Chapter XX: Sulphur, Selenium, and Tellurium (4)

[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.

[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.