Chapter XIX: Phosphorus and the Other Elements of the Fifth Group (2)

[15] Silver orthophosphate, Ag_{3}PO_{4}, is yellow, sp. gr. 7·32, and
insoluble in water. When heated it fuses like silver chloride, and
if kept fused for some length of time it gives a white
pyrophosphate (the decomposition which causes this is not known).
It is soluble in aqueous solutions of phosphoric, nitric, and even
acetic acids, of ammonia, and many of its salts. If silver nitrate
acts on a dimetallic orthophosphate--for instance,
Na_{2}HPO_{4}--it still gives Ag_{3}PO_{4}, nitric acid being
disengaged: Na_{2}HPO_{4} + 3AgNO_{3} = Ag_{3}PO_{4} + 2NaNO_{3} +
HNO_{3}. When alcohol is added to silver orthophosphate,
Ag_{3}PO_{4}, dissolved in syrupy phosphoric acid, it precipitates
a white salt (the alcohol takes up the free phosphoric acid)
having the composition Ag_{2}HPO_{4}, which is immediately
decomposed by water into the normal salt and phosphoric acid.

[16] The researches of Thomsen showed that in very dilute aqueous
solutions the majority of monobasic acids--nitric, acetic,
hydrochloric, &c. (but hydrofluoric acid more and hydrocyanic
less)--HX evolve the following amounts of heat (in thousands of
calories) with caustic soda: NaHO + 2HX = 14; NaHO + HX = 14;
2NaHO + HX = 14; that is, if _n_ be a whole number _n_NaHO + HX =
14 and NaHO + _n_HX = 14. Hence reaction here only takes place
between one molecule of NaHO and one molecule of acid, and the
remaining quantity of acid or alkali does not enter into the
reaction. In the case of bibasic acids, H_{2}R´´ (sulphuric,
dithionic, oxalic, sulphuretted hydrogen, &c.), NaHO + 2H_{2}R´´ =
14; NaHO + H_{2}R´´ = 14; 2NaHO + H_{2}R´´ = 28; _n_NaHO +
H_{2}R´´ = 28; that is, with an excess of acid (NaHO + 2H´_{2}R´´)
14 thousand units of heat are developed, and with an excess of
alkali 28. When phosphoric acid is taken (but not all tribasic
acids--for instance, not citric) the general character of the
phenomenon is similar to the preceding, namely, NaHO +
2H_{3}PO_{4} = 14·7; NaHO + H_{3}PO_{4} = 14·8; 2NaHO +
H_{3}PO_{4} = 27·1; 3NaHO + H_{3}PO_{4} = 34·0; 6NaHO +
H_{3}PO_{4} = 35·3; or, in general terms, NaHO + _n_H_{3}PO_{4} =
14 (approximately) and _n_NaHO + H_{3}PO_{4} = 35 and not 42,
which shows a peculiarity of phosphoric acid. In the case of
energetic acids, when one equivalent (23 grams) of sodium (in the
form of hydroxide) replaces one equivalent (1 gram) of hydrogen
(with the formation of water and in dilute solutions), 14,000 heat
units are evolved; and this is true for phosphoric acid when in
H_{3}PO_{4}, Na or Na_{2} replaces H or H_{2}, but when Na_{3}
replaces H_{3} less heat is developed. This will be seen from the
following scheme based on the preceding figures: H_{3}PO_{4} +
NaHO = 14·8; NaH_{2}PO_{4} + NaHO = 12·3; Na_{2}HPO_{4} + NaHO =
5·9; with Na_{3}PO_{4} + NaHO, a very small amount of heat is
evolved, as may be judged from the fact that Na_{3}PO_{4} + 3NaHO
= 1·3, but still heat is evolved. It must be supposed that in
acting on phosphoric acid in the presence of a large quantity of
water, a certain portion of the sodium hydroxide remains as alkali
uncombined with the acid. Thus, on increasing the mass of the
alkali, heat is still evolved, and a fresh interchange between Na
and H takes place. Hence water shows a decomposing action on the
alkali phosphates. The same decomposing action of water is seen,
but to a less extent, with Na_{2}HPO_{4}, as may be judged both
from the reactions of this salt and from the amount of heat
developed by NaH_{2}PO_{4} with NaHO. Such an explanation is in
accordance with many facts concerning the decomposition of salts
by water already known to us. Recent researches made by Berthelot
and Louguinine have confirmed the above deductions made by me in
the first edition (1871) of this work. At the present time views
of this nature are somewhat generally accepted, although they are
not sufficiently strictly applied in other cases. As regards
PH_{3}O_{4} it may be said that: on the substitution of the first
hydrogen this acid acts as a powerful acid (like HCl, HNO_{3},
H_{2}SO_{4}); on the substitution of the second hydrogen as a
weaker acid (like an organic acid); and on the substitution of the
third, as an alcohol, for instance phenol, having the properties
of a feeble acid.

Phosphoric anhydride, or any of its hydrates, when ignited with an excess of sodium hydroxide, carbonate, &c., forms normal or _trisodium orthophosphate_, Na_{3}PO_{4}, but when a solution of sodium carbonate is decomposed by orthophosphoric acid, only the salt Na_{2}HPO_{4} is formed; and when an excess of sodium chloride is ignited with orthophosphoric acid, hydrochloric acid is evolved, and the acid salt H_{2}NaPO_{4} alone is formed. These facts clearly indicate the small energy of phosphoric acid with respect to the formation of the tri-metallic salt, which is seen further from the fact that the salt Na_{3}PO_{4} has an alkaline reaction, decomposes in the presence of water and carbonic acid, forming Na_{2}HPO_{4}, corrodes glass vessels in which it is boiled or evaporated, just like solutions of the alkalis, disengages, like them, ammonia from ammonium chloride, and crystallises from solutions, as Na_{3}PO_{4},12H_{2}O, only in the presence of an excess of alkali. At 15° the crystals of this salt require five parts of water for solution; they fuse at 77°.

_Disodium orthophosphate_, or common sodium phosphate, Na_{2}HPO_{4}, is more stable both in solution and in the solid state. As it is used in medicine and in dyeing, it is prepared in considerable quantities, most frequently from the impure phosphoric acid obtained by the action of sulphuric acid on bone ash. The solution thus formed--which contains, besides phosphoric and sulphuric acids, salts of sodium, calcium, and magnesium--is heated, and sodium carbonate added so long as carbonic anhydride is disengaged. A precipitate is formed containing the insoluble salts of magnesium and calcium, whilst the solution contains sodium phosphate, Na_{2}HPO_{4}, with a small quantity of other salts, from which it may be easily purified by crystallisation. At the ordinary temperature its solutions, especially in the presence of a small amount of sodium carbonate, give finely-formed inclined prismatic crystals, Na_{2}HPO_{4},12H_{2}O; when the crystallisation takes place above 30° they only contain 7H_{2}O. The former crystals even lose a portion of their water of crystallisation at the ordinary temperature (the salt effloresces), and form the second salt with 7H_{2}O; whilst under the receiver of an air-pump and over sulphuric acid they also part with this water.[17] When ignited they lose the last molecule of water of constitution, and give sodium pyrophosphate, Na_{4}P_{2}O_{7}.

[17] Na_{2}HPO_{4},12H_{2}O has a sp. gr. 1·53. Poggiale determined the
solubility in 100 parts of water (1) of the anhydrous ortho-salt
Na_{2}HPO_{4}, and (2) of the corresponding pyro-salt
Na_{4}P_{2}O_{7}:--

0° 20° 40° 80° 100°
I. 1·5 11·1 30·9 81 108
II. 3·2 6·2 13·5 30 40

At temperatures of 20° to 100° the ortho-salt is so very much less
soluble that this difference alone already indicates the
deeply-seated alteration in constitution which takes place in the
passage from the ortho- to the pyro-salts.

_Monosodium orthophosphate_, NaH_{2}PO_{4}, crystallises with one equivalent of water; its solution has an acid reaction. At 100° the salt only loses this water of crystallisation, and at about 200° it parts with all its water, forming the metaphosphate NaPO_{3}. It is prepared from ordinary sodium phosphate by adding phosphoric acid until the solution does not give a precipitate with barium chloride, and then evaporating and crystallising the solution. The solution of this salt does not absorb carbonic anhydride, and does not give a precipitate with salts of calcium, barium, &c.[18]

[18] The _ammonium orthophosphates_ resemble the sodium salts in many
respects, but the instability of the di- and tri-metallic salts is
seen in them still more clearly than in the sodium salts; thus
(NH_{4})_{3}PO_{4}, and even (NH_{4})_{2}HPO_{4}, lose ammonia in
the air (especially when heated, even in solutions);
NH_{4}H_{2}PO_{4} alone does not disengage ammonia and has an acid
reaction. The crystals of the first salt contain 3H_{2}O, and are
only formed in the presence of an excess of ammonia; both the
others are anhydrous, and may be obtained like the sodium salts.
When ignited these salts leave metaphosphoric acid behind; for
example, (NH_{4})_{2}HPO_{4} = 2NH_{3 + H_{2}O + HPO_{3}. Ammonia
also enters into the composition of many double phosphates.
Ammonium sodium orthophosphate, or simply phosphate,
NH_{4}NaHPO_{4},4H_{2}O, crystallises in large transparent
crystals from a mixture of the solutions of disodium phosphate and
ammonium chloride (in which case sodium chloride is obtained in
the mother liquid), or, better still, from a solution of
monosodium phosphate saturated with ammonia. It is also formed
from the phosphates in urine when it ferments. This salt is
frequently used in testing metallic compounds by the blow-pipe,
because when ignited it leaves a vitreous metaphosphate, NaPO_{3},
which, like borax, dissolves metallic oxides, forming
characteristic tinted glasses.

When a solution of trisodium phosphate is added to a solution of a
magnesium salt it gives a white precipitate of the normal
orthophosphate Mg_{2}(PO_{4})_{2},7H_{2}O. If the trisodium salt
be replaced by the ordinary salt, Na_{2}HPO_{4}, a precipitate is
also formed, and MgHPO_{4},7H_{2}O is obtained. It might be
thought that the normal salt Mg_{3}(PO_{4})_{2} would be
precipitated if disodium phosphate was added to ammonia and a salt
of magnesium, but in reality _ammonium magnesium orthophosphate_,
MgNH_{4}PO_{4},6H_{2}O, is precipitated as a crystalline powder,
which loses ammonia and water when ignited, and gives a
pyrophosphate, Mg_{2}P_{2}O_{7}. This salt occurs in nature as the
mineral struvite, and in various products of the changes of animal
matter. If we consider that the above salt parts with ammonia with
difficulty, and that the corresponding salt of sodium is not
formed under the same conditions (MgNaPO_{4},9H_{2}O is obtained
by the action of magnesia on disodium phosphate), if we turn our
attention to the fact that the salts of calcium and barium do not
form double salts as easily as magnesium, and remember that the
salts of magnesium in general easily form double ammonium salts,
we are led to think that this salt is not really a normal, but an
acid salt, corresponding with Na_{2}HPO_{4}, in which Na_{2} is
replaced by the equivalent group NH_{3}Mg.

The common normal _calcium phosphate_, Ca_{3}(PO_{4})_{2}, occurs
in minerals, in animals, especially in bones, and also probably in
plants, although the ash of many portions of plants, as a rule,
contains less lime than the formation of the normal salt requires.
Thus 100 parts of the ash (from 5,000 parts of grain) of rye grain
contain 47·5 of phosphoric anhydride and only 2·7 of lime, and
even the ash of the whole of the rye (including the straw)
contains twice as much phosphoric anhydride as lime, and the
normal salt contains almost equal weights of these substances.
Only the ash of grasses, and especially of clover, and of trees,
contains in the majority of cases more lime than is required for
the formation of Ca_{3}P_{2}O_{8}. This salt, which is insoluble
in water, dissolves even in such feeble acids as acetic and
sulphurous, and even in water containing carbonic acid. The latter
fact is of immense importance in nature, since by reason of it
rain water is able to transfer the calcium phosphates in the soil
into solutions which are absorbed by plants. The solubility of the
normal salt in acids takes place by virtue of the formation of an
acid salt, which is evident from the quantity of acid required for
its solution, and more especially from the fact that the acid
solutions when evaporated give crystalline scales of the acid
calcium phosphate, CaH_{4}(PO_{4})_{2}, soluble in water. This
solubility of the acid salt forms the basis of the treatment by
acids of bones, phosphorites, guano, and other natural products
containing the normal salt and employed for fertilising the soil.
The perfect decomposition requires at least 2H_{2}SO_{4} to
Ca_{3}(PO_{4})_{2}, but in reality less is taken, so that only a
portion of the normal salt is converted into the acid salt.
Hydrochloric acid is sometimes used. (In practice such mixtures
are known as _superphosphates_). Certain experiments, however,
show that a thorough grinding, the presence of organic, and
especially of nitrogenous, substances, and the porous structure of
some calcium phosphates (for example, in burnt bones), render the
treatment of phosphoric manures by acids superfluous--that is, the
crop is not improved by it.

As a hydrate, orthophosphoric acid should be expressed, after the fashion of other hydrates, as containing three water residues (hydroxyl groups), _i.e._ as PO(OH)_{3}. This method of expression indicates that the type PX_{5}, seen in PH_{4}I, is here preserved, with the substitution of X_{2} by oxygen and X_{3} by three hydroxyl groups. The same type appears in POCl_{3}, PCl_{5}, PF_{5}, &c. And if we recognise phosphoric acid as PO(OH)_{3}, we should expect to find three anhydrides corresponding with it: (1) [PO(OH)_{2}]_{2}O, in which two of the three hydroxyls are preserved; this is pyrophosphoric acid, H_{4}P_{2}O_{7}. (2) PO(OH)O, where only one hydroxyl is preserved. This is metaphosphoric acid. (3) (PO)_{2}O_{3} or P_{2}O_{5}, that is, perfect phosphoric anhydride. Therefore, _pyro- and metaphosphoric acids are imperfect anhydrides_ (or anhydro-acids) _of orthophosphoric acid_.[19]

[19] In this sense the ortho-acid itself might be regarded as an
anhydro-acid, counting P(HO)_{5} as the perfect hydrate, if PH_{5}
existed; but as in general the normal hydrates correspond with the
existing hydrogen compounds with the addition of up to 4 atoms of
oxygen, therefore PH_{3}O_{4} is the normal acid, just as
SH_{2}O_{4} and ClHO_{4}; while NHO_{3}, CH_{2}O_{3} are
meta-acids, or higher normal acids (NH_{3}O_{4} and CH_{4}O_{4})
with the loss of a molecule of water.

In order to see the relation between the ortho-, pyro-, and
metaphosphoric acids, the first thing to remark in them is that
the anhydride P_{2}O_{5} is combined with 3, 2, and 1 molecules of
water. In the absence of data for the molecular weight of ortho-
and pyrophosphoric acids it is necessary to mention that all
existing data for metaphosphoric acid indicate (Note 21) that its
molecule is much more complex and contains at least
H_{3}P_{3}O_{9} or H_{6}P_{6}O_{18}. The explanation of the
problems which here present themselves can, it seems to me, be
only looked for after a detailed study of the phenomena of the
polymerisations of mineral substances, and of those complex acids,
such as phosphomolybdic, which we shall hereafter describe
(Chapter XXI.) A similar instance is exhibited in the solubility
of hydrate of silica (produced by the action of silicon fluoride
on water) in fused metaphosphoric acid, with the formation, on
cooling, of an octahedral compound (sp. gr., 3·1) containing
SiO_{2},P_{2}O_{5}. A certain indication (but no proof) that
ordinary orthophosphoric acid is polymerised is given by
Staudenmaier (1893), who obtained a salt, K_{5}H_{4}P_{3}O_{12},
by the action of a solution of KH_{2}PO_{4} upon K_{2}CO_{3}; and
a compound, KH_{3}P_{2}O_{8}, corresponding to the doubled
molecule of H_{3}PO_{4}, by the action of KH_{2}PO_{4} upon
H_{3}PO_{4} itself.

_Pyrophosphoric acid_, H_{4}P_{2}O_{7}, is formed by heating orthophosphoric acid to 250° when it loses water.[19 bis] Its normal salts are formed by igniting the dimetallic salts of orthophosphoric acid of the types HM_{2}PO_{4}. Thus from the disodium salt we obtain sodium pyrophosphate, Na_{4}P_{2}O_{7} (it crystallises from water with 10H_{2}O, is very stable, fuses when heated, has an alkaline reaction, and does not form ortho-salts when its solution is boiled): and from the monosodium salt NaH_{2}PO_{4} the acid salt Na_{2}H_{2}P_{2}O_{7} (easily soluble in water) is formed; this has an acid reaction, and when ignited further gives the meta-salt.[20]

[19 bis] According to Watson (1893) the ortho-acid is partially
transformed into the pyro-acid at 230°, whilst at 260° the latter
begins to volatilise. At 300° the meta-acid only is formed.

[20] The method of preparation of the acid itself consists in
converting the sodium salt, Na_{4}P_{2}O_{7}, by double
decomposition with water and a salt of lead, into insoluble lead
pyrophosphate, Pb_{2}P_{2}O_{7}, which is then suspended in water
and decomposed by sulphuretted hydrogen; lead sulphide is thus
precipitated, and pyrophosphoric acid remains in solution. This
solution cannot be heated, or the pyro-acid will pass into the
ortho-, but must be evaporated under the receiver of an air-pump.
It concentrates to a syrup and crystallises, and when ignited in
this form loses water, and forms metaphosphoric acid. It resembles
orthophosphoric acid in many respects; its salts with the alkalis
are also soluble, and the others insoluble in water but soluble in
acids. When heated in solution with acid it gives orthophosphoric
acid, as well as when fused with an excess of alkali.

Witt heated ammonium chloride with phosphoric acid (hydrochloric
acid was evolved), ignited the residue to drive off ammonia, and
obtained pyrophosphoric acid in the residue.

_Metaphosphoric acid_, HPO_{3} (the analogue of nitric acid), is formed by the ignition of the pyro- and ortho-acids (or, better, of their ammonium salts), as a vitreous, hygroscopic, fused mass (glacial phosphoric acid, _acidum phosphoricum glaciale_), soluble in water and volatilising without decomposition. It is also formed in the first slow action of cold water on the anhydride, but metaphosphoric acid gradually changes into the ortho-acid when its solution is boiled, or when it is kept for any length of time, especially in the presence of acids.[21]

[21] As when using phenolphthalein as an indicator in neutralising by
an alkali metaphosphoric acid is monobasic, and orthophosphoric
acid is bibasic, it is possible by means of this difference to
follow the transition of meta- into orthophosphoric acid. Sabatier
(1888) carried on an investigation of this nature, and found that
the rate of transformation is dependent on the temperature, and is
subject to the general laws of the rate of chemical
transformations which belongs to physical chemistry.

Metaphosphoric acid has a particular interest in respect to the
variations to which its salts are subject. The metaphosphates are
formed by the ignition of the acid orthophosphates, MH_{2}PO_{4},
or MNH_{4}HPO_{4}, or of the acid pyrophosphates,
M_{2}H_{2}P_{2}O_{7}, or M_{2}(NH_{4})_{2}P_{2}O_{7}, water and
ammonia being given off in the process. The properties of the
metaphosphates, which have a similar composition to nitrates--for
instance, NaPO_{3}, or Ba(PO_{3})_{2}--vary according to the
duration of the ignition to which the ortho-, or pyrophosphates
from which they are prepared have been subjected. When the salts
NaH_{2}PO_{4} or NH_{4}NaHPO_{4} are strongly ignited, a salt
NaPO_{3} is formed, which deliquesces in the air, and gives a
gelatinous precipitate with salts of the alkaline earths. But, as
Graham (in 1830-40), and many others, especially Fleitmann and
Henneberg (in 1840-50), and Tamman (in the nineties), observed,
under other conditions the salts of the same composition acquire
other properties. The above chemists recognise five polymeric
forms of metaphosphates, (HPO_{3})_{_n_}. We will follow the
nomenclature and researches of Fleitmann.

_Monometaphosphoric acid._ The salts are distinguished for their
insolubility in water; even the salts NaPO_{3}, KPO_{3}, are
insoluble. They are obtained by igniting the monometallic
orthophosphates--for example, RH_{2}PO_{4}--up to the temperature
at which all water is evolved (316°), but not to fusion. No double
salts are known.

_Dimetaphosphoric acid_, on the contrary, easily forms double
salts--for example, KNaP_{2}O_{6}, and also the copper potassium
salt, &c. The copper salt is obtained by evaporating a solution of
copper oxide in orthophosphoric acid. A blue ortho-salt,
CuRHO_{4}, first separates from the solution, then a light-blue
pyro-salt, Cu_{2}P_{2}O_{7}; and above 350°, when metaphosphoric
acid itself begins to volatilise, the dimetaphosphate,
CuP_{2}O_{6}, is formed. The residue is washed with water, and
decomposed with a hot solution of sodium sulphide, when the sodium
salt, Na_{2}P_{2}O_{6}, is obtained in solution. This salt, when
evaporated with alcohol, gives crystals containing 2 mol. H_{2}O,
which, however, retain their solubility (in 7 parts of water)
after the water is driven off at 100°. When fused, these crystals
give a deliquescent salt (hexa-metaphosphate). The solution of the
salt has a neutral reaction, which only after prolonged boiling
becomes acid, owing to the formation of orthophosphate,
NaH_{2}PO_{4}. The soluble salts of dimetaphosphoric acid give the
insoluble silver salt, Ag_{2}P_{2}O_{6}, with silver nitrate, and
a precipitate of BaP_{2}O_{6}2H_{2}O with barium chloride.

_Trimetaphosphoric acid_ is obtained as the sodium salt
Na_{3}P_{3}O_{9} when any other metaphosphate of sodium is fused
and _slowly_ cooled, then dissolved in a slight excess of warm
water, and the resultant solution evaporated. The crystals contain
6 mol. H_{2}O, and dissolve in four parts of water. An acid
reaction is only obtained, as with the preceding salt, after
prolonged boiling with water. The acid is a true analogue of
nitric acid, because _all its metallic salts are soluble_.

_Hexametaphosphoric acid._ Fleitmann so named the ordinary
metaphosphoric acid (glacial) which attracts moisture. The
deliquescent sodium salt is obtained, like the trimetaphosphate,
only by _rapid_ cooling. It is also formed by fusing silver oxide
with an excess of phosphoric acid. The sodium salt is soluble in
water, and gives viscous, elastic precipitates with salts of Ba,
Ca, and Mg. Lubert (1893) obtained salts of Ag, Pb, &c.

Jawein and Thillot (1889), who investigated the sodium salts of
metaphosphoric acid by Raoult's method, came to the conclusion
that the salts of di- and tri-metaphosphoric acid behave in such a
manner that their molecule must be represented as non-polymerised
NaPO_{3}, whilst those of hexametaphosphoric acid behave as
(NaPO_{3})_{4}. At all events, the series of salts which Fleitmann
and Henneberg regard as monometaphosphates--_i.e._ as
non-polymerised--are most probably the most polymerised, because
they are insoluble.

According to Tamman's researches, vitreous metaphosphoric acid
contains a mixture consisting chiefly of two varieties, differing
in the solubility and degree of stability of their salts. The
least stable corresponds to Fleitmann's hexa-acid, and gives three
isomeric salts. Tamman came to the conclusion that there exist
polymers also in the form of penta-, ortho-, and
deca-metaphosphoric acids. Without going into details upon this
subject, I do not think it superfluous to point out that the
undoubted capability of metaphosphoric acid to polymerise should
be connected with its faculty of combining with water, whilst the
degree of polymerisation and the number of polymeric forms cannot
yet be considered as sufficiently explained.

In order to see the relation between phosphoric acid and the lower acids of phosphorus, it is simplest to imagine the substitution of hydroxyl in H_{3}PO_{4} or PO(OH)_{3} by hydrogen. Then from orthophosphoric acid, PO(OH)_{3}, we shall obtain phosphorous acid, POH(OH)_{2}, and hypophosphorous acid, POH(OH); and, furthermore, phosphorous acid should be bibasic if orthophosphoric acid was tribasic, and hypophosphorous acid should be monobasic. This conclusion[21 bis] is, in fact, true, and hence all the acids of phosphorus may be referred to one common type, PX_{5}, whose representatives are PH_{4}I and PCl_{5}, POCl_{3}, PCl_{2}F_{3}, &c.

[21 bis] The bibasity of H_{3}PO_{3}, established by Würtz, has been
proved by many direct experiments (see, for instance, Note 22),
among which we may mention that Amat (1892) took a mixture of the
aqueous solutions of Na_{2}HPO_{3} and NaHO and added absolute
alcohol to it. Two layers were formed; the upper, alcoholic,
contained all the excess of NaHO, whilst the lower only contained
the salt Na_{2}HPO_{3}, which was therefore unable to react with
the excess of NaHO. Amat also obtained NaH_{2}PO_{3} by saturating
H_{3}PO_{3} with soda until he obtained a neutral reaction with
methyl-orange. The replacement of one atom of H by sodium here, as
in phosphoric acid (Note 16), gives more heat than the replacement
of the second atom. For the third atom there is no formation of a
salt, and therefore no evolution of heat. The monometallic
salts--for example, NaH_{2}PO_{3}--or the ammonia salts, when
heated to 160°, give, as Amat had previously shown, a salt of
bibasic pyrophosphorous acid, Na_{2}H_{2}P_{2}O_{5}.

_Phosphorous acid_, PH_{3}O_{3}, is generally obtained from phosphorus trichloride, PCl_{3}, by the action of water: PCl_{3} + 3H_{2}O = 3HCl + PH_{3}O_{3}. Both acids formed are soluble in water, but are easily separated, because hydrochloric acid is volatile whilst phosphorous acid volatilises with difficulty, and if a small amount of water be originally taken the hydrochloric acid nearly all passes off directly. Concentrated solutions of phosphorous acid give crystals of H_{3}PO_{3}, which fuse at 70°, attract moisture from the air, and deliquesce when ignited, giving phosphine and phosphoric acid,[22] and are oxidised into orthophosphoric acid by many oxidising agents. In its salts only two hydrogen atoms are replaced by metals (Würtz); the salts of the alkaline metals are soluble, and give precipitates with salts of the majority of other metals.

[22] Phosphorous acid, when subjected to the action of nascent hydrogen
(zinc and sulphuric acid), evolves phosphine, and when boiled with
an excess of alkali it evolves hydrogen (PH_{3}O_{3} + 3KHO =
PK_{3}O_{4} + 2H_{2}O + H_{2}); owing to its liability to
oxidation, it is a reducing agent--for instance, it reduces cupric
chloride to cuprous chloride, and precipitates silver from the
nitrate and mercury from its salts.

These reactions are perhaps connected with the fact that in this
acid one atom of hydrogen should be considered as in the same
condition as in phosphuretted hydrogen, which is expressed by the
formula PHO(OH)_{2}, if we represent it as PH_{4}X, with the
substitution of two of the hydrogen atoms by oxygen and of HX by
two of hydroxyl. The direct passage of phosphorous chloride into
phosphorous acid would, however, indicate that all the three atoms
of hydrogen in it occur in the form of hydroxyl, because no
difference is known between the three atoms of chlorine in
PCl_{3}--they all react alike, as a rule. However, Menschutkin, by
acting on alcohol, C_{2}H_{5}OH, with phosphorous chloride,
obtained hydrochloric acid and a substance P(C_{2}H_{5}O)Cl_{2},
and from it by the action of bromine he obtained ethyl bromide,
C_{2}H_{5}Br, and a compound PBrOCl_{2}, which proves, to a
certain extent, the existence of a difference between the three
atoms of chlorine in phosphorous chloride. If we turn our
attention to the formation of phosphine by the ignition of
phosphorous acid, we see that 4PH_{3}O_{3} only evolve 3H in the
form of PH_{3}, and therefore the residue--that is,
3PH_{3}O_{4}--will still contain one hydrogen of the same nature
as in phosphine, because in 4PH_{3}O_{3} we should recognise four
such hydrogens as in phosphine. We arrive at the same conclusion
by examining the decomposition of hypophosphorous acid,
2PH_{3}O_{2} = PH_{3} + PH_{3}O_{4}. In the two molecules of the
monobasic hypophosphorous acid taken, there are only two atoms of
hydrogen replaceable by metals, whilst in the molecule of the
resultant phosphoric acid there are three. Perhaps relations of
this nature determine the relative stability of the dimetallic
salts of orthophosphoric acid.

The monobasic _hypophosphorous acid_, PH_{3}O_{2}, gives salts PH_{2}O_{2}Na, (PH_{2}O_{2})_{2}Ba, &c.; the two remaining atoms of hydrogen (which exist in the same form as in phosphine, PH_{3}) are not replaceable by metals, and this determines the property of these salts of evolving phosphuretted hydrogen when heated (especially with alkalis). In acting on substances liable to reduction it is this hydrogen which acts, and, for example, _reduces_ gold and mercury from the solutions of their salts, or converts cupric into cuprous salts. In all these instances the hypophosphorous acid is converted into phosphoric acid. Under the action of zinc and sulphuric acid it gives phosphine, PH_{3}. Nevertheless, neither hypophosphorous acid nor its dry salts absorb oxygen from the air. The salts of hypophosphorous acid are more soluble than those of the preceding acids of phosphorus. Thus the sodium salt PNaH_{2}O_{2} does not give a precipitate with barium chloride, and the salts of calcium, barium, and many other metals are soluble.[23] The hypophosphites are prepared by boiling an alkali with phosphorus so long as phosphuretted hydrogen is evolved. The acid itself is obtained from barium hypophosphite (prepared in the same manner by boiling phosphorus in baryta water), by decomposing its solution with sulphuric acid. By concentration of the solution of hypophosphorous acid (it must not be heated above 130°, at which temperature it decomposes) a syrup is formed which is able to crystallise. In the solid state hypophosphorous acid fuses at +17°, and has the properties of a clearly defined acid.

[23] Calcium hypophosphite is used in medicine. According to Cavazzi, a
mixture of sodium hypophosphite, NaH_{2}PO_{2}, and sodium nitrate
explodes violently.

The types PX_{3} and PX_{5}, which are evident for the hydrogen and oxygen compounds of phosphorus, are most clearly seen in its halogen compounds,[24] to the consideration of which we will proceed, fixing our attention more especially on the chlorine compounds, as being the most important from the historical, theoretical, and practical point of view.

[24] Fluorine and bromine give PX_{3} and PX_{5}, like chlorine. With
respect to iodine PI_{5} is, in a chemical sense, a very unstable
substance, and generally _phosphorus tri-iodide_ only is formed
(from yellow or red phosphorus and iodine in the requisite
proportions). It is a red crystalline substance, fuses at 55°, is
easily decomposed by water, forming phosphorous and hydriodic
acids, and when heated it evolves iodine vapours and forms
phosphorus di-iodide, PI_{2}. This substance may be obtained in
the same manner as the preceding by taking a smaller proportion of
iodine (8 parts of iodine to 1 part of phosphorus, whilst the
tri-iodide requires 12·3); it also forms red crystals, which melt
at 110°. When decomposed by water it not only gives phosphorous
and hydriodic acids, but also phosphine and a yellow substance (a
lower oxide of phosphorus). In its composition di-iodide of
phosphorus corresponds with liquid phosphuretted hydrogen, PH_{2},
and probably its molecular weight is much higher: P_{2}I_{4} or
P_{3}I_{6}, &c. As the iodine compounds of phosphorus give
hydriodic and phosphorous acids with water, and as both these
substances are reducing agents in the presence of water (and
hydrates), iodide of phosphorus also acts as a reducing agent.

Phosphorus burns in chlorine, forming phosphorous chloride, PCl_{3}, and with an excess of chlorine, phosphoric chloride, PCl_{5}. The oxychloride, POCl_{3}, as the simplest chloranhydride according to the type PX_{5}, and also phosphoric chloride, correspond with orthophosphoric acid, PO(OH)_{3}, while phosphorous chloride, PCl_{3}, corresponds with phosphorous acid and the type PX_{3}. Phosphoric oxychloride, POCl_{3}, is a colourless liquid, boiling at 110°. Phosphorus trichloride is also a colourless liquid, boiling at 76°,[25] whilst phosphoric chloride is a solid yellowish substance, which volatilises without melting at about 168°. They are all heavier than water, and form types of the _chloranhydrides_ or chlorine compounds of the non-metallic elements whose hydrates are acids, just as NaCl or BaCl_{2} are types of halogen metallic salts.

[25] In a liquid state the density of phosphorous chloride at 10° =
1·597, and therefore its molecular volume = 137·5/1·597 = 86·0,
and that of phosphorus oxychloride is equal to 153·5/1·693 = 90·7;
hence the addition of oxygen has produced considerable increase in
volume, just as in the conversion of sulphur dichloride, SCl_{2},
into sulphuryl chloride, SOCl_{2}, the volume changes from 64 to
71. It is the same with the boiling-points; phosphorus trichloride
boils at 70°, the oxychloride at 100°, sulphur dichloride at 64°,
and sulphuryl chloride at 78°--that is, the addition of oxygen
raises the boiling points.

_The vapour density_ of phosphorus trichloride and oxychloride
corresponds with their formulæ (Cahours, Würtz)--namely, is equal
to half the molecular weight referred to hydrogen. But it is not
so with phosphorus pentachloride. Cahours showed that the vapour
density of phosphorus pentachloride referred to air = 3·65, to
hydrogen = 52·6, whilst according to the formula PCl_{5} it should
be = 104·2. Hence this formula corresponds with four, and not with
two, molecules. This shows that the vapour of phosphoric chloride
contains two and not one molecule, that in a state of vapour it
splits up, like sal-ammoniac, sulphuric acid, &c. The products of
disruption must here be phosphorous chloride, PCl_{3}, and
chlorine, Cl_{2}, bodies which easily re-form phosphoric chloride,
PCl_{5}, at a lower temperature. This decomposition of phosphoric
chloride in its conversion into vapour is confirmed by the fact
that the vapour of this almost colourless substance shows the
greenish-yellow colour proper to chlorine. This dissociation of
phosphoric chloride has been considered by some chemists as a sign
that phosphorus, like nitrogen, does not give volatile compounds
of the type PX_{5}, and that such substances are only obtained as
unstable molecular compounds which break up when distilled; for
example, PH_{3},HI, PCl_{3},Cl_{2}, NH_{3},HCl, &c. To prove that
the molecule PCl_{5} actually exists, Würtz in 1870 observed that
when mixed with the vapour of phosphorous chloride the vapour of
phosphoric chloride distils over (from 160° to 190°) perfectly
colourless, and has a density which is really near to the
formula--namely, to 104--and the same density was determined for
the pentachloride in an atmosphere of chlorine. Hence at low
temperatures and in admixture with one of the products of
dissociation, there is no longer that decomposition which occurs
at higher temperatures--that is, we have here a case of
dissociation proceeding at moderate temperatures.

An important proof in favour of the type PX_{5} is exhibited by
phosphorus pentafluoride PF_{5}, obtained by Thorpe as a
colourless gas which only corrodes glass after the lapse of time;
it may be kept over mercury, and has a normal density. It is
formed when liquid arsenic trifluoride, AsF_{3}, is added to
phosphoric chloride surrounded by a freezing mixture: 3PCl_{5} +
5AsF_{3} = 3PF_{5} + 5AsCl_{3}.

In general, fluorine and phosphorus give stable compounds: PF_{3},
POF_{3}, and PF_{5}, as would be expected from the fact that in
passing from Cl to I (_i.e._ as the atomic weight of the halogen
increases) the stability of the compounds with P and the tendency
to give PX_{5} (Note 24) decreases. _Phosphorus trifluoride_ is
obtained by heating a mixture of ZnF_{2} and PBr_{3}, by the
action of AsF_{3} upon PCl_{3}, by heating phosphide of copper
with PbF_{2}, &c. It is a strong-smelling gas, which liquefies at
-10° under a pressure of 40 atmospheres, giving a colourless
liquid. It dissolves easily in (is absorbed by, reacts with)
water, and acts upon glass; when mixed with Cl_{2} it combines
with it (Poulenc, 1891), forming PCl_{2}F_{3}, a colourless gas of
normal density, which is transformed into a liquid at 8°,
decomposes into PF_{3} + Cl_{2} at 250°, and, with a small amount
of water, gives _oxy-fluoride_ of phosphorus, POF_{3} (with a
large amount of water it gives PH_{3}O_{4}), which Moissan (1891)
obtained by the action of dry HF upon P_{2}O_{5}, and Thorpe and
Tutton (1890) by heating a mixture of cryolite and P_{2}O_{5}. It
is a gas of normal density, like PF_{3}, and was obtained by
Moissan by the action of fluorine upon PF_{3} (PSF_{3}, _see_