Chapter XX: , Note 20). Thus the forms PX{3} and PX{5} not only (1)

exist in many solid and non-volatile substances, but also as
vapours.

If a piece of phosphorus be dropped into a flask containing chlorine, it burns when touched with a red-hot wire, and combines with the chlorine. If the phosphorus be in excess, liquid _phosphorus trichloride_, PCl_{3}, is always formed, but if the chlorine be in excess the solid pentachloride is obtained. The trichloride is generally prepared in the following manner. Dry chlorine (passed through a series of Woulfe's bottles containing sulphuric acid) is led into a retort containing sand and phosphorus. The retort is heated, the phosphorus melts, spreads through the sand, and gradually forms the trichloride, which distils over into a receiver, where it condenses. _Phosphoric chloride_ or _phosphorus pentachloride_, PCl_{5}, is prepared by passing dry chlorine into a vessel containing phosphorus trichloride (purified by distillation). Phosphorous chloride combines directly with oxygen, but more rapidly with ozone or with the oxygen of potassium chlorate (3PCl_{3} + KClO_{3} = 3POCl_{3} + KCl), forming _phosphorus oxychloride_, POCl_{3} (Brodie). This compound is also formed by the first action of water on phosphoric chloride; for example, if two vessels, one containing phosphoric chloride and the other water, are placed under a bell jar, after a certain time the crystals of the chloride disappear and hydrochloric acid passes into the water. The aqueous vapour acts on the pentachloride, and the following reaction occurs: PCl_{5} + H_{2}O = POCl_{3} + 2HCl, the result being that liquid phosphorus oxychloride is found in one vessel, and a solution of hydrochloric acid in the other. However, an excess of water directly transforms phosphoric chloride into orthophosphoric acid, PCl_{5} + 4H_{2}O = PH_{3}O_{4} + 5HCl,[26] since POCl_{3} reacts with water (3H_{2}O), forming 3HCl and phosphoric acid PO(OH)_{3}.

[26] Phosphorus oxychloride is obtained by the action of phosphoric
chloride on hydrates of acids (because alkalis decompose
phosphorus oxychloride), according to the equation PCl_{5} + RHO =
POCl_{3} + RCl + HCl, where RHO is an acid. The reaction only
proceeds according to this equation with monobasic acids, but then
RCl is volatile, and therefore a mixture is obtained of two
volatile substances, the acid chloride and phosphorus oxychloride,
which are sometimes difficult to separate; whilst if the hydrate
be polybasic the reaction frequently proceeds so that an anhydride
is formed: RH_{2}O_{2} + PCl_{5} = RO + POCl_{3} + 2HCl. If the
anhydride be non-volatile (like boric), or easily decomposed (like
oxalic), it is easy to obtain pure oxychloride. Thus phosphorus
oxychloride is often prepared by acting on boric or oxalic acid
with phosphoric chloride. It is also formed when the vapour of
phosphoric chloride is passed over phosphoric anhydride,
P_{2}O_{5} + 3PCl_{5} = 5POCl_{3}. This forms an excellent example
in proof of the fact that the formation of one substance from two
does not necessarily show that the resultant compound contains the
molecules of these substances in its molecule. But other
oxychlorides of phosphorus are also formed by the interaction of
phosphoric anhydride and chloride; thus at 200° the
chloranhydride, PO_{2}Cl, or chloranhydride of metaphosphoric
acid, is formed (Gustavson). The chloranhydride of pyrophosphoric
acid, P_{2}O_{3}Cl_{4}, was obtained (Hayter and Michaelis),
together with NOCl, &c., by the action of NO upon cold PCl_{3}, as
a fuming liquid boiling at 210°.

The above chlorine compounds serve not only as a type of the chloranhydrides, but also as a means for the preparation of other _acid chloranhydrides_. Thus the conversion of acids XHO into chloranhydrides, XCl, is generally accomplished by means of _phosphorus pentachloride_. This fact was discovered by Chancel, and adopted by Gerhardt as an important method for studying organic acids. By this means organic acids, containing, as we know, RCOOH (where R is a hydrocarbon group, and where carboxyl may repeat itself several times by replacing the hydrogen of hydrocarbon compounds), are converted into their chloranhydrides, RCOCl. With water they again form the acid, and resemble the chloranhydrides of mineral acids in their general properties.

Since carbonic acid, CO(OH)_{2}, contains two hydroxyl groups, its perfect chloranhydride, COCl_{2}, _carbonic oxychloride_, _carbonyl chloride_ or _phosgene gas_, contains two atoms of chlorine, and differs from the chloranhydrides of organic acids in that in them one atom of chlorine is replaced by the hydrocarbon radicle RCOCl, if R be a monatomic radicle giving a hydrocarbon RH. It is evident, on the one hand, that in RCOCl the hydrogen is replaced by the radicle COCl, which is also able to replace several atoms of hydrogen (for example, C_{2}H_{4}(COCl)_{2} corresponds with the bibasic succinic acid); and, on the other hand, that the reactions of the chloranhydrides of organic acids will answer to the reactions of carbonyl chloride, as the reactions of the acids themselves answer to those of carbonic acid. Carbonyl chloride is obtained directly from dry carbon monoxide and chlorine[27] exposed to the action of light, and forms a colourless gas, which easily condenses into a liquid, boiling at +8°, specific gravity 1·43, and having the suffocating odour belonging to all chloranhydrides. Like all chloranhydrides, it is immediately decomposed by water, forming carbonic anhydride, according to the equation COCl_{2} + H_{2}O = CO_{2} + 2HCl, and thus expresses the type proper to all chloranhydrides of both mineral and organic acids.[28]

[27] The direct action of the sun's rays, or of magnesium light, is
necessary to start the reaction between carbonic oxide and
chlorine, but when once started it will proceed rapidly in
diffused light. An excess of chlorine (which gives its coloration
to the colourless phosgene) aids the completion of the reaction,
and may afterwards be removed by metallic antimony. Porous
substances, like charcoal, aid the reaction. Phosgene may be
prepared by passing a mixture of carbonic anhydride and chlorine
over incandescent charcoal. Lead or silver chloride, when heated
in a current of carbonic oxide, also partially form phosgene gas.
Carbon tetrachloride, CCl_{4}, also forms it when heated with
carbonic anhydride (at 400°), with phosphoric anhydride (200°),
and most easily of all with sulphuric anhydride (2SO_{3} + CCl_{4}
= COCl_{2} + S_{2}O_{5}Cl_{2}, this is pyrosulphuryl chloride).
Chloroform, CHCl_{3}, is converted into carbonyl chloride when
heated with SO_{2}(OH)Cl (the first chloranhydride of sulphuric
acid); CHCl_{3} + SO_{3}HCl = COCl_{2} + SO_{2} + 2HCl (Dewar),
and when oxidised by chromic acid.

Among the reactions of phosgene we may mention the formation of
urea with ammonia, and of carbonic oxide when heated with metals.

[28] We are already acquainted with some of the chloranhydrides of the
inorganic acids--for instance, BCl_{3}, and SiCl_{4}--and here we
shall describe those which correspond with sulphuric acid in the
following chapter. It may be mentioned here that when hydrochloric
acts on nitric acid (aqua regia, Vol. I. p. 467) there is formed,
besides chlorine, the oxychlorides NOCl and NO_{2}Cl, which may be
regarded as chloranhydrides of nitric and nitrous acids (nitrogen
chloride, Vol. I. p. 476). The former boils at -5°, the latter at
+5°, the specific gravity of the first at -12° = 1·416, and at
-18° = 1·433 (Geuther), and of the second = 1·3; the first is
obtained from nitric oxide and chlorine, the second from nitric
peroxide and chlorine, and also by the action of phosphoric
chloride on nitric acid. If the gases evolved by aqua regia be
passed into cold and strong sulphuric acid, they form crystals of
the composition NHSO_{3} (like chamber crystals), which melt at
86°, and with sodium chloride form acid sodium sulphate and the
oxychloride NOCl. This chloranhydride of nitric acid is termed
_nitrosyl chloride_.

_Cyanogen chloride_, CNCl, is the gaseous chloranhydride of cyanic
acid; it is formed by the action of chlorine on aqueous mercury
cyanide, Hg(CN)_{2} + 2Cl_{2} = HgCl_{2} + 2CNCl. When chlorine
acts on cyanic acid, it forms not only this cyanogen chloride, but
also polymerides of it--a liquid, boiling at 18°, and a solid,
boiling at 190°. The latter corresponds with cyanuric acid, and
consequently contains C_{3}N_{3}Cl_{3}. Details concerning these
substances must be looked for in works on organic chemistry.

In order to show the general method for the preparation of acid chloranhydrides, we will take that of acetic acid, CH_{3}·COOH, as an example. Phosphorus pentachloride is placed in a glass retort, and acetic acid poured over it; hydrochloric acid is then evolved, and the substance distilling over directly after is a very volatile liquid, boiling at 50°, and having all the properties of the chloranhydrides. With water it forms hydrochloric and acetic acids. The reaction here taking place may be explained thus: the substitution of the oxygen taken from the acetic acid (from its carboxyl) by two atoms of chlorine from the PCl_{5} should be as follows: CH_{3}·COOH + PCl_{5} = CH_{3}·COHCl_{2} + POCl_{3}. But the compound CH_{3}·COHCl_{2} does not exist in a free state (because it would indicate the possibility of the formation of compounds of the type CX_{6}, and carbon only gives those of the type CX_{4}); it therefore splits up into HCl and the chloranhydride CH_{3}·COCl. The general scheme for the reaction of phosphorus pentachloride with hydrates ROH is exactly the same as with water; namely, ROH with PCl_{5}, gives POCl_{3} + HCl + RCl--that is a chloranhydride.[28 bis]

[28 bis] This reaction indeed proceeds very easily and completely with
a number of hydroxides, if they do not react on hydrochloric acid
and phosphorus oxychloride, which is the case when they have
alkaline properties. When the hydroxide is bibasic and is present
in excess, it not unfrequently happens that the elements of water
are taken up: R(OH)_{2} + PCl_{5} = RO + 2HCl + POCl_{3}. The
anhydride RO may then be converted into chloranhydride, RO +
PCl_{5} = RCl_{2} + POCl_{3}--that is, phosphorus pentachloride
brings about the substitution of O by Cl_{2}. Thus carbonyl
chloride, COCl_{2}, boron chloride, 2BCl_{3}, and succinic
chloride, C_{4}H_{4}O_{2}Cl_{2}, &c., are respectively obtained by
the action of phosphoric chloride on carbonic, boric, and succinic
anhydrides. Phosphorus pentachloride reacts in a similar manner on
the aldehydes, RCHO, forming RCHCl_{2}, and on the chloranhydrides
themselves--for example, with acetic chloride, CH_{3}.COCl (when
heated in a closed tube), it forms a substance having the
composition CH_{3}CCl_{3}.

Phosphorus trichloride and oxychloride act in a similar manner to
phosphoric chloride. When phosphorus trichloride acts on an acid,
3RHO + PCl_{3} = 3RCl + P(HO)_{3}. If a salt is taken, then by the
action of phosphorus oxychloride a corresponding chloranhydride
and salt of orthophosphoric acid are easily formed: 3R(KO) +
POCl_{3} = 3RCl + PO(KO)_{3}. The chloranhydride RCl is always
more volatile than its corresponding acid, and distils over before
the hydrate RHO. Thus acetic acid boils at 117°, and its
chloranhydride at 50°. Phosphoric and phosphorous acids are very
slightly volatile, whilst their chloranhydrides are comparatively
easily converted into vapour. The faculty of the chloranhydrides
to react at the expense of their own chlorine determines their
great importance in chemistry. For instance, suppose we require to
know the molecular formula of some hydrate which does not pass
into a state of vapour and does not give a chloranhydride with
hydrochloric acid--that is, which has not any basic or alkaline
properties; we must then endeavour to obtain this chloranhydride
by means of phosphoric chloride, and it frequently happens that
the corresponding chloranhydride is volatile. The resultant
chloranhydride is then converted into vapour, and its composition
is determined; and if we know its composition we are able to
decide that of its corresponding hydrate. Thus, for example, from
the formula of silicon chloride, SiCl_{4}, or of boron chloride,
BCl_{3}, we can judge the composition of their corresponding
hydrates, Si(HO)_{4}, B(HO)_{3}. Having obtained the
chloranhydride RCl or RCl_{_n_}, it is possible by its means to
obtain many other compounds of the same radicle R according to the
equation MX + RCl = MCl + RX. M may be = H, K, Ag, or other metal.
The reaction proceeds thus if M forms a stable compound with
chlorine--for example, silver chloride, hydrochloric acid, and R,
an unstable substance. Hence, a chloranhydride is frequently
employed for the formation of other compounds of a given radicle;
for instance, with ammonia they form amides RNH_{2}, and with
salts ROK, with anhydrides R_{2}O, &c.

Containing, as they do, chlorine, which easily reacts with hydrogen, phosphorus pentachloride, trichloride, and oxychloride enter into reaction with ammonia, and give a series of amide and nitrile compounds of phosphorus. Thus, for example, when ammonia acts on the oxychloride we obtain sal-ammoniac (which is afterwards removed by water) and an orthophosphoric triamide, PO(NH_{2})_{3}, as a white insoluble powder on which dilute acids and alkalis do not act, but which, when fused with potassium hydroxide, gives potassium phosphate and ammonia like other amides. When ignited, the triamide liberates ammonia and forms the nitrile PON, just as urea, CO(NH_{2})_{2}, gives off ammonia and forms the nitrile CONH. This nitrile, called _monophosphamide_, PON, naturally corresponds with metaphosphoric acid, namely, with its ammonium salt. NH_{4}PO_{3} - H_{2}O = PO_{2}·NH_{2}, an as yet unknown amide, and PO_{2}·NH_{2} - H_{2}O gives the nitrile PON. This relation is confirmed by the fact that PON, moistened with water, gives metaphosphoric acid when ignited. It is the analogue of nitrous oxide, NON. It is a very stable compound, more so than the preceding.[29]

[29] The reaction of ammonia on phosphorus pentachloride is more
complex than the preceding. This is readily understood: to the
oxychloride, POCl_{3}, tere corresponds a hydrate PO(OH)_{3}, and
a salt PO(NH_{4}O)_{3}, and consequently also an amide
PO(NH_{2})_{3}, whilst the pentachloride, PCl_{5}, has no
corresponding hydrate P(OH)_{5}, and therefore there is no amide
P(NH_{2})_{5}. The reaction with ammonia will be of two kinds:
either instead of 5 mol. NH_{3}, only 3 mol. NH_{3} or still less
will act; _i.e._ PCl_{2}(NH_{2})_{3}, PCl_{3}(NH_{2})_{2}, &c. are
formed; or else the pentachloride will act like a mixture of
chlorine with the trichloride, and then as the result there will
be obtained the products of the action of chlorine on those amides
which are formed from phosphorus trichloride and ammonia. It would
appear that both kinds of reaction proceed simultaneously, but
both kinds of products are unstable, at all events complex, and in
the result there is obtained a mixture containing sal-ammoniac,
&c. The products of the first kind should react with water, and we
should obtain, for example, PCl_{3}(NH_{2})_{2} + 2H_{2}O = 3HCl
and PO(HO)(NH_{2})_{2}. This substance has not actually been
obtained, but the compound PONH(NH_{2}) derived from it by
elimination of the elements of water is known, and is termed
_diphosphamide_; it is, however, more probable that it is a
nitrile than an amide, because only amides contain the group
NH_{2}. It is a colourless, stable, insoluble powder, which
possibly corresponds with pyrophosphoric acid, more especially
since when heated it evolves ammonia and gives and leaves
phosphoryl nitride, PON--that is, the nitrile of metaphosphoric
acid. The amide corresponding with the pyrophosphate
P_{2}O_{3}(NH_{4}O)_{4} should be P_{2}O_{3}(NH_{2})_{4}, and the
nitriles corresponding to the latter would be
P_{2}O_{2}N(NH_{2})_{3}, P_{2}ON_{2}(NH_{2})_{2}, and
P_{2}N_{3}(NH_{2}). The composition of the first is the same as
that of the above diphosphamide. The third pyrophosphoric nitrile
has a formula P_{2}N_{4}H_{2}, and this is the composition of the
body known as _phospham_, PHN_{2} (in a certain sense this is the
analogue of N_{3}H polymerised, Chapter VI.) Indeed, phospham has
been obtained by heating the products of the action of ammonia on
phosphoric chloride, as an insoluble and alkaline powder, which
gives ammonia and phosphoric acid when subjected to the action of
water. The same substance is obtained by the action of ammonium
chloride on phosphoric chloride (PNCl_{2} is first formed, and
reacts further with ammonia, forming phospham), and by igniting
the mass which is formed by the action of ammonia on phosphorus
trichloride. Formerly the composition of phospham was supposed to
be PHN_{2}, now there is reason to think that its molecular weight
is P_{3}H_{3}N_{6}.

The above compounds correspond with normal salts, but nitriles and
amides corresponding to acid salts are also possible, and they
will be acids. For example, the amide PO(HO)_{2}(NH_{2}), and its
nitrile, will be either PN(HO)_{2} or PO(HO)(NH), but at all
events of the composition PNH_{2}O_{2}, and having acid
properties. The ammonium salt of this _phosphonitrilic acid_ (it
is called phosphamic acid), PNH(NH_{4})O_{2}, is obtained by the
action of ammonia on phosphoric anhydride, P_{2}O_{5} + 4NH_3 =
H_{2}O + 2PNH(NH_{4})O_{2}. A non-crystalline soluble mass is thus
formed, which is dissolved in a dilute solution of ammonia and
precipitated with barium chloride, and the resultant barium salt
is then decomposed with sulphuric acid, and thus a solution of the
acid of the above composition is obtained.

It is evident from the theory of the formation of amides and
nitriles (Chapter IX.) that very many compounds of this kind can
correspond with the acids of phosphorus; but as yet only a few are
known. The easy transitions of the ortho-, meta-, and
pyrophosphoric acids, by means of the hydrogen of ammonia, into
the lower acids, and conversely, tend to complicate the study of
this very large class of compounds, and it is rarely that the
nature of a product thus obtained can be judged from its
composition; and this all the more that instances of isomerism and
polymerism, of mixture between water of crystallisation and of
constitution, &c., are here possible. Many data are yet needed to
enable us to form a true judgment as to the composition and
structure of such compounds. As the best proof of this we will
describe the very interesting and most fully investigated compound
of this class, PNCl_{2}, called _chlorophosphamide_, or nitrogen
chlorophosphorite. It is formed in small quantities when the
vapour of phosphoric chloride is passed over ignited sal-ammoniac.
Besson (1892) heated the compound PCl_{5}8NH_{3} (which is easily
and directly formed from PCl_{5} and NH_{3}) under a pressure of
about 50 mm. (of mercury) to 200°, and obtained brilliant crystals
of PNCl_{2}, which melted at 106° (in the residue after the
distillation of sal-ammoniacal phospham). The chlorine in it is
very stable--quite different from that in phosphoric chloride.
Indeed, the resultant substance is not only insoluble in water
(though soluble in alcohol and ether), but it is not even
moistened by it, and distils over, together with steam, without
being decomposed. In a free state it easily crystallises in
colourless prisms, fuses at 114°, boils at 250° (Gladstone,
Wichelhaus), and when fused with potash gives potassium chloride
and the amidonitrile of phosphoric acid. Judging from its formula
and the simplicity of its composition and reactions, it might be
thought that the molecular weight of this substance would be
expressed by the formula PCl_{2}N, that it corresponds with PON
and with PCl_{5} (like POCl_3), with the substitution of Cl_3 by
N, just as in POCl_3 two atoms of chlorine are replaced by oxygen;
but all these surmises are incorrect, because its vapour density
(referred to hydrogen--Gladstone, Wichelhaus) = 182--that is, the
molecular formula must be three times greater, P_{3}N_{3}Cl_{6}.
The polymerisation (tripling) is here of exactly the same kind as
with the nitriles.

The most important analogue of phosphorus is _arsenic_, the metallic aspect of which and the general character of its compounds of the types AsX_{3} and AsX_{5} at once recall the metals. The hydrate of its highest oxide, arsenic acid (ortho-arsenic acid), H_{3}AsO_{4}, is an oxidising agent, and gives up a portion of its oxygen to many other substances; but, nevertheless, it is very like phosphoric acid. Mitscherlich established the conception of isomorphism by comparing the salts of these acids.[30]

[30] It is necessary to remark that, although arsenic is so closely
analogous to phosphorus (especially in the higher forms of
combination, RX_{3} and RX_{5}), at the same time it exhibits a
certain resemblance and even isomorphism with the corresponding
compounds of sulphur (especially the metallic compounds of the
type MAs, corresponding with MS). Thus compounds containing
metals, arsenic, and sulphur are very frequently met with in
nature. Sometimes the relative amounts of arsenic and sulphur
vary, so that an isomorphous substitution between the arsenides
and sulphides must be recognised. Besides FeS_{2} (ordinary
pyrites), and FeAs_{2}, iron forms an arsenical pyrites containing
both sulphur and arsenic, which from its composition, FeAsS or
FeS_{2}FeAs_{2}, resembles the two preceding.

Arsenic occurs _in nature_, not only combined with metals, but also, although rarely, native and also in combination with sulphur in two minerals--one red, _realgar_, As_{2}S_{2}, and the other yellow, _orpiment_, As_{2}S_{3} (Chapter XX., Note 29). Arsenic occurs, but more rarely, in the form of salts of arsenic acid--for instance, the so-called cobalt and nickel blooms, two minerals which are found accompanying other cobalt ores, are the arsenates of these metals. Arsenic is also found in certain clays (ochres) and has been discovered in small quantities in some mineral springs, but it is in general of rarer occurrence in nature than phosphorus. Arsenic is most frequently extracted from arsenical pyrites, FeSAs, which, when roasted without access of air, evolves the vapour of arsenic, ferrous sulphide being left behind. It is also obtained by heating arsenious anhydride with charcoal, in which case carbonic oxide is evolved. In general, the oxides and other compounds are very easily reduced. Solid _arsenic_ is a steel-grey brittle _metal_, having a bright lustre and scaly structure. Its specific gravity is 5·7. It is opaque and infusible, but volatilises as a yellow vapour which on cooling deposits rhombohedral crystals.[30 bis] The vapour density of arsenic is 150 times greater than that of hydrogen--that is, its molecule, like that of phosphorus, contains 4 atoms, As_{4}. When heated in the air, arsenic easily oxidises into white arsenious anhydride, As_{2}O_{3}, but even at the ordinary temperature it loses its lustre (becomes dull), owing to the formation of a coating of a lower oxide. The latter appears to be as volatile as arsenious anhydride, and it is probable that it is owing to the presence of this compound that the vapours of arsenious compounds, when heated with charcoal (for example, in the reducing flame of a blow-pipe), have the characteristic smell of garlic, because the vapour of arsenic itself has not this odour.

[30 bis] According to Retgers (1893) the arsenic mirror (see further
on) is an unstable variety of metallic arsenic, whilst the brown
product which is formed together with it in Marsh's apparatus is a
lower hydride AsH. Schuller and McLeod (1894), however, recognise
a peculiar yellow variety of arsenic.

Arsenic easily combines with bromine and chlorine;[31] nitric acid and aqua regia also oxidise it into the higher oxide, or rather its hydrate, arsenic acid.[32] As far as is known, it does not decompose steam, and it acts exceedingly slowly on those acids, like hydrochloric, which are not capable of oxidising.

[31] Hydrochloric acid dissolves arsenious anhydride in considerable
quantities, and this is probably owing to the formation of
unstable compounds in which the arsenious anhydride plays the part
of a base. A compound called _arsenious oxychloride_, having the
composition AsOCl, is even known. It is formed when arsenious
anhydride is added little by little to boiling arsenic
trichloride, As_{2}O_{3} + AsCl_{3} = 3AsOCl. It is a transparent
substance, which fumes in air, and combines with water to form a
crystalline mass having the composition As_{2}(OH)_{4}Cl_{2}. When
heated it decomposes into arsenious chloride and a fresh
oxychloride of a more complex composition, As_{6}O_{8}Cl_{2}·
Arsenic trichloride, when treated with a small quantity of water,
forms the crystalline compound, As_{2}(HO)_{4}Cl_{2}, mentioned
above. These compounds resemble the basic salts of bismuth and
aluminium. The existence of these compounds shows that arsenic is
of a more metallic or basic character than phosphorus.
Nevertheless _arsenic trichloride_, AsCl_{3}, resembles phosphorus
trichloride in many respects. It is obtained by the direct action
of chlorine on arsenic, or by distilling a mixture of common salt,
sulphuric acid, and arsenious anhydride. The latter mode of
preparation already indicates the basic properties of the oxide.
Arsenious chloride is a colourless oily liquid, boiling at 130°,
and having a sp. gr. of 2·20. It fumes in air like other
chloranhydrides, but it is much more slowly and imperfectly
decomposed by water than phosphorus trichloride. A considerable
quantity of water is required for its complete decomposition into
hydrochloric acid and arsenious anhydride. It forms an excellent
example of the transition from true metallic chlorides to true
chloranhydrides of the acids. It hardly combines with chlorine,
_i.e._ if AsCl_{5} is formed it is very unstable. _Arsenic
tribromide_, AsBr_{3}, is formed as a crystalline substance,
fusing at 20° and boiling at 220°, by the direct action of
metallic arsenic on a solution of bromine in carbon bisulphide,
the latter being then evaporated. The specific gravity of arsenic
tribromide is 3·36. Crystalline arsenic tri-iodide, AsI_{3},
having a sp. gr. 4·39, may be obtained in a like manner; it may be
dissolved in water, and on evaporation separates out from the
solution in an anhydrous state--that is, it is not decomposed--and
consequently behaves like metallic salts. _Arsenic trifluoride_,
AsF_{3}, is obtained by heating fluor spar and arsenious anhydride
with sulphuric acid. It is a fuming, colourless, and very
poisonous liquid, which boils at 63° and has a sp. gr. of 2·73. It
is decomposed by water. It is very remarkable that fluorine forms
a pentafluoride of arsenic also, although this compound has not
yet been obtained in a separate state, but only in combination
with potassium fluoride. This compound, K_{3}AsF_{8}, is formed as
prismatic crystals when potassium arsenate, K_{3}AsO_{4}, is
dissolved in hydrofluoric acid.

[32] _Arsenic acid_, H_{3}AsO_{4}, corresponding with orthophosphoric
acid, is formed by oxidising arsenious anhydride with nitric acid,
and evaporating the resultant solution until it attains a sp. gr.
of 2·2; on cooling it separates in crystals having the above
composition. This hydrate corresponds with the normal salts of
arsenic acid; but on dissolving in water (without heating), and on
cooling a strong solution, crystals containing a greater amount of
water, namely, (AsH_{3}O_{4})_{2},H_{2}O, separate. This water,
like water of crystallisation, is very easily expelled at 100°. At
120° crystals having a composition identical with that of
pyrophosphoric acid, As_{2}H_{4}O_{7}, separate, but water, on
dissolving this hydrate with the development of heat, forms a
solution in no way differing from a solution of ordinary arsenic
acid, so that it is not an independent pyroarsenic acid that is
formed. Neither is there any true analogue of metaphosphoric acid,
although the compound AsHO_{3} is formed at 200°, and on
solidifying forms a mass having a pearly lustre and sparingly
soluble in cold water; but on coming into contact with warm water
it becomes very hot, and gives ordinary orthoarsenic acid in
solution. Arsenic acid forms three series of salts, which are
perfectly analogous to the three series of orthophosphates. Thus
the normal salt, K_{3}AsO_{4}, is formed by fusing the other
potassium arsenates with potassium carbonate; it is soluble in
water and crystallises in needles which do not contain water.
Di-potassium arsenate, K_{2}HAsO_{4}, is formed in solution by
mixing potassium carbonate and arsenic acid until carbonic
anhydride ceases to be evolved; it does not crystallise, and has
an alkaline reaction; hence it corresponds perfectly with the
sodium phosphate. As was mentioned above, arsenic acid itself acts
as an oxidising agent; for example, it is used in the manufacture
of aniline dyes for oxidising the aniline, and it is prepared in
large quantities for this purpose. When sulphuretted hydrogen is
passed through its solution, sulphuric acid and arsenious
anhydride are obtained in solution. Arsenic acid is very easily
soluble in water, and its solution has an exceedingly acid
reaction, and when boiled with hydrochloric acid evolves chlorine,
like selenic, chromic, manganic, and certain other higher metallic
acids.

_Arsenic anhydride_, As_{2}O_{5}, is produced when arsenic acid is
heated to redness. It must be carefully heated, as at a bright red
heat it decomposes into oxygen and arsenious anhydride. Arsenic
anhydride is an amorphous substance almost entirely insoluble in
water, but it attracts moisture from the air, deliquesces, and
passes into the acid. Hot water produces this transformation with
great ease.

_Arseniuretted hydrogen_, _arsine_, AsH_{3}, resembles phosphuretted hydrogen in many respects. This colourless gas, which liquefies into a mobile liquid at -40°, has a disagreeable garlic-like odour, is only slightly soluble in water, and is exceedingly poisonous. Even in a small quantity it causes great suffering, and if present to any considerable amount in air it even causes death. The other compounds of arsenic are also poisonous, with the exception of the insoluble sulphur compound and some compounds of arsenic acid. Arseniuretted hydrogen, AsH_{3}, is obtained by the action of water on the alloy of arsenic and sodium, sodium hydroxide and arseniuretted hydrogen being formed. It is also formed by the action of sulphuric acid on the alloy of arsenic and zinc: Zn_{3}As_{2} + 3H_{2}SO_{4} = 2AsH_{3} + 3ZnSO_{4}.[33] The oxygen compounds of arsenic are very easily reduced by the action of hydrogen at the moment of its evolution from acids, and the reduced arsenic then combines with the hydrogen; hence, if a certain amount of an oxygen compound of arsenic be put into an apparatus containing zinc and sulphuric acid (and thus serving for the evolution of hydrogen), the hydrogen evolved will contain arseniuretted hydrogen. In this case it is diluted with a considerable amount of hydrogen. But its presence in the most minute quantities may be easily recognised from the fact that it is _easily decomposed_ by heat (200° according to Brunn) into metallic arsenic and hydrogen, and therefore if such impure hydrogen he passed through a moderately-heated tube metallic arsenic will be deposited as a bright layer on the part of the tube which was heated (_see_ Note 30 bis). This reaction is so sensitive that it enables the most minute traces of arsenic to be discovered; hence it is employed in medical jurisprudence, as a test in poisoning cases. It is easy to discover the presence of arsenic in common zinc, copper, sulphuric and hydrochloric acids, _&c._ by this method. It is obvious that in testing for poison by Marsh's apparatus it is necessary to take zinc and sulphuric acid quite free from arsenic. The arsenic deposited in the tube may be driven as a volatile metal from one place to another in the current of hydrogen evolved, owing to its volatility. This forms a distinction between arseniuretted and antimoniuretted hydrogen, which is decomposed by heat in just the same way as arseniuretted hydrogen, but the mirror given by Sb is not so volatile as that formed by As.

[33] The formation of arseniuretted hydrogen is accompanied by the
absorption of 37,000 heat units, while phosphine evolves 18,000
(Ogier), and ammonia 27,000. Sodium (0·6 p.c.) amalgam, with a
strong solution of As_{2}O_{3}, gives a gas containing 86 vols. of
arsenic and 14 vols. of hydrogen (Cavazzi).

If hydrogen contains arseniuretted hydrogen, it also gives metallic arsenic when it burns, because in the reducing flame of hydrogen the oxygen attracted combines entirely with the hydrogen and not with the arsenic, so that if a cold object, such as a piece of china, be held in the hydrogen flame the arsenic will be deposited upon it as a metallic spot.[34]

[34] This spot, or the metallic ring which is deposited on the heated
tube, may easily be tested as to whether it is really due to
arsenic or proceeds from some other substance reduced in the
hydrogen flame--for instance, carbon or antimony. The necessity
for distinguishing arsenic from antimony is all the more
frequently encountered in medical jurisprudence, from the fact
that preparations of antimony are very frequently used as
medicine, and antimony behaves in the hydrogen apparatus just like
arsenic, and therefore in making an investigation for poisoning by
arsenic it is easy to mistake it for antimony. The best method to
distinguish between the metallic spots of arsenic and antimony is
to test them with a solution of sodium hypochlorite, free from
chlorine, because this will dissolve arsenic and not antimony.
Such a solution is easily obtained by the double decomposition of
solutions of sodium carbonate and bleaching powder. A solution of
potassium chlorate acts in the same manner, only more slowly.
Further particulars must be looked for in analytical works.

Arseniuretted hydrogen, like phosphuretted hydrogen, is only
slightly soluble in water, has no alkaline properties--that is, it
does not combine with acids--and acts as a reducing agent. When
passed into a solution of silver nitrate it gives a blackish brown
precipitate of metallic silver, the arsenic being oxidised. If
acting on copper sulphate and similar salts, arseniuretted
hydrogen sometimes forms arsenides--_i.e._ it reduces the metallic
salt with its hydrogen, and is itself reduced to arsenic.
Sulphuric, and even hydrochloric, acid reduces arseniuretted
hydrogen to arsenic, and it is still more easily decomposed by
arsenious chloride, and with phosphorous chloride it gives the
compound PAs. Arseniuretted hydrogen gives metallic arsenic with
an acid solution of arsenious anhydride (Tivoli).

The most common compound of arsenic is the solid and volatile _arsenious anhydride_, As_{2}O_{3}, which corresponds with phosphorous and nitrous anhydrides. This very poisonous, colourless, and sweet-tasting substance is generally known under the name of arsenic, or _white arsenic_. The corresponding hydrate is as yet unknown; its solutions, when evaporated, yield crystals of arsenious anhydride. It is chiefly prepared for the dyer, and is also used as a vermin killer, and sometimes in medicine; it is a product from which all other compounds of arsenic can be prepared. It is obtained as a by-product in roasting cobalt and other ores containing arsenic. Arsenical pyrites are sometimes purposely roasted for the extraction of arsenious anhydride. When arsenical ores are burnt in the air, the sulphur and arsenic are converted into the oxides As_{2}O_{3} and SO_{2}. The former is a solid at the ordinary temperature, and the latter gaseous, and therefore the arsenious anhydride is deposited as a sublimate in the cooler portion of the flues through which the vapours escape from the furnace. It collects in condensing chambers especially constructed in the flues. The deposit is collected, and after being distilled gives arsenious anhydride in the form of a vitreous non-crystalline mass. This is one of the varieties of arsenious anhydride, which is also known in two crystalline forms. When sublimed--_i.e._ when it rapidly passes from the state of vapour to the solid state--it appears in the regular system in the form of octahedra.[35] It is obtained in the same form when it is crystallised from acid solutions. The specific gravity of the crystals is 3·7. The other crystalline form (in prisms) belongs to the rhombohedral system, and is also formed by sublimation when the crystals are deposited on a heated surface, or when it is crystallised from alkaline solutions.[36]

[35] According to Mitscherlich's determination, the vapour density of
arsenious anhydride is 199 (H = 1)--that is, it answers to the
molecular formula As_{4}O_{6}. Probably this is connected with the
fact that the molecule of free arsenic contains As_{4}. V. Meyer
and Biltz, however, showed (1889) that at a temperature of about
1,700° the vapour density of arsenic corresponds with the molecule
As_{2}, and not As_{4}, as at lower temperatures.

[36] Arsenious anhydride is obtained in an amorphous form after
prolonged heating at a temperature near to that at which it
volatilises, or, better still, by heating it in a closed vessel.
It then fuses to a colourless liquid, which on cooling forms a
transparent vitreous mass, whose specific gravity is only slightly
less than that of the crystalline anhydride. On cooling, this
vitreous mass undergoes an internal change, in which it
crystallises and becomes opaque, and acquires the appearance of
porcelain. The following difference between the vitreous and
opaque varieties is very remarkable: when the vitreous variety is
dissolved in strong and hot hydrochloric acid it gives crystals of
the anhydride on cooling, and this crystallisation _is accompanied
by the emission of light_ (which is visible in the dark), and the
entire liquid glows as the crystals begin to separate. The opaque
variety does not emit light when the crystals separate from its
hydrochloric acid solution. It is also remarkable that the
vitreous variety passes into the opaque form when it is
pounded--that is, under the action of a series of blows. Thus,
several varieties of arsenious anhydride are known, but as yet
they are not characterised by any special chemical distinctions,
and even differ but little in their specific gravities, so that it
cannot be said that the above differences are due to any isomeric
transformation--that is, to an arrangement of the atoms in the
molecule--but probably only depend on a difference in the
distribution of the molecules, or, in other terms, are physical
and not chemical variations. One part of the vitreous anhydride
requires twelve parts of boiling water for its solution, or
twenty-five parts at the ordinary temperature. The opaque variety
is less soluble, and at the ordinary temperature requires about
seventy parts of water for its solution.

Solutions of arsenious anhydride have a sweet metallic taste, and give _a feeble acid reaction_. Its solubility increases with the admixture of acids and alkalis. This shows the property of arsenious anhydride of forming salts with acids and alkalis. And in fact compounds of it with hydrochloric acid (Note 31), sulphuric anhydride (_see_ further on), and with the alkali oxides are known.[37] If silver nitrate be added to a solution of arsenious anhydride, it does not give any precipitate unless a certain amount of the arsenious anhydride is saturated with an alkali--for instance, ammonia. It then gives a precipitate of silver arsenite, Ag_{3}AsO_{3}. This is yellow, soluble in an excess of ammonia, and anhydrous; it distinctly shows that arsenious acid is tribasic, and that it differs in this respect from phosphorous acid, in which only two atoms of hydrogen can be replaced by metals.[38] The feeble acid character of arsenious anhydride is confirmed by the formation of saline compounds with acids. In this respect the most remarkable example is the anhydrous compound with sulphuric acid, having the composition As_{2}O_{3},SO_{3}. It is formed in the roasting of arsenical pyrites in those spaces where the arsenious anhydride condenses, a portion of the sulphurous anhydride being converted into sulphuric anhydride, SO_{3}, at the expense of the oxygen of the air. The compound in question forms colourless tabular crystals, which are decomposed by water with formation of sulphuric acid and arsenious anhydride.[39]

[37] Arsenious anhydride does not oxidise in air, either in a dry state
or in solution, but in the presence of alkalis it absorbs oxygen
from the air, and acts as an excellent reducing agent. This
probably is connected with the fact that arsenic acid is much more
energetic than arsenious acid, and that it is arsenic acid which
is formed by the oxidation of the latter in the presence of
alkalis. Arsenious anhydride is easily reduced to arsenic by many
metals, even by copper.

[38] The feebleness of the acid properties of arsenious anhydride is
seen in the fact that if it be dissolved in ammonia water, and
then a still stronger solution of ammonia be added, prismatic
crystals separate having the composition of ammonium metarsenite,
NH_{4}AsO_{3}. This ammonium salt deliquesces in air, and loses
all its ammonia. The magnesium salt is tri-metallic,
Mg_{3}(AsO_{3})_{2}; it is insoluble in water, and is formed by
mixing an ammoniacal solution of arsenious anhydride with an
ammoniacal solution of a magnesium salt. It is insoluble even in
ammonia, although it dissolves in an excess of acids. Magnesium
hydroxide gives the same salt with arsenious solutions, and hence
magnesia is one of the best antidotes for arsenic poisoning. _The
arsenites of copper_ are much used in the manufacture of colours,
more especially of pigments. They are distinguished by their
insolubility in water and by their remarkably vivid green colour,
but at the same time by their poisonous character. Not only do
such pigments applied to wall papers or other materials easily
dust off from them, but they give exhalations containing AsH_{3}.
The cupric salts, CuX_{2}, when mixed with an alkaline solution of
arsenious acid, give a green precipitate of a copper salt called
_Scheele's green_. Its composition is probably CuHAsO_{3}. Ammonia
dissolves it, and gives a colourless solution, containing cuprous
arsenate--that is, the cupric compound is reduced and the arsenic
subjected to a further oxidation. The so-called _Schweinfurt
green_ was still more used, especially in former times; it is an
insoluble green cupric salt, which resembles the preceding in many
respects, but has a different tint. It is prepared by mixing
boiling solutions of arsenious acid and cupric acetate. Arsenious
acid forms an insoluble compound with ferric hydroxide, resembling
the phosphate; and this is the reason why freshly precipitated
oxide of iron is employed as an _antidote for arsenic_. The
freshly precipitated oxide of iron, taken immediately after
poisoning by arsenic, converts the arsenious acid into an
insoluble state, by forming a compound on which the acids of the
stomach have no action, so that the poisoning cannot proceed. It
is remarkable that the inhabitants of certain mountainous
countries accustom themselves to taking arsenic, as a means which,
according to their experience, helps to overcome the fatigue of
mountain ascents. Arsenious anhydride and certain of its salts are
also used in medicine, naturally only in small quantities. When
taken internally arsenic passes into the blood, and is mainly
excreted by the urine.

[39] Adie (1889) obtained compounds of As_{2}O_{3} with 1, 2, 4, and 8
SO_{3} by the direct action of ordinary and Nordhausen sulphuric
acid upon As_{2}O_{3}. Weber had previously obtained
As_{2}O_{3}SO_{3} (which disengages SO_{3} at 225°), and also
other As_{2}O_{3}_n_SO_{3} (where _n_ = 3, 6, and 8), by the
action of the vapours of SO_{3} upon As_{2}O_{3} at a definite
temperature. The compound As_{2}O_{3},8SO_{3} loses SO_{3} at
100°. Oxide of antimony, Sb_{2}O_{3}, gives similar compounds.
Adie (1891) also obtained (by the action of SO_{3} upon
H_{3}PO_{4}) a compound H_{3}PO_{4}3SO_{3} in the form of a
viscous liquid decomposed by water.

_Antimony_ (stibium), Sb = 120, is another analogue of phosphorus. In its external appearance and the properties of its compounds it resembles the metals still more closely than arsenic. In fact, antimony has the appearance, lustre, and many of the characteristic properties of the metals. Its oxide, Sb_{2}O_{3}, exhibits the earthy appearance of rust or of lime, and has distinctly basic properties, although it corresponds with nitrous and phosphorous anhydride, and is able, like them, to give saline compounds with bases. At the same time antimony presents, in the majority of its compounds, an entire analogy with phosphorus and arsenic. Its compounds belong to the type SbX_{3} and SbX_{5}. It is found in nature chiefly in the form of sulphide, Sb_{2}S_{3}. This substance sometimes occurs in large masses in mineral veins and is known in mineralogy under the name of antimony glance or _stibnite_, and commercially as _antimony_ (Chapter XX., Note 29). The most abundant deposits of antimony ore occur in Portugal (near Oporto on the Douro). Besides which antimony partially or totally replaces arsenic in some minerals; thus, for example, a compound of antimony sulphide and arsenic sulphide with silver sulphide is found in red silver ore. But in every case antimony is a rather rare metal found in few localities. In Russia it is known to occur in Daghestan in the Caucasus. It is extracted chiefly for the preparation of alloys with lead and tin, which are used for casting printing type.[40] Some of its compounds are also used in medicine, the most important in this respect being antimony pentasulphide, Sb_{2}S_{5} (_sulfur auratum antimonii_), and tartar emetic, which is a double salt derived from tartaric acid and has the composition C_{4}H_{4}K(SbO)O_{6}. Even the native antimony sulphide is used in large quantities as a purgative for horses and dogs. Metallic antimony is extracted from the glance, Sb_{2}S_{2}, by roasting, when the sulphur burns away and the antimony oxidises, forming the oxide Sb_{2}O_{3}, which is then heated with charcoal, and thus reduced to a _metallic state_. The reduction may be carried on in the laboratory on a small scale by fusing the sulphide with iron which takes up the sulphur.[40 bis]

[40] Printers' type consists of an alloy known as 'type-metal,'
containing usually about 15 parts of antimony to 85 parts of lead;
sometimes (for example, for stereotypes) from 10 to 15 per cent.
Bi or 8 per cent. Sn and even Cu is added. The hardness of the
alloy, which is essential for printing, evidently depends upon the
presence of antimony, but an excess must be avoided, since this
renders the alloy brittle, and the type after a time loses its
sharpness.

[40 bis] Antimony is prepared in a state of greater purity by heating
with charcoal the oxide obtained by the action of nitric acid on
the impure commercial metallic antimony. This is based on the fact
that by the action of the acid, antimony forms the oxide
Sb_{2}O_{3}, which is but slightly soluble in water. The arsenic,
which is nearly always present, forms soluble arsenious and
arsenic acids, and remains in solution. The purest antimony is
easily obtained from tartar emetic, by heating it with a small
quantity of nitre. Metallic antimony also occurs, although rarely,
native; and as it is very easily obtained, it was known to the
alchemists of the fifteenth century. Very pure metallic antimony
may be deposited by the electric current from a solution of
antimonious sulphide in sodium sulphide after the addition of
sodium chloride to the solution.

Metallic antimony has a white colour and a brilliant lustre; it remains untarnished in the air, for the metal does not oxidise at the ordinary temperature. It crystallises in rhombohedra, and always shows a distinctly crystalline structure which gives it quite a different aspect from the majority of the metals yet known. It is most like tellurium in this respect. Antimony is brittle, so that it is very easily powdered; its specific gravity is 6·7, it melts at about 432°, but only volatilises at a bright red heat. When heated in the air--for instance, before the blow-pipe--it burns and gives white odourless fumes, consisting of the oxide. This oxide is termed antimonious oxide, although it might as well be termed antimonious anhydride. It is given the first name because in the majority of cases its compounds with acids are used, but it forms compounds with the alkalis just as easily.

Antimonious oxide, like arsenious anhydride, crystallises either in regular octahedra or in rhombic prisms; its specific gravity is 5·56; when heated it becomes yellow and then fuses, and when further heated in air it oxidises, forming an oxide of the composition Sb_{2}O_{4}. Antimonious oxide is insoluble in water and in nitric acid, but it easily dissolves in strong hydrochloric acid and in alkalis, as well as in tartaric acid or solutions of its acid salts. When dissolved in the latter it forms tartar emetic. It is precipitated from its solutions in alkalis and acids (by the action of acids on the former and alkalis on the latter). It occurs native but rarely. As a base it gives salts of the type SbOX (as if the basic salts = SbX_{3}, Sb_{2}O_{3}) and hardly ever forms salts, SbX_{3}. In the antimonyl salts, SbOX, the group SbO is univalent, like potassium or silver. The oxide itself is (SbO)_{2}O, the hydroxide, SbO(OH), &c.; tartar emetic is a salt in which one hydrogen of tartaric acid is replaced by potassium and the other by antimonyl, SbO. Antimonious oxide is very easily separated from its salts by any base, but it must be observed that this separation does not take place in the presence of tartaric acid, owing to the property of tartaric acid of forming a soluble double salt--_i.e._ tartar emetic.[41]

[41] As antimonious oxide answers to the type SbX_{3}, it is evident
that compounds may exist in which antimony will replace three
atoms of hydrogen; such compounds have been to some extent
obtained, but they are easily converted by water into substances
corresponding with the ordinary formulæ of the compounds of
antimony. Thus tartar emetic, C_{4}H_{4}(SbO)KO_{6}, loses water
when heated, and forms C_{4}H_{2}SbKO_{6}--that is, tartaric acid,
C_{2}H_{6}O_{6}, in which one atom of hydrogen is replaced by
potassium and three by antimony. But this substance is reconverted
into tartar emetic by the action of water.

A similar compound is seen in that _intermediate oxide of
antimony_ which is formed when antimonious oxide is heated in air:
its composition is SbO_{2} or Sb_{2}O_{4}. This oxide may be
regarded as orthantimonic acid, SbO(HO)_{3}, in which three atoms
of hydrogen are replaced by antimony in that state in which it
occurs in oxide of antimony--_i.e._ SbO(SbO_{3}) = Sb_{2}O_{4}.
Oxide of antimony is also formed when antimonic acid is ignited;
it then loses water and oxygen, and gives this intermediate oxide
as a white infusible powder, of sp. gr. 6·7. It is somewhat
soluble in water, and gives a solution which turns litmus paper
red.

If metallic antimony, or antimonious oxide, be oxidised by an excess of nitric acid and the resultant mass be carefully evaporated to dryness, _metantimonic acid_, SbHO_{3}, is formed. Its corresponding potassium salt, 2SbKO_{3},5H_{2}O, is prepared by fusing metallic antimony with one-fourth its weight of nitre and washing the resultant mass with cold water. This potassium salt is only slightly soluble in water (in 50 parts) and the sodium salt is still less so. An ortho-acid, SbH_{3}O_{4}, also appears to exist;[41 bis] it is obtained by the action of water on antimony pentachloride, but it is very unstable, like the pentachloride, SbCl_{5}, itself, which easily gives up Cl_{2}, leaving antimony trichloride, SbCl_{3}, and this is decomposed by water, forming an oxychloride--SbOCl, only slightly soluble in water. When antimonic acid is heated to an incipient red heat, it parts with water and forms the anhydride, Sb_{2}O_{5}, of a yellow colour and specific gravity 6·5.[42]

[41 bis] Beilstein and Blaese (1889), after preparing many salts of
antimonic acid, came to the conclusion that it is monobasic, but
all the salts still contain water, so that their general type is
mostly: MSbO_{3}3H_{2}O, for example, M = Li, Hg (salts of the
suboxide), 1/2 Pb, &c. The type of the ortho-salts, M_{2}SbO_{4},
is quite unknown, although it is reproduced in the thio-compounds,
for instance, Schlippe's salt, Na_{2}SbS_{4}, but this salt also
contains water of crystallisation, 9H_{2}O (Chapter XX., Note 29).

[42] Among the other compounds of antimony, _antimoniuretted hydrogen_,
SbH_{3}, resembles arseniuretted hydrogen in its mode of formation
and properties (it splits up at 150°, Brunn 1890; when liquified,
it boils at -65° and solidifies at -92°), whilst the halogen
compounds differ in many respects from those of arsenic. When
chlorine is passed over an excess of antimony powder, it forms
_antimony trichloride_, SbCl_{3}, but if the chlorine be in excess
it forms the _pentachloride_, SbCl_{5}. The trichloride is a
crystalline substance which melts at 72° and distils at 230°,
whilst the pentachloride is a yellow liquid, which splits up into
chlorine and the trichloride when heated; at 140° it begins to
give off chlorine abundantly, carrying away the vapour of the
trichloride with it; and at 200° the decomposition is complete,
and pure antimonious chloride only passes over. This property of
antimony pentachloride has caused it to be applied in many cases
for the transference of chlorine; all the more that when it has
given up its chlorine, it leaves the trichloride, which is able to
absorb a fresh amount of chlorine; and therefore many substances
which are unable to react directly with gaseous chlorine do so
with antimony pentachloride, and in the presence of a small
quantity of it chlorine will act on them, just as oxygen is able,
in the presence of nitrogen oxides, to oxidise substances which
could not be oxidised by means of free oxygen. Thus carbon
bisulphide is not acted on by chlorine at low temperatures--this
reaction requires a high temperature--but in the presence of
antimony pentachloride its conversion into carbon tetrachloride
takes place at low temperatures. Antimony tri- and pentachloride,
having the character of chloranhydrides, fume in air, attract
moisture, and are decomposed by water, forming antimonious and
antimonic acids. But in the first action of water the trichloride
does not evolve all its chlorine as hydrochloric acid, which is
intelligible in view of the fact that antimonious anhydride is
also a base, and is therefore able to react with acids; indeed
antimony sulphide dissolved in an excess of hydrochloric acid
(hydrogen sulphide is evolved) gives an aqueous solution of
antimony trichloride, which, when carefully distilled, even gives
the anhydrous compound. Antimony trichloride is only decomposed by
an excess of water, and then not completely, for with a large
quantity of water it forms _powder of algaroth_--_i.e._ antimony
oxychloride. The first action of water consists in the formation
of _oxychloride_, SbOCl--that is, a salt corresponding to oxide of
antimony as a base. If antimony oxide or antimony chloride be
dissolved in an excess of hydrochloric acid, and the solution
diluted with a considerable amount of water, then this same powder
of algaroth is precipitated. The composition varies with the
relative amount of water; namely, between the limits SbOCl and
Sb_{4}O_{5}Cl_{2}. The latter compound is, as it were, a basic
salt of the former, because its composition = 2(SbOCl)Sb_{2}O_{3}.

With bromine and iodine, antimony forms compounds similar to those
with chlorine. Antimonious bromide, SbBr_{3}, crystallises in
colourless prisms, melts at 94°, and boils at 270°; antimonious
iodide, SbI_{3}, forms red crystals of sp. gr. 5·0; antimony
trifluoride, SbF_{3} separates from a solution of antimonious
oxide in hydrofluoric acid, and SbF_{5} is formed by a similar
treatment of antimonic acid. The latter gives easily-soluble
double salts with the fluorides of the metals of the alkalis.

De Haën (1887) obtained very stable double soluble salts,
SbF_{3},KCl (100 parts of water dissolve 57 parts of salt),
SbF_{3},K_{2}SO_{4}, &c., which he proposed to make use of in the
arts as very easily crystallisable and soluble salts of antimony.

Engel, by passing hydrochloric acid gas into a saturated solution
of antimonious chloride at 0°, obtained a compound
HCl,2SbCl_{3},2H_{2}O, and with the pentachloride a compound
SbCl_{5},5HCl,10H_{2}O. Bismuth trichloride, BiCl_{3}, gives a
similar compound.