount of dilution it requires to bring it
to exact deci-normal strength is calculated ; i c.c. deci-normal
thiocyanate should be equivalent to 0*0108 gram Ag, or 0*00355
gram Cl.
Estimation of Silver in a Silver Alloy. A weighed piece
of the alloy (such as a small silver coin) is dissolved in nitric
acid, and the solution made up to 250 c.c. ; 25 c.c. of this are
transferred to a small flask, the ferric sulphate indicator added,
and the deci-normal thiocyanate run in from a burette until the
red coloration is obtained. From this titration the percentage
of silver in the alloy is calculated.
ABRIDGED TABLE OF ATOMIC WEIGHTS (APPROXI-
MATE VALUES)
Aluminium ... Al ... 27
Antimony Sb ... 120
Arsenic ... As ... 75
Barium Ba ... 137
Bismuth Bi ... 208
Boron B ... 11
Bromine Br ... 80
Cadmium Cd ... 112
Calcium ... Ca ... 40
Carbon C ... 12
Chlorine Cl ... 35-5
Chromium Cr ... 52
Cobalt Co ... 59
Copper (cuprum} Cu ... 63*5
Fluorine F ... 19
Hydrogen H ... i
Iodine I ... 127
Iron (ferruni) Fe ... 56
Lead (plumbum} Pb ... 207
Lithium Li ... 7
Magnesium Mg ... 24
Manganese Mn ... 55
Mercury (hydrargyi-um} Hg ... 200
Nickel Ni ... 59
Nitrogen N ... 14
Oxygen O ... 16
Phosphorus P ... 31
Potassium (kalium) ... ... ... K ... 39
Silicon Si ... 28
Silver Ag ... 108
Sodium Na ... 23
Strontium Sr ... 87-6
Sulphur S ... 32
Tin Sn ... 119
Zinc Zn ... 65-4
144
INDEX
ACID, boric, 105
, carbonic, 101
, chloric, 89
v hydriodic, 87
, hydrobromic, 86
, hydrochloric, 84
, hydrofluoric, 90
, metaphosphoric, 101
, metastannic, 72
, nitric, 96
, nitrous, 98
, orthophosphoric, 100
, permanganic, 106
, phosphoric, 51, 100
, pyrophosphoric, 100
radicals, 83
, silicic, 1 02
, sulphuric, 93
, , standard solution, 128
Acids, analytical classification of, 1 1 5
Aluminium reactions, 32
Ammonium phosphomolybdate, 51
reactions, 21
thiocyanate, deci-normal solu-
tion, 142
Analytical classification, 16
groups, 20
tables
General table, 81
Group I., 79
II., Division I, 63
II., Division 2, 76
IIlA, 41
Analytical tables continued
Group IIlB, 50
III. (phosphate table), 52
IV., 31
V., 27
Antimoniuretted hydrogen, 72
Antimony reactions, 70
Arsenic reactions, 64
, volumetric estimation, 140
Atomic weights, table of, 144
BARIUM reactions, 28
Bismuth reactions, 58
Bleaching powder, available chlorine
in, 141
Blowpipe flame, 10
Borates, 105
Borax, 105
beads, 12
Boron fluoride, 91, 106
Bromides, 86
, iodides, and chlorides, detec-
tion in solution together, 89
Burettes, 123
CADMIUM reactions, 60
Calcium reactions, 29
Carbonates, IOI
Chlorates, 89
and nitrates, detection together,
98
, separation from chlorides, 90
Chlorides, 84
145
146
Index
Chlorides, detection in presence of
bromide or iodide, 85
Chlorine, volumetric estimation by
precipitation, 142
, in bleaching powder,
141
Chromium reactions, 33
Chromyl chloride, 85
Cobalt reactions, 48
Cobalticyanides, 49
Cobaltocyanides, 49
Copper reactions, 60
Cupric salts, 61
Cuprous salts, 61
ETCHING glass, 91
Evaporating to dryness, 4
Evaporation, 4
FILTRATION, i
Flame, oxidising and reducing, 10
Fleitmann's test, 70
Fluorides, 90
Fusion, 5
mixture, 42
with borax, exercises, 12
GENERAL reagents, 18
table for the separation of
the metals, 81
Group reagents, 19
IGNITION, 9
Indicators, use of, 125
Insoluble substances, treatment of,
112
Iodides, 87
, detection in solution with
bromides, 89
Iodine, deci-normal solution, 138
, estimations by means of, 140
Ions, 19, 83
Iron reactions, 36
, volumetric estimation of, in
ores, 137
LEAD reactions, 57
MAGNESIUM reactions, 25
Manganese reactions, 43
Marsh's test, 67
Meniscus, 124
Mercuric compounds, reactions, 54
Mercurous compounds, 56
Mercury reactions, 54
Meta-phosphates, 101
Methyl orange, 125
NEGATIVE radicals, 83
Neutralisation, exercises, 13
Nickel reactions, 46
Nitrates, 96
and chlorates, detection to-
gether, 98
and nitrites, detection together,
99
, reduction by ferrous salts, 97
, by sulphurous acid, 97
Nitrites, distinction from nitrates, 99
ORTHOPHOSPHATES, 100
Oxidation and reduction, 13
PERMANGANATES, 106
Phosphates, reactions of, 100
Phosphoric acid, removal of, in
Group III., 51
Pipettes, 122
Positive radicals, 18
Potassioscope, 23
Potassium dichromate, deci-normal,
135
, analyses by means of, 137
permanganate, deci-normal
solution, 133
, typical analysis by means
of, 135 f
reactions, 23
Precipitation, 7
Preliminary examination for acid
radicals, 114
Index
147
Preliminary examination for metallic
radicals, 107
exercises, i
Purple of Cassius, 73
Pyrophosphates, 100
RADICALS, 18
Reactions of the metals of Group I.,
77
II., Division I, 54
II., Division 2, 64
IIlA, 32
IIIB, 43
IV., 28
. V., 21
Reactions, wet and dry, 1 6
Reagents, 17
, group or general, 19
Reinsch's test, 66
SILICA, 103
Silicates, 102
, treatment of insoluble, 104
Silicon fluoride, 91
Silver nitrate, deci-normal, 141
reactions, 77
, volumetric estimation of, 142
Sodium chloride, deci-normal, 142
reactions, 22
thiosulphate, deci-normal, 133
Solution, 3
Solvent, 3
Spontaneous evaporation, 4
Standard solutions, 121
Stannic compounds, 73
Stannous compounds, 72
Strontium reactions, 29
Sulphates, 93
Sulphides, 92
Sulphites, 94
Sulphuretted hydrogen, 92
Sulphuric acid, standard solution,
128
Systematic detection of the acids, 115
TIN reactions, 72
VOLUMETRIC methods of analysis,
119
based on oxidation, 132
on precipitation, 141
WEIGHING, 119
ZINC reactions, 45
THE END
PRINTED BY WILLIAM CLOWES AND SONS, LIMITED, LONDON AND BECCLES.
14 DAY USE
RETURN TO DESK FROM WHICH BORROWED
LOAN DEPT.
This book is due on the last date stamped below, or
on the date to which renewed.
Renewed books are subject to immediate recall.
1 1 !un'58RH|
^3
, -
General Library
LD 21A-50m-8,'57 University of California
(C8481slO)476B Berkeley
YB 17026
889771
THE UNIVERSITY OF CALIFORNIA LIBRARY