water; one volume at the ordinary temperature dissolves as much as
eight volumes of the gas. The solutions in water and alcohol
undergo change, especially in open vessels, owing to the fact that
the water and alcohol dissolve oxygen from the atmosphere, which,
acting on the sulphuretted hydrogen, forms water and sulphur. The
solution may be so altered in this manner that every trace of
sulphuretted hydrogen disappears. Solutions of sulphuretted
hydrogen in glycerine change much more slowly, and may therefore
be kept for a long time as reagents. De Forcrand obtained a
hydrate, H_{2}S,16H_{2}O, resembling the hydrates given by many
gases.
Sulphuretted hydrogen is very easily _decomposed_ into its component
parts by the action of heat or a series of electric sparks. Hence it is
not surprising that sulphuretted hydrogen undergoes change under the
action of many substances having a considerable affinity for hydrogen and
oxygen. Very many metals[17] evolve hydrogen with sulphuretted hydrogen,
so that in this respect it presents the property of an acid; for
instance, 2H_{2}S + Sn = 2H_{2} + SnS_{2}. This may be taken advantage of
for determining the composition of sulphuretted hydrogen, because a given
volume then leaves the same volume of hydrogen. On the other hand,
oxygen,[18] chlorine,[19] and even iodine decompose sulphuretted
hydrogen, removing the hydrogen from it and leaving free sulphur, so that
in this reaction the sulphur is replaced by the above-named elements; for
example, H_{2}S + Br_{2} = 2HBr + S. In no other hydrogen compound is it
so easy to show the _substitution_, both of hydrogen and of the element
combined with it, as in hydrogen sulphide. This clearly proves the feeble
union between the elements forming this gas. Compounds containing a
considerable amount of oxygen, with which they easily part, can
accomplish the separation of the sulphur very easily. Such are, for
instance, nitrous acid, chromic acid, and even ferric oxide and the
higher oxides like it. Thus, if sulphuretted hydrogen be passed into a
solution of chromic acid or an acid solution of ferric oxide, water is
formed, _and the sulphur is separated in a free state_. Thus,
sulphuretted hydrogen acts as a _reducing agent_, in virtue of the
hydrogen it contains. Salts of iodic, chlorous, chloric, and other acids
are reduced by sulphuretted hydrogen, their oxygen acting mainly on its
hydrogen; but in the presence of an excess of a powerful oxidising agent
a portion of the sulphur may also be oxidised to sulphurous anhydride.
The reducing action of sulphuretted hydrogen is frequently applied in
chemical manipulations for the preparation of lower oxides, and for the
conversion of certain oxygen compounds into hydrogen compounds: thus, the
higher oxides of nitrogen are converted into ammonia by it, and in the
presence of alkalis the nitro-compounds are converted into ammonia
derivatives. The reaction of sulphuretted hydrogen on sulphurous
anhydride belongs to this class of phenomena, the chief products of which
are sulphur and water, 2H_{2}S + SO_{2} = 2H_{2}O + S_{3}.
[17] Some metals evolve hydrogen from sulphuretted hydrogen at the
ordinary temperature. For example, the light metals, and copper
and silver (especially with the access of air?) among the heavy
metals. Hence articles made of silver turn black in the presence
of vapours containing sulphuretted hydrogen, because silver
sulphide is black. Zinc and cadmium act at a red heat, but not
completely.
[18] If sulphuretted hydrogen escapes from a fine orifice into the air,
it will burn when lighted, and be transformed into sulphurous
anhydride and water. But if it burns in a limited supply of
air--for instance, when a cylinder is filled with it and
lighted--then only the hydrogen burns, which has, judging from the
amount of heat developed in its combustion and from all its
properties, a greater affinity for oxygen than sulphur. In this
respect the combustion of sulphuretted hydrogen resembles that of
hydrocarbons.
[19] Hence bleaching powder and chlorine destroy the disagreeable smell
of sulphuretted hydrogen. (For the reaction of hydrogen sulphide
and iodine, _see_ Chapter XI. p. 504.)
The acid character of sulphuretted hydrogen is clearly seen in its
action on alkalis and salts.[19 bis] Thus lead oxide and its salts in the
presence of sulphuretted hydrogen form water or an acid, and sulphide of
lead: PbX_{2} + H_{2}S = PbS + 2HX. This reaction takes place even in the
presence of powerful acids, because lead sulphide is one of those
sulphides which are unacted on by acids, and in solutions the reaction is
a complete one. This reaction is taken advantage of for the preparation
of many acids, by first converting into a lead salt, and then submitting
this salt to the action of sulphuretted hydrogen. For example, lead
formate with sulphuretted hydrogen gives formic acid. Sulphuretted
hydrogen in acting on a number of metallic acid substances in solution or
in an anhydrous state also forms corresponding sulphates: (1) if it does
not reduce the acid; (2) if the sulphur compound corresponding with the
anhydride of the acid be insoluble in water, the reaction proceeds in
solutions; (3) if the sulphuretted hydrogen and the acid taken do not
come in contact with an alkali, on which they would be able to act first;
and (4) if the sulphur compound be not decomposed by water. Thus
solutions of arsenious acid give a precipitate of arsenious sulphide,
As_{2}S_{3}, with sulphuretted hydrogen. This reaction proceeds not only
in the presence of water, but also of acids, because the latter do not
decompose the resultant sulphur compounds. The type of the decomposition
is the same as with bases--that is, the sulphur and oxygen change places:
RO_{_n_} + _n_H_{2}S = RS_{_n_} + _n_H_{2}O. Some sulphides corresponding
with acid anhydrides are decomposed by water, and therefore are not
formed in the presence of water. Such, for example, are the sulphides of
phosphorus.[20]
[19 bis] Perfectly dry H_{2}S (Hughes 1892) has no action upon
perfectly dry salts, just as dry HCl does not react with dry
NH_{3} or metals (Chapter IX., Note 29).
[20] The sulphide P_{4}S is obtained by cautiously fusing the requisite
proportions of common phosphorus and sulphur under water; it is a
liquid which solidifies at 0°, and may be distilled without
undergoing change, but it fumes in air and easily takes fire. The
higher sulphide, P_{2}S, has similar properties. But little heat
is evolved in the formation of these compounds, and it may be
supposed that they are formed by the direct conjunction of whole
molecules of phosphorus and sulphur; but if the proportion of
sulphur be increased, the reaction is accompanied by so
considerable a rise of temperature that an explosion takes place,
and for the sake of safety red phosphorus must be used, mixed as
intimately as possible with powdered sulphur and heated in an
atmosphere of carbonic anhydride. The higher compounds are
decomposed by water. By increasing the proportion of sulphur, the
following compounds have been obtained: P_{4}S_{3} as prisms
(fuses at 165°, Rebs), soluble in carbon bisulphide, and unaltered
by air and water; _phosphorus trisulphide_, P_{2}S_{3}, is the
analogue of P_{2}O_{3}; it is a light yellow crystalline compound
only slightly soluble in carbon bisulphide, fusible and volatile,
decomposed into hydrogen sulphide and phosphorous acid by water,
and, like the highest compound of sulphur and phosphorus,
P_{2}S_{5}, it forms thio-salts with potassium sulphide, &c. This
_phosphorus pentasulphide_ corresponds with phosphoric anhydride;
like the trisulphide it gives hydrogen sulphide and phosphoric
acid with an excess of water. It reacts in many respects like
phosphoric chloride. The sulphide PS_{2} is also known; the vapour
density of this compound seems to indicate a molecule P_{3}S_{6}.
_Phosphorus sulphochloride_, PSCl_{3}, corresponds with phosphorus
oxychloride. It is a colourless, pleasant-smelling liquid, boiling
at 124°, and of sp. gr. 1·63; it fumes in air and is decomposed by
water: PSCl_{3} + 4H_{2}O = PH_{3}O_{4} + H_{2}S + 3HCl. It is
obtained when phosphoric chloride is treated with hydrogen
sulphide, hydrochloric acid being also formed; it is also produced
by the action of phosphoric chloride on certain sulphides--for
example, on antimonious sulphide, also by the (cautious) action of
phosphorus on sulphur chloride: 2P + 3S_{2}Cl_{2} = 2PSCl_{3} +
4S, by the action of PCl_{5} upon certain sulphides, for example,
Sb_{2}S_{3}, by the reaction: 3MCl + P_{2}S_{5} = PSCl_{3} +
M_{3}PS_{4} (Glatzel, 1893), and in the reaction 3PCl_{3} +
SOCl_{2} = PCl_{5} + POCl_{3} + PSCl_{3}, showing the reducing
action of phosphorus trichloride, which is especially clear in the
reaction SO_{3} + PCl_{3} = SO_{2} + POCl_{3}. Thorpe and Rodger
(1889), by heating 3PbF_{2} or BiF_{3} with phosphorus
pentasulphide (and also by heating AsF_{3} and PSCl_{3} to 150°),
obtained thiophosphoryl fluoride as a colourless, spontaneously
inflammable gas (see further on, Note 74 bis, and Chapter XIX.,
Note 25). The action of PSCl_{3} upon NaHO gives a salt of
monothiophosphoric acid (Würtz, Kubierschky), H_{3}PSO_{3}, which
gives soluble salts of the alkalis.
The metallic sulphides corresponding with the metallic oxides have
either a feeble alkaline or a feeble acid character, according to the
character of the corresponding oxide, and therefore by combining together
they are able to form saline substances--that is, salts in which the
oxygen is replaced by sulphur. Thus sulphuretted hydrogen having the
properties of a feeble acid[21] has, at the same time, the properties of
water, and forms the type of the sulphur derivatives, which may also be
formed by means of sulphuretted hydrogen, just as the oxides may be
formed by the aid of water. But as sulphuretted hydrogen has acid
properties, it combines more easily with the basic metallic sulphides.
Hence, for instance, there exists a compound of sulphuretted hydrogen
with potassium sulphide, potassium hydrosulphide, 2KHS = K_{2}S + H_{2}S,
just as there are potassium hydroxides; but there are scarcely any
compounds of sulphuretted hydrogen with the sulphides corresponding with
acids. Thus the sulphides of the metals may be regarded either as salts
of sulphuretted hydrogen or as oxides of the metals in which the oxygen
is replaced by sulphur. In general terms the sulphides exhibit the same
degrees of difference with respect to their solubility in water as do the
oxides. Thus the oxides of the alkali metals, and of some of the metals
of the alkaline earths, are soluble in water, whilst those of nearly all
the other metals are insoluble. The same may be said as to the sulphides;
the sulphides of the metals of the alkalis and certain of the alkaline
earths are soluble in water, whilst those of the other metals are
insoluble. Those metals, like aluminium, whose oxides--for example,
Al_{2}O_{3}--have intermediate properties and do not form compounds with
feeble acids, at least in a wet way, also do not form sulphides by this
method, although these may be obtained indirectly. And in general the
sulphides of the metals are easily formed in a wet way, and with
particular ease if they are insoluble in water. In this case their salts
enter into double decomposition with sulphuretted hydrogen, or with
soluble sulphides, and give an insoluble sulphide--for instance, a salt
of lead gives lead sulphide with sulphuretted hydrogen. By the action of
sulphuretted hydrogen on a salt of a metal, a free acid must be formed
besides the metallic sulphide. Thus if a metal M be in a state of
combination MX_{2}, then by the action of sulphuretted hydrogen there
will be formed, besides MS,[22] an acid 2HX. It is evident that
sulphuretted hydrogen will not precipitate an insoluble sulphide from the
salts of those metals whose sulphides react with free acid, such as zinc,
iron, manganese, &c. The reaction FeCl_{2} + H_{2}S = FeS + 2HCl, and the
like, do not take place because the acid acts on the ferrous sulphide.
Antimonious sulphide is not acted on by dilute hydrochloric acid, but it
is decomposed by strong acid, and therefore in presence of an excess of
hydrochloric acid antimonious chloride does not entirely react with
hydrogen sulphide, whilst the reaction 2SbCl_{3} + 3H_{2}S = Sb_{2}S_{3}
+ 6HCl is a complete one in a dilute solution and with a small quantity
of acid. Those metallic sulphides which are decomposed by acids may be
obtained in a wet way by the double decomposition of the salts of the
metals, not with hydrogen sulphide, but with soluble metallic sulphides,
such as sulphide of ammonium or of potassium, because then no free acid
is formed, but a salt of the metal (potassium or ammonium) which was
taken as a soluble sulphide. So, for example, FeCl_{2} + K_{2}S = FeS +
2KCl.[23]
[21] Sulphuretted hydrogen does not saturate the alkaline properties of
alkali hydroxides, so that a solution of potassium hydroxide will
not under any circumstances give a neutral liquid with
sulphuretted hydrogen. In this case the sulphuretted hydrogen
forms in solution only an acid salt with the potassium: KHO +
H_{2}S = KHS + H_{2}O. It must be supposed that the normal salt is
not formed in the solution--that is, that the reaction 2KHO +
H_{2}S = K_{2}S + 2H_{2}O does not take place. This is seen from
the fact that a development of heat, depending on the formation of
potassium hydrosulphide, KHS, is remarked when as much hydrogen
sulphide is passed into a solution of potassium hydroxide as it
will absorb. But if a further quantity of potassium hydroxide be
added to the resultant solution, heat is not developed, whilst if
alkali be added to potassium acid sulphate or sodium acid
carbonate, heat is developed. It must not be concluded from this
that H_{2}S is a monobasic acid, for here there is a question of
the decomposing action of water upon K_{2}S; K_{2}S and H_{2}O in
reacting on each other should absorb heat if the reaction of KHS
upon KHO evolves heat. Furthermore, it must be taken into account
that potassium oxide, K_{2}O, and the anhydrous oxides like it,
also do not exist in solutions, for whenever they are formed they
immediately react with the water, forming caustic potash, KHO, &c.
In the same way, directly potassium sulphide, K_{2}S, is formed in
water it is decomposed into potassium hydroxide and hydrosulphide:
K_{2}S + H_{2}O = KHO + KHS. Potassium sulphide, K_{2}S, in a
solid state corresponds with K_{2}O, although neither can exist in
solution.
[22] During recent years (beginning with Schulze, 1882) it has been
found that many metallic sulphides which were considered totally
insoluble do, under certain circumstances, form very unstable
solutions in water, as already mentioned in Chapter I., Note 57.
Arsenic sulphide is very easily obtained in the form of a solution
(hydrosol). Solutions of copper and cadmium sulphides may also be
easily obtained by precipitating their salts CuX_{2}, or CdX_{2},
with ammonium sulphide, and washing the precipitate; but they are
re-precipitated by the addition of foreign salts.
[23] In reality the preceding reaction should be expressed thus:
FeCl_{2} + 2KHS = FeS + 2KCl + H_{2}S (Note 21), because in the
presence of water not K_{2}S but KHS reacts. But as the
sulphuretted hydrogen takes no part in the reaction, it is usual
to express the formation of such sulphides without taking the
hydrogen sulphide proceeding from the potassium or ammonium
hydrosulphides into account. It is not usual to employ potassium
sulphide but ammonium sulphide--or, to speak more accurately,
ammonium hydrosulphide--in order to avoid the formation of a
non-volatile salt of potassium and to have, together with the
formation of the sulphide, a salt of ammonium which can always be
driven off by evaporating the solution and igniting the
residue--for instance: FeCl_{2} + (NH_{4})_{2}S = FeS + 2NH_{4}Cl.
Thus the metallic sulphides may be divided into three chief
classes: (1) _those soluble in water_, (2) _those insoluble in
water but reacting with acids_, and (3) _those insoluble both in
water and acids_. The third class may be easily subdivided into
two groups; to the first group belong those sulphides which
correspond with bases or basic oxides, and are therefore unable to
play the part of an acid with the sulphides of the alkalis, and
are insoluble in NH_{4}HS, whilst the sulphides of the second
group are of an acid character, and give soluble thio-salts with
the sulphides of the alkaline metals, in which they play the part
of an acid. To this group belong those metals whose corresponding
oxides have acid properties. It must be observed, however, that
not all metallic acids have corresponding sulphides, partly owing
to the fact that certain acids are reducible by sulphuretted
hydrogen, especially when their lower degrees of oxidation are of
a basic character. Such are, for instance, the acids of chromium,
manganese, &c. Sulphuretted hydrogen converts them into lower
oxides, having the properties of bases. Those bases which do not
combine with feeble acids, such as carbonic acid and hydrogen
sulphide, give a precipitate of hydroxide with ammonium
sulphide--for example, aluminium salts react in this manner. This
difference of the metals in their behaviour towards sulphuretted
hydrogen gives a very valuable means of separating them from each
other, and _is taken advantage of in analytical chemistry_. If,
for instance, the metals of the first and third groups occur
together, it is only necessary to convert them into soluble salts,
and to act on the solution of the salts with sulphuretted
hydrogen; this will precipitate the metals of the third group in
the form of sulphides, whilst the metals of the first group will
not be in the least acted on. Such a method of separating the
metals is considered more fully in analytical chemistry, and we
will therefore limit ourselves here to pointing out to which
groups the most common metals belong, and the colour which is
proper to the sulphide precipitated.
_Metals which are precipitated by sulphuretted hydrogen_, as
sulphides from a solution of their salts, even in the presence of
free acid:
The precipitate is soluble in ammonium sulphide:
_Platinum_ (dark brown) | _Antimony_ (orange)
_Gold_ (dark brown) | _Arsenic_ (yellow)
_Tin_ (yellow and brown) |
The precipitate is insoluble in ammonium sulphide:
_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).
[42] Thus when alkali waste, which contains calcium sulphide, undergoes
oxidation in the air it first forms a calcium polysulphide, and
then calcium thiosulphate, CaS_{2}O_{3}. When iron or zinc acts on
a solution of sulphurous acid, besides the hyposulphurous acid
first formed, a mixture of sulphite and thiosulphate is obtained
(Note 39), 3SO_{2} + Zn_{2} = ZnSO_{3} + ZnS_{2}O_{3}. In this
case, as in the formation of hyposulphurous acid, there is no
hydrogen liberated. One of the most common methods for preparing
thiosulphates consists in the _action of sulphur on the alkalis_.
The reaction is accomplished by the formation of sulphides and
thiosulphates, just as the reaction of chlorine on alkalis is
accompanied by the formation of hypochlorites and chlorides; hence
in this respect the thiosulphates hold the same position in the
order of the compounds of sulphur as the hypochlorites do among
the chlorine compounds. The reaction of caustic soda on an excess
of sulphur may be expressed thus: 6NaHO + 12S = 2Na_{2}S_{5} +
Na_{2}S_{2}O_{3} + 3H_{2}O. Thus sulphur is soluble in alkalis. On
a large scale sodium thiosulphate, Na_{2}S_{2}O_{3}, is prepared
by first heating sodium sulphate with charcoal, to form sodium
sulphide, which is then dissolved in water and treated with
sulphurous anhydride. The reaction is complete when the solution
has become slightly acid. A certain amount of caustic alkali is
added to the slightly acid solution; a portion of the sulphur is
thus precipitated, and the solution is then boiled and evaporated
when the salt crystallises out. The saturation of the solution of
sodium sulphide by sulphurous anhydride is carried on in different
ways--for example, by means of coke-towers, by causing the
solution of sulphide to trickle over the coke, and the sulphurous
anhydride, obtained by burning sulphur, to pass up the coke-tower
from below. An excess of sulphurous anhydride must be avoided, as
otherwise sodium trithionate is formed. Sodium thiophosphate is
also prepared by the double decomposition of the soluble calcium
thiosulphate with sodium sulphate or carbonate, in which case
calcium sulphate or carbonate is precipitated. The calcium
thiosulphate is prepared by the action of sulphurous anhydride on
either calcium sulphide or alkali waste. A dilute solution of
calcium thiosulphate may be obtained by treating alkali waste
which has been exposed to the action of air with water. On
evaporation, this solution gives crystals of the salt containing
CaS_{2}O_{3},5H_{2}O. A solution of calcium thiosulphate must be
evaporated with great care, because otherwise the salt breaks up
into sulphur and calcium sulphide. Even the crystallised salt
sometimes undergoes this change.
The crystals of sodium thiosulphate are stable, do not effloresce
and at 0° dissolve in one part of water, and at 20° in 0·6 part.
The solution of this salt does not undergo any change when boiled
for a short time, but after prolonged boiling it deposits sulphur.
The crystals fuse at 56°, and lose all their water at 100°. When
the dry salt is ignited it gives sodium sulphide and sulphate.
With acids, a solution of the thiosulphate soon becomes cloudy and
deposits an exceedingly fine powder of sulphur (Note 10). If the
amount of acid added be considerable, it also evolves sulphurous
anhydride: H_{2}S_{2}O_{3} = H_{2}O + S + SO_{2}. Sodium
thiosulphate has many practical uses; it is used in photography
for dissolving silver chloride and bromide. Its solvent action on
silver chloride may be taken advantage of in extracting this metal
as chloride from its ores. In dissolving, it forms a double salt
of silver and sodium: AgCl + Na_{2}S_{2}O_{3} = NaCl +
AgNaS_{2}O_{3}. Sodium thiosulphate is an _antichlor_--that is, a
substance which hinders the destructive action of free chlorine
owing to its being very easily oxidised by chlorine into sulphuric
acid and sodium chloride. The reaction with iodine is different,
and is remarkable for the accuracy with which it proceeds. The
iodine takes up half the sodium from the salt and converts it into
a tetrathionate; 2Na_{2}S_{2}O_{3} + I_{2} = 2NaI +
Na_{2}S_{4}O_{6}, and hence this reaction is employed for the
determination of free iodine. As iodine is expelled from potassium
iodide by chlorine, it is possible also to determine the amount of
chlorine by this method if potassium iodide be added to a solution
containing chlorine. And as many of the higher oxides are able to
evolve iodine from potassium iodide, or chlorine from hydrochloric
acid (for example, the higher oxides of manganese, chromium, &c.),
it is also possible to determine the amounts of these higher
oxides by means of sodium thiosulphate and liberated iodine. This
forms the basis of the iodometric method of volumetric analysis.
The details of these methods will be found in works on analytical
chemistry.
On adding a solution of a _lead salt_ gradually to a solution of
sodium thiosulphate a white precipitate of lead thiosulphate,
PbS_{2}O_{3}, is formed (a soluble double salt is first formed,
and if the action be rapid, lead sulphide). When this substance is
heated at 200°, it undergoes a change and takes fire. Sodium
thiosulphate in solution rapidly reduces cupric salts to cuprous
salts by means of the sulphurous acid contained in the
thiosulphate, but the resultant cuprous oxide is not precipitated,
because it passes into the state of a thiosulphate and forms a
double salt. These double cuprous salts are excellent reducing
agents. The solution when heated gives a black precipitate of
copper sulphide.
The following formulæ sufficiently explain the position held by
thiosulphuric acid among the other acids of sulphur:
Sulphurous acid SO_{2}H(OH)
Sulphuric acid SO_{2}OH(OH)
Thiosulphuric acid SO_{2}SH(OH)
Hyposulphurous acid SO_{2}H(SO_{2}H)
Dithionic acid SO_{2}OH(SO_{2}OH)
At one time it was thought that all the salts of thiosulphuric
acid only existed in combination with water, and it was then
supposed that their composition was H_{4}S_{2}O_{4}, or
H_{2}SO_{2}, but Popp obtained the anhydrous salts.
Although sulphur, oxidising at a high temperature, only forms a small
quantity of sulphuric anhydride, SO_{3}, and nearly all passes into
sulphurous anhydride, still the latter may be converted into the higher
oxide, or _sulphuric anhydride_, SO_{3}, by many methods. Sulphuric
anhydride is a solid crystalline substance at the ordinary temperature;
it is easily fusible (15°), and volatile (46°), and rapidly attracts
moisture. Although it is formed by the combination of sulphurous
anhydride with oxygen, it is capable of further combination. Thus it
combines with water, hydrochloric acid, ammonia, with many hydrocarbons,
and even with sulphuric acid, boric and nitrous anhydrides, &c., and also
with bases which burn directly in its vapour, forming sulphates in the
presence of traces of moisture (_see_ Chapter IX., Note 29). The
oxidation of sulphurous anhydride, SO_{2}, into sulphuric anhydride,
SO_{3}, is effected by passing a mixture of the former and dry oxygen or
air over incandescent spongy platinum. An increase of pressure
accelerates the reaction (Hanisch). If the product be passed into a cold
vessel, crystalline sulphuric anhydride is deposited upon the sides of
the vessel, but as it is difficult to avoid all traces of moisture it
always contains compounds of its hydrates: H_{2}S_{2}O_{7} and
H_{2}S_{4}O_{13}, whose presence so modifies the properties of the
anhydride (Weber) that formerly two modifications of the anhydride were
recognised. The same sulphuric anhydride may be obtained from certain
anhydrous sulphates, or those which are almost so, which are decomposed
by heat, whilst an impure but perfectly anhydrous anhydride is formed by
distillation over phosphoric anhydride. For instance, acid sodium
sulphate, NaHSO_{4}, and the pyro- or di-sulphate, Na_{2}S_{2}O_{7}
(Chapter XII.) formed from it, when ignited evolve sulphuric anhydride.
Green vitriol--that is, ferrous sulphate, FeSO_{4}--belongs to the number
of those sulphates which easily give off sulphuric anhydride under the
action of heat. It contains water of crystallisation and parts with it
when it is heated, but the last equivalent of water is driven off with
difficulty, just as is the case with magnesium sulphate, MgSO_{4}7H_{2}O;
however, when thoroughly heated, this evolution of sulphuric anhydride
does take place, although not completely, because at a high temperature a
portion of it is decomposed by the ferrous oxide (SO_{3} + 2FeO), which
is converted into ferric oxide, Fe_{2}O_{3}, and in consequence part of
the sulphuric anhydride is converted into sulphurous anhydride. Thus the
products of the decomposition of ferrous sulphate will be: ferric oxide,
Fe_{2}O_{3}, sulphurous anhydride, SO_{2}, and sulphuric anhydride,
SO_{3}, according to the equation: 2FeSO_{4} = Fe_{2}O_{3} + SO_{2} +
SO_{3}. As water still remains with the ferrous sulphate when it is
heated, the result will partially consist of the hydrate H_{2}SO_{4},
with anhydride, SO_{3}, dissolved in it. Sulphuric acid was for a long
time prepared in this manner; the process was formerly carried on on a
large scale in the neighbourhood of Nordhausen, and hence the sulphuric
acid prepared from ferrous sulphate is called _fuming Nordhausen acid_.
At the present time the fuming acid is prepared by passing the volatile
products of the decomposition of ferrous sulphate through strong
sulphuric acid prepared by the ordinary method. The sulphurous anhydride
is insoluble in it, but it absorbs the sulphuric anhydride. Sulphuric
anhydride may be prepared not only by igniting FeSO_{4} or sodium
pyrosulphate, Na_{2}S_{2}O_{7} (the decomposition proceeds at 600°), but
also by heating a mixture of the latter and MgSO_{4} (Walters); in the
former case a stable double salt MgNa_{2}(SO_{4})_{2} finally remains. It
is also obtained by the direct combination of SO_{2} and O under the
action of spongy platinum or asbestos coated with platinum black (C.
Winkler's process). Nordhausen sulphuric acid fumes in air, owing to its
containing and easily giving off sulphuric anhydride, and it is therefore
also called _fuming sulphuric acid_; these fumes are nothing but the
vapour of sulphuric anhydride combining with the moisture in the air and
forming non-volatile sulphuric acid (hydrate).[43]
[43] Nordhausen sulphuric acid may serve as a very simple means for the
preparation of sulphuric anhydride. For this purpose the
Nordhausen acid is heated in a glass retort, whose neck is firmly
fixed in the mouth of a well-cooled flask. The access of moisture
is prevented by connecting the receiver with a drying-tube. On
heating the retort the vapours of sulphuric anhydride will pass
over into the receiver, where they condense; the crystals of
anhydride thus prepared will, however, contain traces of sulphuric
acid--that is, of the hydrate. By repeatedly distilling over
phosphoric anhydride, it is possible to obtain the pure anhydride,
SO_{3}, especially if the process be carried on without access of
air in a closed vessel.
The ordinary sulphuric anhydride, which is imperfectly freed from
the hydrate, is a snow-white, exceedingly volatile substance,
which crystallises (generally by sublimation) in long silky
prisms, and only gives the pure anhydride when carefully distilled
over P_{2}O_{5}. Freshly prepared crystals of almost pure
anhydride fuse at 16° into a colourless liquid having a specific
gravity at 26° = 1·91, and at 47° = 1·81; it volatilises at 46°.
After being kept for some time the anhydride, even containing only
small traces of water, undergoes a change of the following nature:
A small quantity of sulphuric acid combines by degrees with a
large proportion of the anhydride, forming polysulphuric acids,
H_{2}SO_{4},_n_SO_{3}, which fuse with difficulty (even at 100°,
Marignac), but decompose when heated. In the entire absence of
water this rise in the fusing point does not occur (Weber), and
then the anhydride long remains liquid, and solidifies at about
+15°, volatilises at 40°, and has a specific gravity 1·94 at 16°.
We may add that Weber (1881), by treating sulphuric anhydride with
sulphur, obtained a blue lower oxide of sulphur, S_{2}O_{3}.
Selenium and tellurium also give similar products with SO_{3},
SeSO_{3}, and TeSO_{3}. Water does not act upon them.
Nordhausen sulphuric acid contains a peculiar compound of SO_{3} and
H_{2}SO_{4}, or _pyrosulphuric acid_; an imperfect anhydride of sulphuric
acid, H_{2}S_{2}O_{7}, analogous in composition with the salts
Na_{2}S_{2}O_{7}, K_{2}Cr_{2}O_{7}, and bearing the same relation to
H_{2}SO_{4} that pyrophosphoric acid does to H_{3}PO_{4}. The bond
holding the sulphuric acid and anhydride together is unstable. This is
obvious from the fact that the anhydride may easily be separated from
this compound, by the action of heat. In order to obtain the definite
compound, the Nordhausen acid is cooled to 5°, or, better still, a
portion of it is distilled until all the anhydride and a certain amount
of sulphuric acid have passed over into the distillate, which will then
solidify at the ordinary temperature, because the compound
H_{2}SO_{4},SO_{3} fuses at 35°. Although this substance reacts on water,
bases, &c., like a mixture of SO_{3} + H_{2}SO_{4}, still since a
definite compound, H_{2}S_{2}O_{7}, exists in a free state and gives
salts and a chloranhydride, S_{2}O_{5}Cl_{2},[44] we must admit the
existence of a definite pyrosulphuric acid, like pyrophosphoric acid,
only that the latter has a far greater stability and is not even
converted into a perfect hydrate by water. Further, the salts
M_{2}S_{2}O_{7} dissolved in water react in the same manner as the acid
salts MHSO_{4}, whilst the imperfect hydrates of phosphoric acid (for
example, PHO_{3}, H_{4}P_{2}O_{7}) have independent reactions even in an
aqueous solution which distinguish them and their salts from the perfect
hydrates.
[44] Pyrosulphuric chloranhydride, or _pyrosulphuryl chloride_,
S_{2}O_{5}Cl_{2}, corresponds to pyrosulphuric acid, in the same
way that sulphuryl chloride, SO_{2}Cl_{2}, corresponds to
sulphuric acid. The composition S_{2}O_{5}Cl_{2} = SO_{2}Cl_{2} +
SO_{3}. It is obtained by the action of the vapour of sulphuric
anhydride on sulphur chloride: S_{2}Cl_{2} + 5SO_{3} = 5SO_{2} +
S_{2}O_{5}Cl_{2}. It is also formed (and not sulphuryl chloride,
SO_{2}Cl_{2}, Michaelis) by the action of phosphorus pentachloride
in excess on sulphuric acid (or its first chloranhydride,
SHO_{3}Cl). It is an oily liquid, boiling at about 150°, and of
sp. gr. 1·8. According to Konovaloff (Chapter VII.), its vapour
density is normal. It should be noticed that the same substance is
obtained by the action of sulphuric anhydride on sulphur
tetrachloride, and also on carbon tetrachloride, and this
substance is the last product of the metalepsis of CH_{4}, and
therefore the comparison of SCl_{2} and S_{2}Cl_{2} with products
of metalepsis (_see_ later) also finds confirmation in particular
reactions. Rose, who obtained pyrosulphuryl chloride,
S_{2}O_{5}Cl_{2}, regarded it as SCl_{6},5SO_{3}, for at that time
an endeavour was always made to find two component parts of
opposite polarity, and this substance was cited as a proof of the
existence of a hexachloride, SCl_{6}. Pyrosulphuryl chloride is
decomposed by cold water, but more slowly than chlorosulphuric
acid and the other chloranhydrides.
The relation between pyrosulphuric acid and the normal acid will
be obvious if we express the latter by the formula OH(SO_{3}H),
because the sulphonic group (SO_{3}H) is then evidently equivalent
to OH, and consequently to H, and if we replace both the hydrogens
in water by this radicle we shall obtain (SO_{3}H)_{2}O--that is,
pyrosulphuric acid.
[Illustration: FIG. 87.--Concentration of sulphuric acid in glass
retorts. The neck of each retort is attached to a bent glass tube, whose
vertical arm is lowered into a glass or earthenware vessel acting as a
receiver for the steam which comes over from the acid, as the former
still contains a certain amount of acid.]
_Sulphuric acid_, H_{2}SO_{4}, is formed by the combination of its
anhydride, SO_{3}, and water, with the evolution of a large amount of
heat; the reaction SO_{3} + H_{2}O develops 21,300 heat units. The method
of its preparation on a large scale, and most of the methods employed for
its formation, are dependent on the oxidation of sulphurous anhydride,
and the formation of sulphuric anhydride, which forms sulphuric acid
under the action of water. The technical method of its manufacture has
been described in Chapter VI. The acid obtained _from the lead chambers_
contains a considerable amount of water, and is also impure owing to the
presence of oxides of nitrogen, lead compounds, and certain impurities
from the burnt sulphur which have come over in a gaseous and vaporous
state (for example, arsenic compounds). For practical purposes, hardly
any notice is taken of the majority of these impurities, because they do
not interfere with its general qualities. Most frequently endeavours are
only made to remove, as far as possible, all the water which can be
expelled.[45] That is, the object is to obtain the hydrate, H_{2}SO_{4},
from the dilute acid (60 per cent.), and this is effected by evaporation
by means of heat. Every given mixture of water and sulphuric acid begins
to part with a certain amount of aqueous vapour when heated to a certain
definite temperature. At a low temperature either there is no evaporation
of water, or there can even be an absorption of moisture from the air. As
the removal of the water proceeds, the vapour tension of the residue
decreases for the same temperature, and therefore the more dilute the
acid the lower the temperature at which it gives up a portion of its
water. In consequence of this, the removal of water from dilute solutions
of sulphuric acid may be easily carried on (up to 75 p.c. H_{2}SO_{4}) in
lead vessels, because at low temperatures dilute sulphuric acid does not
attack lead. But as the acid becomes more concentrated the temperature at
which the water comes over becomes higher and higher, and then the acid
begins to act on lead (with the evolution of sulphuretted hydrogen and
conversion of the lead into sulphate), and therefore lead vessels cannot
be employed for the complete removal of the water. For this purpose the
evaporation is generally carried on in glass or platinum retorts, like
those depicted in figs. 87 and 88.
[45] The removal of the water, or concentration to almost the real
acid, H_{2}SO_{4}, is effected for two reasons: in the first place
to avoid the expense of transit (it is cheaper to remove the water
than to pay for its transit), and in the second place because many
processes--for instance, the refining of petroleum--require a
strong acid free from an excess of water, the weak acid having no
action. When in the manufacture of chamber acid, both the
Gay-Lussac tower (cold, situated at the end of the chambers) and
the Glover tower (hot, situated at the beginning of the plant,
between the chambers and ovens for the production of SO_{2}) are
employed, a mixture of nitrose (_i.e._ the product of the
Gay-Lussac tower) and chamber acid containing about 60 p.c.
H_{2}SO_{4}, is poured into the Glover tower, where under the
action of the hot furnace gases containing SO_{2}, and the water
held in the chamber acid (1) N_{2}O_{3} is evolved from the
nitrose; (2) water is expelled from the chamber acid; (3) a
portion of the SO_{2} is converted into H_{2}SO_{4}; and (4) the
furnace gases are cooled. Thus, amongst other things, the Glover
tower facilitates the concentration of the chamber acid (removal
of H_{2}O), but the product generally contains many impurities.
[Illustration: FIG. 88.--Concentration of sulphuric acid in platinum
retorts.]
_The concentration of sulphuric acid_ in glass retorts is not a
continuous process, and consists of heating the dilute 75 per cent. acid
until it ceases to give off aqueous vapour, and until acid containing
93-98 per cent. H_{2}SO_{4} (66° Baumé) is obtained--and this takes place
when the temperature reaches 320° and the density of the residue reaches
1·847 (66° Baumé).[46] The platinum vessels designed for the continuous
concentration of sulphuric acid consist of a still _B_, furnished with a
still head _E_, a connecting pipe _E F_, and a syphon tube _H R_, which
draws off the sulphuric acid concentrated in the boiler. A stream of
sulphuric acid previously concentrated in lead retorts to a density of
about 60° Baumé--_i.e._ to 75 per cent. or a sp. gr. of 1·7--runs
continuously into the retort through a syphon funnel _E´_. The apparatus
is fed from above, because the acid freshly supplied is lighter than that
which has already lost water, and also because the water is more easily
evaporated from the freshly supplied acid at the surface. The platinum
retort is heated, and the steam coming off[47] is condensed in a worm _F
G_, whilst as fresh dilute acid is supplied to the boiler the acid
already concentrated is drawn off through the syphon tube _H B_, which is
furnished with a regulating cock by means of which the outflow of the
concentrated acid from the bottom of the retort can be so regulated that
it will always present one and the same specific gravity, corresponding
with the strength required. For this purpose the acid flowing from the
syphon is collected in a receiver _R_, in which a hydrometer, indicating
its density, floats; if its density be less than 66° Baumé, the
regulating cock is closed sufficiently to retard the outflow of sulphuric
acid, so as to lengthen the time of its evaporation in the retort.[48]
[46] The difficulty with which the last portions of water are removed
is seen from the fact that the boiling becomes very irregular,
totally ceasing at one moment, then suddenly starting again, with
the rapid formation of a considerable amount of steam, and at the
same time bumping and even overturning the vessel in which it is
held. Hence it is not a rare occurrence for the glass retorts to
break during the distillation; this causes platinum retorts to be
preferred, as the boiling then proceeds quite uniformly.
[47] According to Regnault, the vapour tensions (in millimetres of
mercury) of the water given off by the hydrates of sulphuric acid,
H_{2}SO_{4},_n_H_{2}O, are--
_t_=5° 15° 30°
_n_ = 1 0·1 0·1 0·2
2 0·4 0·7 1·5
3 0·9 1·6 4·1
4 1·3 2·8 7·0
5 2·1 4·2 10·7
7 3·2 6·2 15·6
9 4·1 8·0 19·6
11 4·4 9·0 22·2
17 5·5 10·6 26·1
According to Lunge, the vapour tension of the aqueous vapour given
off from solutions of sulphuric acid containing _p_ per cent.
H_{2}SO_{4}, at _t_°, equals the barometric pressure 720 to 730
mm.
_p_= 10 20 30 40 50 60 70 80 85 90 95
_t_= 102° 105° 108° 114° 124° 141° 170° 207° 233° 262° 295°
The latter figures give the temperature at which water is easily
expelled from solutions of sulphuric acid of different strengths.
But the evaporation begins sooner, and concentration may be
carried on at lower temperatures if a stream of air be passed
through the acid. Kessler's process is based upon this (Note 48).
[48] The greatest part of the sulphuric acid is used in the soda
manufacture, in the conversion of the common salt into sulphate.
For this purpose an acid having a density of 60° Baumé is amply
sufficient. Chamber acid has a density up to 1·57 = 50° to 51°
Baumé; it contains about 35 per cent. of water. About 15 per cent.
of this water can be removed in leaden stills, and nearly all the
remainder may be expelled in glass or platinum vessels. Acid of
66° Baumé, = 1·847, contains about 96 per cent. of the hydrate
H_{2}SO_{4}. The density falls with a greater or less proportion
of water, the maximum density corresponding with 97-1/2 per cent.
of the hydrate H_{2}SO_{4}. The concentration of H_{2}SO_{4} in
platinum retorts has the disadvantage that sulphuric acid, upwards
of 90 per cent. in strength, does corrode platinum, although but
slightly (a few grams per tens of tons of acid). The retorts
therefore require repairing, and the cost of the platinum exceeds
the price obtained for concentrating the acid from 90 per cent. to
98 per cent. (in factories the acid is not concentrated beyond
this by evaporation in the air). This inconvenience has lately
(1891, by Mathey) been eliminated by coating the inside of the
platinum retorts with a thin (0·1 to 0·02 mm.) layer of gold which
is 40 times less corroded by sulphuric acid than platinum. Négrier
(1890) carries on the distillation in porcelain dishes, Blond by
heating a thin platinum wire immersed in the acid by means of an
electric current, but the most promising method is that of Kessler
(1891), which consists in passing hot air over sulphuric acid
flowing in a thin stream in stone vessels, so that there is no
boiling but only evaporation at moderate temperatures: the
transference of the heat is direct (and not through the sides of
the vessels), which economises the fuel and prevents the
distilling vessels being damaged.
When, by evaporation of the water, sulphuric acid attains a
density of 66° Baumé (sp. gr. 1·84), it is impossible to
concentrate it further, because it then distils over unchanged.
_The distillation of sulphuric acid_ is not generally carried on
on a large scale, but forms a laboratory process, employed when
particularly pure acid is required. The distillation is effected
either in platinum retorts furnished with corresponding condensers
and receivers, or in glass retorts. In the latter case, great
caution is necessary, because the boiling of sulphuric acid itself
is accompanied by still more violent jerks and greater
irregularity than even the evaporation of the last portions of
water contained in the acid. If the glass retort which holds the
strong sulphuric acid to be distilled be heated directly from
below, it frequently jerks and breaks. For greater safety the
heating is not effected from below, but at the sides of the
retort. The evaporation then does not proceed in the whole mass,
but only from the upper portions of the liquid, and therefore goes
on much more quietly. The acid may be made to boil quietly also by
surrounding the retort with good conductors of heat--for example,
iron filings, or by immersing a bunch of platinum wires in the
acid, as the bubbles of sulphuric acid vapour then form on the
extremities of the wires.
Strictly speaking, _sulphuric acid is not volatile_, and at its
so-called boiling-point it really decomposes into its anhydride and
water; its boiling-point (338°) being nothing else but its temperature of
decomposition. The products of this decomposition are substances boiling
much below the temperature of the decomposition of sulphuric acid. This
conclusion with regard to the process of the distillation of sulphuric
acid may be deduced from Bineau's observations on the vapour-density of
sulphuric acid. This density referred to hydrogen proved to be half that
which sulphuric acid should have according to its molecular weight,
H_{2}SO_{4}, in which case it should be 49, whilst the observed density
was equal to 24·5. Besides which, Marignac showed that the first portions
of the sulphuric acid distilling over contain less of the elements of
water than the portion which remains behind, or which distils over
towards the end. This is explained by the fact that on distillation the
sulphuric acid is decomposed, but a portion of the water proceeding from
its decomposition is retained by the remaining mass of sulphuric acid,
and therefore at first a mixture of sulphuric acid and sulphuric
anhydride--_i.e._ fuming sulphuric acid--is obtained in the distillate.
It is possible by repeating the distillation several times and only
collecting the first portions of the distillate, to obtain a distinctly
fuming acid. To obtain the definite hydrate H_{2}SO_{4} it is necessary
to refrigerate a highly concentrated acid, of as great a purity as
possible, to which a small quantity of sulphuric anhydride has been
previously added. Sulphuric acid containing a small quantity (a fraction
of a per cent. by weight) of water only freezes at a very low
temperature, while the pure normal acid, H_{2}SO_{4}, solidifies when it
is cooled below 0°, and therefore the normal acid first crystallises out
from the concentrated sulphuric acid. By repeating the refrigeration
several times, and pouring off the unsolidified portion, it is possible
to obtain a pure _normal hydrate_, H_{2}SO_{4}, which melts at 10°·4.
Even at 40° it gives off distinct fumes--that is, it begins to evolve
sulphuric anhydride, which volatilises, and therefore even in a dry
atmosphere the hydrate H_{2}SO_{4} becomes weaker, until it contains
1-1/2 p.c. of water.[49]
[49] Thus it appears that so common, and apparently so stable, a
compound as sulphuric acid decomposes even at a low temperature
with separation of the anhydride, but this decomposition is
restricted by a limit, corresponding to the presence of about
1-1/2 p.c. of water, or to a composition of nearly
H_{2}O,12H_{2}SO_{4}.
Now there is no reason for thinking that this substance is a
definite compound; it is an equilibrated system which does not
decompose under ordinary circumstances below 338°. Dittmar carried
on the distillation under pressures varying between 30 and 2,140
millimetres (of mercury), and he found that the composition of the
residue hardly varies, and contains from 99·2 to 98·2 per cent. of
the normal hydrate, although at 30 mm. the temperature of
distillation is about 210° and at 2,140 mm. it is 382°.
Furthermore, it is a fact of practical importance that under a
pressure of two atmospheres the distillation of sulphuric acid
proceeds very quietly.
Sulphuric acid may be _purified_ from the majority of its
impurities by distillation, if the first and last portions of the
distillate be rejected. The first portions will contain the oxides
of nitrogen, hydrochloric acid, &c., and the last portions the
less volatile impurities. The oxides of nitrogen may be removed by
heating the acid with charcoal, which converts them into volatile
gases. Sulphuric acid may be freed from arsenic by heating it with
manganese dioxide and then distilling. This oxidises all the
arsenic into non-volatile arsenic acid. Without a preliminary
oxidation it would partially remain as volatile arsenious acid,
and might pass over into the distillate. The arsenic may also be
driven off by first reducing it to arsenious acid, and then
passing hydrochloric acid gas through the heated acid. It is then
converted into arsenious chloride, which volatilises.
In a concentrated form sulphuric acid is commercially known as _oil of
vitriol_, because for a long time it was obtained from green vitriol and
because it has an oily appearance and flows from one vessel into another
in a thick and somewhat sluggish stream, like the majority of oily
substances, and in this clearly differs from such liquids as water,
spirit, ether, and the like, which exhibit a far greater mobility. Among
its reactions the first to be remarked is its faculty for the formation
of many compounds. We already know that it combines with its anhydride,
and with the sulphates of the alkali metals; that it is soluble in water,
with which it forms more or less stable compounds. Sulphuric acid, when
mixed with water, develops a very considerable amount of heat.[50]
[50] The amount of heat developed by the mixture of sulphuric acid with
water is expressed in the diagram on p. 77, Volume I., by the
middle curve, whose abscissæ are the percentage amounts of acid
(H_{2}SO_{4}) in the resultant solution, and ordinates the number
of units of heat corresponding with the formation of 100 cubic
centimetres of the solution (at 18°). The calculations on which
the curve is designed are based on Thomsen's determinations, which
show that 98 grams or a molecular amount of sulphuric acid, in
combining with _m_ molecules of water (that is, with _m_=18 grams
of water), develop the following number of units of heat, R:--
_m_ = 1 2 3 5 9
R = 6379 9418 11137 13108 14952
_c_ = 0·432 0·470 0·500 0·576 0·701
T = 127° 149° 146° 121° 82°
_m_ = 19 49 100 200
R = 16256 16684 16859 17066
_c_ = 0·821 0·914 0·954 0·975
T = 145° 19° 9° 5°
_c_ stands for the specific heat of H_{2}SO_{4}_m_H_{2}O
(according to Marignac and Pfaundler), and T for the rise in
temperature which proceeds from the mixture of H_{2}SO_{4} with
_m_H_{2}O. The diagram shows that contraction and rise of
temperature proceed almost parallel with each other.
Besides the normal hydrate H_{2}SO_{4}, _another definite
hydrate_, H_{2}SO_{4},H_{2}O (84·48 per cent. of the normal
hydrate, and 15·52 per cent. of water) is known; it
crystallises[50 bis] extremely easily in large six-sided prisms,
which form above 0°--namely, at about +8°·5; when heated to 210°
it loses water.[51] If the hydrates H_{2}SO_{4} and
H_{2}SO_{4},H_{2}O exist at low temperatures as definite
crystalline compounds, and if pyrosulphuric acid,
H_{2}SO_{4}SO_{3}, has the same property, and if they all
decompose with more or less ease on a rise of temperature, with
the disengagement of either SO_{3} or H_{2}O, and in their
ordinary form present all the properties of simple solutions, it
follows that between sulphuric anhydride, SO_{3}, and water,
H_{2}O, there exists a consecutive series of homogeneous liquids
or solutions, among which we must distinguish _definite
compounds_, and therefore it is quite justifiable to look for
other definite compounds between SO_{3} and H_{2}O, beyond the
conditions for a change of state. In this respect we may be guided
by the variation of properties of any kind, proceeding
concurrently with a variation in the composition of a solution.
[50 bis] Pickering (1890) showed (_a_) that dilute solutions of
sulphuric acid containing up to H_{2}SO_{4} + 10H_{2}O deposit ice
(at -0°·12 when there is 2,000H_{2}O per H_{2}SO_{4}, at -0°·23
when there is 1,000H_{2}O, at -1°·04 when there is 200H_{2}O, at
-2°·12 when there is 100H_{2}O, at -4°·5 when there is 50H_{2}O,
at -15°·7 when there is 20H_{2}O, and at -61° when the composition
of the solution is H_{2}SO_{4} + 10H_{2}O); (_b_) that for higher
concentrations crystals separate out at a considerable degree of
cold, having the composition H_{2}SO_{4}4H_{2}O, which melt at
-24°·5, and if either water or H_{2}SO_{4} be added to this
compound the temperature of crystallisation falls, so that a
solution of the composition 12H_{2}SO_{4} + 100H_{2}O gives
crystals of the above hydrate at -70°, 15H_{2}SO_{4} + 100H_{2}O
at -47°, 30H_{2}SO_{4} + 100H_{2}O at -32°, 40H_{2}SO_{4} +
100H_{2}O at -52°; (_c_) that if the amount of H_{2}SO_{4} be
still greater, then a hydrate H_{2}SO_{4}H_{2}O separates out and
melts at +8°·5, while the addition of water or sulphuric acid to
it lowers the temperature of crystallisation so that the
crystallisation of H_{2}SO_{4}H_{2}O from a solution of the
composition H_{2}SO_{4} + 1·73H_{2}O takes place at -22°,
H_{2}SO_{4} + 1·5H_{2}O at -6°·5, H_{2}SO_{4} + 1·2H_{2}O at
+3°·7, H_{2}SO_{4} + 0·75H_{2}O at +2°·8, H_{2}SO_{4} + 0·5H_{2}O
at -16°; (_d_) that when there is less than 40H_{2}O per
100H_{2}SO_{4}, refrigeration separates out the normal hydrate
H_{2}SO_{4}, which melts at +10°·35, and that a solution of the
composition H_{2}SO_{4} + 0·35H_{2}O deposits crystals of this
hydrate at -34°, H_{2}SO_{4} + 0·1H_{2}O at -4°·1, H_{2}SO_{4} +
0·05H_{2}O at +4°·9, while fuming acid of the composition
H_{2}SO_{4} + 0·05SO_{3} deposits H_{2}SO_{4} at about +7°. Thus
the temperature of the separation of crystals clearly
distinguishes the above four regions of solutions, and in the
space between H_{2}SO_{4} + H_{2}O and +25H_{2}O a particular
hydrate H_{2}SO_{4}4H_{2}O separates out, discovered by Pickering,
the isolation of which deserves full attention and further
research. I may add here that the existence of a hydrate
H_{2}SO_{4}4H_{2}O was pointed out in my work, _The Investigation
of Aqueous Solutions_, p. 120 (1887), upon the basis that it has
at all temperatures a smaller value for the coefficient of
expansion _k_ in the formula S_{_t_} = S_{0}/(1 - _kt_) than the
adjacent (in composition) solutions of sulphuric acid. And for
solutions approximating to H_{2}SO_{4}10H_{2}O in their
composition, _k_ is constant at all temperatures (for more dilute
solutions the value of _k_ increases with _t_ and for more
concentrated solutions it decreases). This solution (with
10H_{2}O) forms the point of transition between more dilute
solutions which deposit ice (water) when refrigerated and those
which give crystals of H_{2}SO_{4}4H_{2}O. According to R. Pictet
(1894) the solution H_{2}SO_{4}10H_{2}O freezes at -88° (but no
reference is made as to what separates out), _i.e._ at a lower
temperature than all the other solutions of sulphuric acid.
However, in respect to these last researches of R. Pictet (for
88·88 p.c. H_{2}SO_{4} -55°, for H_{2}SO_{4}H_{2}O +3·5°, for
H_{2}SO_{4}2H_{2}O -70°, for H_{2}SO_{4}4H_{2}O -40°, &c.) it
should be remarked that they offer some quite improbable data; for
example, for H_{2}SO_{4}75H_{2}O they give the freezing point as
0°, for H_{2}SO_{4}300H_{2}O +4°·5, and even for
H_{2}SO_{4}1000H_{2}O +0°·5, although it is well known that a
small amount of sulphuric acid lowers the temperature of the
formation of ice. I have found by direct experiment that a frozen
solidified solution of H_{2}SO_{4} + 300H_{2}O melted completely
at 0°.
[51] With an excess of snow, the hydrate H_{2}SO_{4},H_{2}O, like the
normal hydrate, gives a freezing mixture, owing to the absorption
of a large amount of heat (the latent heat of fusion). In melting,
the molecule H_{2}SO_{4} absorbs 960 heat units, and the molecule
H_{2}SO_{4}H_{2}O 3,680 heat units. If therefore we mix one gram
molecule of this hydrate with seventeen gram molecules of snow,
there is an absorption of 18,080 heat units, because 17H_{2}O
absorbs 17 × 1,430 heat units, and the combination of the
monohydrate with water evolves 9,800 heat units. As the specific
heat of the resultant compound H_{2}SO_{4},18H_{2}O = 0·813, the
fall of temperature will be -52°·6. And, in fact, a very low
temperature may be obtained by means of sulphuric acid.
But only a few properties have been determined with sufficient accuracy.
In those properties which have been determined for many solutions of
sulphuric acid, it is actually seen that the above-mentioned definite
compounds are distinguished by distinctive marks of change. As an example
we may cite the variation of the specific gravity with a variation of
temperature (namely K = _ds/dt_, if _s_ be the sp. gr. and _t_ the
temperature). For the normal hydrate, H_{2}SO_{4}, this factor is easily
determined from the fact that--
_s_ = 18528 - 10·65_t_ + 0·013_t_^2,
where _s_ is the specific gravity at _t_ (degrees Celsius) if the sp.
gr. of water at 4° = 10,000. Therefore K = 10·65 - 0·026_t_. This means
that at 0° the sp. gr. of the acid H_{2}SO_{4} decreases by 10·65 for
every rise of a degree of temperature, at 10° by 10·39, at 20° by 10·13,
at 30° by 9·87.[52] And for solutions containing slightly more anhydride
than the acid H_{2}SO_{4} (_i.e._ for fuming sulphuric acid), as well as
for solutions containing more water, K is greater than for the acid
H_{2}SO_{4}. Thus for the solution SO_{3},2H_{2}SO_{4}, at 10° K = 11·0.
On diluting the acid H_{2}SO_{4} K again increases until the formation of
the solution H_{2}SO_{4},H_{2}O (K = 11·1 at 10°), and then, on further
dilution with water, it again decreases. Consequently both hydrates
H_{2}SO_{4} and H_{2}SO_{4},H_{2}O are here expressed by an alteration of
the magnitude of K.
[52] For example, if it be taken that at 19° the sp. gr. of pure
sulphuric acid is 1·8330, then at 20° it is 1·8330 - (20 -
19)10·13 = 1·8320.
[Illustration: FIG. 89.--Diagram showing the variation of the factor
(_ds/dp_) of the specific gravity of solutions of sulphuric acid. The
percentage quantities of the acid, H_{2}SO_{4}, are laid out on the axes
of abscissæ. The ordinates are the factors or rises in sp. gr. (water
at 4 = 10,000) with the increase in the quantity of H_{2}SO_{4}.]
This shows that in liquid solutions it is possible by studying the
variation of their properties (without a change of physical state) to
recognise the presence or formation of definite hydrate compounds, and
therefore an exact investigation of the properties of solutions, of their
specific gravity for instance, should give direct indications of such
compounds.[53] The mean result of the most trustworthy determinations of
this nature is given in the following tables. The first of these tables
gives the specific gravities (in vacuo, taking the sp. gr. of water at 4°
= 1), at 0° (column 3), 15° (column 4), and 30° (column 5),[53 bis] for
solutions having the composition H_{2}SO_{4} + _n_H_{2}O (the value of
_n_ is given in the first column), and containing _p_ (column 2) per
cent. (by weight in vacuo) of H_{2}SO_{4}.[53 tri]
_n_ _p_ 0° 15° 30°
100 5·16 1·0374 1·0341 1·0292
50 9·82 1·0717 1·0666 1·0603
25 17·88 1·1337 1·1257 1·1173
15 26·63 1·2040 1·1939 1·1837
10 35·25 1·2758 1·2649 1·2540
8 40·50 1·3223 1·3110 1·2998
6 47·57 1·3865 1·3748 1·3622
5 52·13 1·4301 1·4180 1·4062
4 57·65 1·4881 1·4755 1·4631
3 64·47 1·5635 1·5501 1·5370
2 73·13 1·6648 1·6500 1·6359
1 84·48 1·7940 1·7772 1·7608
0·5 91·59 1·8445 1·8284 1·8128
H_{2}SO_{4} 100 1·8529 1·8372 1·8221
[53] Unfortunately, notwithstanding the great number of fragmentary and
systematic researches which have been made (by Parks, Ure, Bineau,
Kolbe, Lunge, Marignac, Kremers, Thomsen, Perkin, and others) for
determining the relation between the sp. gr. and composition of
solutions of sulphuric acid, they contain discrepancies which
amount to, and even exceed, 0·002 in the sp. gr. For instance, at
15°·4 the solution of composition H_{2}SO_{4}3H_{2}O has a sp. gr.
1·5493 according to Perkin (1886), 1·5501 according to Pickering
(1890), and 1·5525 according to Lunge (1890). The cause of these
discrepancies must be looked for in the methods employed for
determining the composition of the solutions--_i.e._ in the
inaccuracy with which the percentage amount of H_{2}SO_{4} is
determined, for a difference of 1 p.c. corresponds to a difference
of from 0·0070 (for very weak solutions) to 0·0118 (for a solution
containing about 73 p.c.) in the specific gravity (that is the
factor _ds/dp_) at 15°. As it is possible to determine the
specific gravity with an accuracy even exceeding 0·0002, the
specific gravities given in the adjoining tables are only averages
and most probable data in which the error, especially for the
30-80 p.c. solutions cannot be less than 0·0010 (taking water at
4° as 1).
[53 bis] Judging from the best existing determinations (of Marignac,
Kremers, and Pickering) for solutions of sulphuric acid
(especially those containing more than 5 p.c. H_{2}SO_{4}) within
the limits of 0° and 30° (and even to 40°), the variation of the
sp. gr. with the temperature _t_ may (within the accuracy of the
existing determinations) be perfectly expressed by the equation
S_{_t_} = S_{_0_} + A_t_ + B_t_^2. It must be added that (1) three
specific gravities fully determine the variation of the density
with _t_; (2) _ds/dt_ = A + 2B_t_--_i.e._ the factor of the
temperature is expressed by a straight line; (3) the value of A
(if _p_ be greater than 5 p.c.) is negative, and numerically much
greater than B; (4) the value of B for dilute solutions containing
less than 25 p.c. is negative; for solutions approximating to
H_{2}SO_{4}3H_{2}O in their composition it is equal to 0, and for
solutions of greater concentration B is positive; (5) the factor
_ds/dp_ for all temperatures attains a maximum value about
H_{2}SO_{4}H_{2}O; (6) on dividing _ds/dt_ by S_{_0_}, and so
obtaining the coefficient of expansion _k_ (_see_ Note 53), a
minimum is obtained near H_{2}SO_{4} and H_{2}SO_{4}4H_{2}O, and a
maximum at H_{2}SO_{4}H_{2}O for all temperatures.
[53 tri] These data (as well as those in the following table) have been
recalculated by me chiefly upon the basis of Kremer's,
Pickering's, Perkin's, and my own determinations; all the
requisite corrections have been introduced, and I have reason for
thinking that in each of them the probable error (or difference
from the true figures, now unknown) of the specific gravity does
not exceed ±0·0007 (if water at 4° = 1) for the 25-80 p.c.
solutions, and ±0·0002 for the more dilute or concentrated
solutions.
In the second table the first column gives the percentage amount _p_ (by
weight) of H_{2}SO_{4}, the second column the weight in grams (S_{15}) of
a litre of the solution at 15° (at 4° the weight of a litre of water =
1,000 grams), the third column, the variation (_d_S/_dt_) of this weight
for a rise of 1°, the fourth column, the variation _d_S/_dp_ of this
weight (at 15°) for a rise of 1 per cent. of H_{2}SO_{4}, the fifth
column, the difference between the weight of a litre at 0° and 15° (S_{0}
- S_{15}), and the sixth column, the difference between the weight of a
litre at 15° and 30° (S_{15} - S_{30}).
_p_ _S__{15} _dS__{15}/_dt_ _dS__{15}/_dp_ _S__{0}- _S__{15}-
_S__{15} _S__{30}
0 999·15 0·148 7·0 0·7 3·4
5 1033·0 0·27 6·8 3·1 5·0
10 1067·7 0·38 7·1 5·2 6·4
20 1141·9 0·58 7·7 8·6 8·9
30 1221·3 0·69 8·2 10·4 10·4
40 1306·6 0·75 8·8 11·3 11·2
50 1397·9 0·79 9·9 11·9 11·8
60 1501·2 0·86 10·8 13·0 12·7
70 1613·1 0·93 11·6 14·1 13·8
80 1731·4 1·04 11·0 15·8 15·4
90 1819·9 1·08 5·4 16·4 16·0
95 1837·6 1·03 +1·7 15·8 15·1
100 1837·2 1·03 -1·9[54] 15·7 15·1
The figures in these tables give the means of finding the amount of
H_{2}SO_{4} contained in a solution from its specific gravity,[55] and
also show that 'special points' in the lines of variation of the specific
gravity with the temperature and percentage composition correspond to
certain definite compounds of H_{2}SO_{4} with OH_{2}. This is best seen
in the variation of the factors (_d_S/_dt_ and _d_S/_dp_) with the
temperature and composition (columns 3, 4, second table). We have already
mentioned how the factor of temperature points to the existence of
hydrates, H_{2}SO_{4} and H_{2}SO_{4},H_{2}O. As regards the factor
_d_S/_dp_ (giving the increase of sp. gr. with an increase of 1 per cent.
H_{2}SO_{4}) the following are the three most salient points: (1) In
passing from 98 per cent. to 100 per cent. the factor is negative, and at
100 per cent. about -0·0019 (_i.e._ at 99 per cent. the sp. gr. is about
1·8391, and at 100 per cent. about 1·8372, at 15°, the amount of
H_{2}SO_{4} has increased whilst the sp. gr. has decreased), but as soon
as a certain amount of SO_{3} is added to the definite compound
H_{2}SO_{4} (and 'fuming' acid formed) the specific gravity rises (for
example, for H_{2}SO_{4} 0·136 SO_{3} the sp. gr. at 15° = 1·866), that
is the factor becomes positive (and, in fact, greater by +0·01), so that
the formation of the definite hydrate H_{2}SO_{4} is accompanied by a
distinct and considerable break in the continuity of the factor[55 bis];
(2) The factor (_d_S/_dp_) in increasing in its passage from dilute to
concentrated solutions, attains a maximum value (at 15° about 0·012)
about H_{2}SO_{4}2H_{2}O, _i.e._ at about the hydrate corresponding to
the form SX_{6}; proper to the compounds of sulphur, for S(OH)_{6} =
H_{2}SO_{4}2H_{2}O; the same hydrate corresponds to the composition of
gypsum CaSO_{4}2H_{2}O, and to it also corresponds the greatest
contraction and rise of temperature in mixing H_{2}SO_{4} with H_{2}O
(_see_ Chapter I., Note 28); (3) The variation of the factor (_d_S/_dp_)
under certain variations in the composition proceeds so uniformly and
regularly, and is so different from the variation given under other
proportions of H_{2}SO_{4} and H_{2}O, that the sum of the variations of
_d_S/_dp_ is expressed by a series of straight lines, if the values of
_p_ be laid along the axis of abscissæ and those of _d_S/_dp_ along the
ordinates.[56] Thus, for instance, for 15°, at 10 per cent. _d_S/_dp_ =
0·0071, at 20 per cent. = 0·0077, at 30 per cent. = 0·0082, at 40 per
cent. = 0·0088, that is, for each 10 per cent. the factor increases by
about 0·0006 for the whole of the above range, but beyond this it becomes
larger, and then, after passing H_{2}SO_{4}2H_{2}O, it begins to fall
rapidly. Such changes in the variation of the factor take place
apparently about definite hydrates,[56 bis] and especially about
H_{2}SO_{4}4H_{2}O, H_{2}SO_{4}2H_{2}O and H_{2}SO_{4}H_{2}O. All this
indicating as it does the special chemical affinity of sulphuric acid for
water, although of no small significance for comprehending the nature of
solutions (_see_ Chapter I. and Chapter VII.), contains many special
points which require detailed investigation, the chief difficulty being
that it requires great accuracy in a large number of experimental data.