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.
[Illustration: FIG. 84.--Formation and decomposition of arseniuretted
hydrogen. Hydrogen is evolved in the Woulfe's bottle, and when the gas
comes off, a solution containing arsenic is poured through the funnel.
The presence of AsH_{3} is recognised from the deposition of a mirror of
arsenic when the gas-conducting tube is heated. If the escaping hydrogen
be lighted, and a porcelain dish be held in the flame, a film of arsenic
is deposited on it. The gas is dried by passing through the tube
containing calcium chloride. This apparatus is used for the detection of
arsenic by Marsh's test.]
[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.
Saunders (1892) obtained 5RbCl,3SbCl_{3} and RbCl,SbCl_{3}. Ditte
and Metzner (1892) showed that Sb and Bi dissolve in hydrochloric
acid only owing to the participation of the oxygen of the air or
of that dissolved in the acid.
The heaviest analogue of nitrogen and phosphorus is _bismuth_, Bi = 208.
Here, as in the other groups, the basic, metallic, properties increase
with the atomic weight. Bismuth does not give any hydrogen compound and
the highest oxide, Bi_{2}O_{5}, is a very feeble acid oxide. Bismuthous
oxide, Bi_{2}O_{3}, is a base, and bismuth itself a perfect metal. To
explain the other properties of bismuth it must further be remarked that
in the eleventh series it follows mercury, thallium and lead, whose
atomic weights are near to that of bismuth, and that therefore it
resembles them and more especially its nearest neighbour, lead. Although
PbO and PbO_{2}, represent types different from Bi_{2}O_{3} and
Bi_{2}O_{5}, they resemble them in many respects, even in their external
appearance, moreover the lower oxides both of Pb and Bi are basic and the
higher acid, which easily evolve oxygen. But judging by the formula,
Bi_{2}O_{3} is a more feeble base than PbO. They both easily give basic
salts.
Bismuth forms compounds of two types, BiX_{3} and BiX_{5},[43] which
entirely recall the two types we have already established for the
compounds of lead. Just as in the case of lead, the type PbX_{2}, is
basic, stable, easily formed, and passes with difficulty into the higher
and lower types, which are unstable, so also in the case of bismuth the
type of combination BiX_{3} is the usual basic form. The higher type of
combination, BiX_{5},[44] in fact behaves toward this stable type,
BiX_{3}, in exactly the same manner as lead dioxide does to the monoxide;
and bismuthic acid is obtained by the action of chlorine on bismuth oxide
suspended in water, in exactly the same way as lead dioxide is obtained
from lead oxide. It is an oxidising agent like lead dioxide, and even the
acid character in bismuthic acid is only slightly more developed than in
lead dioxide. Here, as in the case of lead (minium), intermediate
compounds are easily formed in which the bismuth of the lower oxide plays
the part of a base combined with the acid which is formed by the higher
form of the oxidation of bismuth.
[43] Metallic bismuth is very easily obtained when the compounds of the
oxide are reduced by powerful reducing agents, but when less
powerful reducing agents--for example, stannous oxide--are taken,
bismuth suboxide is formed as a black crystalline powder. It is a
compound of the type BiX_{2}, its composition being BiO; it is
decomposed by acids into the metal and oxide, which passes into
solution.
[44] The type BiX_{5} is represented by the pentoxide, Bi_{2}O_{5}, its
metahydrate, Bi_{2}O_{5},H_{2}O, or BiHO_{3}, known as bismuthic
acid, and the pyrohydrate, Bi_{2}H_{4}O_{7}. _Bismuth pentoxide_
is obtained by the prolonged passage of chlorine through a boiling
solution of potassium hydroxide (sp. gr. 1·38), containing bismuth
oxide in suspension; the precipitate is washed with water, with
boiling nitric acid (but not for long, as otherwise the bismuthic
acid is decomposed), then again with water, and finally the
resultant bright red powder of the hydrate BiHO_{3} is dried at
125°. The prolonged action of nitric acid on bismuthic anhydride,
Bi_{2}O_{5}, results in the formation of the compound
Bi_{2}O_{4},H_{2}O, which decomposes in moist air, forming
Bi_{2}O_{3}. The density of bismuthic anhydride is 5·10, of the
tetroxide, Bi_{2}O_{4}, 3·60, and of bismuthic acid, BiHO_{3},
5·75. _Pyrobismuthic acid_, Bi_{2}H_{4}O_{7}, forms a brown
powder, which loses a portion of its water at 150°, and decomposes
on further heating, with the evolution of oxygen and water. It is
obtained by the action of potassium cyanide on a solution of
bismuth nitrate. The meta-salts of bismuthic acid are known, for
example KBiO_{3}. They generally occur, however, in combinations
with metabismuthic acid itself. Thus André (1891) took a solution
of the double salt of BiBr_{3} and KBr, treated it with bromine
after adding ammonia, and obtained a red-brown precipitate, which
after being washed (for several weeks) had the composition
KBiO_{3},HBiO_{3} When washed with dilute nitric acid this salt
gave bismuthic acid.
[Illustration: FIG. 85.--Furnace used for the extraction of bismuth from
its ores.]
In nature, bismuth occurs in only a few localities and in small
quantities, most frequently in a native state, and more rarely as oxide
and as a compound of bismuth sulphide with the sulphides of other metals,
and sometimes in gold ores. It is extracted from its native ores by
simple fusion in the furnace shown in fig. 85. This furnace contains an
inclined iron retort, into the upper extremity of which the ore is
charged, and the molten _metal_ flows from the lower extremity. It is
refined by re-melting, and the pure metal may be obtained by dissolving
in nitric acid, decomposing the resultant salt with water, and reducing
the precipitate by heating it with charcoal. Bismuth is a metal which
crystallises very well from a molten state. Its specific gravity is 9·8;
it melts at 269°, and if it be melted in a crucible, allowed to cool
slowly, and the crust broken and the remaining molten liquid poured out,
perfect rhombohedral crystals of bismuth are obtained on the sides of the
crucible.[44 bis] It is brittle, has a grey-coloured fracture with a
reddish lustre, is not hard, and is but very slightly ductile and
malleable; it volatilises at a white heat and easily oxidises. It recalls
antimony and lead in many of its properties. When oxidised in air, or
when the nitrate is ignited, bismuth forms the _oxide_, Bi_{2}O_{3}, as a
white powder which fuses when heated and resembles massicot. The addition
of an excess of caustic potash to a solution of a bismuthous salt gives a
white precipitate of the hydroxide, BiO(OH), which loses its water and
gives the anhydrous oxide when boiled with a solution of caustic potash.
Both the hydroxide and oxide easily dissolve in acids and form bismuthous
salts.
[44 bis] Hérard (1889) obtained a peculiar variety of bismuth by
heating pure crystalline bismuth to a bright red heat in a stream
of nitrogen. A greenish vapour was deposited in the cooler
portions of the apparatus in the form of a grey powder, which
under the microscope had the appearance of minute globules. An
atmosphere of nitrogen is necessary for this transformation, other
gases such as hydrogen and carbonic oxide do not favour the
transition. The resultant amorphous bismuth fuses at 410° (the
crystalline variety at 269°), sp. gr. 9·483. (Does it not contain
a nitride?)
_Bismuthous oxide_, Bi_{2}O_{3}, is a feeble and unenergetic base. The
normal hydroxide of the oxide Bi_{2}O_{3} is Bi(OH)_{3}; it parts with
water and forms a metahydroxide (bismuthyl hydroxide), BiO(OH). Both of
these hydroxides have their corresponding saline compounds of the
composition BiX_{3} and BiOX. And the form BiOX is nothing else but the
type of the basic salt, because 3ROX = RX + R_{2}O_{3}. It is evident
that in the type BiX_{3} the bismuth replaces three atoms of hydrogen.
And indeed with phosphoric acid solutions of the bismuthous salts give a
precipitate of the composition BiPO_{4}. On the other hand, in the form
of compounds BiOX or Bi(OH)_{2}X, the univalent group (BiO) or
(BiH_{2}O_{2}) is combined with X. Many bismuth salts are formed
according to the type BiOX. For instance the carbonate, (BiO)_{2}CO_{3},
which corresponds with the other carbonates M_{2}CO_{3}. It is obtained
as a white precipitate when a solution of sodium carbonate is added to a
solution of a bismuth salt.[45] The compound radicle BiO is not a special
natural grouping, as it was formerly represented to be; it is simply a
mode of expression for showing the relation between the compound in
question and the compounds of other oxides.
[45] Basic bismuth carbonate is employed for whitening the skin
(veloutine, &c.)
Three _salts of nitric acid_ are known containing bismuthous oxide. If
metallic bismuth or its oxide be dissolved in nitric acid, it forms a
colourless transparent solution containing a salt which separates in
large transparent crystals containing Bi(NO_{3})_{3},5H_{2}O. When heated
at 80° these crystals melt in their water of crystallisation, and in so
doing lose a portion of their nitric acid together with water, forming a
salt whose empirical formula is Bi_{2}N_{2}H_{2}O_{9}. If the preceding
salt belongs to the type BiX_{3}, this one should belong to the form
BiOX, because it = BiO(NO_{3}) + Bi(H_{2}O_{2})(NO_{3}). This salt may be
heated to 150° without change. When the first colourless crystalline salt
dissolves in water _it is decomposed_. There is no decomposition if an
excess of acid be added to the water--that is to say, the salt is able to
exist in an acid solution without decomposing, without separation of the
so-called basic salt--but by itself it cannot be kept in solution; water
decomposes this salt, acting on it like an alkali. In other words the
basic properties of bismuthic oxide are so feeble that even water acts by
taking up a portion of the acid from it. Here we see one of the most
striking facts, long since observed, confirming that action of water on
salts about which we have spoken in Chapter X. and elsewhere. This action
on water may be expressed thus:--BiX_{3} + 2H_{2}O = Bi(OH)_{2}X + 2XH. A
salt of the type Bi(OH)_{2}X is obtained in the precipitate. But if the
quantity of acid, HX, be increased, the salt BiX_{3} is again formed and
passes into solution. The quantity of the salt BiOX which passes into
solution on the addition of a given quantity of acid depends indisputably
on the amount (mass) of water (Muir). The solution, which is perfectly
transparent with a small amount of water, becomes cloudy and deposits the
salt of the type BiOX, when diluted. The white flaky precipitate of
Bi(OH)_{2}NO_{3} formed from the normal salt Bi(NO_{3})_{3} by mixing it
with five parts of water, and in general with a small amount of water, is
used in medicine under the name of magistery of bismuth.[46]
[46] With an excess of water a further quantity of acid is separated
and a still more basic salt formed. The ultimate product, on which
an excess of water has apparently no action whatever, is a
substance having the composition BiO(NO_{3}).BiO(OH). In the
latter salt we see the limit of change, and this limit appears to
show that the type of the saline compounds of bismuthic oxide is
of the form Bi_{2}X_{6}, and not BiX_{3}; but it is very probable,
on the basis of the examples which we considered in the case of
lead, that this type should be still further polymerised in order
to give a correct idea of the type of the bismuthous compounds. If
we refer all the bismuthous compounds to this type, Bi_{2}X_{6},
we shall obtain the following expression for the composition of
the nitrates: normal salt, Bi_{2}(NO_{3})_{6}, first basic salt,
Bi_{2}O(OH)_{2}(NO_{3})_{2}, magistery of bismuth,
Bi_{2}(OH)_{4}(NO_{3})_{2}, and the limiting form
Bi_{2}O_{2}(OH)(NO_{3}).
The general character of bismuthous oxide in its compounds is well
exemplified in the nitrate; bismuthous chloride, BiCl_{3}, which
is obtained by heating bismuth in chlorine, or by dissolving it in
aqua regia, and then distilling without access of air, is also
decomposed by water in exactly the same manner, and forms basic
salts--for instance, first, BiOCl, like the above salt of nitric
acid. Bismuth chloride boils at 447° and probably its formula is
BiCl_{3}. Polymerisation may take place in some compounds and not
in others. A volatile compound of the composition
Bi(C_{2}H_{5})_{3} is also known as a liquid which is insoluble in
water and decomposes with explosion when heated at 130°. Double
salts containing chloride of bismuth are: 2(KCl)BiCl_{3}2H_{2}O
(from a solution of Bi_{2}O_{3} and KCl in hydrochloric acid) and
KClBiCl_{3}H_{2}O. Bigham (1892) also obtained KBr(SO_{4})_{2} in
tabular crystals by treating the above-named double salt with
strong sulphuric acid. The composition of this salt recalls that
of alum.
Metallic bismuth is used in the preparation of fusible alloys. The
addition of bismuth to many metals renders them very hard, and at the
same time generally lowers their melting point to a considerable extent.
Thus Wood's metal, which contains one part of cadmium, one part of tin,
two parts of lead, and four parts of bismuth, fuses at about 60°, and in
general many alloys composed of bismuth, tin, lead, and antimony melt
below or about the boiling point of water.[47]
[47] As the metals contained in alloys like the above (bismuth, lead,
tin, cadmium) are difficultly volatile and their alloys are
fusible, they may be employed in the place of mercury in many
physical experiments conducted at or above 70°, and they offer the
advantage that they do not give any vapour having an appreciable
tension (mercury at 100°, 0·75 mm.) Bismuth expands in passing
into a molten state, but it has a temperature of maximum density.
According to Luedeking the mean coefficient of expansion of liquid
bismuth is 0·0000442 (between 270° and 303°), and of solid bismuth
0·0000411.
Just as in group II., side by side with the elements zinc, cadmium, and
mercury in the uneven series, we found calcium, strontium, and barium in
the even series; and as in group IV., parallel to silicon, germanium,
tin, and lead, we noticed thallium, zirconium, cerium, and thorium; so
also in group V. we find, beside those elements of the uneven series just
considered by us, a series of analogues in the even series, which, with a
certain degree of similarity (mainly quantitative, or relative to the
atomic weights), also present a series of particular (qualitative)
independent points of distinction. In the even series are known
_vanadium_, which stands between titanium and chromium, _niobium_,
between zirconium and molybdenum, and _tantalum_, situated near tungsten
(an element of group VI. like chromium and molybdenum). Just as bismuth
is similar in many respects to its neighbour lead, so also do these
neighbouring elements resemble each other, even in their external
appearance, not to mention the quality of their compounds, naturally
taking into account the differences of type corresponding with the
different groups. The occurrence in group V. determines the type of the
oxides, R_{2}O_{3} and R_{2}O_{5}, and the development of an acid
character in the higher oxides. The occurrence in the even series
determines the absence of volatile compounds, RH_{3}, for these metals,
and a more basic character of the oxides of a given composition than in
the uneven series, &c.[48] Vanadium, niobium, and tantalum belong to the
category of rare metals, and are exceedingly difficult to obtain pure,
more especially owing to their similarity to, and occurrence with,
chromium, tungsten and other metals, and also in combination among
themselves; therefore it is natural that they have been far from
completely studied, although since 1860 chemists have devoted not a
little time to their investigation. The researches carried out by
Marignac, at Geneva, on niobium, and by Sir Henry Roscoe, at Manchester,
on vanadium deserve special attention. The undoubted external resemblance
of the compounds of chromium and vanadium, as well as the want of
completeness in the knowledge of the compounds of vanadium, long caused
its oxides to be considered analogous in atomic composition to those
formed by chromium. The higher oxide of vanadium was therefore supposed
to have the formula VO_{3}. But the fact of the matter is, that the
chemical analogy of the elements does not hold in one direction only;
vanadium is at one and the same time the analogue of chromium, and
consequently of the elements like sulphur of group VI, and also the
analogue of phosphorus, arsenic, and antimony; just as bismuth stands in
respect to lead and antimony. Investigation has shown that the compounds
of vanadium are always accompanied by those of phosphorus as well as of
iron, and that it is even more difficult to separate it from the
compounds of phosphorus than from those of iron and tungsten. We should
have to extend our description considerably if we wished to give the
complete history, even of vanadium alone, not to mention niobium and
tantalum, all the more as questions would not unfrequently arise
concerning the compounds of these elements which have not yet been fully
elucidated. We shall therefore limit ourselves to pointing out the most
important features in the history of these elements, the more so since
the minerals themselves in which they occur are exceedingly rare and only
accessible to a few investigators.
[48] Although, guided by Brauner, who showed that didymium gives a
higher oxide, Di_{2}O_{5}, I place this element in the fifth
group, still I am not certain as to its position, because I
consider that the questions relating to this metal are still far
from being definitely answered.
An important point in the history of the members of this group is the
circumstance that they form volatile compounds with chlorine, similar to
the compounds of the elements of the phosphorus group, namely, of the
type RX_{5}. The vapour densities of the compounds of this kind were
determined, and served as the most important basis for the explanation of
the atomic composition of these molecules. In this we see the power of
general and fundamental laws, like the law of Avogadro-Gerhardt. An
oxychloride, VOCl_{3}, is known for vanadium, which is the perfect
analogue of phosphorus oxychloride. It was formerly considered to be
vanadium chloride, for just as in the case of uranium (Chapter XXI.), its
lower oxide, VO, was considered to be the metal, because it is
exceedingly difficultly reduced--even potassium does not remove all the
oxygen, besides which it has a metallic appearance, and decomposes acids
like a metal; in a word, it simulates a metal in every respect. _Vanadium
oxychloride_ is obtained by heating the trioxide, V_{2}O_{3}, mixed with
charcoal, in a current of hydrogen; the lower oxide of vanadium is then
formed, and this, when heated in a current of dry chlorine, gives the
oxychloride VOCl_{3} as a reddish liquid which does not act on sodium and
may be purified by distillation over this metal. It fumes in the air,
giving reddish vapours; it reacts on water, forming hydrochloric and
vanadic acids; hence, on the one hand it is very similar to phosphorus
oxychloride, and on the other hand to chromium oxychloride, CrO_{2}Cl_{2}
(Chapter XXI.). It is of a yellow colour, its specific gravity is 1·83,
it boils at 120°, and its vapour density is 86 with respect to hydrogen;
therefore the above formula expresses its molecular weight.[49]
[49] When the vapours of vanadium oxychloride are heated with zinc in a
closed tube at 400°, they lose a portion of their chlorine and
form a green crystalline mass of sp. gr. 2·88, which is
deliquescent in air and has the composition VOCl_{2}. Only its
vapour density is unknown, and it would be extremely important to
determine whether its molecular composition is that given above,
or whether it corresponds with the formula V_{2}O_{2}Cl_{4}.
Another less volatile oxychloride, VOCl, is formed with it as a
brown insoluble substance, which is, however, soluble in nitric
acid like the preceding. Roscoe obtained a still less chlorinated
substance, namely, (VO)_{2}Cl; but it may only consist of a
mixture of VO and VOCl. At all events, we here find a graduated
series such as is met with in the compounds of very few other
elements.
_Vanadic anhydride_, V_{2}O_{5}, is obtained either in small quantities
from certain clays where it accompanies the oxides of iron (hence some
sorts of iron contain vanadium) and phosphoric acid, or from the rare
minerals: _volborthite_, CuHVO_{4}, or basic vanadate of copper;
_vanadinite_, PbCl_{2}3Pb_{3}(VO_{4})_{2}; lead vanadate,
Pb_{3}(VO_{4})_{2}, &c. The latter salts are carefully ignited for some
time with one-third of their weight of nitre; the fused mass thus formed
is powdered and boiled in water: the yellow solution obtained contains
potassium vanadate. The solution is neutralised with acid, and barium
chloride added; a meta-salt, Ba(VO_{3})_{2}, is then precipitated as an
almost insoluble white powder, which gives a solution of vanadic acid
when boiled with sulphuric acid. (The precipitate is at first yellow, as
long as it remains amorphous, but it afterwards becomes crystalline and
white.) The solution thus obtained is neutralised with ammonia, which
thus forms ammonium (meta) vanadate, NH_{4}VO_{3}, which, when
evaporated, gives colourless crystals, insoluble in water containing
sal-ammoniac; hence this salt is precipitated by adding solid
sal-ammoniac to the solution. Ammonium vanadate, when ignited, leaves
vanadic acid behind. In this it differs from the corresponding chromium
salt, which is deoxidised into chromium oxide when ignited. In general,
vanadic acid has but a small oxidising action. It is reduced with
difficulty, like phosphoric or sulphuric acid, and in this differs from
arsenic and chromic acids. Vanadic acid, like chromic acid, separates
from its solution as the anhydride V_{2}O_{5}, and not in a hydrous
state. Vanadic anhydride, V_{2}O_{5}, forms a reddish-brown mass, which
easily fuses and re-solidifies into transparent crystals having a violet
lustre (another point of resemblance to chromic acid); it dissolves in
water, forming a yellow solution with a slightly acid reaction.[50]
[50] Strong acids and alkalis dissolve vanadic anhydride in
considerable quantities, forming yellow solutions. When it is
ignited, especially in a current of hydrogen, it evolves oxygen
and forms the lower oxides; V_{2}O_{4} (acid solutions of a green
colour, like the salts of chromic oxide), V_{2}O_{3}, and the
lowest oxide, VO. The latter is the metallic powder which is
obtained when the vanadium oxychloride is heated in an excess of
hydrogen, and was formerly mistaken for metallic vanadium. When a
solution of vanadic acid is treated with metallic zinc it forms a
blue solution, which seems to contain this oxide. It acts as a
reducing agent (and forms a close analogue to chromous oxide,
CrO). Metallic _vanadium_ can only be obtained from vanadium
chloride which is quite free from oxygen. Moissan (1893) obtained
it by reducing the oxide with carbon in the electric furnace, and
considered it to be most infusible of the metals in the series Pt,
Cr, Mo, U, W, and V (he also obtained a compound of vanadium and
carbon). The specific gravity of this metal is 5·5. It is of a
grey-white colour, is not decomposed by water, and is not oxidised
in air, but burns when strongly heated, and can be fused in a
current of hydrogen (forming perhaps a compound with hydrogen). It
is insoluble in hydrochloric acid, but easily dissolves in nitric
acid, and when fused with caustic soda it forms sodium vanadate.
As regards the salts of vanadic acid, three different classes are
known; the first correspond with metavanadic acid, VMO_{3} =
M_{2}OV_{2}O_{5}, the second correspond with the dichromates--that
is, have the composition V_{4}M_{2}O_{11}, which is equal to
M_{2}O + 2V_{2}O_{5}--and the third correspond with orthovanadic
acid, VM_{3}O_{4} or 3M_{2}O + V_{2}O_{5}. The latter are formed
when vanadic anhydride is fused with an excess of an alkaline
carbonate.
Vanadic acid gives the so-called 'complex' acids (which are
considered more fully in Chapter XXI. in speaking of Mo and
W)--_i.e._ acids formed of two acids assimilated into one. Thus
Friedheim (1890) obtained phosphor-vanadic acid, and
Schmitz-Dumont (1890) a similar arseno-vanadic acid. The former is
obtained by heating V_{2}O_{5} with sirupy phosphoric acid. The
resultant golden-yellow tabular crystals have the composition
H_{2}OV_{2}O_{5}P_{2}O_{5}9H_{2}O, and there are corresponding
salts--for example, (NH_{4})_{2}V_{2}O_{5}P_{2}O_{5} with 3 and
7H_{2}O, &c. These salts cannot be separated by crystallisation,
so that there are 'complexes' of these acids in a whole series of
salts (and also in nature). It may be supposed (Friedheim) that
V_{2}O_{5} here, as it were, plays the part of a base, or that
those acids may be looked upon as double salts. Among the true
double salts of vanadium (Nb and Ta) very many are known among the
fluorides, such as VF_{3}2NH_{4}F, VOF_{2}2NH_{4}F,
VO_{2}F,3NH_{4}F, &c. (Pettersson, Piccini, and Georgi, 1890-92).
Vanadium was discovered at the beginning of this century by
Del-Rio, and afterwards investigated by Sefström, but it was only
in 1868 that Roscoe established the above formulæ of the vanadic
compounds.
_Niobium and tantalum_[51] occur as acids in rare minerals, and are
mainly extracted from _tantalite_ and _columbite_, which are found in
Bavaria, Finland, North America, and in the Urals. These minerals are
composed of the ferrous salts of niobic and tantalic acids; they contain
about 15 per cent. of ferrous oxide in isomorphous mixture with manganous
oxide, in combination with various proportions of tantalic and niobic
anhydrides. These minerals are first fused with a considerable amount of
potassium bisulphate, and the fused mass is boiled in water, which
dissolves the ferrous and potassium salts and leaves an insoluble residue
of impure niobic and tantalic acids. This raw product is then treated
with ammonium sulphide, in order to extract the tin and tungsten, which
pass into solution. The residue containing the acids (according to
Marignac) is then treated with hydrofluoric acid, in which it entirely
dissolves, and potassium fluoride is added to the resultant hot solution;
on cooling, a sparingly soluble double fluoride of potassium and tantalum
separates out in fine crystals, while the much more soluble niobium salt
remains in solution. The difference in the solubility of these double
salts in water acidified with hydrofluoric acid (in pure water the
solution becomes cloudy after a certain time) is so great that the
tantalum compound requires 150 parts of water for its solution, and the
niobium compound only 13 parts. The Greenland columbite (specific gravity
5·36) only contains niobic acid, and that from Bodenmais, Bavaria
(specific gravity 6·06) almost equal quantities of tantalic and niobic
acids. Having isolated tantalic and niobic salts, Marignac found that the
relation between the potassium and fluorine in them is very
variable--that is, that there exist various double salts of fluoride of
potassium, and of the fluorides of the metals of this group, but that
with an excess of hydrofluoric acid both the tantalum and niobium
compounds contain seven atoms of fluorine to two of potassium, whence it
must be concluded that the simplest formula for these double salts will
be K_{2}RF_{7} = RF_{5},2KF; that is, that the type of the higher
compounds of niobium and tantalum is RX_{5}, and hence is similar to
phosphoric acid. A chloride, TaCl_{5}, may be obtained from pure tantalic
acid by heating it with charcoal in a current of chlorine. This is a
yellow crystalline substance, which melts at 211°, and boils at 241°; its
vapour density with respect to hydrogen is 180, as would follow from the
formula TaCl_{5}. It is completely decomposed by water into tantalic and
hydrochloric acids. _Niobium pentachloride_ may be prepared in the same
manner; it fuses at 194°, and boils at 240°. When treated with water this
substance gives a solution containing niobic acid, which only separates
out on boiling the solution. Delafontaine and Deville found its vapour
density to be 9·3 (air = 1), as is shown by its formula NbCl_{5}.[52]
[51] The researches made by Roscoe were preceded by those of Marignac
in 1865, on the _compounds_ of _niobium_ and _tantalum_, to which
were also ascribed different formulæ from those now recognised.
Tantalum was discovered simultaneously with vanadium by Hatchett
and Ekeberg, and was afterwards studied by Rose, who in 1844
discovered niobium in it. Notwithstanding the numerous researches
of Hermann (in Moscow), Kobell, Rose, and Marignac, still there is
not yet any certainty as to the purity of, and the properties
ascribed to, the compounds of these elements. They are difficult
to separate from each other, and especially from the cerite metals
and titanium, &c., which accompany them. Before the investigations
of Rose the highest oxide of tantalum was supposed to belong to
the type TaX_{6}--that is, its composition was taken as TaO_{3},
and to the lower oxide was ascribed a formula TaO_{2}. Rose gave
the formula TaO_{2} to the higher oxide, and discovered a new
element called niobium in the substance previously supposed to be
the lower oxide. He even admitted the existence of a third element
occurring together with tantalum and niobium, which he named
pelopium, but he afterwards found that pelopic acid was only
another oxide of niobium, and he considered it probable that the
higher oxide of this element is NbO_{2}, and the lower
Nb_{2}O_{3}. Hermann found that niobic acid which was considered
pure contained a considerable quantity of tantalic acid, and
besides this he admitted the existence of another special metallic
acid, which he called ilmenic acid, after the locality (the Ilmen
mountains of the Urals) of the mineral from which he obtained it.
V. Kobell recognised still another acid, which he called dianic
acid, and these diverse statements were only brought into
agreement in the sixties by Marignac. He first of all indicated an
accurate method for the separation of tantalic and niobic
compounds, which are always obtained in admixture.
[52] If niobic acid be mixed with a small quantity of charcoal and
ignited in a stream of chlorine, a difficultly-fusible and
difficultly-volatile oxychloride, NbOCl_{3} separates. The vapour
density of this compound with respect to air is 7·5, and this
vapour density perfectly confirms the accuracy of the formulæ
given by Marignac, and indicates the quantitative analogy between
the compounds of niobium and tantalum, and those of phosphorus and
arsenic, and consequently also of vanadium. In their qualitative
relations (as is evident also from the correspondence of the
atomic weights), the compounds of tantalum and niobium exhibit a
great analogy with the compounds of molybdenum and tungsten. Thus
zinc, when acting on acid solutions of tantalic and niobic
compounds, gives a blue coloration, exactly as it does with those
of tungsten and molybdenum (also titanium). These acids form the
same large number of salts as those of tungsten and molybdenum.
The anhydrides of the acids are also insoluble in water, but as
colloids are sometimes held in solution, just like those of
titanic and molybdic acids. Furthermore, niobium is in every
respect the nearest analogue of molybdenum, and tantalum of
tungsten. _Niobium_ is obtained by reducing the double fluoride of
niobium and sodium, with sodium. It is difficult to obtain in a
pure state. It is a metal on which hydrochloric acid acts with
some energy, as also does hydrofluoric acid mixed with nitric
acid, and also a boiling solution of caustic potash. _Tantalum_,
which is obtained in exactly the same way, is a much heavier
metal. It is infusible, and is only acted on by a mixture of
hydrofluoric and nitric acids. Rose in 1868 showed that in the
reduction of the double fluoride, NbF_{5},2KF, by sodium, a
greyish powder is obtained after treating with water. The specific
gravity of this powder is 6·8, and he considers it to be niobium
hydride, NbH. Neither did he obtain metallic niobium when he
reduced with magnesium and aluminium, but an alloy, Al_{3}Nb,
having a sp. gr. of 4·5.
Niobium, so far as is known, unites in three proportions with
oxygen. NbO, which is formed when NbOF_{3},2KF is reduced by
sodium; NbO_{2}, which is formed by igniting niobic acid in a
stream of hydrogen, and niobic anhydride, Nb_{2}O_{5}, a white
infusible substance, which is insoluble in acids, and has a
specific gravity of 4·5. Tantalic anhydride closely resembles
niobic anhydride, and has a specific gravity of 7·2. _The
tantalates and niobates_ present the type of ortho-salts--for
example, Na_{2}HNbO_{4},6H_{2}O, and also of pyro-salts, such as
K_{3}HNb_{2}O_{7},6H_{2}O, and of meta-salts--for example,
KNbO_{3},2H_{2}O. And, besides these, they give salts of a more
complex type, containing a larger amount of the elements of the
anhydride; thus, for instance, when niobic anhydride is fused with
caustic potash it forms a salt which is soluble in water, and
crystallises in monoclinic prisms, having the composition
K_{8}Nb_{6}O_{19},16H_{2}O. There is a perfectly similar
isomorphous salt of tantalic acid. Tantalite is a salt of the type
of metatantalic acid, Fe(TaO_{3})_{2}. The composition of
Yttrotantalite appears to correspond with orthotantalic acid.
CHAPTER XX
SULPHUR, SELENIUM, AND TELLURIUM
The acid character of the higher oxides RO_{3} of the elements of group
VI. is still more clearly defined than that of the higher oxides of the
preceding groups, whilst feeble basic properties only appear in the
oxides RO_{3} of the elements of the even series, and then only for those
elements having a high atomic weight--that is, under those two conditions
in which, as a rule, the basic characters increase. Even the lower types
RO_{2} and R_{2}O_{3}, &c., formed by the elements of group VI., are acid
anhydrides in the uneven series, and only those of the elements of the
even series have the properties of peroxides or even of bases.
_Sulphur_ is the typical representative of group VI., both on account of
the fact that the acid properties of the group are clearly defined in it,
and also because it is more widely distributed in nature than any of the
other elements belonging to this group. As an element of the uneven
series of group VI., sulphur gives H_{2}S, sulphuretted hydrogen, SO_{3},
sulphuric anhydride, and SO_{2}, sulphurous anhydride. And in all of them
we find acid properties--SO_{3} and SO_{2} are anhydrides of acids, and
H_{2}S is an acid, although a feeble one. As an element sulphur has all
the properties of a true non-metal; it has not a metallic lustre, does
not conduct electricity, is a bad conductor of heat, is transparent, and
combines directly with metals--in short it has all the properties of the
non-metals, like oxygen and chlorine. Furthermore, sulphur exhibits a
great qualitative and quantitative _resemblance to oxygen_, especially in
the fact that, like oxygen, it combines _with two atoms of hydrogen_, and
forms compounds resembling oxides with metals and non-metals. From this
point of view sulphur is bivalent, if the halogens are univalent.[1] The
chemical character of sulphur is expressed by the fact that it forms a
very slightly stable and feebly energetic acid with hydrogen. The salts
corresponding with this acid are the sulphides, just as the oxides
correspond to water and the chlorides to hydrochloric acid. However, as
we shall afterwards see more fully, the sulphides are more analogous to
the former than to the latter. But although combining with metals, like
oxygen, sulphur also forms chemically stable compounds with oxygen, and
this fact impresses a peculiar character on all the relations of this
element.[2]
[1] The character of sulphur is very clearly defined in the
organo-metallic compounds. Not to dwell on this vast subject, which
belongs to the province of organic chemistry, I think it will be
sufficient for our purpose to compare the physical properties of
the ethyl compounds of mercury, zinc, sulphur and oxygen. The
composition of all of them is expressed by the general formula
(C_{2}H_{5})_{2}R, where R = Hg, Zn, S, or O. They are all
volatile: mercury ethyl, Hg(C_{2}H_{5})_{2}, boils at 159°, its sp.
gr. is 2·444, molecular volume = 106; zinc ethyl boils at 118°, sp.
gr. 1·882, volume 101; ethyl sulphide, S(C_{2}H_{5})_{2}, boils at
90°, sp. gr. 0·825, volume 107; common ether, or ethyl oxide,
O(C_{2}H_{5})_{2}, boils at 35°, sp. gr. 0·736, volume 101, in
addition to which diethyl itself, (C_{2}H_{5})_{2} = C_{4}H_{10},
boils about 0°, sp. gr. about 0·62, volume about 94. Thus the
substitution of Hg, S, and O scarcely changes the volume,
notwithstanding the difference of the weights; the physical
influence, if one may so express oneself, of these elements, which
are so very different in their atomic weights, is almost alike.
[2] Therefore in former times sulphur was known as an amphid element.
Although the analogy between the compounds of sulphur and oxygen
has been recognised from the very birth of modern chemistry (owing,
amongst other things, to the fact that the oxides and sulphides are
the most widely spread metallic ores in nature), still it has only
been clearly expressed by the periodic system, which places both
these elements in group VI. Here, moreover, stands out that
parallelism which exists between SO_{2} and ozone OO_{2}, between
K_{2}SO_{3} and peroxide of potassium K_{2}O_{4} (Volkovitch in
1893 again drew attention to this parallelism).
Sulphur belongs to the number of those elements which _are very widely
distributed in nature_, and occurs both free and combined in various
forms. The atmosphere, however, is almost entirely free from compounds of
sulphur, although a certain amount of them should be present, if only
from the fact that sulphurous anhydride is emitted from the earth in
volcanic eruptions, and in the air of cities, where much coal is burnt,
since this always contains FeS_{2}. Sea and river water generally contain
more or less sulphur in the form of sulphates. The beds of gypsum, sodium
sulphate, magnesium sulphate, and the like are formations of undoubtedly
aqueous origin. The sulphates contained in the soil are the source of the
sulphur found in plants, and are indispensable to their growth. Among
vegetable substances, the proteïds always contain from one to two per
cent. of sulphur. From plants the albuminous substances, together with
their sulphur, pass into the animal organism, and therefore the
decomposition of animal matter is accompanied by the odour of
sulphuretted hydrogen, as the product into which the sulphur passes in
the decomposition of the albuminous substances. Thus a rotten egg emits
sulphuretted hydrogen. Sulphur occurs largely in nature, as the various
insoluble sulphides of the metals. Iron, copper, zinc, lead, antimony,
arsenic, &c., occur in nature combined with sulphur. These _sulphides_
frequently have a metallic lustre, and in the majority of cases occur
crystallised, and also very often several sulphides occur combined or
mixed together in these crystalline compounds. If they are yellow and
have a metallic lustre they are called pyrites. Such are, for example,
copper pyrites, CuFeS_{2}, and iron pyrites, FeS_{2}, which is the
commonest of all. They are all also known as glances or blendes if they
are greyish and have a metallic lustre--for example, zinc blende, lead
glance, PbS, antimony glance, Sb_{2}S_{3}, &c. And, lastly, sulphur
occurs _native_. It occurs in this form in the most recent geological
formations in admixture with limestone and gypsum, and most frequently in
the vicinity of active or extinct volcanoes. As the gases of volcanoes
contain sulphur compounds--namely, sulphuretted hydrogen and sulphurous
anhydride, which by reacting on one another may produce sulphur, which
also frequently appears in the craters of volcanoes as a sublimate--it
might be imagined that the sulphur was of volcanic origin. But on a
nearer acquaintance with its mode of occurrence, and more especially
considering its relation to gypsum, CaSO_{4}, and limestone, the present
general opinion leads to the conclusion that the 'native' sulphur has
been formed by the reduction of the gypsum by organic matter and that its
occurrence is only indirectly connected with volcanic agencies. Near
Tetush, on the Volga, there are beds containing gypsum, sulphur, and
asphalt (mineral tar). In Europe the most important deposits of sulphur
are in the south of Sicily from Catania to Girgenti.[3] There are very
rich deposits of sulphur in Daghestan near Cherkai and Cherkat in Khyut,
near Mount Kanabour-bam, near Petrovsk, and in the Kira Koumski steppes
in the Trans-Caspian provinces, which are able to supply the whole of
Russia with this mineral. Abundant deposits of sulphur have also been
found in Kamtchatka in the neighbourhood of the volcanoes. The method of
separation of the sulphur from its earthy impurities is based on the fact
that sulphur melts when it is heated. The fusion is carried on at the
expense of a portion of the sulphur, which is burnt, so that the
remainder may melt and run from the mass of the earth. This is carried on
in special furnaces called calcaroni, built up of unhewn stone in the
neighbourhood of the mines.[4]
[3] When in Sicily, I found, near Caltanisetta, a specimen of sulphur
with mineral tar. In the same neighbourhood there are naphtha
springs and mud volcanoes. It may be that these substances have
reduced the sulphur from gypsum.
The chief proof in favour of the origin of sulphur from gypsum is
that in treating the deposits for the extraction of the sulphur it
is found that the proportion of sulphur to calcium carbonate never
exceeds that which it would be had they both been derived from
calcium sulphate.
[4] Naturally only those ores of sulphur which contain a considerable
amount of sulphur can be treated by this method. With poor ores it
is necessary to have recourse to distillation or mechanical
treatment in order to separate the sulphur, but its price is so low
that this method in most cases is not profitable.
The sulphur obtained by the above-described method still contains
some impurities, but it is frequently made use of in this form for
many purposes, and especially in considerable quantities for the
manufacture of sulphuric acid, and for strewing over grapes. For
other purposes, and especially in the preparation of gunpowder, a
purer sulphur is required. Sulphur may be purified by distillation.
The crude sulphur is called _rough_, and the distilled sulphur
_refined_. The arrangement given in fig. 86 is employed for
refining sulphur. The rough sulphur is melted in the boiler _d_,
and as it melts it is run through the tube F into an iron retort B
heated by the naked flame of the furnace. Here the sulphur is
converted into vapour, which passes through a wide tube into the
chamber G, surrounded by stone walls and furnished with a
safety-valve S.
[Illustration: FIG. 86.--Refining sulphur by sublimation.]
Sulphur is purified by distillation in special retorts (see fig. 86) by
passing the vapour into a chamber G built of stone. The first portions of
the vapour entering into the condensing chamber are condensed straightway
from the vapour into a solid state, and form a fine powder known as
_flowers of sulphur_.[5] But when the temperature of the receiver attains
the melting point of sulphur, it passes into a liquid state and is cast
into moulds (like sealing wax), and is then known under the name of _roll
sulphur_.[6]
[5] Flowers of sulphur always contain a certain amount of the oxides of
sulphur.
[6] Sulphur may be extracted by various other means. It may be
extracted from iron pyrites, FeS_{2}, which is very widely
distributed in nature. From 100 parts of iron pyrites about half
the sulphur contained, namely, about 25 parts, may be extracted by
heating without the access of air, a lower sulphide of iron, which
is more stable under the action of heat, being left behind. Alkali
waste (Chapter XII.), containing calcium sulphide and gypsum,
CaSO_{4}, may be used for the same purpose, but native sulphur is
so cheap that recourse can only be had to these sources when the
calcium sulphide appears as a worthless by-product. The most simple
process for the extraction of sulphur from alkali waste, in a
chemical sense, consists in evolving sulphuretted hydrogen from the
calcium sulphide by the action of hydrochloric acid. The
sulphuretted hydrogen when burnt gives water and sulphurous
anhydride, which reacts on fresh sulphuretted hydrogen with the
separation of sulphur. The combustion of the sulphuretted hydrogen
may be so conducted that a mixture of 2H_{2}S and SO_{2} is
straightway formed, and this mixture will deposit sulphur (Chapter
XII., Note 14). Gossage and Chance treat alkali waste with carbonic
anhydride, and subject the sulphuretted hydrogen evolved to
incomplete combustion (this is best done by passing a mixture of
sulphuretted hydrogen and air, taken in the requisite proportions,
over red-hot ferric oxide), by which means water and the vapour of
sulphur are formed: H_{2}S + O = H_{2}O + S.
In an uncombined state sulphur exists in _several modifications_, and
forms a good example of the facility with which an alteration of
properties can take place without a change of composition--that is, as
regards the material of a substance. Common sulphur has the well-known
yellow colour. This colour fades as the temperature falls, and at -50°
sulphur is almost colourless. It is very brittle, so that it may be
easily converted into a powder, and it presents a crystalline structure,
which, by the way, shows itself in the unequal expansion of lumps of
sulphur by heat. Hence when a piece of sulphur is heated by the warmth of
the hand, it emits sounds and sometimes cracks, which probably also
depends on the bad heat-conducting power of this substance. It is easily
obtained in a crystalline form by artificial means, because although
insoluble in water it dissolves in carbon bisulphide, and in certain
oils.[7] Solutions of sulphur in carbon bisulphide when evaporated at the
ordinary temperature yield well-formed transparent crystals of sulphur in
the form of rhombic octahedra, in which form it occurs native. The
specific gravity of these crystals is 2·045. Fused sulphur, cast into
moulds and cooled, has, after being kept a long time, a specific gravity
2·066; almost the same as that of the crystalline sulphur of the above
form, which shows that common sulphur is the same as that which
crystallises in octahedra. The specific heat of octahedral sulphur is
0·17; it melts at 114°, and forms a bright yellow mobile liquid. On
further heating, the fused sulphur undergoes an alteration, which we
shall presently describe, first observing that the above octahedral state
of sulphur is its most stable form. Sulphur may be kept at the ordinary
temperature in this form for an indefinite length of time, and many other
modifications of sulphur pass into this form after being left for a
certain time at ordinary temperature.
[7] One hundred parts of liquid carbon bisulphide, CS_{2}, dissolve
16·5 parts of sulphur at -11°, 24 parts at 0°, 37 parts at 15°, 46
parts at 22°, and 181 parts at 55°. The saturated solution boils at
55°, whilst pure carbon bisulphide boils at 47°. The solution of
sulphur in carbon bisulphide reduces the temperature, just as in
the solution of salts in water. Thus the solution of 20 parts of
sulphur in 50 parts of carbon bisulphide at 22° lowers the
temperature by 5°; 100 parts of benzene, C_{6}H_{6}, dissolves
0·965 part of sulphur at 26°, and 4·377 parts at 71°; chloroform,
CHCl_{3}, dissolves 1·2 part of sulphur at 22°, and 16·35 parts at
174°.
If sulphur be melted and then slightly cooled, so that it forms a crust
on the surface and over the sides of the crucible, while the internal
mass remains liquid, then the sulphur takes another crystalline form as
it solidifies. This may be seen by breaking the crust, and pouring out
the remaining molten sulphur.[8] It is then found that the sides of the
crucible are covered with _prismatic crystals_ of the monoclinic system;
they have a totally different appearance from the above-described
crystals of rhombic sulphur. The prismatic crystals are brown,
transparent, and less dense than the crystals of rhombic sulphur, their
specific gravity being only 1·93, and their melting point higher--about
120°. These crystals of sulphur cannot be kept at the ordinary
temperature, which is indeed evident from the fact that in time they turn
yellow; the specific gravity also changes, and they pass completely into
the ordinary modification. This is accompanied by a considerable
development of heat, so that the temperature of the mass may rise 12°.
Thus sulphur is _dimorphous_--that is, it exists in two crystalline
forms, and in both forms it has independent physical properties. However,
no chemical reactions are known which distinguish the two modifications
of sulphur, just as there are none distinguishing aragonite from
calcspar.[9]
[8] If the experiment be made in a vessel with a narrow capillary tube,
the sulphur fuses at a lower temperature (occurs, as it were, in a
supersaturated state), and solidifying at 90°, appears in a rhombic
form (Schützenberger).
[9] If sulphur be cautiously melted in a U tube immersed in a salt
bath, and then gradually cooled, it is possible for all the sulphur
to remain liquid at 100°. It will now be in a state of superfusion;
thus also by careful refrigeration water may be obtained in a
liquid state at -10°, and a lump of ice then causes such water to
form ice, and the temperature rises to 0°. If a prismatic crystal
of sulphur be thrown into one branch of the U tube containing the
liquid sulphur at 100°, and an octahedral crystal be thrown into
the other branch, then, as Gernez showed, the sulphur in each
branch will crystallise in the corresponding form, and both forms
are obtained at the same temperature; therefore it is not the
influence of temperature only which causes the molecules of sulphur
to distribute themselves in one or another form, but also the
influence of the crystalline parts already formed. This phenomenon
is essentially analogous to the phenomena of supersaturated
solutions.
If molten sulphur be heated to 158° it loses its mobility and becomes
thick and very dark-coloured, so that the crucible in which it is heated
may be inverted without the sulphur running out. When heated above this
temperature the sulphur again becomes liquid, and at 250° it is very
mobile, although it does not acquire its original colour, and at 440° it
boils. These modifications in the properties of sulphur depend not only
on the variations of temperature, but also on a change of structure. If
sulphur, heated to about 350°, be poured in a thin stream into cold
water, it does not solidify into a solid mass, but retains its brown
colour and _remains soft_, may be stretched out into threads, and is
elastic, like guttapercha. But in this soft and ductile state, also, it
does not remain for a long time. After the lapse of a certain period this
soft transparent sulphur hardens, becomes opaque, passes into the
ordinary yellow modification of sulphur, and in so doing develops heat,
just as in the conversion of the prismatic into the octahedral variety.
The soft sulphur is characterised by the fact that a certain portion of
it is insoluble in carbon bisulphide. When soft sulphur is immersed in
this liquid, only a portion of common sulphur passes into solution,
whilst a certain portion is quite insoluble and remains so for a long
time. The maximum proportion of insoluble sulphur is obtained by heating
slightly above 170°. It melts at 114°. An exactly similar _insoluble
amorphous sulphur_ is obtained in certain reactions in the wet way, when
sulphur separates out from solutions. Thus sodium thiosulphate,
Na_{2}S_{2}O_{3}, when treated with acids, gives a precipitate of
sulphur, which is insoluble in carbon bisulphide. The action of water on
sulphur chloride also gives a similar modification of sulphur. Certain
sulphides, when treated with nitric acid, also yield sulphur in this
form.[10]
[10] A certain amount of insoluble sulphur remains for a long time in
the mass of soft sulphur, changing into the ordinary variety.
Freshly-cooled soft sulphur contains about one-third of insoluble
sulphur, and after the lapse of two years it still contains about
15 p.c. Flowers of sulphur, obtained by the rapid condensation of
sulphur from a state of vapour, also contains a certain amount of
insoluble sulphur. _Rapidly distilled and condensed sulphur_ also
contains some insoluble sulphur. Hence a certain amount of
insoluble sulphur is frequently found in roll sulphur. The action
of light on a solution of sulphur converts a certain portion into
the insoluble modification. Insoluble sulphur is of a lighter
colour than the ordinary variety. It is best prepared by
vaporising sulphur in a stream of carbonic anhydride, hydrochloric
acid, &c., and collecting the vapour in cold water. When condensed
in this manner it is nearly all insoluble in carbon bisulphide. It
then has the form of hollow spheroids, and is therefore lighter
than the common variety: sp. gr. 1·82. An idea of the
modifications taking place in sulphur between 110° and 250° may be
formed from the fact that at 150° liquid sulphur has a coefficient
of expansion of about 0·0005, whilst between 150° and 250° it is
less than 0·0003.
Engel (1891), by decomposing a saturated solution of hyposulphite
of sodium (Note 42) with HCl in the cold (the sulphur is not
precipitated directly in this case), obtained, after shaking up
with chloroform and evaporation, crystals of sulphur (sp. gr.
2·135), which, after several hours, passed into the insoluble (in
CS_{2}) state, and in so doing became opaque, and increased in
volume. But if a mixture of solution of Na_{2}S_{2}O_{3} and HCl
be allowed to stand, it deposits sulphur, which, after sufficient
washing, is able to dissolve in water (like the colloid varieties
of the metallic sulphides, alumina, boron, and silver), but this
colloid _solution of sulphur_ soon deposits sulphur insoluble in
CS_{2}.
When a solution of sulphuretted hydrogen in water is decomposed by
an electric current the sulphur is deposited on the positive pole,
and has therefore an electro-negative character, and this sulphur
is soluble in carbon bisulphide. When a solution of sulphurous
acid is decomposed in the same manner, the sulphur is deposited on
the negative pole, and is therefore electro-positive, and the
sulphur so deposited is insoluble in carbon bisulphide. The
sulphur which is combined with metals must have the properties of
the sulphur contained in sulphuretted hydrogen, whilst the sulphur
combined with chlorine is like that which is combined with oxygen
in sulphurous anhydride. Hence Berthelot recognises the presence
of soluble sulphur in metallic sulphides, and of the insoluble
modification of amorphous sulphur in sulphur chloride. Cloez
showed that the sulphur precipitated from solutions is either
soluble or insoluble, according to whether it separates from an
alkaline or acid solution. If sulphur be melted with a small
quantity of iodine or bromine, then on pouring out the molten mass
it forms amorphous sulphur, which keeps so for a very long time,
and is insoluble, or nearly so, in carbon bisulphide. This is
taken advantage of in casting certain articles in sulphur, which
by this means retain their tenacity for a long time; for example,
the discs of electrical machines.
At temperatures of 440° to 700° the vapour density of sulphur is 6·6
referred to air--_i.e._ about 96 referred to hydrogen. Hence, at these
temperatures _the molecule of sulphur contains six atoms_, it has the
composition S_{6}. The agreement between the observations of Dumas,
Mitscherlich, Bineau, and Deville confirms the accuracy of this result.
But in this respect the properties of sulphur were found to be variable.
When heated to higher temperatures, that is to say, _above_ 800°, the
vapour density of sulphur is found to be one-third of this quantity,
_i.e._ about 32 referred to hydrogen. At this temperature _the molecule
of sulphur_, like that of hydrogen, oxygen, nitrogen, and chlorine,
_contains two atoms_; hence the molecular formula is then S_{2}. This
variation in the vapour density of sulphur evidently corresponds with a
polymeric modification, and may be likened to the transformation of
ozone, O_{3}, into oxygen, O_{2}, or better still, of benzene,
C_{6}H_{6}, into acetylene, C_{2}H_{2}.[11]
[11] Here, however, it is very important to remark that both benzene
and acetylene can exist at the ordinary temperature, whilst the
sulphur molecule S_{2} only exists at high temperatures; and if
this sulphur be allowed to cool, it passes first into S_{6} and
then into a liquid state. Were it possible to have sulphur at the
ordinary temperature in both the above modifications, then in all
probability the sulphur in the state S_{2} would present totally
different properties from those which it has in the form S_{6},
just as the properties of gaseous acetylene are far from being
similar to those of liquid benzene. Sulphur, in the form of S_{2},
is probably a substance which boils at a much lower temperature
than the variety with which we are now dealing. Paterno and Nasini
(1888), following the method of depression or fall of the
freezing-point in a benzene solution, found that the molecule of
sulphur in solution contains S_{6}.
One must here call attention to the fact that sulphur, with all
its analogy to oxygen (which also shows itself in its faculty to
give the modification S_{2}), is also able to give a series of
compounds containing more atoms of sulphur than the analogous
oxygen compounds do of oxygen. Thus, for instance, compounds of 5
atoms of sulphur with 1 atom of barium, BaS_{5}, are known,
whereas with oxygen only BaO_{2} is known. On every side one
cannot but see in sulphur a faculty for the union of a greater
number of atoms than with oxygen. With oxygen the form of ozone,
O_{3}, is very unstable, the stable form is O_{2}; whilst with
sulphur S_{6} is the stable form, and S_{2} is exceedingly
unstable. Furthermore, it is remarkable that sulphur gives a
higher degree of oxidation, H_{2}SO_{4}, corresponding, as it
were, with its complex composition, if we suppose that in S_{6}
four atoms of sulphur are replaced by oxygen and one by two atoms
of hydrogen. The formulæ of its compounds, K_{2}SO_{4},
K_{2}S_{2}O_{3}, K_{2}S_{5}, BaS_{5}, and many others, have no
analogues among the compounds of oxygen. They all correspond with
the form S_{6} (one portion of the sulphur being replaced by
oxygen and another by metals), which is not attained by oxygen. In
this faculty of sulphur to hold many atoms of other substances the
same forces appear which cause many atoms of sulphur to form one
complex molecule.
_In its faculty for combination_, sulphur most closely resembles oxygen
and chlorine; like them, it combines with nearly all elements, with the
development of heat and light, forming sulphur compounds, but as a rule
this only takes place at a high temperature. At the ordinary temperature
it does not enter into reactions, owing, amongst other things, to the
fact that it is a solid. In a molten state it acts on most metals and on
the halogens. It burns in air at about 300°, and with carbon at a red
heat, but it does not combine with nitrogen.
Fine wires, or the powders of the greater number of metals, burn in the
vapour of sulphur. The direct combination of hydrogen with sulphur is
restricted by a limit--that is, at a given temperature and under other
given conditions it does not proceed unrestrictedly; there is no
explosion or recalescence. Sulphuretted hydrogen, H_{2}S, decomposes at
its temperature of combination--that is, it is easily dissociated.[12]
The same phenomenon is repeated here as with water, except that the
temperatures at which the attraction of hydrogen for sulphur begins and
ceases are much lower than in the case of oxygen and hydrogen. The
temperature at which combination takes place is here, as in many other
instances, nearly the same as that at which dissociation begins. Hence
_sulphuretted hydrogen_ is formed in a small quantity by the direct
ignition of a mixture of the vapour of sulphur and hydrogen. However, the
temperature must not be high, because otherwise the whole of the
sulphuretted hydrogen is decomposed; but at lower temperatures a small
amount of sulphuretted hydrogen is formed by direct combination.[13]
Sulphuretted hydrogen however, like all other hydrogen compounds, may be
easily obtained by the double decomposition of its corresponding metallic
compounds, the replacement of the metal by hydrogen being effected by the
action of acids on the sulphides. The metallic sulphides are, as a rule,
easily formed. A sulphide, when mixed with a non-volatile acid, may give,
by double decomposition, a salt of the acid taken and sulphuretted
hydrogen, M_{2}S + H_{2}SO_{4} = H_{2}S + M_{2}SO_{4}. However, it is not
all sulphides nor solutions of all acids that will evolve sulphuretted
hydrogen, which fact is exceedingly characteristic, because, for example,
all carbonates evolve carbonic anhydride when treated with any acid.
Sulphuric acid will only evolve sulphuretted hydrogen from those
sulphides which contain a metal capable of decomposing the acid with the
evolution of hydrogen. Thus zinc, iron, calcium, magnesium, manganese,
potassium, sodium, &c., form sulphides which evolve sulphuretted hydrogen
when treated with sulphuric acid, and the metals themselves evolve
hydrogen with acids.[14] The sulphides of those metals which do not
liberate hydrogen from acids do not generally act on acids--that is, do
not form sulphuretted hydrogen with them; such are, for example, the
sulphides of lead, silver, copper, mercury, tin, &c. Therefore, the
_modus operandi_ of the formation of sulphuretted hydrogen by the action
of acids on metallic sulphides may be looked on as a phenomenon of the
combination of hydrogen, at the moment of its evolution, with the
sulphur, which is combined with the metal. Such a representation is all
the more simple as all the circumstances under which sulphuretted
hydrogen is formed are exactly similar to the conditions of the formation
of hydrogen itself. Thus the usual mode of preparing sulphuretted
hydrogen is by the action of _sulphuric acid on ferrous sulphide_, in
which the same apparatus and method are employed as in the preparation of
hydrogen, only replacing the metallic iron or zinc by ferrous sulphide or
zinc sulphide. The reaction between sulphide of iron and sulphuric acid
takes place at the ordinary temperature, and is accompanied by just as
small a development of heat as in the liberation of hydrogen itself, FeS
+ H_{2}SO_{4} = FeSO_{4} + H_{2}S.[15]
[12] In the formation of potassium sulphide, K_{2}S (that is, in the
combination of 32 parts of sulphur with 78 parts of potassium),
about 100 thousand heat units are developed. Nearly as much heat
is developed in the combination of an equivalent quantity of
sodium; about 90 thousand heat units in the formation of calcium
or strontium sulphide; about 40 thousand for zinc or cadmium
sulphide, and about 20 thousand for iron, cobalt, or nickel
sulphide. Less heat is evolved in the combination of sulphur with
copper, lead, and silver. According to Thomsen, sulphur develops
heat with hydrogen in solutions. The reaction I_{2},Aq,H_{2}S =
21,830 calories. But, as the reaction I_{2} + H_{2} + Aq develops
26,842 calories, it follows that the reaction H_{2} + S develops
4,512 calories.
[13] If sulphur be melted in a flask and heated nearly to its boiling
point, as Lidoff showed, the addition, drop by drop (from a funnel
with a stopcock) of heavy (0·9) naphtha oil (of lubricating
oleonaphtha), &c., is followed by a regular evolution of
sulphuretted hydrogen. This is analogous to the action of bromine
or iodine on paraffin and other oils, because hydrobromic or
hydriodic acid is then formed (Chapter XI.) A certain amount of
hydrogen sulphide is even formed when sulphur is boiled with
water.
[14] However, the matter is really much more complicated. Thus zinc
sulphide evolves sulphuretted hydrogen with sulphuric or
hydrochloric acids, but does not react with acetic acid and is
oxidised by nitric acid. Ferrous sulphide evolves sulphuretted
hydrogen with acids, whilst the bisulphide, FeS_{2}, does not
react with acids of ordinary strength. This absence of action
depends, among other things, on the form in which the native iron
pyrites occurs; it is a crystalline, compact, and very dense
substance; and acids in general react with great difficulty on
such metallic sulphides. This is seen very clearly in the case of
zinc sulphide; if this substance is obtained by double
decomposition, it separates as a white precipitate, which evolves
sulphuretted hydrogen with great ease when treated with acids.
Zinc sulphide is obtained in the same form when zinc is fused with
sulphur, but native zinc sulphide--which occurs in compact masses
of zinc blende, and has a metallic lustre--is not decomposed or
scarcely decomposed by sulphuric acid.
Another source of complication in the behaviour of the metallic
sulphides towards acids depends on the action of water, and is
shown in the fact that the action varies with different degrees of
dilution or proportion of water present. The best known example of
this is antimonious sulphide, Sb_{2}S_{3}, for strong hydrochloric
acid, containing not more water than corresponds with HCl,6H_{2}O,
even decomposes native antimony glance, with evolution of
sulphuretted hydrogen, whilst dilute acid has no action, and in
the presence of an excess of water the reaction 2SbCl_{3} +
3H_{2}S = Sb_{2}S_{3} + 6HCl occurs, whilst in the presence of a
small amount of water the reaction proceeds in exactly the
opposite direction. Here the participation of water in the
reaction and its affinity are evident.
The facts that lead sulphide is insoluble in acids, that zinc
sulphide is soluble in hydrochloric acid but insoluble in acetic
acid, that calcium sulphide is even decomposed by carbonic acid,
&c.--all these peculiarities of the sulphides are in correlation
with the amount of heat evolved in the reaction of the oxides with
hydrogen sulphide and with acids, as is seen from the observations
of Favre and Silberman, and from the comparisons made by Berthelot
in the Proceedings of the Paris Academy of Sciences, 1870, to
which we refer the reader for further details.
[15] _Ferrous sulphide_ is formed by heating a piece of iron to an
incipient white heat, and then removing it from the furnace and
bringing it into contact with a piece of sulphur. Combination then
proceeds, accompanied by the development of heat, and the ferrous
sulphide formed fuses. The sulphide of iron thus formed is a
black, easily-fusible substance, insoluble in water. When damp it
attracts oxygen from the air, and is converted into green vitriol,
FeSO_{4}. If all the iron does not combine with the sulphur in the
method described above, the action of sulphuric acid will evolve
hydrogen as well as hydrogen sulphide.
We will not describe the details of the preparation of
sulphuretted hydrogen employed as a reagent in the laboratory,
because, in the first place, the methods are essentially the same
as in the preparation of hydrogen, and, in the second place,
because the apparatus and methods employed are always described in
text-books of analytical chemistry. Ferrous sulphide may be
advantageously replaced by calcium sulphide or a mixture of
calcium and magnesium sulphides. A solution of magnesium
hydrosulphide, MgS,H_{2}S, is very convenient, as at 60° it
evolves a stream of pure hydrogen sulphide. A paste, consisting of
CuS with crystals of MgCl_{2} and water, may also be employed,
since it only evolves H_{2}S when heated (Habermann).
_In nature_ sulphuretted hydrogen is formed in many ways. The most usual
mode of its formation is by the decomposition of albuminous substances
containing sulphur, as mentioned above. Another method is by the reducing
action of organic matter on sulphates, and by the action of water and
carbonic acid on the sulphides formed by this reduction. Volcanic
eruptions are a third source of sulphuretted hydrogen in nature. Although
sulphuretted hydrogen is formed in small quantities everywhere, it
nevertheless soon disappears from the atmosphere, owing to its being
easily decomposed by oxidising agencies. Many mineral waters contain
sulphuretted hydrogen, and smell of it; they are called 'sulphur waters.'
Sulphuretted hydrogen, at the ordinary temperature, is a colourless gas,
having a very unpleasant odour. It has, as its composition H_{2}S shows,
a specific gravity seventeen times greater than hydrogen, and therefore
it is somewhat heavier than air. Sulphuretted hydrogen _liquefies_ at
about -74°, and at the ordinary temperature when subjected to a pressure
of 10 to 15 atmospheres; at -85° it is converted into a solid crystalline
mass.[15 bis] The easy liquefaction of sulphuretted hydrogen is evidently
allied to its solubility. One volume of water at 0° dissolves 4·37
volumes of sulphuretted hydrogen, at 10° 3·58 volumes, and at 20° 2·9
volumes.[16] The solutions impart a very feeble red coloration to litmus
paper. This gas is poisonous. One part in fifteen hundred parts of air
will kill birds. Mammalia die in an atmosphere containing 1/200 of this
gas.
[15 bis] Liquid sulphuretted hydrogen is most easily obtained by the
decomposition of hydrogen polysulphide, which we shall presently
describe, by the action of heat, and in the presence of a small
amount of water. If poured into a bent tube, like that described
for the liquefaction of ammonia (Chapter VI.), the hydrogen
polysulphide is decomposed by heat, in the presence of water, into
sulphur and sulphuretted hydrogen, which condenses in the cold end
of the tube into a colourless liquid.