gravity attains a maximum value at about 98 p.c. This was
discovered by Kohlrausch, and confirmed by Chertel, Pickering, and
others.
[55] Naturally under the condition that there is no other ingredient
besides water, which is sufficiently true. For commercial acid,
whose specific gravity is usually expressed in degrees of Baumé's
hydrometer, we may add that at 15°
Specific gravity 1 1·1 1·2 1·3 1·4 1·5 1·6 1·7 1·8
Degree Baumé 0 13 24 33·3 41·2 48·1 54·1 59·5 64·2
66° Baumé (the strongest commercial acid or oil of vitriol)
corresponds to a sp. gr. 1·84.
By employing the second table (by the method of interpolation) the
specific gravity, at a given temperature (from 0° to 30°) can be
found for any percentage amount of H_{2}SO_{4}, and therefore
conversely the percentage of H_{2}SO_{4} can be found from the
specific gravity.
[55 bis] Whether similar (even small) breaks in the continuity of the
factor _dS_/_dp_ exist or not, for other hydrates (for instance,
for H_{2}SO_{4}H_{2}O and H_{2}SO_{4}4H_{2}O) cannot as yet be
affirmed owing to the want of accurate data (Note 53). In my
investigation of this subject (1887) I admit their possibility,
but only conditionally; and now, without insisting upon a similar
opinion, I only hold to the existence of a distinct break in the
factor at H_{2}SO_{4}, being guided by C. Winkler's observations
ond the specific gravities of fuming sulphuric acid.
[56] In 1887, on considering all the existent observations for a
temperature 0°, I gave the accompanying scheme (p. 243) of the
variation of the factor _ds_/_dp_ at 0°.
I did not then (1887) give this scheme an absolute value, and now
after the appearance of two series of new determinations (Lunge
and Pickering in 1890), which disagree in many points, I think it
well to state quite clearly: (1) that Lunge's and Pickering's new
determinations have not added to the accuracy of our data
respecting the variation of the specific gravity of solutions of
sulphuric acid; (2) that the sum total of existing data does not
negative (within the limit of experimental accuracy) the
possibility of a rectilinear and broken form for the factors
_ds_/_dp_; (3) that the supposition of 'special points' in
_ds_/_dp_, indicating definite hydrates, finds confirmation in all
the latest determinations; (4) that the supposition respecting the
existence of hydrates determining a break of the factor _ds_/_dp_
is in in way altered if, instead of a series of broken straight
lines, there be a continuous series of curves, nearly approaching
straight lines; and (5) that this subject deserves (as I mentioned
in 1887) new and careful elaboration, because it concerns that
foremost problem in our science--solutions--and introduces a
special method into it--that is, the study of differential
variations in a property which is so easily observed as the
specific gravity of a liquid.
[56 bis] These hydrates are: (_a_) H_{2}SO_{4} = SO_{3}H_{2}O (melts
at + 10°·4); (_b_) H_{2}SO_{4}H_{2}O = SO_{3}2H_{2}O
(crystallo-hydrate, melts at +8°·5); (_c_) H_{2}SO_{4}2H_{2}O (is
apparently not crystallisable); (_d_) one of the hydrates between
H_{2}SO_{4}6H_{2}O and H_{2}SO_{4}3H_{2}O, most probably
H_{2}SO_{4}4H_{2}O = SO_{3}5H_{2}O, for it crystallises at -24°·5
(Note 50 bis); and (_e_) a certain hydrate with a large proportion
of water, about H_{2}SO_{4}150H_{2}O. The existence of the last is
inferred from the fact that the factor _ds_/_dp_ first falls,
starting from water, and then rises, and this change takes place
when _p_ is less than 5 p.c. Certainly a change in the variation
of _ds_/_dp_ or _ds_/_dt_ does take place in the neighbourhood of
these five hydrates (Pickering, 1890, recognised a far greater
number of hydrates). I think it well to add that if the
composition of the solutions be expressed by the percentage amount
of molecules--_r__{1}SO_{3} + (100 - _r__{1})H_{2}O we find that
for H_{2}SO_{4}, _r__{1} = 50, for H_{2}SO_{4}2H_{2}O _r__{1} = 25
= 50/2, for H_{2}SO_{4}H_{2}O, _r__{1} = 33·333 = 50 · 2/3, while
for H_{2}SO_{4}4H_{2}O, _r__{1} = 16·666 = 50 · 1/3--_i.e._ that
the chief hydrates are distributed symmetrically between H_{2}O
and H_{2}SO_{4}. Besides which I may mention that my researches
(1887) upon the abrupt changes in the factor for solutions of
sulphuric acid, and upon the correspondence of the breaks of
_ds_/_dp_ with definite hydrates, received an indirect
confirmation not only in the solutions of HNO_{3}, HCl,
C_{2}H_{6}O, C_{3}H_{8}O, &c., which I investigated (in my work
cited in Chapter I., Note 19), but also in the careful
observations made by Professor Cheltzoff on the solutions of
FeCl_{3} and ZnCl_{2} (Chapter XVI., Note 4) which showed the
existence in these solutions of an almost similar change in
_ds_/_dp_ as is found in sulphuric acid. The detailed researches
(1893) made by Tourbaba on the solutions of many organic
substances are of a similar nature. Besides which, H. Crompton
(1888), in his researches on the electrical conductivity of
solutions of sulphuric acid, and Tammann, in his observations on
their vapour tension, found a correlation with the hydrates
indicated as above by the investigation of their specific
gravities. The influence of mixtures of a definite composition
upon the chemical relations of solutions is even exhibited in such
a complex process as electrolysis. V. Kouriloff (1891) showed that
mixtures containing about 3 p.c., 47 p.c. and 73 p.c. of sulphuric
acid--_i.e._ whose composition approaches that of the hydrates
H_{2}SO_{4}150H_{2}O, H_{2}SO_{4}6H_{2}O and
H_{2}SO_{4}2H_{2}O--exhibit certain peculiarities in respect to
the amount of peroxide of hydrogen formed during electrolysis.
Thus a 3 p.c. solution gives a maximum amount of peroxide of
hydrogen at the negative pole, as compared with that given by
other neighbouring concentrations. Starting from 3 p.c., the
formation of peroxide of hydrogen ceases until a concentration of
47 p.c. is reached.
The great affinity of sulphuric acid for water is also seen from the
fact that when the strong acid acts on the majority of _organic
substances_ containing hydrogen and oxygen (especially on heating) it
very frequently _takes up these elements in the form of water_. Thus
strong sulphuric acid acting on alcohol, C_{2}H_{6}O, removes the
elements of water from it, and converts it into olefiant gas, C_{2}H_{4}.
It acts in a similar manner on wood and other vegetable tissues, which it
chars. If a piece of wood be immersed in strong sulphuric acid it turns
black. This is owing to the fact that the wood contains carbohydrates
which give up hydrogen and oxygen as water to the sulphuric acid, leaving
charcoal, or a black mass very rich in it. For example, cellulose,
C_{6}H_{10}O_{5}, acts in this manner.[57]
[57] Cellulose, for instance unsized paper or calico, is dissolved by
strong sulphuric acid. Acid diluted with about half its volume of
water converts it (if the action be of short duration) into
vegetable parchment (Chapter I., Note 18). The action of dilute
solutions of sulphuric acid converts it into hydro-cellulose, and
the fibre loses its coherent quality and becomes brittle. The
prolonged action of strong sulphuric acid chars the cellulose
while dilute acid converts it into glucose. If sulphuric acid be
kept in an open vessel, the organic matter of the dust held in the
atmosphere falls into it and blackens the acid. The same thing
happens if sulphuric acid be kept in a bottle closed by a cork;
the cork becomes charred, and the acid turns black. However, the
chemical properties of the acid undergo only a very slight change
when it turns black. Sulphuric acid which is considerably diluted
with water does not produce the above effects, which clearly shows
their dependence on the affinity of the sulphuric acid for water.
It is evident from the preceding that strong sulphuric acid will
act as a powerful poison; whilst, on the other hand, when very
dilute it is employed in certain medicines and as a fertiliser for
plants.
We have already had frequent occasion to notice the very _energetic acid
properties_ of sulphuric acid, and therefore we will now only consider a
few of their aspects. First of all we must remember that, with calcium,
strontium, and especially with barium and lead, sulphuric acid forms very
slightly soluble salts, whilst with the majority of other metals it gives
more easily soluble salts, which in the majority of cases are able, like
sulphuric acid itself, to combine with water to form crystallo-hydrates.
Normal sulphuric acid, containing two atoms of hydrogen in its molecule,
is able for this reason alone to form two classes of salts, _normal_ and
_acid_, which it does with great facility _with the alkali metals_. The
metals of the alkaline earths and the majority of other metals, if they
do form acid sulphates, do so under exceptional conditions (with an
excess of strong sulphuric acid), and these salts when formed are
decomposable by water--that is, although having a certain degree of
physical stability they have no chemical stability. Besides the acid
salts RHSO_{4}, sulphuric acid also gives other forms of acid salts. An
entire series of salts having the composition RHSO_{4},H_{2}SO_{4}, or
for bivalent metals RSO_{4},3H_{2}SO_{4},[58] has been prepared. Such
salts have been obtained for potassium, sodium, nickel, calcium, silver,
magnesium, manganese. They are prepared by dissolving the sulphates in an
excess of sulphuric acid and heating the solution until the excess of
sulphuric acid is driven off; on cooling, the mass solidifies to a
crystalline salt. Besides which, Rose obtained a salt having the
composition Na_{2}SO_{4},NaHSO_{4}, and if HNaSO_{4} be heated it easily
forms a salt Na_{2}S_{2}O_{7} = Na_{2}SO_{4},SO_{3}; hence it is clear
that sulphuric anhydride combines with various proportions of bases, just
as it combines with various proportions of water.
[58] Weber (1884) obtained a series of salts R_{2}O,8SO_{3}_n_H_{2}O
for K, Rb, Cs, and Tl.
We have already learned that sulphuric acid displaces the acid from the
salts of nitric, carbonic, and many other volatile acids. Berthollet's
laws (Chapter X.) explain this by the small volatility of sulphuric acid;
and, indeed, in an aqueous solution sulphuric acid displaces the much
less soluble boric acid from its compounds--for instance, from borax, and
it also displaces silica from its compounds with bases; but both boric
anhydride and silica, when fused with sulphates, decompose them,
displacing sulphuric anhydride, SO_{3}, because they are less volatile
than sulphuric anhydride. It is also well known that with metals,
sulphuric acid forms salts giving off hydrogen (Fe, Zn, &c.), or sulphur
dioxide (Cu, Hg, &c.).[58 bis]
[58 bis] Ditte (1890) divides all the metals into two groups with
respect to sulphuric acid; the first group includes silver,
mercury, copper, lead, and bismuth, which are only acted upon by
hot concentrated acid. In this case sulphurous anhydride is
evolved without any by-reactions. The second group contains
manganese, nickel, cobalt, iron, zinc, cadmium, aluminium, tin,
thallium, and the alkali metals. They react with sulphuric acid of
any concentration at any temperature. At a low temperature
hydrogen is disengaged, and at higher temperatures (and with very
concentrated acid) hydrogen and sulphurous anhydride are
simultaneously evolved.
The reactions of sulphuric acid _with respect to organic substances_ are
generally determined by its acid character, when the direct extraction of
water, or oxidation at the expense of the oxygen of the sulphuric
acid,[59] or disintegration does not take place. Thus the majority of the
saturated hydrocarbons, C_{_n_}H_{2_m_}, form with sulphuric acid a
special class of _sulphonic acids_, C_{_n_}H_{2_m_-1}(HSO_{3}); for
example, benzene, C_{6}H_{6}, forms benzenesulphonic acid,
C_{6}H_{5}.SO_{3}H, water being separated, for the formation of which
oxygen is taken up from the sulphuric acid, for the product contains less
oxygen than the sulphuric acid. It is evident from the existence of these
acids that the hydrogen in organic compounds is replaceable by the group
SO_{3}H, just as it may be replaced by the radicles Cl, NO_{2}, CO_{2}H
and others. As the radicle of sulphuric acid or _sulphoxyl_, SO_{2}OH or
SHO_{3}, contains, like carboxyl (Vol. I., p. 395), one hydrogen
(hydroxyl) of sulphuric acid, the resultant substances are acids whose
basicity is equal to the number of hydrogens replaced by sulphoxyl. Since
also sulphoxyl takes the place of hydrogen, and itself contains hydrogen,
the sulpho-acids are equal to a hydrocarbon + SO_{3}, just as every
organic (carboxylic) acid is equal to a hydrocarbon + CO_{2}. Moreover,
here this relation corresponds with actual fact, because many sulphonic
acids are obtained by the direct combination of sulphuric anhydride:
C_{6}H_{5},(SO_{3}H) = C_{6}H_{6} + SO_{3}. The sulphonic acids give
soluble barium salts, and are therefore easily distinguished from
sulphuric acid. They are soluble in water, are not volatile, and when
distilled give sulphurous anhydride (whilst the hydroxyl previously in
combination with the sulphurous anhydride remains in the hydrocarbon
group; thus phenol, C_{6}H_{5}.OH, is obtained from benzenesulphonic
acid), and they are very energetic, because the hydrogen acting in them
is of the same nature as in sulphuric acid itself.[60]
[59] For example, the action of hot sulphuric acid on nitrogenous
compounds, as applied in Kjeldahl's method for the estimation of
nitrogen (Volume I. p. 249). It is obvious that when sulphuric
acid acts as an oxidising agent it forms sulphurous anhydride.
The action of sulphuric acid on the alcohols is exactly similar to
its action on alkalis, because the alcohols, like alkalis, react
on acids; a molecule of alcohol with a molecule of sulphuric acid
separates water and forms an _acid_ ethereal salt--that is there
is produced an ethereal compound corresponding with acid salts.
Thus, for example, the action of sulphuric acid, H_{2}SO_{4}, on
ordinary alcohol, C_{2}H_{5}OH, gives water and sulphovinic acid,
C_{2}H_{5}HSO_{4}--that is, sulphuric acid in which one atom of
hydrogen is replaced by the radicle C_{2}H_{5} of ethyl alcohol,
SO_{2}(OH)(OC_{2}H_{5}), or, what is the same thing, the hydrogen
in alcohol is replaced by the radicle (sulphoxyl) of sulphuric
acid, C_{2}H_{5}O.SO_{2}(OH).
[60] We will mention the following difference between the sulphonic
acids and the ethereal acid sulphates (Note 59): the former
re-form sulphuric acid with difficulty and the latter easily. Thus
sulphovinic acid when heated with an excess of water is
reconverted into alcohol and sulphuric acid. This is explained in
the following manner. Both these classes of acids are produced by
the substitution of hydrogen by SO_{3}H, or the univalent radicle
of sulphuric acid, but in the formation of ethereal acid sulphates
the SO_{3}H replaces the hydrogen of the hydroxyl in the alcohol,
whilst in the formation of the sulphonic acids the SO_{3}H
replaces the hydrogen of a hydrocarbon. This difference is clearly
evidenced in the existence of two acids of the composition
SO_{4}C_{2}H_{6}. The one, mentioned above, is sulphovinic acid or
alcohol, C_{2}H_{5}.OH, in which the hydrogen of the hydroxyl is
replaced by sulphoxyl = C_{2}H_{5}.OSO_{3}H, whilst the other is
alcohol, in which one atom of the hydrogen in ethyl, C_{2}H_{5},
is replaced by the sulphonic group--that is =
(C_{2}H_{4})SO_{3}H·OH. The latter is called isethionic acid. It
is more stable than sulphovinic acid. The details as to these
interesting compounds must be looked for in works on organic
chemistry, but I think it necessary to note one of the general
methods of formation of these acids. The sulphites of the
alkalis--for example, K_{2}SO_{3}--when heated with the halogen
products of metalepsis, give a halogen salt and a salt of a
sulphonic acid. Thus methyl iodide, CH_{3}I, derived from marsh
gas, CH_{4}, when heated to 100° with a solution of potassium
sulphite, K_{2}SO_{3}, gives potassium iodide, KI, and potassium
methylsulphonate, CH_{3}SO_{3}K--that is a salt of the sulphonic
acid. This shows that the sulphonic acid may be referred to
sulphurous acid, and that there is a resemblance between sulphuric
and sulphurous acid, which clearly reveals itself here in the
formation of one product from them both.
Sulphuric acid, as containing a large proportion of oxygen, is a
substance which frequently acts as an oxidising agent: in which case it
is _deoxidised, forming sulphurous anhydride_ and water (or even,
although more rarely, sulphuretted hydrogen and sulphur). Sulphuric acid
acts in this manner on charcoal, copper, mercury, silver, organic and
other substances, which are unable to evolve hydrogen from it directly,
as we saw in describing sulphurous anhydride.
Although the hydrate of a higher saline form of oxidation (Chapter XV.),
sulphuric anhydride is capable of further oxidation, and forms a kind of
peroxide, just as hydrogen gives hydrogen peroxide in addition to water,
or as sodium and potassium, besides the oxides Na_{2}O and K_{2}O, give
their peroxides, compounds which are in a chemical sense unstable,
powerfully oxidising, and not directly able to enter into saline
combinations. If the oxides of potassium, barium, &c., be compared to
water, then their peroxides must in like manner correspond to hydrogen
peroxide,[61] not only because the oxygen contained in them is very
mobile and easily liberated, and because their reactions are similar, but
also because they can be mutually transformed into each other, and are
able to form compounds with each other, with bases and with water, and
indeed form a kind of peroxide salts.[62] This is also the character of
_persulphuric acid_, discovered in 1878 by Berthelot, and its
corresponding anhydride or peroxide of sulphur S_{2}O_{7}. It is formed
from 2SO_{3} + O with the absorption of heat (-27 thousand heat units),
like ozone from O_{2} + O (-29 thousand units of heat), or hydrogen
peroxide from H_{2}O + O (-21 thousand heat units).
[61] The reaction BaO + O develops 12,000 heat units, whilst the
reaction H_{2}O + O absorbs 21,000 heat units.
[62] Schöne obtained a compound of peroxide of barium with peroxide of
hydrogen. If barium peroxide be dissolved in hydrochloric (or
acetic) acid, or if a solution of hydrogen peroxide be diluted
with a solution of barium hydroxide, a pure hydrate is
precipitated having the composition BaO_{2},8H_{2}O (sometimes the
composition is taken as BaO_{2},6H_{2}O). This fact was already
known to Thénard. Schöne showed that if hydrogen peroxide be in
excess, a crystalline compound of the two peroxides,
BaO_{2}H_{2}O_{2}, is precipitated. Schöne also obtained small
well-formed crystals of the same composition by adding a solution
of ammonia to an acid solution of barium peroxide (containing a
barium salt and hydrogen peroxide or a compound of BaO_{2} with
the acid). Thus barium peroxide combines with both water and
hydrogen peroxide. This is a very important fact for the
comprehension of the composition of other peroxides. Moreover, if
the peroxides are able to give hydrates they can also form
corresponding salts, _i.e._ they can combine with bases and acids,
as was afterwards found to be the case on further research into
this subject.
Peroxide of sulphur is produced by the action of a silent discharge
upon a mixture of oxygen and sulphurous anhydride.[63] With water
S_{2}O_{7} gives persulphuric acid, H_{2}S_{2}O_{8}. The latter is
obtained more simply by mixing strong sulphuric acid (not weaker than
H_{2}SO_{4},2H_{2}O) directly with hydrogen peroxide, or by the action of
a galvanic current on sulphuric acid mixed with a certain amount of
water, and cooled, the electrodes being platinum wires, when persulphuric
acid naturally appears at the positive pole.[64] When an acid of the
strength H_{2}SO_{4},6H_{2}O is taken, at first the hydrate of the
sulphuric peroxide, S_{2}O_{7},H_{2}O only is formed; but when the
concentration about the positive pole reaches H_{2}SO_{4},3H_{2}O, a
mixture of hydrogen peroxide and the hydrate of sulphuric peroxide begins
to be formed. Dilute solutions of sulphuric peroxide can be kept better
than more concentrated solutions, but the latter may be obtained
containing as much as 123 grams of the peroxide to a litre. It is a very
instructive fact that hydrogen peroxide is always formed when strong
solutions of persulphuric acid break up on keeping. So that the bond
between the two peroxides is established both by analysis and synthesis:
hydrogen peroxide is able to produce S_{2}H_{2}O_{8}, and the latter to
produce hydrogen peroxide. A mixture of sulphuric peroxide with sulphuric
acid or water is immediately decomposed, with the evolution of oxygen,
either when heated or under the action of spongy platinum. The same thing
takes place with a solution of baryta, although at first no precipitate
is formed and the decomposition of the barium salt, BaS_{2}O_{8}, with
the formation of BaSO_{4}, only proceeds slowly, so that the solution may
be filtered (the barium salt of persulphuric acid is soluble in water).
Mercury, ferrous oxide, and the stannous salts, are oxidised by
S_{2}H_{2}O_{8}. These are all distinct signs of true peroxides. The same
common properties (capacity for oxidising, property of forming peroxide
of hydrogen, &c.) are possessed by the alkali salts of persulphuric acid,
which are obtained by the action of an electric current upon certain
sulphates, for instance ammonium or potassium sulphate. The ammonium salt
of persulphuric acid, (NH_{4})_{2}S_{2}O_{8}, is especially easily formed
by this means, and is now prepared on a large scale and used (like
Na_{2}O_{2} and H_{2}O_{2}) for bleaching tissues and fibres.[65]
[63] Anhydrous _sulphuric peroxide_, S_{2}O_{7}, is obtained by the
prolonged (8 to 10 hours) action of a silent discharge of
considerable intensity on a mixture of oxygen and sulphurous
anhydride; the vapour of sulphuric peroxide, S_{2}O_{7}, condenses
as liquid drops, or after being cooled to 0° in the form of long
prismatic crystals, resembling those of sulphuric anhydride. The
anhydrous compound S_{2}O_{7} (and also the hydrated compound)
cannot be preserved long, as it splits up into oxygen and
sulphuric anhydride. Direct experiment shows that a mixture of
equal volumes of sulphurous anhydride and oxygen leaves a residue
of a quarter of the oxygen taken, or half of the whole volume,
which indicates the formula S_{2}O_{7}. This substance is soluble
in water, and it then gives a hydrate, probably having the
composition S_{2}O_{7},H_{2}O = 2SHO_{4}. This solution oxidises
the salts SnX_{2}, potassium iodide, and others, which renders it
possible to prove that the solution actually contains one atom of
oxygen capable of effecting oxidation to two molecules of
sulphuric anhydride.
In order to fully demonstrate the reality of a peroxide form for
acids, it should be mentioned that some years ago Brodie obtained
the so-called _acetic peroxide_, (C_{2}H_{2}O)_{2}O_{2}, by the
action of barium peroxide on acetic anhydride, (C_{2}H_{3}O)_{2}O.
Its corresponding hydrate is also known. This shows that true
peroxides and their hydrates, with reactions similar to those of
hydrogen peroxide, are possible for acids. A similar higher oxide
has long been known for chromium, and Berthelot obtained a like
compound for nitric acid (Chapter VI., Note 26).
[64] When an acid of the strength H_{2}SO_{4}6H_{2}O is taken, at first
only the hydrate of the sulphuric peroxide, S_{2}O_{7}H_{2}O, is
formed, but when the concentration at the positive pole reaches
H_{2}SO_{4}3H_{2}O, a mixture of hydrogen peroxide and the hydrate
of sulphuric peroxide begins to be formed. A state of equilibrium
is ultimately arrived at when the amounts of these substances
correspond to the proportion S_{2}O_{7} : 2H_{2}O_{2}, which, as
it were, answers to a new hydrate, S_{2}O_{9}2H_{2}O. But its
existence cannot be admitted because the sulphuric peroxide can be
easily distinguished from the hydrogen peroxide in the solution
owing to the fact that the former does not act on an acid solution
of potassium permanganate, whilst the hydrogen peroxide disengages
both its own oxygen and that of the permanganic acid, converting
it into manganous oxide. Their common property of liberating
iodine from an acid solution of the potassium iodide enables the
sum of the active oxygen in them both to be determined.
[65] If a solution of sulphuric acid which has been first subjected to
electrolysis be neutralised with potash or baryta, the salt which
is formed begins to decompose rapidly with the evolution of oxygen
(Berthelot, 1890). On saturating with caustic baryta, the solution
of the salt formed may be separated from the sulphate of barium,
and then the composition of the resultant compound, BaS_{2}O_{8},
may be determined from the amount of oxygen disengaged. Marshall
(1891) studied the formation of this class of compounds more
fully; he subjected a saturated solution of bisulphate of
potassium to electrolysis with a current of 3-3-1/2 ampères;
before electrolysis dilute sulphuric acid is added to the liquid
surrounding the negative pole, and during electrolysis the
solution at the anode is cooled. The electrolysis is continued
without interruption for two days, and a white crystalline deposit
separates at the anode. To avoid decomposition, the latter is not
filtered through paper, but through a perforated platinum plate,
and dried on a porous tile. The mother liquor, with the addition
of a fresh solution of bisulphate of potassium, is again subjected
to electrolysis and the crystals formed at the anode are again
collected, &c. The salt so obtained may be recrystallised by
dissolving it in hot water and rapidly cooling the solution after
filtration; a small proportion of the salt is decomposed by this
treatment. Rapid cooling is followed by the formation of small
columnar crystals; slow cooling gives large prismatic crystals.
The composition of the salt is determined either by igniting it,
when it forms sulphate of potassium, or else by titrating the
active oxygen with permanganate: its composition was found to
correspond to the salt of persulphuric acid, K_{2}S_{2}O_{8}. The
solution of the salt has a neutral reaction, and does not give a
precipitate with salts of other metals. K_{2}S_{2}O_{8} is the
most insoluble of the salts of persulphuric acid. With nitrate of
silver it forms persulphate of silver, which gives peroxide of
silver under the action of water according to the equation
Ag_{2}S_{2}O_{8} + 2H_{2}O = Ag_{2}O_{2} + 2H_{2}SO_{4}. With an
alkaline solution of a cupric salt (Fehling's solution) it forms a
red precipitate of peroxide of copper. Manganese and cobalt salts
give precipitates of MnO_{2} and Co_{2}O_{3}. Ferrous salts are
rapidly oxidised, potassium iodide slowly disengages iodine at the
ordinary temperature. All these reactions indicate the powerful
oxidising properties of K_{2}S_{2}O_{8}. In oxidising in the
presence of water it gives a residue of KHSO_{4}. The
decomposition of the dry salt begins at 100° but is not complete
even at 250°. The freshly prepared salt is inodorous, but after
being kept in a closed vessel it evolves a peculiar smell
different from that of ozone. The ammonium salt of persulphuric
acid, (NH_{4})_{2}S_{2}O_{8}, is obtained in a similar manner. It
is soluble to the extent of 58 parts per 100 parts by weight of
water. The decomposition of the ammonium salt by the hydrated
oxide of barium gives the barium salt, BaS_{2}O_{8}4H_{2}O, which
is soluble to the extent of 52·2 parts in 100 parts of water at
0°. The crystals do not deliquesce in the air and decompose in the
course of several days; they decompose most rapidly in perfectly
dry air. Solutions of the pure salt decompose slowly at the
ordinary temperature; on boiling barium sulphate is gradually
precipitated, oxygen being liberated simultaneously. To completely
decompose this salt it is necessary to boil its solution for a
long time. Alcohol dissolves the solid salt; the anhydrous salt
does not separate from the alcoholic solution, but a hydrate
containing one molecule of water, BaS_{2}O_{8}H_{2}O, which is
soluble in water but insoluble in absolute alcohol. Solid barium
persulphate decomposes even when slightly heated. The free acid,
which may serve for the preparation of other salts, is obtained by
treating the barium salt with sulphuric acid. The lead salt,
PbS_{2}O_{8}, has been obtained from the free acid; it
crystallises with two or three molecules of water. It is soluble
in water, deliquesces in the air, and with alkalis gives a
precipitate of the hydrated oxide which rapidly oxidises into the
binoxide.
Traube, before Marshall's researches, thought that the
electrolysis of solutions of sulphuric acid did not give
persulphuric acid but a persulphuric oxide having the composition
SO_{4}. On repeating his former researches (1892) Traube obtained
a persulphuric oxide by the electrolysis of a 70 per cent.
solution of sulphuric acid, and he separated it from the solution
by means of barium phosphate. Analysis showed that this substance
corresponded to the above composition SO_{4}, and therefore Traube
considers it very likely that the salts obtained by Marshall
corresponded to an acid H_{2}SO_{4} + SO_{4}, _i.e._ that the
indifferent oxide, SO_{4}, can combine with sulphuric acid and
form peculiar saline compounds.
In order to understand the relation of sulphuric peroxide to sulphuric
acid we must first remark that hydrogen peroxide is to be considered, in
accordance with the law of substitution, as water, H(OH), in which H is
replaced by (OH). Now the relation of H_{2}S_{2}O_{8} to H_{2}SO_{4} is
exactly similar. The radicle of sulphuric acid, equivalent to hydrogen,
is HSO_{4};[65 bis] it corresponds with the (OH) of water, and therefore
sulphuric acid, H(SHO_{4}), gives (SHO_{4})_{2} or S_{2}H_{2}O_{8}, in
exactly the same manner as water gives (HO)_{2}--_i.e._ H_{2}O_{2}.[66]
[65 bis] Or one of those supposed ions which appear at the positive
pole in the decomposition of sulphuric acid by the action of a
galvanic current.
[66] If this be true one would expect the following peroxide hydrates:
for phosphoric acid, (H_{2}PO_{4})_{2} = H_{4}P_{2}O_{8} = 2H_{2}O
+ 2PO_{3}; for carbonic acid, (HCO_{3})_{2} = H_{2}C_{2}O_{6} =
H_{2}O + C_{2}O_{5}; and for lead the true peroxide will be also
Pb_{2}O_{5}, &c. Judging from the example of barium peroxide (Note
62), these peroxide forms will probably combine together. It seems
to me that the compounds obtained by Fairley for uranium are very
instructive as elucidating the peroxides. In the action of
hydrogen peroxide in an acid solution on uranium oxide, UO_{3},
there is formed a uranium peroxide, UO_{4},4H_{2}O (U = 240), but
hydrogen peroxide acts on uranium oxide in the presence of caustic
soda; on the addition of alcohol a crystalline compound containing
Na_{4}UO_{8},4H_{2}O is precipitated, which is doubtless a
compound of the peroxides of sodium, Na_{2}O_{2}, and uranium,
UO_{4}. It is very possible that the first peroxide,
UO_{4},4H_{2}O, contains the elements of hydrogen peroxide and
uranium peroxide, U_{2}O_{7}, or even U(OH)_{6},H_{2}O_{2}, just
as the peroxide form lately discovered by Spring for tin perhaps
contains Sn_{2}O_{3},H_{2}O_{2}.
The largest part _of the sulphuric acid made_ is used for reacting on
sodium chloride in the manufacture of sodium carbonate; for the
manufacture of the volatile acids, like nitric, hydrochloric, &c., from
their corresponding salts; for the preparation of ammonium sulphate,
alums, vitriols (copper and iron), artificial manures, superphosphate
(Chapter XIX., Note 18) and other salts of sulphuric acid; in the
treatment of bone ash for the preparation of phosphorus, and for the
solution of metals--for example, of silver in its separation from
gold--for cleaning metals from rust, &c. A large amount of oil of vitriol
is also used in treatment of organic substances; it is used for the
extraction of stearin, or stearic acid, from tallow, for refining
petroleum and various vegetable oils, in the preparation of
nitro-glycerine (Chapter VI., Notes 37 and 37 bis), for dissolving indigo
and other colouring matters, for the conversion of paper into vegetable
parchment, for the preparation of ether from alcohol, for the preparation
of various artificial scents from fusel oil, for the preparation of
vegetable acids, such as oxalic, tartaric, citric, for the conversion of
non-fermentable starchy substances into fermentable glucose, and in a
number of other processes. It would be difficult to find another
artificially-prepared substance which is so frequently applied in the
arts as sulphuric acid. Where there are not works for its manufacture,
the economical production of many other substances of great technical
importance is impossible. In those localities which have arrived at a
high technical activity the amount of sulphuric acid consumed is
proportionally large; sulphuric acid, sodium carbonate, and lime are the
most important of the artificially-prepared agents employed in factories.
Besides the normal acids of sulphur, H_{2}SO_{3}, H_{2}SO_{3}S, and
H_{2}SO_{4}, corresponding with sulphuretted hydrogen, H_{2}S, in the
same way that the oxy-acids of chlorine correspond with hydrochloric
acid, HCl, there exists a peculiar series of acids which are termed
_thionic acids_. Their general composition is S_{_n_}H_{2}O_{6}, where
_n_ varies from 2 to 5. If _n_ = 2, the acid is called dithionic acid.
The others are distinguished as trithionic, tetrathionic, and
pentathionic acids. Their composition, existence, and reactions are very
easily understood if they be referred to the class of the sulphonic
acids--that is, if their relation to sulphuric acid be expressed in just
the same manner as the relation of the organic acids to carbonic acid.
The organic acids, as we saw (Chapter IX.), proceed from the hydrocarbons
by the substitution of their hydrogen by carboxyl--that is, by the
radicle of carbonic acid, CH_{2}O_{3} - HO = CHO_{2}. The formation of
the acids of sulphur by means of sulphoxyl may be represented in the same
manner, HSO_{3} = H_{2}SO_{4} - HO. Therefore to hydrogen H_{2}, there
should correspond the acids H.SHO_{3}, sulphurous, and SHO_{3}.SHO_{3} =
S_{2}H_{2}O_{6}, or dithionic; to SH_{2} there should correspond the
acids SH(SHO_{3}) = H_{2}S_{2}O_{3} (thiosulphuric), and S(SHO_{3})_{2} =
H_{2}S_{3}O_{6} (trithionic); to S_{2}H_{2} the acids S_{2}H(SHO_{3}) =
H_{2}S_{3}O_{2} (unknown), and S_{2}(SHO_{3})_{2} = H_{2}S_{4}O_{6}
(tetrathionic); to S_{3}H_{2} the acids S_{3}H(SHO_{3}) and
S_{3}(SHO_{3})_{2} = H_{2}S_{5}O_{6} (pentathionic). We know that iodine
reacts directly with the hydrogen of sulphuretted hydrogen and combines
with it, and if thiosulphuric acid contains the radicle of sulphuretted
hydrogen (or hydrogen united with sulphur) of the same nature as in
sulphuretted hydrogen, it is not surprising that iodine reacts with
sodium thiosulphate and forms sodium tetrathionate. Thus, thiosulphuric
acid, HS(SHO_{3}), when deprived of H, gives a radicle which immediately
combines with another similar radicle, forming the tetrathionate
S_{2}(SO_{2}HO)_{2}. On this view[67] of the structure of the thionic
acids and salts, it is also clear how all the thionic acids, like
thiosulphuric acid, easily give sulphur and sulphides, with the exception
only of dithionic acid, H_{2}S_{2}O_{6}, which, judging from the above,
stands apart from the series of the other thionic acids. Dithionic acid
stands in the same relation to sulphuric acid as oxalic acid does to
carbonic acid. Oxalic acid is dicarboxyl, (CHO_{2})_{2} =
C_{2}H_{2}O_{4}, and so also dithionic acid is disulphoxyl, (SHO_{3})_{2}
= S_{2}H_{2}O_{6}. Oxalic acid when ignited decomposes into carbonic
anhydride and carbonic oxide, CO, and dithionic acid when heated
decomposes into sulphuric anhydride and sulphurous anhydride, SO_{2}, and
SO_{2} stands in the same relation to SO_{3} as CO to CO_{2}. This also
explains the peculiarity of the calcium, barium, and lead, &c. salts of
the thionic acids being easily soluble (although the corresponding salts
of H_{2}SO_{3}, H_{2}SO_{4}, and H_{2}S dissolve with difficulty),
because the former are similar to the salts of the sulphonic acids, which
are also soluble in water. Thus the thionic acids are _disulphonic
acids_, just as many dicarboxylic acids are known--for example,
CH_{2}(CO_{2}H)_{2}, C_{6}H_{4}(CO_{2}H)_{2}.[68]
[67] This view was communicated by me in 1870 to the Russian Chemical
Society.
[68] _Dithionic acid_, H_{2}S_{2}O_{6}, is distinguished among the
thionic acids as containing the least proportion of sulphur. It is
also called hyposulphuric acid, because its supposed anhydride,
S_{2}O_{5}, contains more O than sulphurous oxide, SO_{2} or
S_{2}O_{4}, and less than sulphuric anhydride, SO_{3} or
S_{2}O_{6}. Dithionic acid, discovered by Gay-Lussac and Welter,
is known as a hydrate and as salts, but not as anhydride. The
method for preparing dithionic acid usually employed is by the
action of finely-powdered manganese dioxide on a solution of
sulphurous anhydride. On shaking, the smell of the latter
disappears, and the manganese salt of the acid in question passes
into solution; MnO_{2} + 2SO_{2} = MnS_{2}O_{6}. If the
temperature be raised, the dithionate splits up into sulphurous
anhydride and manganese sulphate, MnSO_{4}. Generally owing to
this a mixture of manganese sulphate and dithionate is obtained in
the solution. They may be separated by mixing the solution of the
manganese salts with a solution of barium hydroxide, when a
precipitate of manganese hydroxide and barium sulphate is
obtained. In this manner barium dithionate only is obtained in
solution. It is purified by crystallisation, and separates as
BaS_{2}O_{6},2H_{2}O; this is then dissolved in water, and
decomposed with the requisite amount of sulphuric acid. Dithionic
acid, H_{2}S_{2}O_{6}, then remains in solution. By concentrating
the resultant solution under the receiver of an air-pump it is
possible to obtain a liquid of sp. gr. 1·347, but it still
contains water, and on further evaporation the acid decomposes
into sulphuric acid and sulphurous anhydride: H_{2}S_{2}O_{6} =
H_{2}SO_{4} + SO_{2}. The same decomposition takes place if the
solution be slightly heated. Like all the thionic acids, dithionic
acid is readily attacked by oxidising agents, and passes into
sulphurous acid. No dithionate is able to withstand the action of
heat, even when very slight, without giving off sulphurous
anhydride: K_{2}S_{2}O_{6} = K_{2}SO_{4} + SO_{2}. The alkali
dithionates have a neutral reaction (which indicates the energetic
nature of the acid) are soluble in water, and in this respect
present a certain resemblance to the salts of nitric acid (their
anhydrides are: N_{2}O_{5} and S_{2}O_{5}). Klüss (1888) described
many of the salts of dithionic acid.
Langlois, about 1840, obtained a peculiar thionic acid by heating
a strong solution of acid potassium sulphite with flowers of
sulphur to about 60°, until the disappearance of the yellow
coloration first produced by the solution of the sulphur. On
cooling, a portion of the sulphur was precipitated, and crystals
of a salt of _trithionic acid_, K_{2}S_{3}O_{6} (partly mixed with
potassium sulphate), separated out. Plessy afterwards showed that
the action of sulphurous acid on a thiosulphate also gives sulphur
and trithionic acid: 2K_{2}S_{2}O_{3} + 3SO_{2} = 2K_{2}S_{3}O_{6}
+ S. A mixture of potassium acid sulphite and thiosulphate also
gives a trithionate. It is very possible that a reaction of the
same kind occurs in the formation of trithionic acid by Langloid's
method, because potassium sulphite and sulphur yield potassium
thiosulphate. The potassium thiosulphate may also be replaced by
potassium sulphide, and on passing sulphurous anhydride through
the solution thiosulphate is first formed and then trithionate:
4KHSO_{3} + K_{2}S + 4SO_{2} = 3K_{2}S_{3}O_{6} + 2H_{2}O. The
sodium salt is not formed under the same circumstances as the
corresponding potassium salt. The sodium salt does not crystallise
and is very unstable: the barium salt is, however, more stable.
The barium and potassium salts are anhydrous, they give neutral
solutions and decompose when ignited, with the evolution of
sulphur and sulphurous anhydride, a sulphate being left behind,
K_{2}S_{3}O_{6} = K_{2}SO_{4} + SO_{2} + S. If a solution of the
potassium salt be decomposed by means of hydrofluosilicic or
chloric acid, the insoluble salts of these acids are precipitated
and trithionic acid is obtained in solution, which however very
easily breaks up on concentration. The addition of salts of
copper, mercury, silver, &c., to a solution of a trithionate is
followed, either immediately or after a certain time, by the
formation of a black precipitate of the sulphides whose formation
is due to the decomposition of the trithionic acid with the
transference of its sulphur to the metal.
_Tetrathionic acid_, H_{2}S_{4}O_{6}, in contradistinction to the
preceding acids, is much more stable in the free state than in the
form of salts. In the latter form it is easily converted into
trithionate, with liberation of sulphur. Sodium tetrathionate was
obtained by Fordos and Gélis, by the action of iodine on a
solution of sodium thiosulphate. The reaction essentially consists
in the iodine taking up half the sodium of the thiosulphate,
inasmuch as the latter contains Na_{2}S_{2}O_{3}, whilst the
tetrathionate contains NaS_{2}O_{3} or Na_{2}S_{4}O_{6}, so that
the reaction is as follows: 2Na_{2}S_{2}O_{3} + I_{2} = 2NaI +
Na_{2}S_{4}O_{6}. It is evident that tetrathionic acid stands to
thiosulphuric acid in exactly the same relation as dithionic acid
does to sulphurous acid; for the same amount of the other elements
in dithionate, KSO_{3}, and tetrathionate, KS_{2}O_{3}, there is
half as much metal as in sulphite, K_{2}SO_{3}, and thiosulphate,
K_{2}S_{2}O_{3}. If in the above reaction the sodium thiosulphate
be replaced by the lead salt PbS_{2}O_{3}, the sparingly-soluble
lead iodide PbI_{2} and the soluble salt PbS_{4}O_{6} are
obtained. Moreover the lead salt easily gives tetrathionic acid
itself (PbSO_{4} is precipitated). The solution of tetrathionic
acid may be evaporated over a water-bath, and afterwards in a
vacuum, when it gives a colourless liquid, which has no smell and
a very acid reaction. When dilute it may be heated to its
boiling-point, but in a concentrated form it decomposes into
sulphuric acid, sulphurous anhydride, and sulphur: H_{2}S_{4}O_{6}
= H_{2}SO_{4} + SO_{2} + S_{2}.
_Pentathionic acid_, H_{2}S_{5}O_{6}, also belongs to this series
of acids. But little is known concerning it, either as hydrate or
in salts. It is formed, together with tetrathionic acid, by the
direct action of sulphurous acid on sulphuretted hydrogen in an
aqueous solution; a large proportion of sulphur being precipitated
at the same time: 5SO_{2} + 5H_{2}S = H_{2}S_{5}O_{6} + 5S +
4H_{2}O.
If, as was shown above, the thionic acids are disulphonic acids,
they may be obtained, like other sulphonic acids, by means of
potassium sulphite and sulphur chloride. Thus Spring demonstrated
the formation of potassium trithionate by the action of sulphur
dichloride on a strong solution of potassium sulphite: 2KSO_{3}K +
SCl_{2} = S(SO_{3}K)_{2} + 2KCl. If sulphur chloride be taken,
sulphur also is precipitated. The same trithionate is formed by
heating a solution of double thiosulphates; for example, of
AgKS_{2}O_{3}. Two molecules of the salts then form silver
sulphide and potassium trithionate. If the thiosulphate be the
potassium silver salt SO_{3}K(AgS), then the structure of the
trithionate must necessarily be (SO_{3}K)_{2}S. Previous to
Spring's researches, the action of iodine on sodium thiosulphate
was an isolated accidentally discovered reaction; he, however,
showed its general significance by testing the action of iodine on
mixtures of different sulphur compounds. Thus with iodine, I_{2},
the mixture Na_{2}S + Na_{2}SO_{3} forms 2NaI + Na_{2}S_{2}O_{3},
whilst the mixture Na_{2}S_{2}O_{3} + Na_{2}SO_{3} + I_{2} gives
2NaI + Na_{2}S_{3}O_{6}--that is, trithionic acid stands in the
same relation to thiosulphuric acid as the latter does to
sulphuretted hydrogen. We adopt the same mode of representation:
by replacing one hydrogen in H_{2}S by sulphuryl we obtain
thiosulphuric acid, HSO_{3}.HS, and by replacing a second hydrogen
in the latter again by sulphuryl we obtain trithionic acid,
(HSO_{3})_{2}S. Furthermore, Spring showed that the action of
sodium amalgam on the thionic acids causes reverse reactions to
those above indicated for iodine. Thus sodium thiosulphate with
Na_{2} gives Na_{2}S + Na_{2}SO_{3}, and Spring showed that the
sodium here is not a simple element taking up sulphur, but itself
enters into double decomposition, replacing sulphur; for on taking
a potassium salt and acting on it with sodium, KSO_{3}(SK) + NaNa
= KSO_{3}Na + (SK)Na. In a similar way sodium dithionate with
sodium gives sodium sulphite: (NaSO_{3})_{2} + Na_{2} =
2NaSO_{3}Na; sodium trithionate forms NaSO_{3}Na and NaSO_{3}.SNa,
and tetrathionate forms sodium thiosulphate,
(NaSO_{3})S_{2}(NaSO_{3}) + Na_{2} = 2(NaSO_{3})(NaS).
In all the oxidised compounds of sulphur we may note the presence
of the elements of sulphurous anhydride, SO_{2}, the only product
of the combustion of sulphur, and in this sense the compounds of
sulphur containing one SO_{2} are--
H HO C_{6}H_{5} HS
SO_{2} SO_{2} SO_{2} SO_{2}
HO HO HO HO
Sulphurous Sulphuric Benzene sulphonic Thiosulphuric
acid acid acid acid
while, according to this mode of representation, the thionic acids
are--
HO HO HO HO
SO_{2} SO_{2} SO_{2} SO_{2}
S S_{2} S_{3}
SO_{2} SO_{2} SO_{2} SO_{2}
HO HO HO HO
Dithionic Trithionic Tetrathionic Pentathionic
Hence it is evident that SO_{2} has (whilst CO_{2} has not) the
faculty for combination, and aims at forming SO_{2}X_{2}. These
X_{2} can = O, and the question naturally suggests itself as to
whether the O_{2} which occurs in SO_{2} is not of the same nature
as this oxygen which adds itself to SO_{2}--that is, whether
SO_{2} does not correspond with the more general type SX_{4}, and
its compounds with the type SX_{6}? To this we may answer 'Yes'
and 'No'--'Yes' in the general sense which proceeds from the
investigation of the majority of compounds, especially metals,
where RO corresponds with RCl_{2}, RX_{2}; 'No' in the sense that
sulphur does not give either SH_{4}, SH_{6}, or SCl_{6}, and
therefore the stages SX_{4} and SX_{6} are only observable in
oxygen compounds. With reference to the type SX_{6} a hydrate,
S(HO)_{6}, might be expected, if not SCl_{6}. And we must
recognise this hydrate from a study of the compounds of sulphuric
acid with water. In addition to what has been already said
respecting the complex acids formed by sulphur, I think it well to
mention that, according to the above view, still more complex
oxygen acids and salts of sulphur may be looked for. For instance,
the salt Na_{2}S_{4}O_{8} obtained by Villiers (1888) is of this
kind. It is formed together with sodium trithionate and sulphur,
when SO_{2} is passed through a cold solution of Na_{2}S_{2}O_{3},
which is then allowed to stand for several days at the ordinary
temperature: 2Na_{2}S_{2}O_{3} + 4SO_{2} = Na_{2}S_{4}O_{8} +
Na_{2}S_{3}O_{6} + S. It may be assumed here, as in the thionic
acids, that there are two sulphoxyls, bound together not only by
S, but also by SO_{2}, or what is almost the same thing, that the
sulphoxyl is combined with the residue of trithionic acid, _i.e._
replaces one aqueous residue in trithionic acid.
Sulphur exhibits an acid character, not only in its compounds with
hydrogen and oxygen, but also in those with other elements. The compound
of sulphur and carbon has been particularly well investigated. It
presents a great analogy to carbonic anhydride, both in its elementary
composition and chemical character. This substance is the so-called
carbon bisulphide, CS_{2}, and corresponds with CO_{2}.
The first endeavours to obtain a compound of sulphur with carbon were
unsuccessful, for although sulphur does combine directly with carbon, yet
the formation of this compound requires distinctly definite conditions.
If sulphur be mixed with charcoal and heated, it is simply driven off
from the latter, and not the smallest trace of carbon bisulphide is
obtained. The formation of this compound requires that the charcoal
should be first heated to a red heat, but not above, and then either the
vapour of sulphur passed over it or lumps of sulphur thrown on to the
red-hot charcoal, but in small quantities, so as not to lower the
temperature of the latter. If the charcoal be heated to a white heat, the
amount of carbon bisulphide formed is less. This depends, in the first
place, on the carbon bisulphide dissociating at a high temperature.[69]
In the second place, Favre and Silberman showed that in the combustion of
one gram of carbon bisulphide (the products will be CO_{2} + 2SO_{2})
3,400 heat units are evolved--that is, the combustion of a molecular
quantity of carbon bisulphide evolves 258,400 heat units (according to
Berthelot, 246,000). From a molecule of carbon bisulphide in grams we may
obtain 12 grams of carbon, whose combustion evolves 96,000 heat units,
and 64 grams of sulphur, evolving by combustion (into SO_{2}) 140,800
heat units. Hence we see that the component elements separately evolve
less heat by their combustion (237,000 heat units) than carbon bisulphide
itself--that is, that heat should be evolved (at the ordinary
temperature) and not absorbed in its decomposition, and therefore that
the formation of carbon bisulphide from charcoal and sulphur is in all
probability accompanied by an absorption of heat.[70] It is therefore not
surprising that, like other compounds produced with an absorption of heat
(ozone, nitrous oxide, hydrogen peroxide, &c.), carbon bisulphide is
unstable and easily converted into the original substances from which it
is obtained. And indeed if the vapour of carbon bisulphide be passed
through a red-hot tube, it is decomposed--that is, it dissociates--into
sulphur and carbon. And this takes place at the temperature at which this
substance is formed, just as water decomposes into hydrogen and oxygen at
the temperature of its formation. In this absorption of heat in the
formation of carbon bisulphide is explained the facility with which it
suffers reactions of decomposition, which we shall see in the sequel, and
its main difference from the closely analogous carbonic anhydride.
[Illustration: FIG. 90.--Apparatus for the manufacture of carbon
bisulphide.
[69] Even light decomposes carbon bisulphide, but not to the extent of
separating carbon; under the action of the sun's rays it is
decomposed into sulphur and solid substance which is considered to
be carbon monosulphide; it is of a red colour, and its sp. gr. is
1·66. (The formation of a red liquid compound C_{3}S_{2} has also
been remarked.) Thorpe (1889) observed a complete decomposition of
carbon bisulphide under the action of a liquid alloy of potassium
and sodium; it is accompanied by an explosion and the deposition
of carbon and sulphur. A similar complete decomposition of carbon
bisulphide is also accomplished by the action of mercury fulminate
(Chapter XVI., Note 26), and is due to the fact that _at the
ordinary temperature_ (at which carbon bisulphide is not produced)
_the decomposition_ of carbon bisulphide takes place with the
development of heat--that is, it presents an exothermal reaction,
like the decomposition of all explosives. It is very possible that
at a higher temperature, when carbon bisulphide is formed, the
_combination_ of carbon with sulphur is also an exothermal
reaction--that is, heat is developed. If this should be the case,
carbon bisulphide would present a most instructive example in
thermochemistry.
[70] The fact should not be lost sight of that sulphur and charcoal are
solids at the ordinary temperature, whilst carbon bisulphide is a
very volatile liquid, and consequently, in the act of combination,
referred to the ordinary temperature (Note 69), there is, as it
were, a passage into a liquid state, and this requires the
absorption of heat. And furthermore, the molecule of sulphur
contains at least six atoms, and the molecule of carbon in all
probability (Chapter VIII.) a very considerable number of atoms;
thus the reaction of sulphur on charcoal may be expressed in the
following manner: 3C_{_n_} + _n_S_{6} = 3_n_CS_{2}--that is, from
_n_ + 3 molecules there proceed 3_n_ molecules, and as _n_ must be
very considerable, 3_n_ must be greater than 3 + _n_, which
indicates a decomposition in the formation of carbon bisulphide,
although the reaction at first sight appears as one of
combination. This decomposition is seen also from the volumes in
the solid and liquid states. Carbon bisulphide has a sp. gr. of
1·29; hence its molecular volume is 59. But the volume of carbon,
even in the form of charcoal, is not more than 6, and the volume
of S_{2} is 30; hence 36 volumes after combination give 59
volumes--an expansion takes place, as in decompositions.
In the laboratory carbon bisulphide is prepared as follows: A porcelain
tube is luted into a furnace in an inclined position, the upper extremity
of the tube being closed by a cork, and the lower end connected with a
condenser. The tube contains charcoal, which is raised to a red heat, and
then pieces of sulphur are placed in the upper end. The sulphur melts,
and its vapour comes into contact with the red-hot charcoal, when
combination takes place; the vapours condense in the condenser, carbon
bisulphide being a liquid boiling at 48°. On a large scale the apparatus
depicted in fig. 90 is employed. A cast-iron cylinder rests on a stand in
a furnace. Wood charcoal is charged into the cylinder through the upper
tube closed by a clay stopper, whilst the sulphur is introduced through a
tube reaching to the bottom of the cylinder. Pieces of sulphur thrown
into this tube fall on to the bottom of the cylinder, and are converted
into vapour, which passes through the entire layer of charcoal in the
cylinder. The vapour of carbon bisulphide thus formed passes through the
exit tube first into a Woulfe's bottle (where the sulphur which has not
entered into the reaction is condensed), and then into a strongly-cooled
condenser or worm.[71]
[71] Carbon bisulphide, as prepared on a large scale, is generally very
impure, and contains not only sulphur, but, more especially, other
impurities which give it a very disagreeable odour. The best
method of purifying this malodorous carbon bisulphide is to shake
it up with a certain amount of mercuric chloride, or even simply
with mercury, until the surface of the metal ceases to turn black.
After this the carbon bisulphide must be poured off and distilled
over a water-bath, after mixing with some oil to retain the
impurities.
Pure carbon bisulphide is a colourless liquid, which refracts light
strongly, and has a pure ethereal smell; at 0° its specific gravity is
1·293, and at 15° 1·271. If kept for a long time it seems to undergo a
change, especially when it is kept under water, in which it is insoluble.
It boils at 48°, and the tension of its vapour is so great that it
evaporates very easily, producing cold,[72] and therefore it has to be
kept in well-stoppered vessels; it is generally kept under a layer of
water, which hinders its evaporation and does not dissolve it.[73]
[72] If carbon bisulphide be evaporated under the receiver of an
air-pump, or by means of a current of air, it is possible to
obtain a temperature as low as -60°, and the carbon bisulphide
does not solidify at this temperature. However, if a series of
air-bubbles be passed through it by means of bellows, a
crystalline white substance remains which volatilises below 0°:
this a hydrate, H_{2}O,2CS_{2}; it easily decomposes into water
and carbon bisulphide. It is formed in the above experiment by the
moisture held in the air passed through the carbon bisulphide, and
the fall of temperature.
[73] Strong alcohol is miscible in all proportions with carbon
bisulphide, but dilute alcohol only in a definite amount, owing to
its diminished solubility from the presence of the water in it.
Ether, hydrocarbons, fatty oils, and many other organic substances
are soluble with great ease in carbon bisulphide. This is taken
advantage of in practice for extracting the fatty oils from
vegetable seeds, such as linseed, palm-nuts, or from bones, &c.
The preparation of vegetable oils is usually done by pressing the
seeds under a press, but the residue always contains a certain
amount of oil. These traces of oil can, however, be removed by
treatment with carbon bisulphide. In this manner a solution is
obtained which when heated easily parts with all the carbon
bisulphide, leaving the non-volatile fatty oil behind, so that the
same carbon bisulphide may be condensed and used over again for
the same purpose. It also dissolves iodine, bromine, indiarubber,
sulphur, and tars.
Carbon bisulphide, especially at high temperatures, very often
acts by its elements in a manner in which carbon and sulphur alone
are not able to react, which will be understood from what has been
said above respecting its endothermal origin. If it be passed over
red-hot metals--even over copper, for instance, not to mention
sodium, &c.--it forms a sulphide of the metal and deposits
charcoal, and if the vapour be passed over incandescent metallic
oxides it forms metallic sulphides and carbonic anhydride (and
sometimes a certain amount of sulphurous anhydride). Lime and
similar oxides give under these circumstances a carbonate and a
sulphide--for example, CS_{2} +3CaO = 2CaS + CaCO_{3}. The
sulphides obtained by this means are often well crystallised, like
those found in nature--for example, lead and antimony sulphides.
Carbon bisulphide enters into many combinations, which are frequently
closely analogous to the compounds of carbonic anhydride. In this respect
it is a _thio-anhydride_--_i.e._ it has the character of the acid
anhydrides,[73 bis] like carbonic anhydride, with the difference that the
oxygen of the latter is replaced by sulphur. By thio-compounds in general
are understood those compounds of sulphur which differ from the compounds
of oxygen as carbon bisulphide does from carbonic anhydride--that is,
which correspond with the oxygen compounds, but with substitution of
sulphur for oxygen. Thus thiosulphuric acid is monothiosulphuric
acid--that is, sulphuric acid in which one atom of sulphur replaces one
atom of oxygen. With the sulphides of the alkalis and alkaline earths, it
forms saline substances corresponding with the carbonates, and these
compounds may be termed _thiocarbonates_. For example, the composition of
the sodium salt Na_{2}CS_{3} is exactly like that of sodium carbonate.
They are formed by the direct solution of carbon bisulphide in aqueous
solutions of the sulphides; but they are difficult to obtain in a
crystalline form, because they are easily decomposable. When the
solutions of these salts are highly concentrated they begin to decompose,
with the evolution of sulphuretted hydrogen and the formation of a
carbonate, water taking part in the reaction--for example, K_{2}CS_{3} +
3H_{2}O = K_{2}CO_{3} + 3H_{2}S.[74]
[73 bis] And just as COCl_{2} corresponds to CO_{2}, so also the
chloranhydride, CSCl_{2}, or _thiophosgene_, corresponds to CS_{2}.
[74] If instead of a sulphide we take an alkali hydroxide, a
thiocarbonate is also formed, together with a carbonate--thus,
3BaH_{2}O_{2} + 3CS_{2} = 2BaCS_{3} + BaCO_{3} + 3H_{2}O. From the
instability of the thiocarbonates of the alkaline metals we can
clearly see the reason of the difficulty with which the salts of
the heavier metals are formed, whose basic properties are
incomparably weaker than those of the alkali metals. However,
these salts may be obtained by double decomposition. Ammonia in
reacting on carbon bisulphide gives, besides products like those
formed by other alkalis, a whole series of products of as complex
a structure as those substances which are produced by the action
of carbonic anhydride on ammonia. In the ninth chapter we examined
the formation of the ammonium carbonates, and saw the transition
from them into the cyanides. It is not surprising after this that
the action of carbon bisulphide on ammonia not only produces the
above-mentioned salts, but also amidic compounds corresponding
with them, in which the oxygen is wholly or partially replaced by
sulphur. Thus ammonium dithiocarbamate is very easily obtained if
carbon bisulphide be added to an alcoholic solution of ammonia,
and the mixture cooled in a closed vessel. The salt then separates
out in minute yellow crystals, CN_{2}H_{6}S_{2}.
Carbon bisulphide not only forms compounds with the metallic
sulphides, but also with sulphuretted hydrogen--that is, it forms
_thiocarbonic acid_, H_{2}CS_{3}. This is obtained by carefully
mixing solutions of thiocarbonates with dilute hydrochloric acid.
It then separates in an oily layer, which easily decomposes in the
presence of water into sulphuretted hydrogen and carbon
bisulphide, just as the corresponding carbonic acid (hydrate)
decomposes into water and carbonic anhydride. Carbon bisulphide
combines not only with sodium sulphide, but also with the
bisulphide, Na_{2}S_{2}, not, however, with the trisulphide,
Na_{2}S_{3}.
The relation of carbon bisulphide to the other carbon compounds
presents many most interesting features which are considered in
organic chemistry. We will here only turn our attention to one of
the compounds of this class. Ethyl sulphide, (C_{2}H_{5})_{2}S,
combines with ethyl iodide, C_{2}H_{5}I, forming a new molecule,
S(C_{2}H_{5})_{3}I. If we designate the hydrocarbon group, for
instance ethyl, C_{2}H_{5}, by Et, the reaction would be expressed
by the following equation : Et_{2}S + EtI = SEt_{3}I. This
compound is of a saline character, corresponds with salts of the
alkalis, and is closely analogous to ammonium chloride. It is
soluble in water; when heated it again splits up into its
components EtI and Et_{2}S, and with silver hydroxide gives a
hydroxide, Et_{3}S·OH, having the property of a distinct and
energetic alkali, resembling caustic ammonia. Thus the compound
group SEt_{3} combines, like potassium or ammonium, with iodine,
hydroxyl, chlorine, &c. The hydroxide SEt_{3}·OH is soluble in
water, precipitates metallic salts, saturates acids, &c. Hence
sulphur here enters into a relation towards other elements similar
to that of nitrogen in ammonia and ammonium salts, with only this
difference, that nitrogen retains, besides iodine, hydroxyl, and
other groups, also H_{4} or Et_{4} (for example, NH_{4}Cl,
NEt_{3}HI, NEt_{4}I), whilst sulphur only retains Et_{3}.
Compounds of the formula SH_{3}X are however unknown, only the
products of substitution SEt_{3}X, &c. are known. The distinctly
alkaline properties of the hydroxide, triethylsulphine hydroxide,
SEt_{3}OH, and also the sharply-defined properties of the
corresponding hydroxide, tetraethylammonium hydroxide, NEt_{4}OH,
depend naturally not only on the properties of the nitrogen and
sulphur entering into their composition, but also on the large
proportion of hydrocarbon groups they contain. Judging from the
existence of the ethylsulphine compounds, it might be imagined
that sulphur forms a compound, SH_{4}, with hydrogen; but no such
compound is known, just as NH_{5} is unknown, although NH_{4}Cl
exists.
A remarkable example[74 bis] of the thio-compounds is found in
_thiocyanic acid_--_i.e._ cyanic acid in which the oxygen is replaced by
sulphur, HCNS. We know (Chapter IX.) that with oxygen the cyanides of the
alkaline metals RCN give cyanates RCNO; but they also combine with
sulphur, and therefore if yellow prussiate of potash be treated as in the
preparation of potassium cyanide, and sulphur be added to the mass,
potassium thiocyanate, KNCS, is obtained in solution. This salt is much
more stable than potassium cyanate; it dissolves without change in water
and alcohol, forming colourless solutions from which it easily
crystallises on evaporation. It may be kept exposed to air even when in
solution; in dissolving in water it absorbs a considerable amount of
heat, and forms a starting-point for the preparation of all the
thiocyanates, RCNS, and organic compounds in which the metals are
replaced by hydrocarbon groups. Such, for example, is volatile mustard
oil, C_{3}H_{5}CSN (allyl thiocyanate),[75] which gives to mustard its
caustic properties. With ferric salts the thiocyanates give an
exceedingly brilliant red coloration, which serves for detecting the
smallest traces of ferric salts in solution. Thiocyanic acid, HCNS, may
be obtained by a method of double decomposition, by distilling potassium
thiocyanate with dilute sulphuric acid. It is a volatile colourless
liquid, having a smell recalling that of vinegar, is soluble in water,
and may be kept in solution without change.[75 bis]
[74 bis] Thorpe and Rodger (1889), by heating a mixture of lead
fluoride and phosphorus pentasulphide to 250° in an atmosphere of
dry nitrogen, obtained gaseous _phosphorus fluosulphide_, or
_thiophosphoryl fluoride_, PSF_{3}, corresponding with POCl_{3}.
This colourless gas is converted into a colourless liquid by a
pressure of eleven atmospheres; it does not act on dry mercury,
and takes fire spontaneously in air or oxygen, forming phosphorus
pentafluoride, phosphoric anhydride, and sulphurous anhydride. It
is soluble in ether, but is decomposed by water: PSF_{3} + 4H_{2}O
= H_{2}S + H_{3}PO_{4} + 3HF (Note 20).
[75] Although mustard oil may be obtained from the thiocyanates, it is
only an isomer of allyl thiocyanate proper, as is explained in
Organic Chemistry.
[75 bis] Sulphur can only replace half the oxygen in CO_{2}, as is seen
in _carbon oxysulphide_, or monothiocarbonic anhydride COS. This
substance was obtained by Than, and is formed in many reactions. A
certain amount is obtained if a mixture of carbonic oxide and the
vapour of sulphur be passed through a red-hot tube. When carbon
tetrachloride is heated with sulphurous anhydride, this substance
is also formed; but it is best obtained in a pure form by
decomposing potassium thiocyanate with a mixture of equal volumes
of water and sulphuric acid. A gas is then evolved containing a
certain amount of hydrocyanic acid, from which it may be freed by
passing it over wool containing moistened mercuric oxide, which
retains the hydrocyanic acid. The reaction is expressed by the
equation: 2KCNS + 2H_{2}SO_{4} + 2H_{2}O = K_{2}SO_{4} +
(NH_{4})_{2}SO_{4} + 2COS. It is also formed by passing the vapour
of carbon bisulphide over alumina or clay heated to redness
(Gautier; silicon sulphide is then formed). COS is also formed by
passing phosgene over a long layer of asbestos mixed with cadmium
sulphide at 270°; CdS + COCl_{3} = CdCl_{2} + COS (Nuricsán,
1892). The pure gas has an aromatic odour, is soluble in an equal
volume of water, which, however, acts on it, so that it must be
collected over mercury. When slightly heated, carbon oxysulphide
decomposes into sulphur and carbonic oxide. It burns in air with a
pale blue flame, explodes with oxygen, and yields potassium
sulphide and carbonate with potassium hydroxide: COS + 4KHO =
K_{2}CO_{3} + K_{2}S + 2H_{2}O.
The sulphur compounds of chlorine Cl_{2}S and Cl_{2}S_{2} may be
regarded on the one hand as products of the metalepsis of the sulphides
of hydrogen, H_{2}S and H_{2}S_{2}; and on the other hand of the oxygen
compounds of chlorine, because chloride of sulphur, Cl_{2}S, resembles
chlorine oxide, Cl_{2}O, whilst Cl_{2}S_{2} corresponds with the higher
oxide of chlorine; or thirdly, we may see in these compounds the type of
the acid chloranhydrides, because they are all decomposed by water,
forming hydrochloric acid, and sulphur tetrachloride, SCl_{4}, is
decomposed with the formation of sulphurous anhydride.[76]
[76] There is no reason for seeing any contradiction or mutual
incompatibility in these three views, because every analogy is
more or less modified by a change of elements. Thus, for instance,
it cannot be expected that the product of the metalepsis of
hydrogen sulphide would resemble the corresponding products of
water in all respects, because water has not the acid properties
of hydrogen sulphide. In the days of dualism and electrical
polarity it was supposed that the sulphur varied in its nature: in
hydrogen sulphide or potassium sulphide it was considered to be
negative, and in sulphurous anhydride or sulphur dichloride
positive. It then appeared evident that sulphur dichloride would
have no point of analogy with potassium sulphide. But metalepsis,
or its expression in the law of substitution, necessitates such
opinions being laid aside. If we can compare CO_{2}, CH_{4},
CCl_{4}, CHCl_{3}, CH_{3}(OH) with each other, we cannot recognise
any difference in the sulphur in SH_{2}, SCl_{2}, SK_{2}, or in
general SX_{2}, for otherwise we should have to acknowledge as
many different states of sulphur, carbon, or hydrogen as there are
compounds of sulphur, carbon, or hydrogen. The essential truth of
the matter is that all the elements in a molecule play their part
in the reactions into which it enters. Often this appears to be
contradicted in the result--for example, hydrogen alone may be
replaced; but it is not this hydrogen alone that has determined
the reaction; all the elements present have participated in it.
This may be made clearer by the following rough illustration.
Supposing two regiments of soldiers were fighting against each
other, and that several men were lost by one of the regiments; no
one could say that it was only these men who took part in the
engagement. The other men fired and the bullets flew over the
heads of their opponents. It was not only those who fell who
fought, although they only were removed from the field of battle;
the fighting proceeded among the masses, but only those few were
disabled who went forward and were more conspicuous &c.; not that
the remainder did not take part in the action; they also fought
and were an object of attack, only they remained sound and unhurt.
Hydrogen is lighter than other elements and its atoms more mobile;
it subjects itself more frequently and easily to reactions; but it
is not it alone which reacts, it is even less liable to attack
than other elements. It participates in exceedingly diverse
reactions, not indeed because the hydrogen itself varies, but
because one atom of it puts itself forward, another is hidden, one
is united with carbon, another feebly held by sulphur, one stands
or moves in the neighbourhood of oxygen, another is joined to a
hydrocarbon. All hydrogen atoms are equal, and equally serve as an
object of attack for the atoms of molecules encountering them, but
those only are removed from the sphere of action which are nearer
the surface of a molecule, which are more mobile, or held by a
less sum of forces. So also sulphur is one and the same in sulphur
dichloride, in sulphurous or sulphuric anhydride, in hydrogen
sulphide, in potassium sulphide, but it reacts differently, and
those elements which are with it also vary in their reactions
because they are with it, and it varies its reactions because it
is with them. It is possible to seize on a character common to
substances quantitatively and qualitatively analogous to each
other. It may be admitted that an element in certain forms is not
able to enter into reactions into which in other forms it enters
willingly, if only the requisite conditions are encountered; but
it must not therefore be concluded that an element changes its
essential quality in these different cases. The preceding remarks
touch on questions which are subject to much argument among
chemists, and I mention them here in order to show the treatment
of those most important problems of chemistry which lie at the
basis of this treatise.
[Illustration: FIG. 91.--Apparatus for the preparation of sulphur
chloride, and similar volatile compounds prepared by combustion in a
stream of chlorine.]
The compounds of sulphur with chlorine are prepared in the apparatus
depicted in fig. 91. As sulphur chloride is decomposed by water, the
chlorine evolved in the flask C must be dried before coming into contact
with the sulphur. It is therefore first passed through a Woulfe's bottle,
B, containing sulphuric acid, and then through the cylinder D containing
pumice stone moistened with sulphuric acid, and then led into the retort
E, in which the sulphur is heated. The compound which is formed distils
over into the receiver R. A certain amount of sulphur passes over with
the sulphur chloride, but if the resultant distillate be re-saturated
with chlorine and distilled no free sulphur remains, the boiling-point
rises to 144°, and pure sulphur chloride, S_{2}Cl_{2}, is obtained. It
has this formula because its vapour density referred to hydrogen is 68.
It is also obtained by heating certain metallic chlorides (stannous,
mercuric) with sulphur; both the metal and chlorine then combine with the
sulphur. Sulphur chloride is a yellowish-brown liquid, which boils at
144°, and has a specific gravity of 1·70 at 0°. It fumes strongly in the
air, reacting on the moisture contained therein, and has a heavy
chloranhydrous odour. It dissolves sulphur, is miscible with carbon
bisulphide, and falls to the bottom of a vessel containing water, by
which it is decomposed, forming sulphurous anhydride and hydrochloric
acid; but it first forms various lower stages of oxidation of sulphur,
because the addition of silver nitrate to the solution gives a black
precipitate. With hydrogen sulphide it gives sulphur and hydrochloric
acid, and it reacts directly with metals--especially arsenic, antimony,
and tin--forming sulphides and chlorides. In the cold, it absorbs
chlorine and gives _sulphur dichloride_, SCl_{2}. The entire conversion
into this substance requires the prolonged passage of dry chlorine
through sulphur chloride surrounded by a freezing mixture. The
distillation of the dichloride must be conducted in a stream of chlorine,
as otherwise it partially decomposes into sulphur chloride and chlorine.
Pure sulphur dichloride is a reddish-brown liquid, which resembles the
lower chloride in many respects; its specific gravity is 1·62; its odour
is more suffocating than that of sulphur chloride; it volatilises at
64°.[77]
[77] The observed vapour density of sulphur dichloride referred to
hydrogen is 53·3, and that given by the formula is 51·5. The
smaller molecular weight explains its boiling point being lower
than that of sulphur chloride, S_{2}Cl_{2}. The reactions of both
these compounds are very similar. Sulphur converts the dichloride,
SCl_{2}, into the monochloride, S_{2}Cl_{2}. In one point the
dichloride differs distinctly from the monochloride--that is, in
its capacity for easily giving up chlorine and decomposing. Even
light decomposes it into chlorine and the monochloride. Hence it
acts on many substances in the same manner as chlorine, or
substances which easily part with the latter, such as phosphoric
or antimonic chloride. In distinction to these, however, sulphur
dichloride would appear to distil without any considerable
decomposition, judging by the vapour density. But this is not a
valid conclusion, for if there be a decomposition, then 2SCl_{2} =
S_{2}Cl_{2} + Cl_{2}; now the density of sulphur chloride = 67·5,
and of chlorine = 35·5, and consequently a mixture of equal
volumes of the two = 51·5, just the same as an equal volume of
sulphur dichloride. _Therefore the distillation of sulphur
dichloride is probably nothing but its decomposition._ Hence the
compound SCl_{2}, which is stable at the ordinary temperature,
decomposes at 64°. In the cold it absorbs a further amount of
chlorine, corresponding to SCl_{4}, but even at -10° a portion of
the absorbed chlorine is given off--that is, dissociation takes
place. Thus the tetrachloride is even less stable than the
dichloride.
_Thionyl chloride_, SOCl_{2}, may be regarded as oxidised sulphur
dichloride; it corresponds with sulphur chloride, S_{2}Cl_{2}, in which
one atom of sulphur is replaced by oxygen. At the same time it is
chlorine oxide (hypochlorous anhydride, Cl_{2}O) combined with sulphur,
and also the chloranhydride of sulphurous acid--that is, SO(HO)_{2}, in
which the two hydroxyl groups are replaced by two atoms of chlorine, or
sulphurous anhydride, SO_{2}, in which one atom of oxygen is replaced by
two atoms of chlorine. All these representations are confirmed by
reactions of formation, or decompositions; they all agree with our
notions of the other compounds of sulphur, oxygen, and chlorine; hence
these definitions are not contradictory to each other. Thus, for
instance, thionyl chloride was first obtained by Schiff, by the action of
dry sulphurous anhydride on phosphorus pentachloride. On distilling the
resultant liquid, thionyl chloride comes over first at 80°, and on
continuing the distillation phosphorus oxychloride distils over at above
100°, PCl_{5} + SO_{2} = POCl_{3} + SOCl_{2}. This mode of preparation is
direct evidence of the oxychloride character of SOCl_{2}. Würtz obtained
the same substance by passing a stream of chlorine oxide through a cold
solution of sulphur in sulphur chloride; the chlorine oxide then combined
directly with the sulphur, S + Cl_{2}O = SOCl_{2}, whilst the sulphur
chloride remained unchanged (sulphur cannot be combined directly with
chlorine oxide, as an explosion takes place). Thionyl chloride is a
colourless liquid, with a suffocating acrid smell; it has a specific
gravity at 0° of 1·675, and boils at 78°. It sinks in water, by which it
is immediately decomposed, like all chloranhydrides--for example, like
carbonyl chloride, which corresponds with it: SOCl_{2} + H_{2}O = SO_{2}
+ 2HCl.[77 bis]
[77 bis] Hartog and Sims (1893) obtained thionyl bromide, SOBr_{2}, by
treating SOCl_{2} with sodium bromide; it is a red liquid, sp. gr.
2·62, and decomposes at 150°.
Normal _sulphuric acid has two corresponding chloranhydrides_; the first,
SO_{2}(OH)Cl, is sulphuric acid, SO_{2}(HO)_{2}, in which one equivalent
of HO is replaced by chlorine; the second has the composition
SO_{2}Cl_{2}--that is, two HO groups are substituted by two of chlorine.
The second chloranhydride, or the compound SO_{2}Cl_{2}, is called
sulphuryl chloride, and the first chloranhydride, SO_{2}HOCl, may be
called chlorosulphonic acid, because it is really an acid; it still
retains one hydroxyl of sulphuric acid, and its corresponding salts are
known. Thus, potassium chloride absorbs the vapour of sulphuric
anhydride, forming a salt, SO_{3}KCl, corresponding with SO_{3}HCl as
acid. In acting on sodium chloride it forms hydrochloric acid and the
salt NaSO_{3}Cl. This first chloranhydride of sulphuric acid, SO_{2}HOCl,
discovered by Williamson, is obtained either by the action of phosphorus
pentachloride on sulphuric acid (PCl_{5} + H_{2}SO_{4} = POCl_{3} + HCl +
HSO_{3}Cl), or directly by the action of dry hydrochloric acid on
sulphuric anhydride, SO_{3} + HCl = HSO_{3}Cl. The most easy and rapid
method of its formation is by direct saturation of cold Nordhausen acid
with dry hydrochloric acid gas (SO_{3} + HCl = HSO_{3}Cl), and
distillation of the resultant solution; the distillate then contains
HSO_{3}Cl. It is a colourless fuming liquid, having an acrid odour; it
boils at 153° (according to my determination, confirmed by Konovaloff),
and its specific gravity at 19° is 1·776. It is immediately decomposed by
water, forming hydrochloric and sulphuric acids, as should be the case
with a true chloranhydride. In the reactions of this chloranhydride we
find the easiest means of introducing the sulphonic group HSO_{3} into
other compounds, because it is here combined with chlorine. The second
chloranhydride of sulphuric acid, or _sulphuryl chloride_, SO_{2}Cl_{2},
was obtained by Regnault by the direct action of the sun's ray on a
mixture of equal volumes of chlorine and sulphurous oxide. The gases
gradually condense into a liquid, combining together as carbonic oxide
does with chlorine. It is also obtained when a mixture of the two gases
in acetic acid is allowed to stand for some time. The first
chloranhydride, SO_{3}HCl, decomposes when heated at 200° in a closed
tube into sulphuric acid and sulphuryl chloride. It boils at 70°, its
specific gravity is 1·7, it gives hydrochloric and sulphuric acids with
water, fumes in the air, and, judging by its vapour density, does not
decompose when distilled.[78]
[78] Pyrosulphuryl chloride, S_{2}O_{5}Cl_{2}. See Note 44. Thorpe and
Kirman, by treating SO_{3} with HF, obtained SO_{2}(OH)F, as a
liquid boiling at 163°, but which decomposed with greater facility
and then gave SO_{2}F_{2}.
The acids of sulphur naturally have their corresponding ammonium
salts, and the latter their amides and nitriles. It will be
readily understood how vast a field for research is presented by
the series of compounds of sulphur and nitrogen, if we only
remember that to carbonic and formic acids there corresponds, as
we saw (Chapter IX.), a vast series of derivatives corresponding
with their ammonium salts. To sulphuric acid there correspond two
ammonium salts, SO_{2}(HO)(NH_{4}O) and SO_{2}(NH_{4}O)_{2}; three
amides: the acid amide SO_{2}(HO)(NH_{2}), or sulphamic acid, the
normal saline compound SO_{2}(NH_{4}O)(NH_{2}), or ammonium
sulphamate, and the normal amide SO_{2}(NH_{2})_{2}, or sulphamide
(the analogue of urea); then the acid nitrile, SON(HO), and two
neutral nitriles, SON(NH_{2}) and SN_{2}. There are similar
compounds corresponding with sulphurous acid, and therefore its
nitriles will be, an acid, SN(HO), its salt, and the normal
compound, SN(NH_{2}). Dithionic and the other acids of sulphur
should also have their corresponding amides and nitriles. Only a
few examples are known, which we will briefly describe. Sulphuric
acid forms salts of very great stability with ammonia, and
ammonium sulphate is one of the commonest ammoniacal compounds. It
is obtained by the direct action of ammonia on sulphuric acid, or
by the action of the latter on ammonium carbonate; it separates
from its solutions in an anhydrous state, like potassium sulphate,
with which it is isomorphous. Hence, the composition of crystals
of ammonium sulphate is (NH_{4})_{2}SO_{4}. This salt fuses at
140°, and does not undergo any change when heated up to 180°. At
higher temperatures it does not lose water, but parts with half
its ammonia, and is converted into the acid salt, HNH_{4}SO_{4};
and this acid salt, on further heating, undergoes a further
decomposition, and splits up into nitrogen, water, and acid
ammonium sulphite, HNH_{4}SO_{3}. At the ordinary temperature the
normal salt is soluble in twice its weight of water and at the
boiling-point of water in an equal weight. In its faculty for
combinations this salt exhibits a great resemblance to potassium
sulphate, and, like it, easily forms a number of double salts; the
most remarkable of which are the ammonia alums,
NH_{4}AlS_{2}O_{8},12H_{2}O, and the double salts formed by the
metals of the magnesium group, having, for example, the
composition (NH_{4})_{2}MgS_{2}O_{8},6H_{2}O. Ammonium sulphate
does not give an amide when heated, perhaps owing to the faculty
of sulphuric anhydride to retain the water combined with it with
great force. But the amides of sulphuric acid may be very
conveniently prepared from sulphuric anhydride. Their formation by
this method is very easily understood because an amide is equal to
an ammonium salt less water, and if the anhydride be taken it will
give an amide directly with ammonia. Thus, if dry ammonia be
passed into a vessel surrounded by a freezing mixture and
containing sulphuric anhydride, it forms a white powdery mass
called sulphatammon, having the composition SO_{3},2H_{3}N, and
resembling the similar compound of carbonic acid, CO_{2},2NH_{3}.
This substance is naturally the ammonium salt of sulphamic acid,
SO_{2}(NH_{4}O)NH_{2}. It is slowly acted on by water, and may
therefore be obtained in solution, in which it slowly reacts with
barium chloride, which proves that with water it still forms
ammonium sulphate. If this substance be carefully dissolved in
water and evaporated, it yields well-formed crystals, whose
solution no longer gives a precipitate with barium chloride. This
is not due to the presence of impurities, but to a change in the
nature of the substance, and therefore Rose calls the crystalline
modification _parasulphatammon_. Platinum chloride only
precipitates half the nitrogen as platinochloride from solutions
of sulphat- and parasulphatammon, which shows that they are
ammonium salts, SO_{2}(NH_{4}O)(NH_{2}). It may be that the reason
of the difference in the two modifications is connected with the
fact that two different substances of the composition
N_{2}H_{4}SO_{2} are possible: one is the amide SO_{2}(NH_{2})_{2}
corresponding with the normal salt, and the other is the salt of
the nitrile acid corresponding with acid ammonium sulphate--that
is, SON(ONH_{4}) corresponds with the acid SON(OH) =
SO_{2}(NH_{4}O)OH - 2H_{2}O. Hence there may here be a difference
of the same nature as between urea and ammonium cyanate. Up to the
present, the isomerism indicated above has been but little
investigated, and might be the subject of interesting researches.
If in the preceding experiment the ammonia, and not the sulphuric
anhydride, be taken in excess, a soluble substance of the
composition 2SO_{2},3NH_{3} is formed. This compound, obtained by
Jacqueline and investigated by Voronin, doubtless also contains a
salt of sulphamic acid--that is, of the amide corresponding with
the acid ammonium sulphate = HNH_{4}SO_{4} - H_{2}O =
(NH_{2})SO_{2}(OH). Probably it is a compound of sulphatammon with
sulphamic acid. Thus it has an acid reaction, and does not give a
precipitate with barium chloride.
With normal sulphate of ammonium, an amide of the composition
N_{2}H_{4}SO_{2} should correspond, which should bear the same
relation to sulphuric acid as urea bears to carbonic acid. This
amide, known as _sulphamide_, is obtained by the action of dry
ammonia on the sulphuryl chloride, SO_{2}Cl_{2}, just as urea is
obtained by the action of ammonia on carbonyl chloride,
SO_{2}Cl_{2} + 4NH_{3} = N_{2}H_{4}SO_{2} + 2NH_{4}Cl. The
ammonium chloride is separated from the resultant sulphamide with
great difficulty. Cold water, acting on the mixture, dissolves
them both; the cold solution does not gives precipitate with
barium chloride. Alkalis act on it slowly, as they do on urea; but
on boiling, especially in the presence of alkalis or acids, it
easily recombines with water, and gives an ammonium salt. V.
Traube (1892) obtained sulphamide by the reaction of sulphuryl,
dissolved in chloroform, upon ammonia. The resultant precipitate
dissolves when shaken up with water, and the solution (after
boiling with the oxides or lead or silver) is evaporated, when a
syrupy liquid remains. With nitrate of silver the latter gives a
solid compound, which, when decomposed by hydrochloric acid, gives
free sulphamide in large colourless crystals, having the
composition SO_{2}(NH_{2})_{2}. This substance fuses at 81°,
begins to decompose below 100°, and is entirely decomposed above
250°; it is soluble in water, and the solution has a neutral
reaction and bitter taste. When heated with acids, sulphamide
gradually decomposes, forming sulphuric acid and ammonia. If the
silver compound obtained by the action of sulphamide on nitrate of
silver be heated at 170°-180° until ammonia is no longer evolved,
and the residue be extracted with water acidulated with nitric
acid, a salt separates out from the solution, answering in its
composition to sulphamide, SO_{2}NAg, which = the amide - NH_{3} =
SO_{2}N_{2}H_{4} - NH_{3} = SO_{2}NH. The action of sulphuryl
chloride (and of the other chloranhydrides of sulphur) on ammonium
carbonate always, as Mente showed (1888), results in the formation
of the salt NH(SO_{3}NH_{4})_{2}.
The nitriles corresponding with sulphuric acid are not as yet
known with any certainty. The most simple nitrile corresponding
with sulphuric acid should have the composition N_{2}H_{8}SO_{4} -
4H_{2}O = N_{2}S. This would be a kind of cyanogen corresponding
with sulphuric acid. On comparing sulphurous acid with carbonic
acid, we saw that they present a great analogy in many respects,
and therefore it might be expected that nitrile compounds having
the composition NHS and N_{2}S_{2} would be found. The latter of
these compounds is well known, and was obtained by Soubeiron, by
the action of dry ammonia on sulphur chloride. This substance
corresponds with cyanogen (paracyanogen), and is known as
_nitrogen sulphide_, N_{2}S_{2}. It is formed according to the
equation 3SCl_{2} + 8NH_{3} = N_{2}S_{2} + S + 6NH_{4}Cl. The free
sulphur and nitrogen sulphide are dissolved by acting on the
product with carbon bisulphide, the nitrogen sulphide being much
less soluble than the sulphur. It is a yellow substance, which is
excessively irritating to the eyes and nostrils. It explodes when
rubbed with a hard substance, being naturally decomposed with the
evolution of nitrogen; but when heated it fuses without
decomposing, and only decomposes with explosion at 157°. It is
insoluble in water, and only slightly so in alcohol, ether, and
carbon bisulphide; 100 parts of the latter dissolve 1·5 part of
nitrogen sulphide at the boiling point. This solution on cooling
deposits it in minute transparent prisms of a golden yellow
colour.
In the group of the halogens we saw four closely analogous
elements--fluorine, chlorine, bromine, and iodine--and we meet with the
same number of closely allied analogues in the oxygen group; for besides
sulphur this group also includes _selenium_ and _tellurium_: O, S, Se,
Te. These two groups are very closely allied, both in respect to the
magnitudes of their atomic weights and also in the faculty of the
elements of both groups for combining with metals. The distinct analogy
and definite degree of variance known to us for the halogens, also repeat
themselves in the same degree for the elements of the oxygen group.
Amongst the halogens fluorine has many peculiarities compared to Cl, Br
and I which are more closely analogous, whilst oxygen differs in many
respects from S, Se, Te, which possess greater similarities. The analogy
in a quantitative respect is perfect in both cases. Thus the halogens
combine with H, and the elements of the oxygen group with H_{2}, forming
H_{2}O, H_{2}S, H_{2}Se, H_{2}Te. The hydrogen compounds of selenium and
tellurium are acids like hydrogen sulphide. Selenium, by simple heating
in a stream of hydrogen, partially combines with it directly, but
seleniuretted hydrogen is more readily decomposable by heat than
sulphuretted hydrogen, and this property is still more developed in
telluretted hydrogen. Hydrogen selenide and telluride are gases like
sulphuretted hydrogen, and, like it, are soluble in water, form saline
compounds with alkalis, precipitate metallic salts, are obtained by the
action of acids on their compounds with metals, &c. Selenium and
tellurium, like sulphur, give two normal grades of combination with
oxygen, both of an acid character, of which only the forms corresponding
to sulphurous anhydride--namely, selenious anhydride, SeO_{2}, and
tellurous anhydride, TeO_{2}[79]--are formed directly. These are both
solids, obtained by the combustion of the elements themselves and by the
action of oxidising agents on them. They form feebly energetic acids,
having distinct bibasic properties; however, a characteristic difference
from SO_{2} is observable both in the physical properties of these
compounds and in their stability and capacity for further oxidation, just
as in the series of the halogens already known to us, only in an inverse
order; in the latter we saw that iodine combines more easily than bromine
or chlorine with oxygen, forming more stable oxygen compounds, whereas
here, on the contrary, sulphurous anhydride, as we know, is difficultly
decomposed, parts with its sulphur with difficulty, and is easily
oxidised and especially in its salts, while selenious and tellurous
anhydrides are oxidised with difficulty and easily reduced, even by means
of sulphurous acid.
[79] _Selenious anhydride_, SeO_{2}, is a volatile solid, which
crystallises in prisms soluble in water. It is best procured by
the action of nitric acid on selenium. The well-known researches
of Nilson (1874) showed that the salts of selenious acid easily
form acid salts, and are so characteristic in many respects that
they may even serve for judging the analogy of types of oxides.
Thus the oxides of the composition RO give normal salts of the
composition RSeO_{3},2H_{2}O, where R = Mn, Co, Ni, Cu, Zn. The
salts of magnesium, barium, and calcium contain a different
quantity of water, as do also the salts of the oxides R_{2}O_{3}.
We here turn attention to the fact that beryllium gives a normal
salt, BeSeO_{3},2H_{2}O, and not a salt analogous to those of
aluminium, scandium, Sc_{2}(SeO_{3})_{3},H_{2}O, yttrium,
Y_{2}(SeO_{3})_{2},12H_{2}O, and other oxides of the form
R_{2}O_{3}, which speaks in favour of the formula BeO.
_Tellurous anhydride_ is also a colourless solid, which
crystallises in octahedra; it also, when heated, first fuses and
then volatilises. It is insoluble in water, and the decomposition
of its salts gives a hydrate, H_{2}TeO_{3}, which is insoluble.
It is a very characteristic circumstance that selenious and
tellurous anhydrides are very easily _reduced_ to selenium and
tellurium. This is not only effected by metals like zinc, or by
sulphuretted hydrogen, which are powerful deoxidisers, but even by
sulphurous anhydride, which is able to precipitate selenium and
tellurium from solutions of the selenites and tellurites, and even
of the acids themselves, which is taken advantage of in obtaining
these elements and separating them from sulphur.
Sulphuric acid, as we know, rarely acts as an oxidising agent. It
is otherwise with selenic and telluric acids, H_{2}SeO_{4} and
H_{2}TeO_{4}, which are powerful oxidising agents--that is, are
easily reduced in many circumstances either into the lower oxide
or even to selenium and tellurium. A powerful oxidising agent is
required in order to convert selenious and tellurous anhydrides
into selenic and telluric anhydrides, and, moreover, it must be
employed in excess. If chlorine be passed through a solution of
potassium selenide, K_{2}Se, telluride, K_{2}Te, selenite,
K_{2}SeO_{3}, or tellurite, K_{2}TeO_{3}, it acts as an oxidiser
in the presence of the water, forming potassium selenate,
K_{2}SeO_{4}, or tellurate, K_{2}TeO_{4}. The same salts are
formed by fusing the lower oxides with nitre. These salts are
isomorphous with the corresponding sulphates, and cannot therefore
be separated from them by crystallisation. The salts of potassium,
sodium, magnesium, copper, cadmium, &c. are soluble like the
sulphates, but those of barium and calcium are insoluble, in
perfect analogy with the sulphates. When copper selenate,
CuSeO_{4}, is treated with sulphuretted hydrogen (CuS is
precipitated), _selenic acid_ remains in solution. On evaporation
and drying in vacuo at 180° it gives a syrupy liquid, which may be
concentrated to almost the pure acid, H_{2}SeO_{4}, having a
specific gravity of 2·6. Cameron and Macallan (1891) showed that
pure H_{2}SeO_{4} only remains liquid in a state of superfusion
whilst the solidified acid melts at +58°, the solid acid
crystallises well, its sp. gr. is then 2·95. The hydrate
H_{2}SeO_{4},H_{2}O melts at +25°. The acid in a superfused state
has a sp. gr. 2·36 and the solid 2·63. Like sulphuric acid strong
selenic acid attracts moisture from the atmosphere; it is not
decomposed by sulphurous acid, but oxidises hydrochloric acid
(like nitric, chromic, and manganic acids), evolving chlorine and
forming selenious acid, H_{2}SeO_{4} + 2HCl = H_{2}SeO_{3} +
H_{2}O + Cl_{2}. _Telluric acid_, H_{2}TeO_{4}, is obtained by
fusing tellurous anhydride with potassium hydroxide and chlorate;
the solution, containing potassium tellurate, is then precipitated
with barium chloride, and the barium tellurate, BaTeO_{4} obtained
in the precipitate is decomposed by sulphuric acid. A solution of
telluric acid is thus obtained, which on evaporation yields
colourless prisms, soluble in water, and containing
TeH_{2}O_{4},2H_{2}O. Two equivalents of water are driven off at
160°; on further heating the last equivalent of water is expelled,
and then oxygen is given off. It also gives chlorine with
hydrochloric acid, like selenic acid. Its salts also correspond
with those of sulphuric acid. It must, however, be remarked that
telluric and selenic acids are able to give poly-acid salts with
much greater ease than sulphuric acid. Thus, for example, there
are known for telluric acid not only K_{2}TeO_{4},5H_{2}O and
KHTeO_{4},3H_{2}O, but also KHTeO_{4},H_{2}TeO_{4},H_{2}O =
K_{2}TeO_{4},3H_{2}TeO_{4},2H_{2}O. This salt is easily obtained
from acid solutions of the preceding salts and is less soluble in
water. As selenious anhydride is volatile and gives similar
poly-salts, it may be surmised that selenious, tellurous, selenic,
and telluric anhydrides are polymeric as compared with sulphurous
and sulphuric anhydrides, for which reason it would be desirable
to determine the vapour density of selenious anhydride. It would
probably correspond with Se_{2}O_{4} or Se_{3}O_{6}.
In order to show the very close analogy of selenium to sulphur, I
will quote two examples. Potassium cyanide dissolves selenium, as
it does sulphur, forming potassium selenocyanate, KCNSe,
corresponding with potassium thiocyanate. Acids precipitate
selenium from this solution, because selenocyanic acid, H_{2}CNSe,
when in a free state is immediately decomposed. A boiling solution
of sodium sulphite dissolves selenium, just as it would sulphur,
forming a salt analogous to thiosulphate of sodium, namely, sodium
selenosulphate, Na_{2}SSeO_{3}. Selenium is separated from a
solution of this salt by the action of acid.
_Selenium_ was obtained in 1817 by Berzelius from the sublimate which
collects in the first chamber in the preparation of sulphuric acid from
Fahlun pyrites. Certain other pyrites also contain small quantities of
selenium. Some native selenides, especially those of lead, mercury, and
copper, have been found in the Hartz Mountains, but only in small
quantities. Pyrites and blendes, in which the sulphur is partially
replaced by selenium, still remain the chief source for its extraction.
When these pyrites are roasted they evolve selenious anhydride, which
condenses in the cooler portions of the apparatus in which the pyrites
are roasted, and is partially or wholly reduced by the sulphurous
anhydride simultaneously formed. The presence of selenium in ores and
sublimates is most simply tested by heating them before the blowpipe,
when they evolve the characteristic odour of garlic. Selenium exhibits
two modifications, like sulphur: one amorphous and insoluble in carbon
bisulphide, the other crystalline and slightly soluble in carbon
bisulphide (in 1,000 parts at 45° and 6,000 at 0°), and separating from
its solutions in monoclinic prisms. If the red precipitate obtained by
the action of sulphurous anhydride on selenious anhydride be dried, it
gives a brown powder, having a specific gravity of 4·26, which when
heated changes colour and fuses to a metallic mass, which gains lustre as
it cools. The selenium acquires different properties according to the
rate at which it is cooled from a fused state; if rapidly cooled, it
remains amorphous and has the same specific gravity (4·28) as the powder,
but if slowly cooled it becomes crystalline and opaque, soluble in carbon
bisulphide, and has a specific gravity of 4·80. In this form it fuses at
214° and remains unchanged, whilst the amorphous form, especially above
80°, gradually passes into the crystalline variety. The transition is
accompanied by the evolution of heat, as in the case of sulphur; thus the
analogy between sulphur and selenium is clearly shown here. In the fused
amorphous form selenium presents a brown mass, slightly translucent, with
a vitreous fracture, whilst in the crystalline form it has the appearance
of a grey metal, with a feeble lustre and a crystalline fracture.[79 bis]
Selenium boils at 700°, forming a vapour whose density is only constant
at a temperature of about 1,400°, when it is equal to 79·4 (referred to
hydrogen)--that is, the molecular formula is then Se_{2}, like sulphur at
an equally high temperature.
[79 bis] Muthmann, in his researches upon the allotropic forms of
selenium, pointed out (1889) a peculiar modification, which
appears, as it were, as a transition between crystalline and
amorphous selenium. It is obtained together with the crystalline
variety by slowly evaporating a solution of selenium in bisulphide
of carbon, and differs from the crystalline variety in the form of
its crystals; it passes into the latter modification when heated.
Schultz also obtained selenium (like Ag, _see_ Chapter XXIV.) in a
soluble form, but these researches are not so conclusive as those
upon soluble silver, and we shall therefore not consider them more
fully.
_Tellurium_ is met with still more rarely than selenium (it is known in
Saxony) in combination with gold, silver, lead, and antimony in the
so-called foliated tellurium ore. Bismuth telluride and silver telluride
have been found in Hungary and in the Altai. Tellurium is extracted from
bismuth telluride by mixing the finely-powdered ore with potassium and
charcoal in as intimate a mixture as possible, and then heating in a
covered crucible. Potassium telluride, K_{2}Te, is then formed, because
the charcoal reduces potassium tellurite. As potassium telluride is
soluble in water, forming a red-brown solution which is decomposed by the
oxygen of the atmosphere (K_{2}Te + O + H_{2}O = 2KHO + Te), the mass
formed in the crucible is treated with boiling water and filtered as
rapidly as possible, and the resultant solution exposed to the air, by
which means the tellurium is precipitated.[80] In a free state tellurium
has a perfectly _metallic appearance_; it is of a silver-white colour,
crystallises very easily in long brilliant needles; is very brittle, so
that it can be easily reduced to powder; but it is a bad conductor of
heat and electricity, and in this respect, as in many others, it forms a
transition from the metals to the non-metals. Its specific gravity is
6·18, it melts at an incipient red heat, and takes fire when heated in
air, like selenium and sulphur, burning with a blue flame, evolving white
fumes of tellurous anhydride, TeO_{2}, and emitting an acrid smell if no
selenium be present; but if it be, the odour of the latter preponderates.
Alkalis dissolve tellurium when boiled with it, potassium telluride,
K_{2}Te, and potassium tellurite, K_{2}TeO_{3}, being formed. The
solution is of a red colour, owing to the presence of the telluride,
K_{2}Te; but the colour disappears when the solution is cooled or
diluted, the tellurium being all precipitated: 2K_{2}Te + K_{2}TeO_{3} +
3H_{2}O = 6KHO + 3Te.[81]
[80] The tellurium thus prepared is impure, and contains a large amount
of selenium. The latter may be removed by converting the mixture
into the salts of potassium, and treating this with nitric acid
and barium nitrate, when barium selenate only is precipitated,
whilst the barium tellurate remains in solution. This method does
not, however, give a pure product, and it appears to be best to
separate the selenium from the tellurium in a metallic form; this
is done by boiling the impure potassium tellurate with
hydrochloric acid, which converts it into potassium tellurite,
from which the tellurium is reduced by sulphurous anhydride. The
metal thus obtained is then fused and distilled in a stream of
hydrogen; the selenium volatilises first, and then the tellurium,
owing to its being much less volatile than the former.
Nevertheless, tellurium is also volatile, and may be separated in
this manner from less volatile metals, such as antimony. Brauner
determined the atomic weight of pure tellurium, and found it to be
125, but showed (1889) that tellurium purified by the usual
method, even after distillation, contains a large amount of
impurities.
[81] The decomposition proceeds in the above order in the cold, but in
a hot solution with an excess of potassium hydroxide it proceeds
inversely. A similar phenomenon takes place when tellurium is
fused with alkalis, and it is therefore necessary in order to
obtain potassium telluride to add charcoal.
Selenium and tellurium form higher compounds with chlorine with
comparative ease. For selenium, SeCl_{2} and SeCl_{4} are known,
and for tellurium TeCl_{2} and TeCl_{4}. The tetrachlorides of
selenium and tellurium are formed by passing chlorine over these
elements. Selenium tetrachloride, SeCl_{4}, is a crystalline,
volatile mass which gives selenious anhydride and hydrochloric
acid with water. Tellurium tetrachloride is much less volatile,
fuses easily, and is also decomposed by water. Both elements form
similar compounds with bromine. Tellurium tetrabromide is red,
fuses to a brown liquid, volatilises, and gives a crystalline
salt, K_{2}TeBr_{6},3H_{2}O, with an aqueous solution of potassium
bromide.
CHAPTER XXI
CHROMIUM, MOLYBDENUM, TUNGSTEN, URANIUM, AND MANGANESE
Sulphur, selenium, and tellurium belong to the uneven series of the sixth
group. In the even series of this group there are known _chromium,
molybdenum, tungsten, and uranium_; these give acid oxides of the type
RO_{3}, like SO_{3}. Their acid properties are less sharply defined than
those of sulphur, selenium, and tellurium, as is the case with all
elements of the even series as compared with those of the uneven series
in the same group. But still the oxides CrO_{3}, MoO_{3}, WO_{3}, and
even UO_{3}, have clearly defined acid properties, and form salts of the
composition MO,_n_RO_{3} with bases MO. In the case of the heavy
elements, and especially of uranium, the type of oxide, UO_{3}, is less
acid and more basic, because in the even series of oxides the element
with the highest atomic weight always acquires a more and more pronounced
basic character. Hence UO_{3} shows the properties of a base, and gives
salts UO_{2}X_{2}. The basic properties of chromium, molybdenum,
tungsten, and uranium are most clearly expressed in the lower oxides,
which they all form. Thus chromic oxide, Cr_{2}O_{3}, is as distinct a
base as alumina, Al_{2}O_{3}.
Of all these elements _chromium_ is the most widely distributed and the
most frequently used. It gives chromic anhydride, CrO_{3}, and chromic
oxide, Cr_{2}O_{3}--two compounds whose relative amounts of oxygen stand
in the ratio 2 : 1. Chromium is, although somewhat rarely, met with in
nature as a compound of one or the other type. The red chromium ore of
the Urals, lead chromate or crocoisite PbCrO_{4}, was the source in which
chromium was discovered by Vauquelin, who gave it this name (from the
Greek word signifying colour) owing to the brilliant colours of its
compounds; the chromates (salts of chromic anhydride) are red and yellow,
and the chromic salts (from Cr_{2}O_{3}) green and violet. The red lead
chromate is, however, a rare chromium ore found only in the Urals and in
a few other localities. Chromic oxide, Cr_{2}O_{3}, is more frequently
met with. In small quantities it forms the colouring matter of many
minerals and rocks--for example, of some serpentines. The commonest ore,
and the chief source of the chromium compounds, is the _chrome iron ore_
or chromite, which occurs in the Urals[1] and Asia Minor, California,
Australia, and other localities. This is magnetic iron ore,
FeO,Fe_{2}O_{3}, in which the ferric oxide is replaced by chromic oxide,
its composition being FeO,Cr_{2}O_{3}. Chrome iron ore crystallises in
octahedra of sp. gr. 4·4; it has a feeble metallic lustre, is of a
greyish-black colour, and gives a brown powder. It is very feebly acted
on by acids, but when fused with potassium acid sulphate it gives a
soluble mass, which contains a chromic salt, besides potassium sulphate
and ferrous sulphate. In practice the treatment of chrome iron ore is
mainly carried on for the preparation of chromates, and not of chromic
salts, and therefore we will trace the history of the element by
beginning with chromic acid, and especially with the working up of the
chrome iron ore into _potassium dichromate_, K_{2}Cr_{2}O_{7}, as the
most common salt of this acid. It must be remarked that chromic
anhydride, CrO_{3}, is only obtained in an anhydrous state, and is
distinguished for its capacity for easily giving anhydro-salts with the
alkalis, containing one, two, and even three equivalents of the anhydride
to one equivalent of base. Thus among the potassium salts there is known
the normal or yellow chromate, K_{2}CrO_{4}, which corresponds to, and is
perfectly isomorphous with, potassium sulphate, easily forms isomorphous
mixtures with it, and is not therefore suitable for a process in which it
is necessary to separate the salt from a mixture containing sulphates. As
in the presence of a certain excess of acid, the dichromate,
K_{2}Cr_{2}O_{7} = 2K_{2}CrO_{4} + 2HX - 2KX - H_{2}O, is easily formed
from K_{2}CrO_{4}, the object of the manufacturer is to produce such a
dichromate, the more so as it contains a larger proportion of the
elements of chromic acid than the normal salt. Finely-ground chrome iron
ore, when heated with an alkali, absorbs oxygen almost as easily (Chapter
III., Note 7) as a mixture of the oxides of manganese with an alkali.
This absorption is due to the presence of chromic oxide, which is
oxidised into the anhydride, and then combines with the alkali
Cr_{2}O_{3} + O_{3} = 2CrO_{3}. As the oxidation and formation of the
chromate proceeds, the mass turns _yellow_. The iron is also oxidised,
but does not give ferric acid, because the capacity of the chromium for
oxidation is incomparably greater than that of the iron.
[1] The working of the Ural chrome iron ore into chromium compounds has
been firmly established in Russia, thanks to the endeavours of P.
K. Ushakoff, who constructed large works for this purpose on the
river Kama, near Elabougi, where as much as 2,000 tons of ore are
treated yearly, owing to which the importation of chromium
preparations into Russia has ceased.
A mixture of lime (sometimes with potash) and chrome iron ore is heated
in a reverberatory furnace, with free access of air and at a red heat for
several hours, until the mass becomes yellow; it then contains normal
calcium chromate, CaCr_O_{4}, which is insoluble in water in the presence
of an excess of lime.[1 bis] The resultant mass is ground up, and treated
with water and sulphuric acid. The excess of lime forms gypsum, and the
soluble calcium dichromate, CaCr_{2}O_{7}, together with a certain amount
of iron, pass into solution. The solution is poured off, and chalk added
to it; this precipitates the ferric oxide (the ferrous oxide is converted
into ferric oxide in the furnace) and forms a fresh quantity of gypsum,
while the chromic acid remains in solution--that is, it does not form the
sparingly-soluble normal salt (1 part soluble in 240 parts of water). The
solution then contains a fairly pure calcium dichromate, which by double
decomposition gives other chromates; for example, with a solution of
potassium sulphate it gives a precipitate of calcium sulphate and a
solution of potassium dichromate, which crystallises when evaporated.[2]
[1 bis] But the calcium chromate is soluble in water in the presence of
an excess of chromic acid, as may be seen from the fact that a
solution of chromic acid dissolves lime.
[2] There are many variations in the details of the manufacturing
processes, and these must be looked for in works on technical
chemistry. But we may add that the chromate may also be obtained by
slightly roasting briquettes of a mixture of chrome iron and lime,
and then leaving the resultant mass to the action of moist air
(oxygen is absorbed, and the mass turns yellow).
_Potassium dichromate_, K_{2}Cr_{2}O_{7}, easily crystallises from acid
solutions in red, well-formed prismatic crystals, which fuse at a red
heat and evolve oxygen at a very high temperature, leaving chromic oxide
and the normal salt, which undergoes no further change: 2K_{2}Cr_{2}O_{7}
= 2K_{2}CrO_{4} + Cr_{2}O_{3} + O_{3}. At the ordinary temperature 100
parts of water dissolve 10 parts of this salt, and the solubility
increases as the temperature rises. It is most important to note that the
dichromate does not contain water, it is K_{2}CrO_{4} + CrO_{3}; the acid
salt corresponding to potassium acid sulphate, KHSO_{4}, does not exist.
It does not even evolve heat when dissolving in water, but on the
contrary produces cold, _i.e._ it does not form a very stable compound
with water. The solution and the salt itself are poisonous, and act as
powerful oxidising agents, which is the character of chromic acid in
general. When heated with sulphur or organic substances, with sulphurous
anhydride, hydrogen sulphide, &c., this salt is deoxidised, yielding
chromic compounds.[2 bis] Potassium dichromate[3] is used in the arts and
in chemistry as a source for the preparation of all other chromium
compounds. It is converted into yellow pigments by means of double
decomposition with salts of lead, barium, and zinc. When solutions of the
salts of these metals are mixed with potassium dichromate (in dyeing
generally mixed with soda, in order to obtain normal salts), they are
precipitated as insoluble normal salts; for example, 2BaCl_{2} +
K_{2}Cr_{2}O_{7} + H_{2}O = 2BaCrO_{4} + 2KCl + 2HCl. It follows from
this that these salts are insoluble in dilute acids, but the
precipitation is not complete (as it would be with the normal salt). The
barium and zinc salts are of a lemon yellow colour; the lead salt has a
still more intense colour passing into orange. Yellow cotton prints are
dyed with this pigment. The silver salt, Ag_{2}CrO_{4}, is of a bright
red colour.
[2 bis] The oxidising action of potassium dichromate on organic
substances at the ordinary temperature is especially marked under
the action of light. Thus it acts on gelatin, as Poutven
discovered; this is applied to photography in the processes of
photogravure, photo-lithography, pigment printing, &c. Under the
action of light this gelatin is oxidised, and the chromic anhydride
deoxidised into chromic oxide, which unites with the gelatin and
forms a compound insoluble in warm water, whilst where the light
has not acted, the gelatin remains soluble, its properties being
unaffected by the presence of chromic acid or potassium dichromate.
[3] Ammonium and sodium dichromates are now also prepared on a large
scale. The sodium salts may be prepared in exactly the same manner
as those of potassium. The normal salt combines with ten
equivalents of water, like Glauber's salt, with which it is
isomorphous. Its solution above 30° deposits the anhydrous salt.
Sodium dichromate crystals contain Na_{2}Cr_{2}O_{7},2H_{2}O. The
_ammonium salts of chromic acid_ are obtained by saturating the
anhydride itself with ammonia. The dichromate is obtained by
saturating one part of the anhydride with ammonia, and then adding
a second part of anhydride and evaporating under the receiver of an
air-pump. On ignition, the normal and acid salts leave chromic
oxide. Potassium ammonium chromate, NH_{4}KCrO_{4}, is obtained in
yellow needles from a solution of potassium dichromate in aqueous
ammonia; it not only loses ammonia and becomes converted into
potassium dichromate when ignited, but also by degrees at the
ordinary temperature. This shows the feeble energy of chromic acid,
and its tendency to form stable dichromates. Magnesium chromate is
soluble in water, as also is the strontium salt. The calcium salt
is also somewhat soluble, but the barium salt is almost insoluble.
The isomorphism with sulphuric acid is shown in the chromates by
the fact that the magnesium and ammonium salts form double salts
containing six equivalents of water, which are perfectly
isomorphous with the corresponding sulphates. The magnesium salt
crystallises in large crystals containing seven equivalents of
water. The beryllium, cerium, and cobalt salts are insoluble in
water. Chromic acid dissolves manganous carbonate, but on
evaporation the solution deposits manganese dioxide, formed at the
expense of the oxygen of the chromic acid. Chromic acid also
oxidises ferrous oxide, and ferric oxide is soluble in chromic
acid.
One of the chromates most used by the dyer is the insoluble yellow
lead chromate, PbCrO_{4} (Chapter XVIII., Note 46), which is
precipitated on mixing solutions of PbX_{2} with soluble chromates.
It easily forms a basic salt, having the composition PbO,PbCrO_{4},
as a crystalline powder, obtained by fusing the normal salt with
nitre and then rapidly washing in water. The same substance is
obtained, although impure and in small quantity, by treating lead
chromate with neutral potassium chromate, especially on boiling the
mixture; and this gives the possibility of attaining, by means of
these materials, various tints of lead chromate, from yellow to
red, passing through different orange shades. The decomposition
which takes place (incompletely) in this case is as follows:
2PbCrO_{4} + K_{2}CrO_{4} = PbCrO_{4},PbO + K_{2}Cr_{2}O_{7}--that
is, potassium dichromate is formed in solution.
When potassium dichromate is mixed with potassium hydroxide or carbonate
(carbonic anhydride being disengaged in the latter case) it forms the
_normal_ salt, K_{2}CrO_{4}, known as _yellow chromate of potassium_. Its
specific gravity is 2·7, being almost the same as that of the dichromate.
It absorbs heat in dissolving; one part of the salt dissolves in 1·75
part of water at the ordinary temperature, forming a yellow solution.
When mixed even with such feeble acids as acetic, and more especially
with the ordinary acids, it gives the dichromate, and Graham obtained a
trichromate, K_{2}Cr_{3}O_{10} = K_{2}CrO_{4},2CrO_{3}, by mixing a
solution of the latter salt with an excess of nitric acid.
_Chromic anhydride_ is obtained by preparing a saturated solution of
potassium dichromate at the ordinary temperature, and pouring it in a
thin stream into an equal volume of pure sulphuric acid.[4] On mixing,
the temperature naturally rises; when slowly cooled, the solution
deposits chromic anhydride in needle-shaped crystals of a red colour
sometimes several centimetres long. The crystals are freed from the
mother liquor by placing them on a porous tile.[4 bis] It is very
important at this point to call attention to the fact that a hydrate of
chromic anhydride is never obtained in the decomposition of chromic
compounds, but always the _anhydride_, CrO_{3}. The corresponding
hydrate, CrO_{4}H_{2}, or any other hydrate, is not even known.
Nevertheless, it must be admitted that chromic acid is bibasic, because
it forms salts isomorphous or perfectly analogous with the salts formed
by sulphuric acid, which is the best example of a bibasic acid. A clear
proof of the bibasicity of CrO_{3} is seen in the fact that the anhydride
and salts give (when heated with sodium chloride and sulphuric acid) a
volatile chloranhydride, CrO_{2}Cl_{2}, containing two atoms of chlorine
as a bibasic acid should.[5] Chromic anhydride is a red crystalline
substance, which is converted into a black mass by heat; it fuses at
190°, and disengages oxygen above 250°, leaving a residue of chromium
dioxide, CrO_{2},[6] and, on still further heating, chromic oxide,
Cr_{2}O_{3}. Chromic anhydride is exceedingly soluble in water, and even
attracts moisture from the air, but, as was mentioned above, it does not
form any definite compound with water. The specific gravity of its
crystals is 2·7, and when fused it has a specific gravity 2·6. The
solution presents perfectly defined acid properties. It liberates
carbonic anhydride from carbonates; gives insoluble precipitates of the
chromates with salts of barium, lead, silver, and mercury.
[4] The sulphuric acid should not contain any lower oxides of nitrogen,
because they reduce chromic anhydride into chromic oxide. If a
solution of a chromate be heated with an excess of acid--for
instance, sulphuric or hydrochloric acid--oxygen or chlorine is
evolved, and a solution of a chromic salt is formed. Hence, under
these circumstances, chromic acid cannot be obtained from its
salts. One of the first methods employed consisted in converting
its salts into volatile _chromium hexafluoride_, CrF_{6}. This
compound, obtained by Unverdorben, may be prepared by mixing lead
chromate with fluor spar in a dry state, and treating the mixture
with fuming sulphuric acid in a platinum vessel: PbCrO_{4} +
3CaF_{2} + 4H_{2}SO_{4} = PbSO_{4} + 3CaSO_{4} + 4H_{2}O + CrF_{6}.
Fuming sulphuric acid is taken, and in considerable excess, because
the chromium fluoride which is formed is very easily decomposed by
water. It is volatile, and forms a very caustic, poisonous vapour,
which condenses when cooled in a dry platinum vessel into a red,
exceedingly volatile liquid, which fumes powerfully in air. The
vapours of this substance when introduced into water are decomposed
into hydrofluoric acid and chromic anhydride: CrF_{6} + 3H_{2}O =
CrO_{3} + 6HF. If very little water be taken the hydrofluoric acid
volatilises, and chromic anhydride separates directly in crystals.
The chloranhydride of chromic acid, CrO_{2}Cl_{2} (Note 5), is also
decomposed in the same manner. A solution of chromic acid and a
precipitate of barium sulphate are formed by treating the insoluble
barium chromate with an equivalent quantity of sulphuric acid. If
carefully evaporated, the solution yields crystals of chromic
anhydride. Fritzsche gave a very convenient method of preparing
chromic anhydride, based on the relation of chromic to sulphuric
acid. At the ordinary temperature the strong acid dissolves both
chromic anhydride and potassium chromate, but if a certain amount
of water is added to the solution the chromic anhydride separates,
and if the amount of water be increased the precipitated chromic
anhydride is again dissolved. The chromic anhydride is almost all
separated from the solution when it contains two equivalents of
water to one equivalent of sulphuric acid. Many methods for the
preparation of chromic anhydride are based on this fact.
[4 bis] They cannot be filtered through paper or washed, because the
chromic anhydride is reduced by the filter-paper, and is dissolved
during the process of washing.
[5] Berzelius observed, and Rose carefully investigated, this
remarkable reaction, which occurs between chromic acid and sodium
chloride in the presence of sulphuric acid. If 10 parts of common
salt be mixed with 12 parts of potassium dichromate, fused, cooled,
and broken up into lumps, and placed in a retort with 20 parts of
fuming sulphuric acid, it gives rise to a violent reaction,
accompanied by the formation of brown fumes of _chromic
chloranhydride_, or _chromyl chloride_, CrO_{2}Cl_{2}, according to
the reaction: CrO_{3} + 2NaCl + H_{2}SO_{4} = Na_{2}SO_{4} + H_{2}O
+ CrO_{2}Cl_{2}. The addition of an excess of sulphuric acid is
necessary in order to retain the water. The same substance is
always formed when a metallic chloride is heated with chromic acid,
or any of its salts, in the presence of sulphuric acid. The
formation of this volatile substance is easily observed from the
brown colour which is proper to its vapour. On condensing the
vapour in a dry receiver a liquid is obtained having a sp. gr. of
1·9, boiling at 118°, and giving a vapour whose density, compared
with hydrogen, is 78, which corresponds with the above formula.
Chromyl chloride is decomposed by heat into chromic oxide, oxygen,
and chlorine: 2CrO_{2}Cl_{2} = Cr_{2}O_{3} + 2Cl_{2} + O; so that
it is able to act simultaneously as a powerful oxidising and
chlorinating agent, which is taken advantage of in the
investigation of many, and especially of organic, substances. When
treated with water, this substance first falls to the bottom, and
is then decomposed into hydrochloric and chromic acids, like all
chloranhydrides: CrO_{2}Cl_{2} + H_{2}O = CrO_{3} + 2HCl. When
brought into contact with inflammable substances it sets fire to
them; it acts thus, for instance, on phosphorus, sulphur, oil of
turpentine, ammonia, hydrogen, and other substances. It attracts
moisture from the atmosphere with great energy, and must therefore
be kept in closed vessels. It dissolves iodine and chlorine, and
even forms a solid compound with the latter, which depends upon the
faculty of chromium to form its higher oxide, Cr_{2}O_{7}. The
close analogy in the physical properties of the chloranhydrides,
CrO_{2}Cl_{2} and SO_{2}Cl_{2}, is very remarkable, although
sulphurous anhydride is a gas, and the corresponding oxide,
CrO_{2}, is a non-volatile solid. It may be imagined, therefore,
that chromium dioxide (which will be mentioned in the following
note) presents a polymerised modification of the substance having
the composition CrO_{2}; in fact, this is obvious from the method
of its formation.