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When a solution of the salt is carefully evaporated to dryness in
a dish and gently warmed, these colour changes will be rendered
evident ; and upon exposing the dried and white residue to the air,
or by gently breathing into the dish, the salt rehydrates itself, and

2i8 Inorganic Chemistry

is converted into the crimson compound having seven molecules
of water.

Many salts can have their combined water withdrawn by power-
ful dehydrating agents ; thus, if a crystal of copper sulphate ("blue
vitriol," CuSO 4 ,5H 2 O) be immersed in strong sulphuric acid, the
acid abstracts four out of the five molecules from the salt, leaving
the nearly white salt CuSO 4 ,H 2 O ; or when alcohol is added to a
solution of cobalt chloride, or to crystals of the salt, CoCl 2 ,6H 2 O,
the alcohol abstracts water, and the solution becomes blue.

When salts containing water of crystallisation are heated, it
frequently happens that a portion of the water is more easily parted
with than the remainder. Thus copper sulphate, CuSO 4 ,5H 2 O,
when heated to 100, parts with four molecules of water, leaving the
salt CuSO 4 ,H 2 O ; and in order to drive off this one remaining mole-
cule, the temperature must be raised above 200. Zinc sulphate
(or white vitriol), ZnSO 4 ,7H 2 O, in like manner loses six molecules of
water at 100, but retains the seventh until a temperature of 240 is
reached. In order, therefore, to distinguish between the water that
is more firmly held and that which is readily parted with, the term
water of constitution is frequently applied to the former, and the
fact is sometimes expressed in notation in the following manner:

CuSO 4 H 2 0,4H 2 O ; ZnSO 4 H 2 O,6H 2 O.

Natural Waters. On account of the great solvent powers of
water, this compound is never found upon the earth in a state of
absolute purity ; even rain, as it falls in regions far removed from
the dirty atmosphere of towns, not only dissolves the gases of the
atmosphere, but also small quantities of those suspended matters
which are always present in the air. As soon as the rain reaches
the earth, the water at once exerts its solvent action upon the
mineral matter constituting the portion of the earth's crust over
which it flows, and through which it percolates, and the liquid is
rapidly rendered less and less pure as it travels on its course to
lake or ocean.

Natural waters may be broadly divided into two classes, based
upon the amount of dissolved impurities they contain. If the sub-
stances in solution are present in excessive quantities, or to such an
extent as to be perceptible to the taste, the water is said to be a
mineral water; while, on the other hand, waters that are not so
rich in dissolved impurities are known as fresh "waters.

Natural Waters 219

Mineral Waters. The most exaggerated examples of mineral
waters are to be found in sea-water and in the waters of certain
lakes, which, having no outlet, are fulfilling the purpose of enormous
evaporating basins, in which the waters that flow into them are
undergoing evaporation and therefore concentration ; such, for
example, as the salt lakes of Egypt, the Elton lake in Russia, and
the Dead Sea. In waters of this description the total quantity of
dissolved solid matter is very considerable, and, as in the case of
the Dead Sea, is often deposited in crystalline masses round the
shores of the lake. The following table gives the total amount of
dissolved saline matter contained in 1000 grammes of certain of
these waters :

Irish Sea . . 33.86 I Dead Sea . . 228.57

Mediterranean Sea . 40.0 Elton Lake . . 271.43

As a typical example of a sea water, the composition of the
water of the English Channel may be quoted ; 1000 grammes of
this water contain

Sodium chloride .
Magnesium chloride
Magnesium sulphate
Calcium sulphate .
Potassium chloride
Calcium carbonate
Magnesium bromide



Water 964-745


Passing from these highly concentrated mineral waters, we find
a large number of spring waters which are classed as mineral, not
because the total quantity of foreign matter in solution is excessive,
but rather because they contain an abnormally large proportion of
a few special substances. Thus, large quantities of magnesium
sulphate and chloride are found in such springs as those at
Epsom and Friedrichshall. Others are found to contain consider-
able quantities of sodium sulphate and sodium carbonate ; while
those known as chalybeate waters contain ferrous carbonate in
solution. Spring waters that are charged with unusual quantities of
soluble gases are likewise placed in the category of mineral waters,
such as the waters of Apollinaris and Seltzer, containing large
quantities of carbon dioxide ; and the sulphur springs at Harrogate
and Aachen, which hold in solution sulphuretted hydrogen as well
as alkaline sulphides.

22O Inorganic Chemistry

Fresh Waters. The purest form of natural water is rain-water.
The average weight of solid matter dissolved in rain-water, col-
lected in the country and in perfectly clean vessels upon which it
exerts no solvent action, is found to be 0.0295 part in 1000 parts
of water. Collected in or near towns, rain-water always contains
a larger amount of dissolved impurities, such as nitrates, sulphates,
ammoniacal salts, and often considerable quantities of sulphuric
acid : it is the acid nature of the rain that causes so much damage
to stcne buildings.

The nature and extent of the contamination that rain-water
suffers after it has fallen must obviously depend very largely upon
geographical and geological circumstances, and therefore there are
no special features that are distinctly characteristic of waters from
rivers, lakes, or springs.

Thus, the total solid impurity in 1000 parts of water from the
river Dee at Aberdeen is 0.057, while that contained in the
Thames is 0.30 parts.

The water- of Loch Katrine only contains 0.032 part of solid
matter dissolved in 1000 parts, while that of Elton lake contains
as much as 271.43.

The same wide differences are also seen in spring waters from
different geological strata. Spring waters from granite and gneiss
rocks contain on an average 0.059 part of dissolved solid matter
in looo parts, while those from magnesian limestone average as
much as 0.665 P ar t- As a broad general rule, river waters are
found to contain less solid matter in solution than spring waters,
and these in their turn less than deep well waters. Thus, com-
paring waters from different sources, and selecting only such
samples as are known to be free from pollution from either sewage
matter or other abnormal impurities, it will be seen that, with
regard to the dissolved solid matter they contain, they fall in the
following order :

Total Solid Impurity Dissolved in 1000 Parts of
Unpolluted Waters.

Rain-water (average of 39 samples) . . .0295

Rivers and lakes (average of 195 samples) . .0967

Spring waters (average of 198 samples) . . .2820

Deep well waters (average of 157 samples) . .4378

Natural Waters 221

Hardness of Water. Certain of the salts that are very fre-
quently present as impurities in natural waters give to these
waters the property that is known as hardness. The chief com-
pounds that produce this effect are the salts of calcium and
magnesium. The term hardness is applied to such waters on
account of the difficulty of obtaining a lather, with soap, in the
ordinary process of washing. Pure soap may be regarded as a
mixture of the sodium salts of certain fatty acids (oleic, stearic,
palmitic, &c.), which are soluble in pure water. In the presence
of salts of calcium and magnesium the soap is decomposed, and
an insoluble curdy precipitate is formed by the union of the fatty
acid of the soap with the calcium and magnesium of the salts.
Until the whole of the hardening salts have in this way been
thrown out of solution, no lather can be obtained, and the soap is
useless as a cleansing agent ; but as soon as this point is reached,
the addition of any further quantity of soap at once raises a lather
on the water, and the soap is capable of acting as a detergent.
This process of precipitating the salts of calcium and magnesium
is known as softening, and in this instance the water is softened at
the expense of the soap.

Hard waters often become less hard after being boiled for a
short time, and this hardness which is so removed is termed the
temporary hardness. The degree of hardness which the water still
possesses after prolonged boiling is distinguished by the term
permanent hardness. The diminution of the total hardness of a
water by boiling is due to the fact that the soluble acid carbonates
of calcium and magnesium are decomposed during this process
into water, carbon dioxide (which escapes as gas), and the prac-
tically insoluble normal carbonates of these metals ; thus, in the
case of the calcium salt

CaH 2 (CO 3 ) 2 = H 2 O + CO 2 + CaCO 3 .

When such a water is boiled, the calcium carbonate is thrown
down as a white precipitate, which gradually collects upon the
bottom of the containing vessel. The " furring " of kettles, and the
formation of calcareous deposits in boilers, is largely due to this

In the case of waters that are highly charged with calcium car-
bonate, held in solution by dissolved carbonic acid, this deposition
of calcium carbonate may even take place at the ordinary tempe-
rature, owing to the diffusion of the dissolved carbon dioxide into
the air. It is in this way that those remarkable, and often beauti-

222 Inorganic Chemistry

fully fantastic formations, known as stalactites, have been produced
in certain subterranean caves. Water charged with the soluble
calcium carbonate, in slowly dropping from the roof of such a cave,
loses a portion of its dissolved carbon dioxide, and, in consequence,
deposits a certain amount of the calcium carbonate which was in
solution. Each drop, as it slowly forms, adds its little share of
calcium carbonate to the deposit, which thereby gradually grows,
much as an icicle grows, as a dependent mass called a stalactite.
Whether the water that drops from the stalactite has deposited
the whole of its calcium carbonate, will depend largely upon the
time occupied by each drop in gathering and dropping ; if, as often
happens, the whole has not been precipitated, the remainder is
deposited upon the floor of the cave, and a growing column of
calcium carbonate, called a stalagmite, gradually rises from the
ground until it ultimately meets the stalactite.

Clark's Process for Softening- Water. Waters whose hard-
ness is due to the presence of the carbonates of calcium and
magnesium can be deprived of their hardness by the addition to
them of lime. The amount of hardness is first estimated, and such
an amount of milk of lime is then added as is demanded by the
following equation :

CaH 2 (CO 3 ) 2 + CaO = H 2 O + 2CaCO 3 .

In this way the soluble calcium salt is converted into the insoluble
normal carbonate, which settles to the bottom of the tank.

The salts, which are mainly instrumental in causing the per-
manent hardness, are the sulphates of calcium and magnesium.
The degree of hardness and its particular order, that is, whether
temporary or permanent, will obviously be determined entirely by
the particular geological formation from which the water is derived.

The Permutit System of water softening is a recently in-
vented process, based upon the fact that if water containing
calcium carbonate in solution is made to filter through a stratum
of certain silicates containing sodium silicate, the following inter-
action takes place :

CaCO 3 + Na 2 SiO 4 = CaSiO 4 +Na 2 CO 3 .

The insoluble calcium silicate remains in the filter while the
soluble sodium carbonate passes into the water. The silicate
actually employed is an artificially produced compound to which
the coined name permutit has been given. When in course of
time this substance becomes exhausted, it may be regenerated

Natural Waters 223

by passing a solution of brine through the filter ; this reacts with
the calcium silicate, reforming the sodium compound, while soluble
calcium chloride passes away. It is claimed that this process
of regeneration may be repeated indefinitely.

Potable Waters, Undoubtedly the most important use to
which water is put is its employment as an article of food to man,
and since it has been proved beyond dispute that many virulent
diseases, such as cholera, typhoid fever, and others, are propagated
through the medium of drinking-water, it becomes a matter of the
greatest sanitary importance that the waters supplied for this pur-
pose should be as pure as possible. Excepting in very rare in-
stances, where poisonous mineral matters accidentally gain access
to drinking-water (as, for example, in the case of certain waters
which are capable of attacking, and to a slight extent dissolving,
the lead of the pipes through which they may be passed)^ the solid
matters that are usually found in waters are not injurious to health.
The living germs or bacilli, through whose agency zymotic diseases
are caused, cannot be detected in a sample of water by any direct
chemical analysis. A specimen of pure distilled water might
be artificially contaminated with such organisms so as to con-
stitute it a most virulent poison, and still chemical analysis
would fail to detect the danger, and the water would be pronounced
pure. Chemical analysis can, however, reveal the presence of
excrementitious matter, and also of the characteristic products re-
sulting from its decomposition : it can with certainty detect in the
water the evidence of recent contamination with sewage matters,
and it can also, with considerable precision, trace the evidences
of its having been so contaminated at an earlier stage of its history.
It cannot, however, distinguish between pollution with healthy and
with infected excreta, and therefore it is necessary to regard with
the greatest suspicion any water to which sewage has at any time
gained access. Waters that are made use of for drinking purposes
may be classified in the following order :

f i. Spring water.

Safe . . < 2. Deep well water.

( 3. Mountain rivers and lakes.

~ . . | 4. Stored rain-water.

( 5. Surface water from cultivated land.

( 6. River water to which sewage gains access.
Dangerous <

7. Shallow well water.

224 Inorganic Chemistry


Formula, H 2 O 2 .

Occurrence. This compound is occasionally found in small
quantities in the atmosphere, and also in dew and rain.

Modes Of Formation. ( i.) Hydrogen peroxide is produced in
small quantities during the burning of hydrogen in the air. If a
jet of burning hydrogen be caused to impinge upon the surface of
water, the temperature of which is not allowed to rise above 20,
the water will be found, after a short time, to contain hydrogen

(2.) This compound is also produced by the decomposition of
barium peroxide by carbonic acid. For this purpose a stream of
carbon dioxide is passed through ice-cold water, into which from
time to time small quantities of barium peroxide are stirred.
Barium carbonate is precipitated, and a dilute aqueous solution
of hydrogen peroxide is obtained.

BaO 2 + H 2 CO 3 =BaCO 3 + H 2 O 2 .

(3.) Barium peroxide may be decomposed by either hydrochloric,
sulphuric, silicofluoric, or phosphoric acid. Whichever acid be
employed, the barium peroxide, previously mixed with a small
quantity of water, is added gradually to the acid ; which, in the
case of either hydrochloric or sulphuric acid, should be diluted
with from five to ten times its volume of water. The temperature
of the mixture is not allowed to rise above 20. Thus, in the case
of hydrochloric acid

the soluble barium chloride is removed by the addition of sulphuric
acid, whereby barium sulphate is precipitated and hydrochloric
acid formed

BaCl 2 + H 2 SO 4 = BaSO 4 + 2HC1.

The hydrochloric acid may be removed by adding a solution of
silver sulphate, which precipitates silver chloride, leaving sulphuric
acid in solution

2HCl + Ag 2 SO 4 = 2AgCl + H 2 SO 4 .

And, lastly, the free sulphuric acid is withdrawn by the addition of
barium carbonate

H 2 SO 4 + BaCO 3 = BaSO 4 + H 2 O + CO 2 .

When sulphuric acid is employed for the decomposition of barium
peroxide, the crystallised, or hydrated peroxide (BaO 2 , 8H 2 O), is
* See "Chemical Lecture Experiments," new ed., p. 74.

Hydrogen Peroxide


most advantageous for the purpose. This salt, made into a paste
with water, is gradually added to the diluted and cooled acid, until
the acid is nearly but not quite neutralised. The slight excess of
acid is remove^-*t>y r ~fhe addition of the exact quantity of barium
hydroxide (b^r^ta=a^t^r)ccc&sarjHo neutralise it, and the insoluble
barium ^raHate is rem^wed by filtraNon. On a large scale silico-
fluoric ac$ or tfhosphcfric acid is ujuQly employed, preferably the

Matter, as it is

in hydrogen j iroxide__griaai^letaird its decomposition.


(4.) H
peroxide is
acid, when p(

Entities of free phosphoric acid

en ^roxide is al<

oxide by means

y obtained by decomposing
'tartaric acid. The potassium

-cooled ^tfong aqueous solution of tartaric
sep^-ates out, and an aqueous solu-
tion of hydrogen peroxide is obtained.

(5.) When small quantitie
for the purpose of illi(stratiffgits p
obtained by adding sodium per
hydrochloric acid, whereS^v so
oxide are formed, both of w

Jgen peroxide are required
ferties, it is most conveniently
dde to dilute and well-cooled
iloride and hydrogen per-

(6.) Hydrogen peroxide is formed in con^K^erabfe.Ojjantity when
ozone is passed through ether floating uporiN^ater.vErobably a
peroxidised compound of ether is first produceu}\wmfcj**' l t5 then
decomposed by the water. This production of hydrag^n r^roxide
may readily be demonstrated by placing a small ofuantit)N^f water
and ether in a beaker, and suspending into the wipour a sphcal of
platinum wire which has been gently heated. The combustion of
the ether vapour upon the wire, whereby the latter is maintained
at a red heat, is attended with the formation of ozone, and this
acting upon the ether, as already described, results in the pro-
duction of hydrogen peroxide, which may be detected in solution
in the water.

(7.) In small quantities, hydrogen peroxide is produced when
moist ether is exposed to the action of oxygen, under the prolonged
influence of sunlight.

Properties. The dilute aqueous solution of hydrogen peroxide,
obtained by the foregoing methods, is concentrated by evaporation
over sulphuric acid in vacuo. In the pure condition it is a colour-
less and odourless, syrupy liquid, having an extremely bitter and

226 Inorganic Chemistry

metallic taste. The specific gravity of the liquid is 1.4532. The
substance is extremely unstable, giving up some of its oxygen even
at temperatures as low as - 20, and decomposing with explosive
violence when heated to 100. Hydrogen peroxide bleaches
organic colours, but less rapidly than chlorine. When placed
upon the skin it destroys the colour, and gives rise to an irritating
blister. When diluted with water, and especially if rendered acid,
the compound is far more stable, and in this condition may be
preserved at the ordinary temperature for a considerable length of
time. When such an aqueous solution is strongly cooled, it deposits
ice, and in this way, by the removal of the frozen water, the solu-
tion may be concentrated. Hydrogen peroxide itself solidifies
between 20 and 23. When heated the solution is decom-
posed into water and oxygen

H 2 O 2 =H 2 O + O.

Owing to the readiness with which hydrogen peroxide gives up
the half of its oxygen and is converted into water, its properties
are generally those of a powerful oxidising agent. It liberates
iodine from potassium iodide ; it converts sulphurous acid into
sulphuric acid, and oxidises lead sulphide into lead sulphate. Its
action upon lead sulphide is made use of in restoring something
of the original brilliancy to oil paintings that have become dis-
coloured. The " white-lead " used in oil paints is gradually con-
verted into lead sulphide when such paintings are exposed to air,
especially the air of towns, which is liable to contain small
quantities of sulphuretted hydrogen. Lead sulphide being black,
the picture slowly assumes a uniformly dark colour. When such
a discoloured picture is washed over with dilute hydrogen peroxide,
the black sulphide is oxidised into the white lead sulphate

PbS + 4H 2 O 2 =4H 2 O + PbSO 4 .

This compound is employed for bleaching articles that would
suffer injury by the use of other bleaching agents, such as ivory,
feathers, and even the teeth.

Hydrogen paroxide is also capable of oxidising hydrogen.
Thus when a dilute acidulated solution of the peroxide is electro-
lysed, oxygen is evolved at the anode, but no gas escapes from
the cathode ; the nascent hydrogen being oxidised to water

Hydrogen Peroxide 227

Hydrogen peroxide, in many of its reactions, appears to act as a
deoxidising agent ; thus, manganese dioxide in contact with this
substance is reduced to manganous oxide

Mn0 2 + H 2 2 = MnO + O 2 + H 2 O.

Similarly, silver oxide is reduced to metallic silver with the
evolution of oxygen

In like manner, when ozone is acted upon by hydrogen per-
oxide, a reaction takes place exactly analogous to that with silver
oxide, which will be the more obvious if the formula for ozone be
written O 2 O instead of O 3 , thus

Although, in a sense, these reactions may be regarded as reduc-
ing^ or deoxidising^ actions, in essence they are not different from
those which have been given as illustrative of the oxidising power
of hydrogen peroxide. It will be seen that they all depend upon
the readiness with which the compound parts with an atom of
oxygen, but that in these latter cases the oxygen that is so given
up is engaged in oxidising another atom of oxygen, contained in the
other compound. Thus, in the case of silver oxide, its atom of
oxygen is oxidised by the liberated oxygen from the hydrogen
peroxide, and converted into the complete molecule of oxygen.
By these reactions Brodie first demonstrated the dual, or di-
atomic, character of the molecule of oxygen.

When hydrogen peroxide is added to a dilute acidulated solution
of potassium dichromate, a deep azure-blue solution is obtained
(see Chromium), which affords a delicate test for this com-
pound. To apply the test, the dilute hydrogen peroxide is shaken
up with ether, a few drops of acidulated potassium dichromate
are then added, and the mixture again shaken. The blue com-
pound being more soluble in ether than in water, the ethereal
liquid will separate as a blue layer. In this way, the presence of
0.00025 grammes of hydrogen peroxide in 20 c.c. of water can be

Hydrogen peroxide is decomposed by contact with many sub-
stances which themselves do not combine with the oxygen ; thus
charcoal, finely divided palladium, platinum, mercury, and notably
silver, when brought into hydrogen peroxide, determine its decom-

228 Inorganic Chemistry

position into water and oxygen, the rapidity of the action being
increased if the liquid be made alkaline. The action is doubtless
catalytic, although in all cases the exact modus operandi is not
clearly understood. In the case of silver it is believed that silver
oxide (perhaps peroxide) is first formed, and then decomposed,

Ag 2 + H 2 O 2 = H 2 O + Ag 2 O
Ag 2 O + H 2 O 2 = H 2 O + O 2 + Ag 2 .

When hydrogen peroxide is added to solutions of the hydroxides
of barium, strontium, or calcium, the peroxide of the metal is

Ba(HO) 2 + H 2 O,=2H 2 O + BaO 2 .

The compound is deposited in crystals having the composition
BaO 2 ,8H 2 O.

With the hydroxides of the alkali metals, the peroxide (which is
soluble in water) may be precipitated by the addition of alcohol ;
when in the case of sodium peroxide, crystals are obtained of
Na 2 O 2 ,8H 2 O.

Hydrogen peroxide is a useful antiseptic ; it possesses the ad-
vantages of being free from smell, without poisonous or injurious
action upon the system, and of leaving as a residue, after having
furnished its available oxygen, only water.

The constitution of hydrogen peroxide is usually expressed by the formula
H-OOH, but the accumulating evidence that oxygen is capable of functioning
as a quadrivalent element has led to the view that its constitution is better

H \
represented by the formula yO : O.

Online LibraryGeorge S NewthA text-book of inorganic chemistry → online text (page 20 of 67)