Chapter XX: Sulphur, Selenium, and Tellurium (2)

_Copper_ (black) | _Mercury_ (black)
_Silver_ (black) | _Lead_ (black)
_Cadmium_ (yellow) |

_Metals which are precipitated by ammonium sulphide_ from neutral
solutions, but not precipitated from acid solutions by
sulphuretted hydrogen:

The sulphide precipitated is soluble in hydrochloric acid:

_Zinc_ (white) | _Manganese_ (rose colour) | _Iron_ (black)

The sulphide precipitated is not soluble in dilute hydrochloric
acid:

_Nickel_ (black) | _Cobalt_ (black)

A hydroxide, and not a sulphide, is precipitated:

_Chromium_ (green) | _Aluminium_ (white)

The metals of the alkalis and of the alkaline earths are not
precipitated either by sulphuretted hydrogen or ammonium sulphide.
The metals of the alkaline earths when in acid solutions in the
form of phosphates and many other salts are precipitated by
ammonium sulphide, because the latter neutralises the free acid,
with formation of an ammonium salt of the acid and evolution of
sulphuretted hydrogen.

Metallic sulphides may be obtained by many other means besides the action of sulphuretted hydrogen on salts and oxides, or by the simple combination of metals with sulphur when heated or fused. Thus they may also be formed by the reduction of sulphates by heating them with charcoal or other means. Charcoal takes up the oxygen from many sulphates, leaving corresponding sulphides. Thus sodium sulphate, Na_{2}SO_{4}, when heated with charcoal, forms sodium sulphide, Na_{2}S. Besides which metallic sulphides are also obtained by heating metals or their oxides in the vapours of many sulphur compounds--for example, in the vapour of carbon bisulphide, CS_{2}, when the carbon takes up the oxygen and the sulphur combines with the metal. The sulphides formed in this manner are often crystalline, and often appear with those properties and in that crystalline form in which they occur in nature. Besides which we must mention that many of the sulphides of the metals are oxidised in air at the ordinary, and especially at a higher, temperature, forming either SO_{2} and the oxide of the metal or sulphates. This oxidation proceeds with particular ease, even at the ordinary temperature, when a metallic sulphide is precipitated from its solutions, as a fine powder containing water. The sulphides of iron and manganese, &c., are very easily oxidised in this manner. But if these hydrates be ignited, they lose their water (the ignition must be carried on in a stream of hydrogen to prevent their oxidation during the process), become denser, and are no longer oxidised at the ordinary temperature. Those sulphides whose corresponding sulphates are decomposed by heat part with their sulphur in the form of sulphurous anhydride when they are ignited in air, and the metal, as a rule, remains behind as oxide. This is taken advantage of in the treatment of sulphurous ores. The process is called _roasting_.

Hydrogen not only forms sulphuretted hydrogen with sulphur, but it also combines with it in several other proportions, just as it combines with oxygen, forming not only water but also hydrogen peroxide. Moreover these _polysulphides of hydrogen_ are also unstable, like hydrogen peroxide, and are also obtained from the corresponding polysulphides of the metals of the alkaline earths, just as hydrogen peroxide is obtained from barium peroxide. Thus calcium forms not only calcium sulphide, CaS, but also as bi-, tri-, and pentasulphide, CaS_{5}, and all these compounds are soluble in water. Sodium also combines with sulphur in the same proportions, forming sulphides from Na_{2}S to Na_{2}S_{5}. If an acid be added to a solution of a polysulphide, it gives sulphur, sulphuretted hydrogen, and a salt of the metal. For instance, MS_{5}, + 2HCl = MCl_{2} + H_{2}S + 4S. If we reverse the operation, and pour a solution of a polysulphide into an acid, sulphur is not precipitated, but an oily liquid is formed which is heavier than water and insoluble in it. This is the polysulphide of hydrogen: MS_{5} + 2HCl = MCl_{2} + H_{2}S_{5}. As Rebs showed (1888), whatever polysulphide be taken--of sodium, for instance--it always gives one and the same _hydrogen pentasulphide_,[24] of specific gravity 1·71 (15°). It can only be preserved in the absence of water and at low temperatures, and then not for long: for, especially in the presence of alkalis and when slightly warmed, it splits up very easily into sulphuretted hydrogen and sulphur.[25]

[24] Rebs took di-, tri-, tetra-, and pentasulphides of sodium,
potassium, and barium, which he prepared by dissolving sulphur in
solutions of the normal sulphides; on adding hydrochloric acid he
always obtained hydrogen pentasulphide, whence it is evident that
4H_{2}S_{n} = (_n_ - 1)H_{2}S_{5} + (5 - _n_)H_{2}S. For example,
if H_{2}S_{2} were formed, it would decompose according to the
equation 4H_{2}S_{2} = H_{2}S_{5} + 3H_{2}S. The hydrogen
pentasulphide formed breaks up into hydrogen sulphide and sulphur
when brought into contact with water. Previous to Rebs' researches
many chemists stated that all polysulphides gave the bisulphide
H_{2}S_{2}, and Hofmann recognised only hydrogen trisulphide,
H_{2}S_{3}.

[25] The formation of the polysulphides of hydrogen, H_{2}S_{n} is
easily understood from the law of substitution, like that of the
saturated hydrocarbons, C_{n}H_{2n + 2}, knowing that sulphur
gives H_{2}S, because the molecule of sulphuretted hydrogen may be
divided into H and HS. This radicle, HS, is equivalent to H. By
substituting this radicle for hydrogen in H_{2}S we obtain (HS)HS
= H_{2}S_{2}, (HS)(HS)S = H_{2}S_{3}, &c., in general H_{2}S_{n}.
The homologues of CH_{4}, C_{n}H_{2n + 2} are formed in this
manner from CH_{4}, and consequently the polysulphides H_{2}S_{n}
are the homologues of H_{2}S. The question arises why in
H_{2}S_{n} the apparent limit of _n_ is 5--that is, why does the
substitution end with the formation of H_{2}S_{5}? The answer
appears to me to be clearly because in the molecule of sulphur,
S_{6}, there are six atoms of sulphur (Note 11). The forces in one
and the other case are the same. In the one case they hold S_{6}
together, in the other S_{5} and H_{2}; and, judging from H_{2}S,
the two atoms of hydrogen are equal in power and significance to
the atom of sulphur. Just as hydrogen peroxide, H_{2}O_{2},
expresses the composition of ozone, O_{3}, in which O is replaced
by H_{2}, so also H_{2}S_{5} corresponds with S_{6}.

The soluble sulphides and polysulphides of the metals of the alkalis and alkaline earths--for example, of ammonium,[26] potassium,[27] and calcium,[28]--have the appearance and properties of salts, just as the hydrated oxides have, whilst the sulphides of the metals of the higher groups resemble their oxides and have not at all the appearance of salts, and this is more especially the case with regard to the crystalline forms in which they frequently occur in nature.[29]

[26] _Ammonium sulphide_, (NH_{4})_{2}S, may be prepared by passing
sulphuretted hydrogen into a vessel full of dry ammonia, or by
passing both dry gases together into a very cold receiver. In the
latter case it is necessary to prevent the access of air, and to
have an excess of ammonia. Under these circumstances, two volumes
of ammonia combine with one volume of sulphuretted hydrogen, and
form a colourless, very volatile, crystalline substance, having a
very unpleasant odour, which is very poisonous and exceedingly
unstable. When exposed to the air it absorbs oxygen and acquires a
yellow colour, and then contains oxygen and polysulphide compounds
(because a portion of the hydrogen sulphide gives water and
sulphur). It is soluble in water and forms a colourless solution,
which, however, in all probability contains free ammonia and the
acid salt--that is, ammonium hydrosulphide, NH_{4}HS, or
(NH_{4})_{2},S,H_{2}S. This salt is formed when dry ammonia is
mixed with an excess of dry sulphuretted hydrogen. The compound
contains equal volumes of the components NH_{3} + H_{2}S =
(NH_{4})HS. It crystallises in an anhydrous state in colourless
plates, and may be easily volatilised (dissociating like ammonium
chloride), even at the ordinary temperature; it has an alkaline
reaction, absorbs oxygen from the air, is soluble in water, and
its solution is usually prepared by saturating an aqueous solution
of ammonia with sulphuretted hydrogen. According to the ordinary
rule, these salts, like other ammonium salts, split up into
ammonia and sulphuretted hydrogen when they are distilled.

A solution of ammonium sulphide is able to dissolve sulphur, and
it then contains compounds of hydrogen polysulphide and ammonia.
Some of these compounds may be obtained in a crystalline form.
Thus Fritzsche obtained a compound of ammonia with hydrogen
pentasulphide, or ammonium pentasulphide, (NH_{4})_{2}S_{5}, in
the following manner: He saturated an aqueous solution of ammonia
with sulphuretted hydrogen, added powdered sulphur to it, and
passed ammonia gas into the solution, which then absorbed a fresh
amount. After this he again passed sulphuretted hydrogen into the
solution, and then added sulphur, and then again ammonia. After
repeating this several times, orange-yellow crystals of
(NH_{4})_{2}S_{5} separated out from the liquid. These crystals
melted at 40° to 50°, and were very unstable.

When a solution of ammonium hydrosulphide, prepared by saturating
a solution of ammonia with sulphuretted hydrogen, is exposed to
the air, it turns yellow, owing to the presence of an ammonium
polysulphide, whose formation is due to the sulphuretted hydrogen
being oxidised by the air and converted into water and sulphur,
which is dissolved by the ammonium sulphide. In certain analytical
reactions it is usual to employ a solution of ammonium sulphide
which has been kept for some time and acquired a yellow colour.
This yellow sulphide of ammonium deposits sulphur when saturated
with acids, whilst a freshly-prepared solution only evolves
sulphuretted hydrogen. The yellow solution furthermore contains
ammonium thiosulphate, which is derived not only from the
oxidation of the ammonium sulphide, but also from the action of
the liberated sulphur on the ammonia, just as an alkaline salt of
thiosulphuric acid and a sulphide are formed by the action of
sulphur on a solution of a caustic alkali.

[27] _Potassium sulphide_, K_{2}S, is obtained by heating a mixture of
potassium sulphate and charcoal to a bright-red heat. It may be
prepared in solution by taking a solution of potassium hydroxide,
dividing it into two equal parts, and saturating one portion with
sulphuretted hydrogen so long as it is absorbed. This portion will
then contain the acid salt KHS (Note 21). The two portions are
then mixed together, and potassium sulphide will then be obtained
in the solution. This solution has a strongly alkaline reaction,
and is colourless when freshly prepared, but it very easily
undergoes change when exposed to the air, forming potassium
thiosulphate and polysulphides. When the solution is evaporated at
low temperatures under the receiver of an air-pump, it yields
crystals containing K_{2}S,5H_{2}O (heated at 150°, they part with
3 mol. H_{2}O, and at higher temperatures they lose nearly all
their water without evolving sulphuretted hydrogen). When they are
ignited in glass vessels they corrode the glass. When a solution
of caustic potash, completely saturated with sulphuretted
hydrogen, is evaporated under the receiver of an air-pump it forms
colourless rhombohedra of _potassium hydrosulphide_,
2(KHS),H_{2}O,K_{2}S,H_{2}S,H_{2}O. These crystals are
deliquescent in the air, but do not change in a vacuum when heated
up to 170°, and at higher temperatures they lose water but do not
evolve sulphuretted hydrogen. The anhydrous compound, KHS, fuses
at a dark-red heat into a very mobile yellow liquid, which
gradually becomes darker in colour and solidifies to a red mass.
It is remarkable that when a solution of the compound KHS is
boiled it somewhat easily evolves half its sulphuretted hydrogen,
leaving potassium sulphide, K_{2}S, in solution; and a solution of
the latter in water is also able to evolve sulphuretted hydrogen
on prolonged boiling, but the evolution cannot be rendered
complete, and, therefore, at a certain temperature, a solution of
potassium sulphide will not be capable of absorbing sulphuretted
hydrogen at all. From this we must conclude that potassium
hydroxide, water, and sulphuretted hydrogen form a system whose
complex equilibrium is subject to the laws of dissociation,
depends on the relative mass of each substance, on the
temperature, and the dissociation pressure of the component
elements. Potassium sulphide is not only soluble in water, but
also in alcohol.

Berzelius showed that in addition to potassium sulphide there also
exist potassium bisulphide, K_{2}S_{2}; trisulphide, K_{2}S_{3};
tetrasulphide, K_{2}S_{4}; and pentasulphide, K_{2}S_{5}.
According to the researches of Schöne, the last three are the most
stable. These different compounds of potassium and sulphur may be
prepared by fusing potassium hydroxide or carbonate with an excess
of sulphur in a porcelain crucible in a stream of carbonic
anhydride. At about 600° potassium pentasulphide is formed; this
is the highest sulphur compound of potassium. When heated to 800°
it loses one-fifth of its sulphur and gives the tetrasulphide,
which at this temperature is stable. At a bright-red heat--namely,
at about 900°--the trisulphide is formed. This compound may be
also formed by igniting potassium carbonate in a stream of carbon
bisulphide, in which case a compound, K_{2}CS_{3}, is first formed
corresponding to potassium carbonate, and carbonic anhydride is
evolved. On further ignition this compound splits up into carbon
and potassium trisulphide, K_{2}S_{3}. The tetrasulphide may also
be obtained in solution if a solution of potassium sulphide be
boiled with the requisite amount of sulphur without access of air.
This solution yields red crystals of the composition
K_{2}S_{4},2H_{2}O when it is evaporated in a vacuum. These
crystals are very hygroscopic, easily soluble in water, but very
sparingly in alcohol; when ignited they give off water,
sulphuretted hydrogen, and sulphur. If a solution of potassium
sulphide be boiled with an excess of sulphur it forms the
pentasulphide, which, however, is decomposed on prolonged boiling
into sulphuretted hydrogen and potassium thiosulphate: K_{2}S_{5}
+ 3H_{2}O = K_{2}S_{2}O_{3} + 3H_{2}S. A substance called _liver
of sulphur_ was formerly frequently used in chemistry and
medicine. Under this name is known the substance which is formed
by boiling a solution of caustic potash with an excess of flowers
of sulphur. This solution contains a mixture of potassium
pentasulphide and thiosulphate, 6KHO + 12S = 2K_{2}S_{5} +
K_{2}S_{2}O_{3} + 3H_{2}O. The substance obtained by fusing
potassium carbonate with an excess of sulphur was also known as
liver of sulphur. If this mixture be heated to an incipient
dark-red heat it will contain potassium thiosulphate, but at
higher temperatures potassium sulphate is formed. In either case a
polysulphide of potassium is also present. The sulphides of
sodium, for example Na_{2}S, NaHS, &c., in many respects closely
resemble the corresponding potassium compounds.

[28] The metals of the alkaline earths, like those of the alkalis, form
several compounds with sulphur; thus calcium forms compounds with
one and with five atoms of sulphur. There are doubtless also
intermediate sulphides. If sulphuretted hydrogen be passed over
ignited lime it forms water and _calcium sulphide_, which may also
be formed by heating calcium sulphate with charcoal, whilst if
sulphur be heated with lime or with calcium carbonate, then
naturally oxygen compounds (calcium thiosulphate and sulphate) are
formed at the same time as calcium sulphide. The prolonged action
of the vapour of carbon bisulphide, especially when mixed with
carbonic anhydride, on strongly ignited calcium carbonate entirely
converts it into sulphide. Calcium sulphide is generally obtained
as an almost colourless, opaque, brittle mass, which is infusible
at a white heat, and is soluble in water. The act of solution (as
with K_{2}S, Note 21) is partly accompanied by a double
decomposition with the water. When heated, dry calcium sulphide
does not absorb oxygen from the air. An excess of water decomposes
it, like many other metallic sulphides, precipitating lime (as a
product of the decomposition the lime hinders the action of the
water upon the CaS; see soda refuse, Chapter XII., Note 12), and
forming a hydrosulphide, CaH_{2}S_{2}, in solution. This compound
is also formed by passing sulphuretted hydrogen through an aqueous
solution of calcium sulphide or lime. Its solution, like that of
calcium sulphide, has an alkaline reaction. It decomposes when
evaporated, and absorbs oxygen from the air. _Calcium
pentasulphide_, CaS_{5}, is not known in a pure state, but may be
obtained in admixture with calcium thiosulphate by boiling a
solution of lime or calcium sulphide with sulphur: 3CaH_{2}O_{2} +
12S = 2CaS_{5} + CaS_{2}O_{3} + 3H_{2}O. A similar compound in an
impure form is formed by the action of air on alkali waste, and is
used for the preparation of thiosulphates.

Many of the sulphides of the metals of the alkaline earths are
phosphorescent--that is, they have the faculty of _emitting
light_, after having been subjected to the action of sunlight, or
of any bright source of light (Canton phosphorus, &c.). The
luminosity lasts some time, but it is not permanent, and gradually
disappears. This phosphorescent property is inherent, in a greater
or less degree, to nearly all substances (Becquerel), but for a
very short time, whilst with calcium sulphide it is comparatively
durable, lasting for several hours, and Dewar (1894) showed that
it is far more intense at very low temperatures (for instance, in
bodies cooled in liquid oxygen to -182°). It is due to the
excitation of the surfaces of substances by the action of light,
and is determined by those rays which exhibit a chemical action.
Hence daylight or the light of burning magnesium, &c., acts more
powerfully than the light of a lamp, &c. Warnerke has shown that a
small quantity of magnesium lighted near the surface of a
phosphorescent substance rapidly excites the greatest possible
intensity of luminosity; this enabled him to found a method of
measuring the intensity of light--_i.e._ to obtain a constant unit
of light--and to apply it to photography. The nature of the change
which is accomplished on the surface of the luminous substance is
at present unknown, but in any case it is a renewable one, because
the experiment may be repeated for an infinite number of times and
takes place in a vacuum. The intensity and tint of the light
emitted depend on the method of preparation of the calcium
sulphide, and on the degree of ignition and purity of the calcium
carbonate taken. According to the observations of Becquerel, the
presence of compounds of manganese, bismuth, &c., sodium sulphide
(but not potassium sulphide), &c., although in minute traces, is
perfectly indispensable. This gives reason for thinking that the
formation (in the dark) and decomposition (in light) of double
salts like MnS,Na_{2}S perhaps form the chemical cause of the
phenomena. Compounds of strontium and barium have this property to
even a greater extent than calcium sulphide. These compounds may
be prepared as in the following example: A mixture of sodium
thiosulphate and strontium chloride is prepared; a double
decomposition takes place between the salts, and, on the addition
of alcohol, strontium thiosulphate, SrS_{2}O_{3}, is precipitated,
which, when ignited, leaves strontium sulphide behind. The
strontium sulphide thus prepared emits (when dry) a greenish
yellow light. It contains a certain amount of sulphur, sodium
sulphide, and strontium sulphate. By ignition at various
temperatures, and by different methods of preparation, it is
possible to obtain mixtures which emit different coloured lights.

[29] As examples, we will describe the sulphides of arsenic, antimony,
and mercury. Arsenic trisulphide, or _orpiment_, As_{2}S_{3},
occurs native, and is obtained pure when a solution of arsenious
anhydride in the presence of hydrochloric acid comes into contact
with sulphuretted hydrogen (there is no precipitate in the absence
of free acid). A beautiful yellow precipitate is then obtained:
As_{2}O_{3} + 3H_{2}S = 3H_{2}O + As_{2}S_{3}; it fuses when
heated, and volatilises without decomposition. As_{2}S_{3} is
easily obtained in a colloid form (Chapter I., Note 57). When
fused it forms a semi-transparent, yellow mass, and it is thus
that it enters the market. The specific gravity of native orpiment
is 3·4, and that of the artificially-fused mass is 2·7. It is used
as a yellow pigment, and owing to its insolubility in water and
acids it is less injurious than the other compounds corresponding
to arsenious acid. According to the type AsX_{2}, realgar, AsS, is
known, but it is probable that the true composition of this
compound is As_{4}S_{4}--that is, it presents the same relation to
orpiment as liquid phosphuretted hydrogen does to gaseous.
_Realgar_ (_Sandaraca_) occurs native as brilliant red crystals of
specific gravity 3·59, and may be prepared artificially by fusing
arsenic and sulphur in the proportions indicated by its formulæ.
It is prepared in large quantities by distilling a mixture of
sulphur and arsenical pyrites. Like orpiment it dissolves in
calcium sulphide, and even in caustic potash. It is used for
signal lights and fireworks, because it deflagrates and gives a
large and very brilliant white flame with nitre.

With antimony, sulphur gives a tri- and a pentasulphide. The
former, Sb_{2}S_{3}, which corresponds with antimonious oxide,
occurs native (Chapter XIX.) in a crystalline form; its sp. gr. is
then 4·9, and it presents brilliant rhombic crystals of a grey
colour, which fuse when heated. A substance of the same
composition is obtained as an amorphous orange powder by passing
sulphuretted hydrogen into an acid solution of antimonious oxide.
In this respect antimonious oxide again reacts like arsenious
acid, and the sulphides of both are soluble in ammonium and
potassium sulphides, and, especially in the case of arsenious
sulphide, are easily obtained in colloidal solutions. By prolonged
boiling with water, antimonious sulphide may be entirely converted
into the oxide, hydrogen sulphide being evolved (Elbers). Native
antimony sulphide, or the orange precipitated trisulphide when
fused with dry, or boiled with dissolved, alkalis, forms a
dark-coloured mass (Kermes mineral) formerly much used in
medicine, which contains a mixture of antimonious sulphide and
oxide. There are also compounds of these substances. A so-called
antimony vermilion is much used as a dye; it is prepared by
boiling sodium thiosulphate (six parts) with antimony trichloride
(five parts) and water (fifty parts). This substance probably
contains an oxysulphide of antimony--that is, a portion of the
oxygen in the oxide of antimony in it is replaced by sulphur. Red
antimony ore, and antimony glass, which is obtained by fusing the
trisulphide with antimonious oxide, have a similar composition,
Sb_{2}OS_{2}. In the arts, the _antimony pentasulphide_,
Sb_{2}S_{5}, is the most frequently used of the sulphur compounds
of antimony. It is formed by the action of acids on the so-called
Schlippe's salt, which is a _sodium thiorthantimonate_,
SbS(NaS)_{3}, corresponding with (Chapter XIX., Note 41 bis)
orthantimonic acid, SbO(OH)_{3}, with the replacement of oxygen by
sulphur. It is obtained by boiling finely-powdered native antimony
trisulphide with twice its weight of sodium carbonate, and half
its weight of sulphur and lime, in the presence of a considerable
quantity of water. The processes taking place are as follows:--The
sodium carbonate is converted into hydroxide by the lime, and then
forms sodium sulphide with the sulphur; the sodium sulphide then
dissolves the antimony sulphide, which in this form already
combines with the greatest amount of sulphur, so that a compound
is formed corresponding with antimony pentasulphide dissolved in
sodium sulphide. The solution is filtered and crystallised, care
being taken to prevent access of air, which oxidises the sodium
sulphide. This salt crystallises in large, yellowish crystals,
which are easily soluble in water and have the composition
Na_{3}SbS_{4},9H_{2}O. When heated they lose their water of
crystallisation and then fuse without alteration; but when in
solution, and even in crystalline form, this salt turns brown in
air, owing to the oxidation of the sulphur and the breaking up of
the compound. As it is used in medicine, especially in the
preparation of antimony pentasulphide, it is kept under a layer of
alcohol, in which it is insoluble. Acids precipitate antimony
pentasulphide from a solution of this salt, as an orange powder,
insoluble in acids and very frequently used in medicine (_sulfur
auratum antimonii_). This substance when heated evolves vapours of
sulphur, and leaves antimony trisulphide behind.

Mercury forms compounds with sulphur of the same types as it does
with oxygen. Mercurous sulphide, Hg_{2}S, easily splits up into
mercury and mercuric sulphide. It is obtained by the action of
potassium sulphide on mercurous chloride, and also by the action
of sulphuretted hydrogen on solutions of salts of the type HgX.
Mercuric sulphide, HgS, corresponding with the oxide, is cinnabar;
it is obtained as a black precipitate by the action of an excess
of sulphuretted hydrogen on solutions of mercuric salts. It is
insoluble in acids, and is therefore precipitated in their
presence. If a certain amount of water containing sulphuretted
hydrogen be added to a solution of mercuric chloride, it first
gives a white precipitate of the composition
Hg_{3}S_{2}Cl_{2}--that is, a compound HgCl,2HgS, a sulphochloride
of mercury like the oxychloride. But in the presence of an excess
of sulphuretted hydrogen, the black precipitate of mercuric
sulphide is formed. In this state it is not crystalline (the red
variety is formed by the prolonged action of polysulphides of
ammonium upon the black HgS), but if it be heated to its
temperature of volatilisation it forms a red crystalline sublimate
which is identical with native cinnabar. In this form its specific
gravity is 8·0, and it forms a red powder, owing to which it is
used as a red pigment (vermilion) in oil, pastel, and other
paints. It is so little attacked by reagents that even nitric acid
has no action on it, and the gastric juices do not dissolve it, so
that it is not poisonous. When heated in air, the sulphur burns
away and leaves metallic mercury. On a large scale cinnabar is
usually prepared in the following manner: 300 parts of mercury and
115 parts of sulphur are mixed together as intimately as possible
and poured into a solution of 75 parts of caustic potash in 425
parts of water, and the mixture is heated at 50° for several
hours. Red mercury sulphide is thus formed, and separates out from
the solution. The reaction which takes place is as follows: A
soluble compound, K_{2}HgS_{2}, is first formed; this compound is
able to separate in colourless silky needles, which are soluble in
the caustic potash, but are decomposed by water, and at 50°; this
solution (perhaps by attracting oxygen from the air) slowly
deposits HgS in a crystalline form.

Spring conducted an interesting research (at Liège, 1894) upon the
conversion of the black amorphous sulphide of mercury, HgS, into
red crystalline cinnabar. This research formed a sequel to
Spring's classical researches on the influence of high pressures
upon the properties of solids and their capacity for mutual
combination. He showed, among other things, that ordinary solids
and even metals (for instance, Pb), after being considerably
compressed under a pressure of 20,000 atmospheres, return on
removal of the pressure to their original density like gases. But
this is only true when the compressed solid is not liable to an
allotropic variation, and does not give a denser variety. Thus
prismatic sulphur (sp. gr. 1·9) passes under pressure into the
octahedral (sp. gr. 2·05) variety. Black HgS (precipitated from
solution) has a sp. gr. 7·6, while that of the red variety is 8·2,
and therefore it might be expected that the former would pass into
the latter under pressure, but experiments both at the ordinary
and a higher temperature did not give the looked-for result,
because even at a pressure of 20,000 atmospheres the black
sulphide was not compressed to the density of cinnabar (a pressure
of as much as 35,000 atmospheres was necessary, which could not be
attained in the experiment). But Spring prepared a black HgS,
which had a sp. gr. of 8·0, and this, under a pressure of 2,500
atmospheres, passed into cinnabar. He obtained this peculiar black
variety of HgS (sp. gr. 8·0) by distilling cinnabar in an
atmosphere of CO_{2}, when the greater portion of the HgS is
redeposited in the form of cinnabar. Under the action of a
solution of polysulphide of ammonium, this variety of HgS passes
more slowly into the red variety than the precipitated variety
does, while under pressure the conversion is comparatively easy.

It is worthy of remark, that Linder and Picton obtained complex
compounds of many of the sulphides of the heavy metals (Ca, Hg,
Sb, Zn, Cd, Ag, Au) with H_{2}S, for example H_{2}S,7CuS (by the
action of H_{2}S upon the hydrate of oxide of copper), H_{2}S,9CuS
(in the presence of acetic acid and with an excess of H_{2}S), &c.
Probably we have here a sort of 'solid' solution of H_{2}S in the
metallic sulphides.

As the acids derived from chlorine, phosphorus, and carbon are the oxidised hydrogen compounds of these elements, so also we can form an idea of the acid hydrates of sulphur, or of _the normal acids of sulphur_, by representing them as the oxidised products of sulphuretted hydrogen--

HCl H_{2}S H_{3}P H_{4}C
HClO H_{2}SO(?) H_{3}PO(?) H_{4}CO
HClO_{2} H_{2}SO_{2}(?) H_{3}PO_{2} H_{4}CO_{2}
HClO_{3} H_{2}SO_{3} H_{3}PO_{3} H_{4}CO_{3}
HClO_{4} H_{2}SO_{4} H_{3}PO_{4} H_{4}CO_{4}[30]

In the case of chlorine, if not all the hydrates, at all events salts of all the normal hydrates are known, whilst in the case of sulphur only the acids H_{2}S, H_{2}SO_{3} and H_{2}SO_{4} are known. But, on the other hand, the latter are obtained not only as hydrates but also as stable anhydrides, SO_{2} and SO_{3}, which are formed with the evolution of heat from sulphur and oxygen; 32 parts of sulphur in combining with 32 parts of oxygen--that is, in forming SO_{2}--evolve 71,000 heat units,[31] and if the oxidation proceeds to the formation of SO_{3}, 103,000 heat units are evolved. These figures may be compared with those which correspond with the passage of carbon into CO and CO_{2}, when 29,000 and 97,000 units of heat are evolved. This determines the stability of the higher oxides of sulphur, and also expresses the peculiarity of sulphur as an element which, although an analogue of oxygen, forms stable compounds with it, and thus fundamentally differs from chlorine. The higher and lower oxides of chlorine are powerful oxidising agents, whilst the higher oxide of sulphur, SO_{3}, has but feeble oxidising powers, and the lower oxide, SO_{2}, frequently acts as a reducing agent, and is formed by the direct combustion of sulphur, just as carbonic anhydride, CO_{2}, proceeds from the combustion of carbon. In the combustion of sulphur, and also in the oxidation (roasting) of the sulphides and polysulphides by their ignition in air, _sulphurous oxide_, or _sulphurous anhydride_, or _sulphur dioxide_, SO_{2},[31 bis] is exclusively formed. It is prepared on a large scale by burning sulphur or roasting iron pyrites or other sulphides[32] for the manufacture of sulphuric acid (Chapter VI.), and for direct application in the manufacture of wine or for bleaching tissues and other purposes. In the latter instances its application is based on the fact that sulphurous anhydride acts on certain vegetable matters, and has the property of a reducing and feeble acid.[32 bis]

[30] CH_{4} gives CH_{4}O or CH_{3}(OH), wood spirit; CH_{4}O_{2} or
CH_{2}(OH)_{2}, which decomposes into water and CH_{2}O--that is,
methylene oxide or formaldehyde; CH_{4}O_{3} = CH(OH)_{3} = H_{2}O
+ CHO(OH), or formic acid; and CH_{4}O_{4} = C(OH)_{4} = 2H_{2}O +
CO_{2}. There are four typical hydrogen compounds, RH, RH_{2},
RH_{3}, and RH_{4}, and each of them has its typical oxide. Beyond
H_{4} and O_{4} combination does not proceed.

[31] Rhombic sulphur, 71,080 heat units; monoclinic sulphur, 71,720
units, according to Thomsen.

[31 bis] However, when sulphur or metallic sulphides burn in an excess
of air, there is always formed a certain, although small, amount
of SO_{3}, which gives sulphuric acid with the moisture of the
air.

[32] The enormous amount of sulphuric acid now manufactured is chiefly
prepared by roasting native pyrites, but a considerable amount of
the SO_{2} for this purpose is obtained by roasting zinc blende
(ZnS) and copper and lead sulphides. A certain amount is also
procured from soda refuse (Note 6) and the residues obtained from
the purification of coal gas.

[32 bis] Sulphurous anhydride is also obtained by the decomposition of
many sulphates, especially of the heavy metals, by the action of
heat; but this requires a very powerful heat. This formation of
sulphurous anhydride from sulphates is based on the decomposition
proper to sulphuric acid itself. When sulphuric acid is strongly
heated (for instance, by dropping it upon an incandescent surface)
it is decomposed into water, oxygen, and sulphurous
anhydride--that is, into those compounds from which it is formed.
A similar decomposition proceeds during the ignition of many
sulphates. Even so stable a sulphate as gypsum does not resist the
action of very high temperatures, but is decomposed in the same
manner, lime being left behind. The decomposition of sulphates by
heat is accomplished with still greater facility in the presence
of sulphur, because in this case the liberated oxygen combines
with the sulphur and the metal is able to form a sulphide. Thus
when ferrous sulphate (green vitriol) is ignited with sulphur, it
gives ferrous sulphide and sulphurous anhydride: FeSO_{4} + 2S =
FeS + 2SO_{2}, and this reaction may even be used for the
preparation of this gas. At 400° sulphuric acid and sulphur give
an extremely uniform stream of pure sulphurous anhydride, so that
it is best prepared on a manufacturing scale by this method. Iron
pyrites, FeS_{2}, when heated to 150° with sulphuric acid (sp. gr.
1·75) in cast-iron vessels also gives an abundant and uniform
supply of sulphurous anhydride.

In the laboratory--that is, on a small scale--sulphurous anhydride is best prepared by deoxidising sulphuric acid by heating it with charcoal, or copper, sulphur, mercury, &c. Charcoal produces this decomposition of sulphuric acid at but moderately high temperatures; it is itself converted into carbonic anhydride,[32 tri] and therefore when sulphuric acid is heated with charcoal it evolves a mixture of sulphurous and carbonic anhydrides: C + 2H_{2}SO_{4} = CO_{2} + 2SO_{2} + 2H_{2}O. The metals which are unable to decompose water, and which do not, therefore, expel hydrogen from sulphuric acid, are frequently capable of decomposing sulphuric acid, with the evolution of sulphurous anhydride, just as they decompose nitric acid, forming the lower oxides of nitrogen. These metals are silver, mercury, copper, lead, and others. Thus, for example, the action of copper on sulphuric acid may be expressed by the following equation: Cu + 2H_{2}SO_{4} = CuSO_{4} + SO_{2} + 2H_{2}O. In the laboratory this reaction is carried on in a flask with a gas-conducting tube, and does not take place unless aided by heat.[33]

[32 tri] Mellitic acid is formed at the same time (Verneuille).

[33] The thermochemical data connected with this reaction are as
follows: A molecule of hydrogen H_{2}, in combining with oxygen (O
= 16) develops about 69,000 heat units, whilst the molecule of
SO_{2}, in combining with oxygen only develops about 32,000 heat
units--that is, about half as much--and therefore those metals
which cannot decompose water may still be able to deoxidise
sulphuric into sulphurous acid. Those metals which decompose water
and sulphuric acid with the evolution of hydrogen, evolve in
combining with sixteen parts by weight of oxygen more heat than
hydrogen does--for example, K_{2}, Na_{2}, Ca develop about or
more than 100,000 heat units; Fe, Zn, Mn about 70,000 to 80,000
heat units; whilst those metals which neither decompose water nor
evolve hydrogen from sulphuric acid, but are still capable of
evolving sulphurous anhydride from it, develop less heat with
oxygen than hydrogen, but nearly the same amount, if not more
than, sulphurous anhydride develops--for example, Cu and Hg
develop about 40,000 and Pb about 50,000 heat units.

In its physical and chemical properties sulphurous anhydride presents a great _resemblance to carbonic anhydride_. It is a heavy gas, somewhat considerably soluble in water, very easily condensed into a liquid; it forms normal and acid salts, does not evolve oxygen under the direct action of heat,[34] although such metals as sodium and magnesium burn in it, just as in carbonic anhydride. It has a suffocating odour, which is well known owing to its being evolved when sulphur or sulphur matches are burnt. In characterising the properties of sulphurous anhydride, it is very important to remember (Chapter II.) also that it is more easily liquefied (at -10°, or at 0° under two atmospheres pressure) than carbonic anhydride (thirty-six atmospheres at 0°),[35] that it is more soluble than carbonic anhydride (Vol. I. p. 79); at 0°, 100 vols. of water dissolve 180 vols. of carbonic anhydride and 688 vols. of sulphuric anhydride), that the molecular weight of SO_{2} = 64 and of CO_{2} = 44, and that the density of liquid sulphurous anhydride at 0° = 1·43 (molecular volume = 45) and of carbonic anhydride = 0·95 (molecular volume = 49). Although sulphur dioxide is the anhydride of an acid, nevertheless, like carbonic anhydride, it does not form any stable compounds with water, but gives a solution from which it may be entirely expelled by the action of heat.[36] The acid character of sulphurous anhydride is clearly expressed by the fact that it is entirely absorbed by alkalis, with which it forms acid and normal salts easily soluble in water. With salts of barium, calcium, and the heavy metals, the normal salts of the alkalis, M_{2}SO_{3}, give precipitates exactly like those formed by the carbonates. In general, the salts of sulphurous acid are closely analogous to the corresponding carbonates.

[34] That is, it only dissociates and re-forms the original product on
cooling.

[35] At a given temperature the pressure of this gas evolved from any
salt will be less than that of carbonic anhydride, if we compare
the separation of a gas from its salts with the phenomenon of
evaporation, as was done in discussing the decomposition of
calcium carbonate.

Liquid sulphurous anhydride is used on a large scale (Pictet) for
the production of cold.

[36] De la Rive, Pierre, and more especially Roozeboom, have
investigated the crystallo-hydrate which is formed by sulphurous
anhydride and water at temperatures below 7° under the ordinary
pressure, and in closed vessels (at temperatures below 12°). Its
composition is SO_{2},7H_{2}O, and density 1·2. This hydrate
corresponds with the similar hydrate CO_{2},8H_{2}O obtained by
Wroblewsky.

_Acid sodium sulphite_, NaHSO_{3}, may be obtained by passing sulphurous anhydride into a solution of sodium hydroxide. It is also formed by saturating a solution of sodium carbonate with the gas (carbonic anhydride is then given off), and as the solubility of the acid sulphite is much greater than that of the carbonate, a further quantity of the latter may be dissolved after the passage of the sulphurous anhydride, so that ultimately a very strong solution of the sulphite may be formed in this manner, from which it may be obtained in a crystalline form, either by cooling and evaporating (without heating, for then the salt would give off sulphurous anhydride) or by adding alcohol to the solution. When exposed to the air this salt loses sulphurous anhydride and attracts oxygen, which converts it into sodium sulphate. The acid sulphites of the alkali metals are able to combine not only with oxygen, but also with many other substances--for example, a solution of the sodium salt dissolves sulphur, forming sodium thiosulphate, gives crystalline compounds with the aldehydes and ketones, and dissolves many bases, converting them into double sulphites. Having the faculty of attracting or absorbing oxygen, acid sodium sulphite is also able to absorb chlorine, and is therefore employed, like sodium thiosulphate, for the removal of chloride (as an antichlor), especially in the bleaching of fabrics, when it is necessary to remove the last traces of the chlorine held in the tissues, which might otherwise have an injurious effect on them. If a solution of an alkali hydroxide be divided into two parts, and one half is saturated with sulphurous anhydride, and then the other half added to it, a normal salt will be obtained in the solution, having an alkaline reaction, like a solution of sodium carbonate. The acid salt has a neutral reaction.[36 bis] Like sodium carbonate, _normal sodium sulphite_ has the composition Na_{2}SO_{3},10H_{2}O, and its maximum solubility is at 33°--in a word, it very closely resembles sodium carbonate. Although this salt does not give off sulphurous anhydride from its solution, it is able, like the acid salt, to absorb oxygen from the air, and is then converted into sodium sulphate.[37]

[36 bis] Schwicker (1889) by saturating NaHSO_{3} with potash, or
KHSO_{3} with soda, obtained NaKSO_{3}, in the first instance with
H_{2}O, and in the second instance with 2H_{2}O, probably owing to
the different media in which the crystals are formed. In general
sulphurous acid easily forms double salts.

[37] The normal salts of calcium and magnesium are slightly, and the
acid salts easily, soluble in water. These acid sulphites are much
used in practice; thus calcium bisulphite is employed in the
manufacture of cellulose from sawdust, for mixing with fibrous
matter in the manufacture of paper.

Besides the acid character we must also point out the reducing character of sulphurous anhydride. The reducing action of sulphurous acid, its anhydride and salts, is due to their faculty of passing into sulphuric acid and sulphates. The reducing action of the sulphites is particularly energetic, so that they even convert nitric oxide into nitrous oxide: K_{2}SO_{3} + 2NO = K_{2}SO_{4} + N_{2}O. The salts of many of the higher oxides are converted into those of the lower--for example, FeX_{3} into FeX_{2}, CuX_{2} into CuX, HgX_{2} into HgX; thus 2FeX_{3} + SO_{2} + 2H_{2}O = 2FeX_{2} + H_{2}SO_{4} + 2HX. In the presence of water, sulphurous anhydride is oxidised by chlorine (SO_{2} + 2H_{2}O + Cl_{2} = H_{2}SO_{4} + 2HCl), iodine, nitrous acid, hydrogen peroxide, hypochlorous acid, chloric acid, and other oxygen compounds of the halogens, chromic, manganic, and many other metallic acids and higher oxides, as well as all peroxides. Free oxygen in the presence of spongy platinum is able to oxidise sulphurous anhydride even in the absence of water, in which case sulphuric anhydride SO_{3} is formed, so that the latter may be prepared by passing a mixture of sulphurous anhydride and oxygen over incandescent spongy platinum, or, as it is now prepared on a large scale in chemical works, by passing this mixture over asbestos or pumice stone moistened with a solution of platinum salt and ignited. Sulphurous anhydride is completely absorbed by certain higher oxides--for instance, by barium peroxide and lead dioxide (PbO_{2} + SO_{2} = PbSO_{4}).[38]

[38] This reaction is taken advantage of in removing sulphurous
anhydride from a mixture of gases. Lead dioxide, PbO_{2}, is
brown, and when combined with sulphurous anhydride it forms lead
sulphate, PbSO_{4}, which is white, so that the reaction is
evident both from the change in colour and development of heat.
Sulphurous anhydride is slowly decomposed by the action of light,
with the separation of sulphur and formation of sulphuric
anhydride. This explains the fact that sulphurous anhydride
prepared in the dark gives a white precipitate of silver sulphite,
Ag_{2}SO_{3}, with silver chlorate, AgClO_{4}, but when prepared
in the light, even in diffused light, it gives a dark precipitate.
This naturally depends on the fact that the sulphur liberated then
forms silver sulphide, which is black.

There are, however, cases where sulphurous anhydride acts as an oxidising agent--that is, it is _deoxidised_ in the presence of substances which are capable of absorbing oxygen with still greater energy than the sulphurous anhydride itself. This oxidising action proceeds with the formation of sulphuretted hydrogen or of sulphides, while the reducing agent is oxidised at the expense of the oxygen of the sulphurous anhydride. In this respect, the action of stannous salts is particularly remarkable. Stannous chloride, SnCl_{2}, in an aqueous solution gives a precipitate of stannic sulphide, SnS_{2}, with sulphurous anhydride--that is, the latter is deoxidised to sulphuretted hydrogen, while SnX_{2} is oxidised into SnX_{4}. A solution of sulphurous anhydride has also an oxidising action on zinc. The zinc passes into solution, but no hydrogen is evolved,[39] because a salt of _hydrosulphurous acid_, ZnS_{2}O_{4}, is formed. The free acid is still less stable than the salt.

[39] Schönebein observed that the liquid turns yellow, and acquires the
faculty of decolorising litmus and indigo. Schützenberger showed
that this depends on the formation of a zinc salt of a peculiar
and very powerfully-reducing acid, for with cupric salts the
yellow solution gives a red precipitate of cuprous hydrate or
metallic copper, and it reduces salts of silver and mercury
entirely. An exactly similar solution is obtained by the action of
zinc on sodium bisulphite without access of air and in the cold.
The yellow liquid absorbs oxygen from the air with great avidity,
and forms a sulphate. If the solution be mixed with alcohol, it
deposits a double sulphite of zinc and sodium,
ZnNa_{2}(SO_{3})_{2}, which does not decolorise litmus or indigo.
The remaining alcoholic solution deposits colourless crystals in
the cold, which absorb oxygen with great energy in the presence of
water, but are tolerably stable when dried under the receiver of
an air-pump. The solution of these crystals has the
above-mentioned decolorising and reducing properties. These
crystals contain a sodium salt of a lower acid; their composition
was at first supposed to be HNaSO_{2}, but it was afterwards
proved that they do not contain hydrogen, and present the
composition Na_{2}S_{2}O_{4} (Bernthsen). The same salt is formed
by the action of a galvanic current on a solution of sodium
bisulphite, owing to the action of the hydrogen at the moment of
its liberation. If SO_{2} resembles CO_{2} in its composition,
then hyposulphurous acid H_{2}S_{2}O_{4} resembles oxalic acid
H_{2}C_{2}O_{4}. Perhaps an analogue of formic acid SH_{2}O_{2}
will be discovered.

The faculty of sulphurous anhydride of combining with various substances is evident from the above-cited reactions, where it combines with hydrogen and with oxygen, and this faculty also appears in the fact that, like carbonic oxide, it combines with chlorine, forming a chloranhydride of sulphuric acid, SO_{2}Cl_{2}, to which we shall afterwards return. The same faculty for combination also appears in the salts of sulphurous acid, in their liability to oxidation and in the exceedingly characteristic formation of a peculiar series of salts obtained by Pelouze and Frémy. At a temperature of -10° or below, nitric oxide NO is absorbed by alkaline solutions of the alkali sulphites, forming a peculiar series of _nitrosulphates_. At a higher temperature these salts are not formed but the nitric oxide is reduced to nitrous oxide. But in the cold the liquid saturated with nitric oxide after a certain time gives prismatic crystals resembling those of nitre. The composition of the potassium salt is K_{2}SN_{2}O_{3}--that is, the salt contains the elements of potassium sulphite and of nitric oxide.[40]

[40] The instability of this salt is very great, and may be compared to
that of the compound of ferrous sulphate with nitric oxide, for
when heated under the contact influence of spongy platinum,
charcoal, &c., it splits up into potassium sulphate and nitrous
oxide. At 130° the dry salt gives off nitric oxide, and re-forms
potassium sulphite. The free acid has not yet been obtained. These
salts resemble the series of _sulphonitrites_ discovered by Frémy
in 1845. They are obtained by passing sulphurous anhydride through
a concentrated and strongly alkaline aqueous solution of potassium
nitrite. They are soluble in water, but are precipitated by an
excess of alkali. The first product of the action has the
composition K_{3}NS_{3}HO_{9}. It is then converted by the further
action of sulphurous anhydride, cold water, and other reagents
into a series of similar complex salts, many of which give
well-formed crystals. One must suppose that the chief cause of the
formation of these very complex compounds is that they contain
unsaturated compounds, NO, KNO_{2}, and KHSO_{3}, all of which are
subject to oxidation and further combination, and therefore easily
combine among each other. The decomposition of these compounds,
with the evolution of ammonia, when their solutions are heated is
due to the fact that the molecule contains the deoxidant,
sulphurous anhydride, which reduces the nitrous acid, NO(OH), to
ammonia. In my opinion the composition of the sulphonitrites may
be very simply referred to the composition of ammonia, in which
the hydrogen is partly replaced by the radicle of the sulphates.
If we represent the composition of potassium sulphate as
KO.KSO_{3}, the group KSO_{3} will be equivalent (according to the
law of substitution) to HO and to hydrogen. It combines with
hydrogen, forming the potassium acid sulphite, KHSO_{3}. Hence the
group KSO_{3} may also replace the hydrogen in ammonia. Judging by
my analysis (1870) the extreme limit of this substitution,
N(HSO_{3})_{3}, agrees with that of the sulphonitrite, which is
easily formed, simultaneously with alkali, by the action of
potassium sulphite on potassium nitrite, according to the equation
3K(KSO_{3}) + KNO_{2} + 2H_{2}O = N(KSO_{3})_{3} + 4HKO. The
researches of Berglund, and especially of Raschig (1887), fully
verified my conclusions, and showed that we must distinguish the
following types of salts, corresponding with ammonia, where X
stands for the sulphonic group, HSO_{3}, in which the hydrogen is
replaced by potassium; hence X = KSO_{3}: (1) NH_{2}X, (2)
NHX_{2}, (3) NH_{3}, (4) N(OH)XH, (5) N(OH)X_{2}, (6) N(OH)_{2}X,
just as NH_{2}(OH) is hydroxylamine, NH(OH)_{2}, is the hydrate of
nitrous oxide, and N(OH)_{3} is orthonitrous acid, as follows from
the law of substitution. This class of compounds is in most
intimate relation with the series of sulphonitrous compounds,
corresponding with 'chamber crystals' and their acids, which we
shall consider later.

There are also several other substances, formed by the oxides of nitrogen and sulphur, which belong to this class of complex and, under some circumstances, unstable compounds. In the manufacture of sulphuric acid, both these classes of oxides come into contact with each other in the lead chambers, and if there be insufficient water for the formation of sulphuric acid they give crystalline compounds, termed _chamber crystals_. As a rule, the composition of the crystals is expressed by the formula NHSO_{3}. This is a compound of the radicles NO_{2} of nitric acid, and HSO_{3} of sulphuric acid, or nitro-sulphuric acid, NO_{2}.SHO_{3}, if sulphuric acid be expressed as OH.SHO_{3} and nitric by NO_{2}.OH. The tabular crystals of this substance fuse at about 70°, are formed both by the direct action of nitrous anhydride or nitric peroxide (but not NO, which is not absorbed by sulphuric acid) on sulphuric acid (Weltzien and others), and especially on sulphuric acid containing an anhydride and the lower oxides of sulphur and nitric acid.[41]

[41] In the sulphuric acid chambers the lower oxides of nitrogen and
sulphur take part in the reaction. They are oxidised by the oxygen
of the air, and form nitro-sulphuric acid--for example, 2SO_{2} +
N_{2}O_{3} + O_{2} + H_{2}O = 2NHSO_{5}. This compound dissolves
in strong sulphuric acid without changing, and when this solution
is diluted (when the sp. gr. falls to 1·5), it splits up into
sulphuric acid and nitrous anhydride, and by the action of
sulphurous anhydride is converted into nitric oxide, which by
itself (in the absence of nitric acid or oxygen) is insoluble in
sulphuric acid. These reactions are taken advantage of in
retaining the oxides of nitrogen in the Gay-Lussac coke-towers,
and for extracting the absorbed oxides of nitrogen from the
resultant solution in the Glover tower. Although nitric oxide is
not absorbed by sulphuric acid, it reacts (Rose, Brüning) on its
anhydride, and forms sulphurous anhydride and a crystalline
substance, N_{2}S_{2}O_{9} = 2NO + 3SO_{3} - SO_{2} =
N_{2}O_{3}2SO_{3}. This may be regarded as the anhydride of
nitro-sulphuric acid, because N_{2}S_{2}O_{9} = 2NHSO_{5} -
H_{2}O; like nitro-sulphuric acid, it is decomposed by water into
nitro-sulphuric acid and nitrous anhydride. Since boric and
arsenious anhydrides, alumina and other oxides of the form
R_{2}O_{3} are able to combine with sulphuric anhydride to form
similar compounds decomposable by water, the above compound does
not present any exceptional phenomenon. The substance NOClSO_{3}
obtained by Weber by the action of nitrosyl chloride upon
sulphuric anhydride belongs to this class of compounds.

_Thiosulphuric acid_, H_{2}S_{2}O_{3}--that is, a compound of sulphurous acid and sulphur--also belongs to the products of combination of sulphurous acid. In the same way that sulphurous acid, H_{2}SO_{3}, gives H_{2}SO_{4} with oxygen, so it gives H_{2}S_{2}O_{3} with sulphur. In a free state it is very unstable, and it is only known in the form of its salts proceeding from the direct action of sulphur on the normal sulphites; if endeavours be made to separate it in a free state, it immediately splits up into those elements from which it might be formed--that is, into sulphur and sulphurous acid. The most important of its salts is the _sodium thiosulphate_ (known as hyposulphite), Na_{2}S_{2}O_{3},5H_{2}O, which occurs in colourless crystals, and is unacted on by atmospheric oxygen either when in a dry state or in solution. Many other salts of this acid are easily formed by means of this salt,[41 bis] although this cannot be done with all bases, for such bases as alumina, ferric oxide, chromium oxide, and others do not give compounds with thiosulphuric acid, just as they do not form stable compounds with carbonic acid. Whenever these salts might be formed, they (like the acid) split up into sulphurous acid and sulphur, and furthermore the elements of thiosulphuric acid in many cases act in a reducing manner, forming sulphuric acid and taking up the oxygen from reducible oxides. Thus when treated with a thiosulphate the soluble ferric salts give a precipitate of sulphur and form ferrous salts. The thiosulphates of the metals of the alkalis are obtained directly by boiling a solution of their sulphites with sulphur: Na_{2}SO_{3} + S = Na_{2}S_{2}O_{3}. The same salts are formed by the action of sulphurous anhydride on solutions of the sulphides; thus sodium sulphide dissolved in water gives sulphur and sodium thiosulphate when a stream of sulphurous anhydride is passed through it: 2Na_{2}S + 3SO_{2} = 2Na_{2}S_{2}O_{3} + S. The polysulphides of the alkali metals when left exposed to the air attract oxygen and also form thiosulphates.[42]

[41 bis] Many double salts of thiosulphuric acid are known, for
instance, PbS_{2}O_{3},3Na_{2}S_{2}O_{3},12H_{2}O;
CaS_{2}O_{3},3K_{2}S_{2}O_{3},5H_{2}O, &c. (Fortman, Schwicker,
Fock, and others).