PHOSPHORUS AND THE OTHER ELEMENTS OF THE FIFTH GROUP
Nitrogen is the lightest and most widely distributed representative of
the elements of the fifth group, which form a higher saline oxide of the
form R_{2}O_{5}, and a hydrogen compound of the form RH_{3}. Phosphorus,
arsenic, bismuth, and antimony belong to the uneven series of this group.
_Phosphorus_ is the most widely distributed of these elements. There is
hardly any mineral substance composing the mass of the earth's crust
which does not contain some--it may be a small--amount of phosphorus
compounds in the form of the salts of phosphoric acid. The soil and
earthy substances in general usually contain from one to ten parts of
phosphoric acid in 10,000 parts. This amount, which appears so small,
has, however, a very important significance in nature. No plant can
attain its natural growth if it be planted in an artificial soil
completely free from phosphoric acid. Plants equally require the presence
of potash, magnesia, lime, and ferric oxide, among basic, and of
carbonic, sulphuric, nitric, and phosphoric anhydrides, among acid
oxides. In order to increase the fertility of a more or less poor soil,
the above-named nutritive elements are introduced into it by means of
fertilisers. Direct experiment has proved that these substances are
undoubtedly necessary to plants, but that they must be all present
simultaneously and in small quantities, and that an excess, like an
insufficiency, of one of these elements is necessarily followed by a bad
harvest, or an imperfect growth, even if all the other conditions (light,
heat, water, air) are normal. The phosphoric compounds of the soil
accumulated by plants pass into the organism of animals, in which these
substances are assimilated in many instances in large quantities. Thus
the chief component part of bones is calcium phosphate, Ca_{3}P_{2}O_{8},
and it is on this that their hardness depends.[1]
[1] Dry bones contain about one-third of gelatinous matter and about
two-thirds of ash, chiefly calcium phosphate. The salts of
phosphoric acid are also found in the mass of the earth as separate
minerals; for example, the _apatites_ contain this salt in a
crystalline form, combined with calcium chloride or fluoride,
CaR_{2},3Ca_{3}(PO_{4})_{2}, where R = F or Cl, sometimes in a
state of isomorphous mixture. This mineral often crystallises in
fine hexagonal prisms; sp. gr. 3·17 to 3·22. Vivianite is a
hydrated ferrous phosphate, Fe_{3}(PO_{4})_{2},8H_{2}O. Phosphates
of copper are frequently found in copper mines; for example,
_tagilite_, Cu_{3}(PO_{4})_{2},Cu(OH)_{2},2H_{2}O. Lead and
aluminium form similar salts. They are nearly all insoluble in
water. The turquoise, for instance, is hydrated phosphate of
alumina, (Al_{2}O_{3})_{2},P_{2}O_{5}5H_{2}O, coloured with a salt
of copper. Sea and other waters always contain a small amount of
phosphates. The ash of sea-plants, as well as of land-plants,
always contains phosphates. Deposits of calcium phosphate are often
met with; they are termed _phosphorites_ and _osteolites_, and are
composed of the fossil remains of the bones of animals; they are
used for manure. Of the same nature are the so-called guano
deposits from Baker's Island, and entire strata in Spain, France,
and in the Governments of Orloff and Kursk in Russia. It is evident
that if a soil destined for cultivation contain very little
phosphoric acid, the fertilisation by means of these minerals will
be beneficial, but, naturally, only if the other elements necessary
to plants be present in the soil.
Phosphorus was first extracted by Brand in 1669, by the ignition of
evaporated urine. After the lapse of a century Scheele, who knew of the
existence of a more abundant source of phosphorus in bones, pointed out
the method which is now employed for the extraction of this element.
Calcium phosphate in bones permeates a nitrogenous organic substance,
which is called ossein, and forms a gelatin. When bones are treated
exclusively for the extraction of phosphorus, neglecting the gelatin,
they are burnt, in which case all the ossein is burnt away. When,
however, it is desired to preserve the gelatin, the bones are immersed in
cold dilute hydrochloric acid, which dissolves the calcium phosphate and
leaves the gelatin untouched; calcium chloride and acid calcium
phosphate, CaH_{4}(PO_{4})_{2}, are then obtained in the solution. When
the bones are directly burnt in an open fire their mineral components
only are left as an ash, containing about 90 per cent. of calcium
phosphate, Ca_{3}(PO_{4})_{2}, mixed with a small amount of calcium
carbonate and other salts. This mass is treated with sulphuric acid, and
then the same substance is obtained in the solution as was obtained from
the unburnt bones immersed in hydrochloric acid--_i.e._ the acid calcium
phosphate soluble in water, in which reaction naturally the chief part of
the sulphuric acid is converted into calcium sulphate:
Ca_{3}(PO_{4})_{2} + 2H_{2}SO_{4} = 2CaSO_{4} + CaH_{4}(PO_{4})_{2}.
Ca_{3}(PO_{4})_{2} + 4HCl = 2CaCl_{2} + CaH_{4}(PO_{4})_{2}.
On evaporating the solution, crystallisable acid calcium phosphate is
obtained. The extraction of the phosphorus from this salt consists in
_heating it with charcoal to a white heat_. When heated, the acid
phosphate, CaH_{4}(PO_{4})_{2}, first parts with water, and forms the
metaphosphate, Ca(PO_{3})_{2}, which for the sake of simplicity may be
regarded, like the acid salt, as composed of pyrophosphate and phosphoric
anhydride, 2Ca(PO_{3})_{2} = Ca_{2}P_{2}O_{7} + P_{2}O_{5}. The latter,
with charcoal, gives phosphorus and carbonic oxide, P_{2}O_{5} + 5C =
P_{2} + 5CO. So that in reality a somewhat complicated process takes
place here, yielding ultimately products according to the following
equation:
2CaH_{4}(PO_{4})_{2} + 5C = 4H_{2}O + Ca_{2}P_{2}O_{7} + P_{2} + 5CO.
After the steam has come over, phosphorus and carbonic oxide distil over
from the retort and calcium pyrophosphate remains behind.[1 bis]
[1 bis] By subjecting the pyrophosphate to the action of sulphuric or
hydrochloric acid it is possible to obtain a fresh quantity of the
acid salt from the residue, and in this manner to extract all the
phosphorus. It is usual to take burnt bones, but mineral
phosphorites, osteolites, and apatites may also be employed as
materials for the extraction of phosphorus. Its extraction for the
manufacture of matches is everywhere extending, and in Russia, in
the Urals, in the Government of Perm, it has attained such
proportions that the district is able to supply other countries
with phosphorus. A great many methods have been proposed for
facilitating the extraction of phosphorus, but none of them differ
essentially from the usual one, because the problem is dependent on
the liberation of phosphoric acid by the action of acids, and on
its ultimate reduction by charcoal. Thus the calcium phosphate may
be mixed directly with charcoal and sand, and phosphorus will be
liberated on heating the mixture, because the silica displaces the
phosphoric anhydride, which gives carbonic oxide and phosphorus
with the charcoal. It has also been proposed to pass hydrochloric
acid over an incandescent mixture of calcium phosphate and
charcoal; the acid then acts just as the silica does, liberating
phosphoric anhydride, which is reduced by the charcoal. It is
necessary to prevent the access of air in the condensation of the
vapours of phosphorus, because they take fire very easily; hence
they are condensed under water by causing the gaseous products to
pass through a vessel full of water. For this purpose the condenser
shown in fig. 83 is usually employed.
[Illustration: FIG. 83.--Preparation of phosphorus. The mixture is
calcined in the retort _c_. The vapours of phosphorus pass through _a_
into water without coming into contact with air. The phosphorus condenses
in the water, and the gases accompanying it escape through _i_.]
As phosphorus melts at about 40°, it condenses at the bottom of the
receiver in a molten liquid mass, which is cast under water in tubes, and
is sold in the form of sticks. This is common or _yellow phosphorus_. It
is a transparent, yellowish, waxy substance, which is not brittle, almost
insoluble in water, and easily undergoes change in its external
appearance and properties under the action of light, heat, and of various
substances. It crystallises (by sublimation or from its solution in
carbon bisulphide) in the regular system, and[2] (in contradistinction to
the other varieties) is easily soluble in carbon bisulphide, and also
partially in other oily liquids. In this it recalls common sulphur. Its
specific gravity is 1·84. It fuses at 44°, and passes into vapour at
290°; it is easily inflammable, and must therefore be handled with great
caution; careless rubbing is enough to cause phosphorus to ignite. Its
application in the manufacture of matches is based on this.[2 bis] It
emits light in the air owing to its slow[3] oxidation, and is therefore
kept under water (such water is phosphorescent in the dark, like
phosphorus itself). It is also very easily oxidised by various oxidising
agents and takes up the oxygen from many substances.[3 bis] Phosphorus
enters into direct combination with many metals and with sulphur,
chlorine, &c., with development of a considerable amount of heat. It is
very poisonous although not soluble in water.
[2] Vernon (1891) observed that ordinary (yellow) phosphorus is
dimorphous. If it be melted and by careful cooling be brought in a
liquid form to as low a temperature as possible, it gives a variety
which melts at 45°·3 (the ordinary variety fuses at 44°·3), sp. gr.
1·827 (that of the ordinary variety is 1·818) at 13°, crystallises
in rhombic prisms (instead of in forms belonging to the cubical
system). This is similar to the relation between octahedral and
prismatic sulphur (Chapter XX.).
[2 bis] According to Herr Irinyi (an Hungarian student), the first
phosphorus matches were made in Austria at Roemer's works in 1835.
[3] The absorption of the oxygen of the atmosphere at a constant
ordinary temperature by a large surface of phosphorus proceeds so
uniformly, regularly, and rapidly, that it may serve, as Ikeda
(Tokio, 1893) has shown, for demonstrating the law of the velocity
(rate) of reaction, which is considered in theoretical chemistry,
and shows that the rate of reaction is proportional to the active
mass of a substance--_i.e._ _dx_/_dt_ = _k_(A - _x_) where _t_ is
the time, A the initial mass of the reacting substance--in this
case oxygen--_x_ the amount of it which has entered into reaction,
and _k_ the coefficient of proportionality. Ikeda took a test-tube
(diameter about 10 mm.), and covered its outer surface with a
coating of phosphorus (by melting it in a test-tube of large
diameter, inserting the smaller test-tube, and, when the phosphorus
had solidified, breaking away the outer test-tube), and introduced
it into a definite volume of air, contained in a Woulfe's bottle
(immersed in a water bath to maintain a constant temperature), one
of whose orifices was connected with a mercury manometer showing
the fall of pressure, _x_. Knowing that the initial pressure of the
oxygen (in air nearly 750 × ·0209) was about 155 mm. = A, the
coefficient of the rate of reaction _k_ is given, by the law of the
variation of the rate of reaction with the mass of the reacting
substance, by the equation: _k_ = (1/_t_)log(A/(A - _x_)), where
_t_ is the time, counting from the commencement, of the experiment
in minutes. When the surface of the phosphorus was about 11 sq.
cm., the following results were actually obtained.
_t_ = 10 20 30 40 50 60 minutes
_x_ = 10·5 21·5 31·1 40·7 49·1 57·3 mm
10,000 _k_ = 32 32 32 33 33 33
The constancy of _k_ is well shown in this case. The determination
takes a comparatively short time, so that it may serve as a lecture
experiment, and demonstrates one of the most important laws of
chemical mechanics.
[3 bis] Not only do oxidising agents like nitric, chromic, and similar
acids act upon phosphorus, but even the alkalis are attacked--that
is, phosphorus acts as a reducing agent. In fact it reduces many
substances, for instance, copper from its salts. When phosphorus is
heated with sodium carbonate, the latter is partially reduced to
carbon. If phosphorus be placed under water slightly warmed, and a
stream of oxygen be passed over it, it will burn under the water.
Besides this, there is a red variety of phosphorus, which differs
considerably from the above. _Red phosphorus_ (sometimes wrongly called
_amorphous phosphorus_) is partially formed when ordinary phosphorus
remains exposed to the action of light for a long time. It is also formed
in many reactions; for example, when ordinary phosphorus combines with
chlorine, bromine, iodine, or oxygen, a portion of it is converted into
red phosphorus. Schrötter, in Vienna, investigated this variety of
phosphorus, and pointed out by what methods it may be produced in
considerable quantities. Red phosphorus is a powdery red-brown opaque
substance of specific gravity 2·14. It does not combine so energetically
with oxygen and other substances as yellow phosphorus, and evolves less
heat in combining with them.[4] Common phosphorus easily oxidises in the
air; red phosphorus does not oxidise at all at the ordinary temperature;
hence it does not phosphoresce in the air, and may be very conveniently
kept in the form of powder. It does not, like yellow phosphorus, fuse at
44°. After being converted into vapour at 290° or 300°, it again passes
into the ordinary variety when slowly cooled. Red phosphorus is not
soluble in carbon bisulphide and other oily liquids, which permits of its
being freed from any admixture of the ordinary phosphorus. It is not
poisonous, and is used in many cases for which the ordinary phosphorus is
unsuitable or dangerous; for example, in the manufacture of matches,
which are then not poisonous or inflammable by accidental friction, and
therefore the red variety has now replaced the ordinary
phosphorus.[4 bis]
[4] The thermochemical determinations for phosphorus and its compounds
date from the last century, when Lavoisier and Laplace burnt
phosphorus in oxygen in an ice calorimeter. Andrews, Despretz,
Favre, and others have studied the same subject. The most accurate
and complete data are due to Thomsen. To determine the heat of
combustion of yellow phosphorus, Thomsen oxidised it in a
calorimeter with iodic acid in the presence of water, and a mixture
of phosphorous and phosphoric acids was thus formed (was not any
hypophosphoric acid formed?--Salzer), and the iodic acid converted
into hydriodic acid. It was first necessary to introduce two
corrections into the calorimetric result obtained, one for the
oxidation of the phosphorous into phosphoric acid, knowing their
relative amounts by analysis, and the other for the deoxidation of
the iodic acid. The result then obtained expresses the conversion
of phosphorous into hydrated phosphoric acid. This must be
corrected for the heat of solution of the hydrate in water, and for
the heat of combination of the anhydride with water, before we can
obtain the heat evolved in the reaction of P_{2} with O_{5} in the
proportion for the formation of P_{2}O_{5}. It is natural that with
so complex a method there is a possibility of many small errors,
and the resultant figures will only present a certain degree of
accuracy after repeated corrections by various methods. Of such a
kind are the following figures determined by Thomsen, which we
express in thousands of calories:--P_{2} + O_{5} = 370; P_{2} +
O_{3} + 3H_{2}O = 400; P_{2} + O_{5} + a mass of water = 405. Hence
we see that P_{2}O_{5} + 3H_{2}O = 30; 2PH_{3}O_{4} + an excess of
water = 5. Experiment further showed that crystallised PH_{3}O_{4},
in dissolving in water, evolves 2·7 thousand calories, and that
fused (39°) PH_{3}O_{4} evolves 5·2 thousand calories, whence the
heat of fusion of H_{3}PO_{4} = 2·5 thousand calories. For
phosphorous acid, H_{3}PO_{3}, Thomsen obtained P_{2} + O_{3} +
3H_{2}O = 250, and the solution of crystallised H_{3}PO_{3} in
water = -0·13, and of fused H_{3}PO_{3} = +2·9. For hypophosphorous
acid, H_{3}PO_{2}, the heats of solution are nearly the same (-0·17
and +2·1), and the heat of formation P_{2} + O + 3H_{2}O = 75;
hence its conversion into 2H_{3}PO_{3} evolves 175 thousand
calories, and the conversion of 2H_{3}PO_{3} into 2H_{3}PO_{4} =
150 thousand calories. For the sake of comparison we will take the
combination of chlorine with phosphorus, also according to Thomsen,
per 2 atoms of phosphorus, P_{2} + 3Cl_{2} = 151, P_{2} + 5Cl_{2} =
210 thousand calories. In their reaction on an excess of water
(with the formation of a solution), 2PCl_{3} = 130, 2PCl_{5} = 247,
and 2POCl_{3} = 142 thousand calories.
Besides which we will cite the following data given by various
observers: heat of fusion for P (that is, for 31 parts of
phosphorus by weight) -0·15 thousand calories; the conversion of
yellow into red phosphorus for P, from +19 to +27 thousand
calories; P + H_{3} = 4·3, HI + PH_{3} = 24, PH_{3} + HBr = 22
thousand calories.
At the ordinary temperature (20° C.) phosphorus is not oxidised by
pure oxygen; oxidation only takes place with a slight rise of
temperature, or the dilution of the oxygen with other gases
(especially nitrogen or hydrogen), or a decrease of pressure.
[4 bis] Ordinary phosphorus takes fire at a temperature (60°) at which
no other known substance will burn. Its application to the
manufacture of matches is based on this property. In order to
illustrate the easy inflammability of common (yellow) phosphorus,
its solution in carbon bisulphide may be poured over paper; this
solvent quickly evaporates, and the free phosphorus spread over a
large surface takes fire spontaneously, notwithstanding the cooling
effect produced by the evaporation of the bisulphide. The majority
of _phosphorus matches_ are composed of common phosphorus mixed
with some oxidising substance which easily gives up oxygen, such as
lead dioxide, potassium chlorate, nitre, &c. For this purpose
common phosphorus is carefully triturated under warm water
containing a little gum; lead dioxide and potassium nitrate are
then added to the resultant emulsion, and the match ends,
previously coated with sulphur or paraffin, are dipped into this
preparation. After this the matches are dipped into a solution of
gum and shellac, in order to preserve the phosphorus from the
action of the air. When such a match containing particles of yellow
phosphorus is rubbed over a rough surface, it becomes (especially
at the point of rupture of the brittle gummy coating) slightly
heated, and this is sufficient to cause the phosphorus to take fire
and burn at the expense of the oxygen of the other ingredients.
The heads of the 'safety' matches do not contain any phosphorus, but
only substances capable of burning and of supporting combustion. Red
phosphorus is spread over a surface on the box, and it is the friction
against this phosphorus which ignites the matches. There is no danger of
the matches taking fire accidentally, nor are they poisonous.[5] This red
phosphorus is prepared by heating the ordinary phosphorus at 230° to
270°; it is evident that this must be done in an atmosphere incapable of
supporting combustion--for example, in nitrogen, carbonic anhydride,
steam, &c. On a large scale, ordinary phosphorus is placed in closed iron
vessels,[5 bis] and immersed in a bath of different proportions of tin
and lead, by which means the temperature of 250° necessary for the
conversion is easily attained. It is kept at this temperature for some
time. The temperature is at first cautiously raised, and the air is thus
partially expelled by the heat, and also by the evolution of steam (the
phosphorus is damp when put in), whilst the remaining oxygen is also
partially absorbed by the phosphorus, so that an atmosphere of nitrogen
is produced in the iron vessel. Red phosphorus enters into all the
reactions proper to yellow phosphorus, only with greater difficulty and
more slowly;[6] and, as its vapour tension (volatility) is less than that
of the yellow variety, it may be supposed that a polymerisation takes
place in the passage of the yellow into the red modification, just as in
the passage of cyanogen into paracyanogen, or of cyanic acid into
cyanuric acid (Chapter IX. Notes 39 bis and 48).
[5] In the so-called 'safety' or Swedish matches (which are not
poisonous, and do not take fire from accidental friction) a mixture
of red phosphorus and glass forms the surface on which the matches
are struck, and the matches themselves do not contain any
phosphorus at all, but a mixture of antimonious sulphide,
Sb_{2}S_{3} (or similar combustible substances) and potassium
chlorate (or other oxidising agents). The combustion, when once
started by contact with the red phosphorus, proceeds by itself at
the expense of the inflammatory and combustible elements contained
in the tip of the match. The mixture applied on the match itself
must not be liable to take fire from a blow or friction. The
mixture forming the heads of the 'safety' matches has the following
approximate composition: 55-60 parts of chlorate of potassium, 5-10
parts of peroxide of manganese (or of K_{2}Cr_{2}O_{7}), about 1
part of sulphur or charcoal, about 1 part of pentasulphide of
antimony, Sb_{2}S_{5}, and 30-40 parts of rouge and powdered glass.
This mixture is stirred up in gum or glue, and the matches are
dipped into it. The paper on which the matches are struck is coated
with a mixture of red phosphorus and trisulphide of antimony,
Sb_{2}S_{3}, stirred up in dextrine.
[5 bis] Phosphorus only acts on iron at a red heat. The boiler is
provided with a safety valve and gas-conducting tube, which is
immersed in mercury or other liquid to prevent the admission of air
into the boiler.
[6] The specific heat of the yellow variety is 0·189--that is, greater
than that of the red variety, which is 0·170. The sp. gr. of the
yellow is 1·84, and of the red prepared at 260° 2·15, and of that
prepared at 580° and above (_i.e._ 'metallic' phosphorus, _see_
below) = 2·34. At 230° the pressure of the vapour of ordinary
phosphorus = 514 millimetres of mercury, and of the red = 0--that
is to say, the red phosphorus does not form any vapour at this
temperature; at 447° the vapour tension of ordinary phosphorus is
at first = 5500 mm., but it gradually diminishes, whilst that of
red phosphorus is equal to 1636 mm.
Hittorf, by heating the lower portion of a closed tube containing
red phosphorus to 530° and the upper portion to 447°, obtained
crystals of the so-called 'metallic' phosphorus at the upper
extremity. As the vapour tensions (according to Hittorf, at 530°
the vapour tension of yellow phosphorus = 8040 mm., of red = 6139
mm., and of metallic = 4130 mm.) and reactions are different,
_metallic phosphorus_ may be regarded as a distinct variety. It is
still less energetic in its chemical reaction than red phosphorus,
and it is denser than the two preceding varieties: sp. gr. = 2·34.
It does not oxidise in the air; is crystalline, and has a metallic
lustre. It is obtained when ordinary phosphorus is heated with lead
for several hours at 400° in a closed vessel, from which the air
has been exhausted. The resultant mass is then treated with dilute
nitric acid, which first dissolves the lead (phosphorus is
electro-negative to lead, and does not, therefore, act on the
nitric acid at first) and leaves brilliant rhombohedral crystals of
phosphorus of a dark violet colour with a slight metallic lustre,
which conduct an electric current incomparably better than the
yellow variety; this also is characteristic of the metallic state
of phosphorus.
The researches of Lemoine partially explain the passage of yellow
(ordinary) phosphorus into its other varieties. He heated a closed
glass globe containing either ordinary or red phosphorus, in the
vapour of sulphur (440°), and then determined the amount of the red
and yellow varieties after various periods of time, by treating the
mixture with carbon bisulphide. It appeared that after the lapse of
a certain time a mixture of definite and equal composition is
obtained from both--that is, between the red and yellow varieties a
state of equilibrium sets in like that of dissociation, or that
observed in double decompositions. But at the same time, the
progress of the transformation appeared to be dependent on the
relative quantity of phosphorus taken per volume of the globe
(_i.e._ upon the pressure). Neglecting the latter, we will cite as
an example the amounts of the red phosphorus transformed into the
ordinary, and of the ordinary not converted into red, per 30 grams
of red or yellow taken per litre capacity of the globe, heated to
440°. When red phosphorus was taken, 4·75 grams of yellow
phosphorus were formed after two hours, four grams after eight
hours, three grams after twenty-four hours, and the last limit
remained constant on further heating. When thirty grams of yellow
phosphorus were taken, five grams remained unaltered after two
hours, four grams after eight hours, and after twenty-four hours
and more three grams as before. Troost and Hautefeuille showed that
liquid phosphorus in general changes more easily into the red than
does phosphorus vapour, which, however, is able, although slowly,
to deposit red phosphorus.
The question presents itself as to whether phosphorus in a state of
vapour is the ordinary or some other variety? Hittorf (1865)
collected many data for the solution of this problem, which leave
no doubt that (as experimental figures show) the density of the
vapour of phosphorus is always the same, although the vapour
tension of the different varieties and their mixtures is very
variable. This shows that the different varieties of phosphorus
only occur in a liquid and solid state, as indeed is implied in the
idea of polymerisation. Strictly speaking, the vapour of phosphorus
is a particular state of this substance, and the molecular formula
P_{4} refers only to it, and not to any other definite state of
phosphorus. But Raoult's solution method showed that in a benzene
solution the fall of the freezing point indicates for ordinary
phosphorus a molecule P_{4}, judging by the determinations of
Paterno and Nasini (1888), Hirtz (1890), and Beckmann (1891), who
obtained for sulphur by the same method a molecular weight = S_{6},
in conformity with the vapour density. Further research in this
direction will perhaps show the possibility of finding the
molecular weight of red phosphorus, if a means be discovered for
dissolving it without converting it into the yellow variety.
I think it will not be out of place here to draw the reader's
attention to the fact that red phosphorus, which we must recognise
as polymeric with the yellow, stands nearer to nitrogen, whose
molecule is N_{2}, in its small inclination towards chemical
reactions, although judging by its small vapour tension it must be
more complex than ordinary (yellow and white) phosphorus.
The vapour of phosphorus is colourless; its density remains constant
between 300° and 1000° (Dumas, 1833; Mitscherlich, Deville, and Troost,
1859, and others). The density with respect to air has been determined as
from 4·3 to 4·5. Hence, referred to hydrogen, it is 4·4 × 14·4 = 63,
corresponding with a molecular weight 124, _i.e._ the molecule of
phosphorus in a state of vapour contains P_{4}. The reader will remember
that the molecule of nitrogen contains N_{2}, of sulphur S_{6} or S_{2},
and of oxygen O_{2} or O_{3}.
The chemical energy of phosphorus in a free state more nearly approaches
that of sulphur than nitrogen. Phosphorus is combustible and inflames at
60°; but having in the act of combination parted with a portion of its
energy in the form of heat it becomes analogous to nitrogen, so long as
there is no question of its reduction back again into phosphorus. Nitric
acid is easily reduced to nitrogen, whilst phosphoric acid is reduced
with very much greater difficulty. All the compounds of phosphorus are
less volatile than those of nitrogen. Nitric acid, HNO_{3}, is easily
distilled; metaphosphoric acid, HPO_{3}, is generally said to be
non-volatile; triethylamine, N(C_{2}H_{5})_{3}, boils at 90°, and
triethylphosphine, P(C_{2}H_{5})_{3}, at 127°.
Phosphorus not only combines easily and directly with oxygen, but also
with chlorine, bromine, iodine, sulphur, and with certain metals, and red
phosphorus when heated combines with hydrogen also.[6 bis] So, for
instance, when fused with sodium under naphtha, phosphorus gives the
compound Na_{3}P_{2}. Zinc, absorbing the vapour of phosphorus, gives the
phosphide Zn_{3}P_{2} (sp. gr. 4·76); tin, SnP; copper, Cu_{2}P; even
platinum combines with phosphorus (PtP_{2}, sp. gr. 8·77).[6 tri] Iron,
when combined even with a small quantity of phosphorus, becomes
brittle.[7] Some of these compounds of phosphorus are obtained by the
action of phosphorus on the solutions of metallic salts, and by the
ignition of metallic oxides in the vapour of phosphorus, or by heating
mixtures of phosphates with charcoal and metals. Phosphides do not
exhibit the external properties of salts, which are so clearly seen in
the chlorides and still distinctly observable in the sulphides. _The
phosphides of the metals_ of the alkalis and of the alkaline earths are
even immediately and very easily decomposed by water, whereas this is
found to be the case with only a very few sulphides, and still more
rarely and indistinctly with the chlorides. We may take calcium phosphide
as an example.[7 bis] Phosphorus is laid in a deep crucible, and covered
with a clay plug, over which lime is strewn. At a red heat the vapours of
phosphorus combine with the oxygen of the lime and form phosphoric
anhydride, which forms a salt with another portion of the lime, whilst
the liberated calcium combines with the phosphorus and forms calcium
phosphide. Its composition is not quite certain; it may be CaP
(corresponding with liquid phosphuretted hydrogen). This substance is
remarkable for the following reaction: if we take water--or, better
still, a dilute solution of hydrochloric acid--and throw calcium
phosphide into it, bubbles of gas are evolved, which take fire
spontaneously in the air and form white rings. This is owing to the fact
that the liquid hydrogen phosphide, PH_{2}, is first formed, thus, CaP +
2HCl = CaCl_{2} + PH_{2}, which, owing to its instability, very easily
splits up into the solid phosphide, P_{2}H, and gaseous phosphide,
PH_{3}; 5PH_{2} = P_{2}H + 3PH_{3}; the latter corresponds with ammonia.
The mixture of the gaseous and liquid phosphides takes fire spontaneously
in the air, forming phosphoric acid. The same hydrogen phosphides are
formed when water acts on sodium phosphide (P_{2}Na_{3}). A similar
mixture of gaseous liquid and solid phosphuretted hydrogen (Retgers 1894)
is formed by heating (in a glass tube) red phosphorus in a stream of dry
hydrogen. Hence we see that there are _three compounds of phosphorus with
hydrogen_. (1) The first or solid yellow phosphide, P_{2}H (more probably
P_{4}H_{2}), is obtained by the action of strong hydrochloric acid on
sodium phosphide; it takes fire when struck or at 175°. (2) The liquid,
PH_{2}, or more correctly expressed as the molecule, P_{2}H_{4}, is a
colourless liquid which takes fire spontaneously in the air, boils at
30°, is very unstable, and is easily decomposed (by light or hydrochloric
acid) into the two other phosphides of hydrogen. It is prepared by
passing the gases evolved by the action of water on calcium phosphide
through a freezing mixture.[8] And, lastly, (3), gaseous hydrogen
phosphide, _phosphine_, PH_{3}, which is distinguished as being the most
stable. It is a colourless gas, which does not take fire in the air. It
has an odour of garlic, and is very poisonous. It resembles ammonia in
many of its properties.[8 bis] It is easily decomposed by heat, like
ammonia, forming phosphorus and hydrogen; but it is very slightly soluble
in water, and does not saturate acids, although it forms compounds with
some of them which resemble ammonium salts in their form and properties.
Among them the _compound with hydriodic acid_, PH_{4}I, analogous to
ammonium iodide, is remarkable. This compound crystallises on sublimation
in well-formed cubes, like sal-ammoniac, which it resembles in many
respects. However, this compound does not enter into those reactions of
double decomposition which are proper to sal-ammoniac, because its saline
properties are very feebly developed. Phosphuretted hydrogen also
combines, like ammonia, with certain chloranhydrides; but they are
decomposed by water, with the evolution of phosphine. Ogier (1880) showed
that hydrochloric acid also combines with phosphine under a pressure of
20 atmospheres at +18°, and under the ordinary pressure at -35°, forming
the crystalline phosphonium chloride PH_{4}Cl, corresponding to
sal-ammoniac. Hydrobromic acid does the same with greater ease, and
hydriodic acid with still greater facility, forming phosphonium iodide,
PH_{4}I.[9]
[6 bis] Retgers (see further on) showed this in 1894, and observed that
As when heated also combines with hydrogen.
[6 tri] The capacity of mercury (Chapter XVI., Note 25 bis) to give
unstable compounds with nitrogen gives rise to the supposition that
similar compounds exist with phosphorus also. Such a compound was
obtained by Granger (1892) by heating mercury with iodide of
phosphorus in a closed tube at 275°-300°. After removing the iodide
of mercury formed, there remain fine rhombic crystals having a
metallic lustre, and composition Hg_{3}P_{2}. This compound is
stable, does not alter at the ordinary temperature and only
decomposes at a red heat; when heated in air it burns with a flame.
Nitric and hydrochloric acids do not act upon it, but it is easily
decomposed by aqua regia. A phosphide of copper, Cu_{2}P_{2}, was
obtained by Granger (1893) by heating a mixture of water, finely
divided copper and red phosphorus in a sealed tube to 130°. The
excess of copper was afterwards washed away by a solution of NH_{3}
in the presence of air.
[7] The metallic compounds of phosphorus possess a great chemical
interest, because they show a transition from metallic alloys (for
instance, of Sb, As) to the sulphides, halogen salts, and oxides,
and on the other hand to the nitrides. Although there are already
many fragmentary data on the subject, the interesting province of
the metallic phosphides cannot yet be regarded as in any way
generalised. The varied applications (phosphor-iron,
phosphor-bronze, &c.), which the phosphides have recently acquired
should give a strong incentive to the complete and detailed study
of this subject, which would, in my opinion, help to the
explanation of chemical relations beginning with alloys (solutions)
and ending with salts and the compounds of hydrogen (hydrides),
because the phosphor-metals, as is proved by direct experiment,
stand in the same relation to phosphuretted hydrogen as the
sulphides do towards sulphuretted hydrogen, or as the metallic
chlorides to hydrochloric acid.
[7 bis] Many other compounds of phosphorus are also capable of forming
phosphuretted hydrogen. Thus BP also gives PH_{3} (_see_ Chapter
XVII., Note 12). According to Lüpke (1890) phosphuretted hydrogen
is formed by phosphide of tin. The latter is prepared by treating
molten tin covered with a layer of carbonate of ammonium, with red
phosphorus; 200-300 c.c. of water are then poured into a flask, 3-5
grams of this phosphide of tin dropped in, and after driving out
the air by a stream of carbonic acid, hydrochloric acid (sp. gr.
1·104) is poured in. The disengagement of phosphuretted hydrogen
takes place on heating the flask in a water bath. The following is
another easy method for preparing PH_{3}. A mixture of 1 part of
zinc dust (fume) and 2 parts of red phosphorus are heated in an
atmosphere of hydrogen (the mixture burns in air). Combination
takes place accompanied by a flash, and a grey mass of Zn_{3}P_{2}
is formed which gives PH_{3} when treated with dilute H_{2}SO_{4}.
[8] The spontaneous inflammability of the hydride PH_{2} in air is very
remarkable, and it is particularly interesting that its analogues
in composition, P(C_{2}H_{5})_{2} (the formula must be doubled) and
Zn(C_{2}H_{5})_{2}, also take fire spontaneously in air.
[8 bis] The analogy between PH_{3} and NH_{3} is particularly clear in
the hydrocarbon derivatives. Just as NH_{2}R, NHR_{2}, and NR_{3},
where R is CH_{3}, and other hydrocarbon radicles, correspond to
NH_{3}, so there are actually similar compounds corresponding to
PH_{3}. These compounds form a branch of organic chemistry.
[9] The periodic law and direct experiment (the molecular weight) show
that PH_{3} is the normal compound of P and H and that it is more
simple than PH_{2} or P_{2}H_{4}, just as methane, CH_{4}, is more
simple than ethane, C_{2}H_{6}, whose empirical composition is
CH_{3}. The formation of liquid phosphuretted hydrogen may be
understood from the law of substitution. The univalent radicle of
PH_{3} is PH_{2}, and if it is combined with H in PH_{3} it
replaces H in liquid phosphuretted hydrogen, which thus gives
P_{2}H_{4}. This substance corresponds with free amidogen
(hydrazine), N_{2}H_{4} (Chapter VI.) Probably P_{2}H_{4} is able
to combine with HI, and perhaps also with 2HI, or other
molecules--that is, to give a substance corresponding to
phosphonium iodide.
_Phosphonium iodide_, PH_{4}I, may be prepared, according to
Baeyer, in large quantities in the following manner:--100 parts of
phosphorus are dissolved in dry carbon bisulphide in a tubulated
retort: when the mixture has cooled, 175 parts of iodide are added
little by little, and the carbon bisulphide is then distilled off,
this being done towards the end of the operation in a current of
dry carbonic anhydride at a moderate temperature. The neck of the
retort is then connected with a wide glass tube, and the tubulure
with a funnel furnished with a stopcock, and containing 50 parts of
water. This water is added drop by drop to the phosphorous iodide,
and a violent reaction takes place, with the evolution of hydriodic
acid and phosphonium iodide. The latter collects as crystals in the
glass tube and the retort itself. It is purified by further
distillations; more than 100 parts may be obtained. Baeyer
expresses the reaction by the equation P_{2}I + 2H_{2}O = PH_{4}I +
PO_{2}; and the compound PO_{2} may be represented as phosphorous
phosphoric anhydride: P_{2}O_{5} + P_{2}O_{3} = 4PO_{2}. As a
better proportion we may take 400 grams of phosphorus, 680 grams of
iodine, and 240 grams of water, and express the formation thus: 13P
+ 9I + 21H_{2}O = 3H_{4}P_{2}O_{7} + 7PH_{4}I + 2HI (Chapter XI.,
Note 77).
Phosphonium iodide and even phosphine act as reducing agents in
solutions of many metallic salts. Cavazzi showed that with a
solution of sulphurous anhydride phosphine gives sulphur and
phosphoric acid.
_Phosphuretted hydrogen, or phosphine_, PH_{3}, is generally prepared by
the action of caustic potash on phosphorus.[10] Small pieces of
phosphorus are dropped into a flask containing a strong solution of
caustic potash and heated. Potassium hypophosphite, H_{2}KPO_{2}, is then
obtained in solution; gaseous phosphuretted hydrogen is evolved:
P_{4} + 3KHO + 3H_{2}O = 3(KH_{2}PO_{2}) + PH_{3}.
Liquid phosphuretted hydrogen (and free hydrogen) is also formed,
together with the phosphine, so that the gaseous product, on escaping
from the water into the air, takes fire spontaneously, forming beautiful
white rings of phosphoric acid. In this experiment, as in that with
calcium phosphide, it is the liquid, P_{2}H_{4}, that takes fire; but the
phosphine set light to by it also burns, PH_{3} + O_{4} = PH_{3}O_{4}.
The same phosphuretted hydrogen, PH_{3}, may be obtained pure, and not
spontaneously combustible, by igniting the hydrates of phosphorous acid
(4PH_{3}O_{3} = PH_{3} + 3PH_{3}O_{4}) and hypophosphorous acid
(2PH_{3}O_{2} = PH_{3} + PH_{3}O_{4}); or, more simply, by the
decomposition of calcium phosphide by hydrochloric acid, because then all
the liquid phosphide, P_{2}H_{4}, is decomposed into non-volatile P_{2}H
and gaseous PH_{3}. Pure phosphine liquefies when cooled to -90°, boils
at -85°, and solidifies at -135° (Olszewski). When phosphorus burns in an
excess[10 bis] of _dry_ oxygen, then only _phosphoric anhydride_,
P_{2}O_{5} is formed. It is prepared by dropping pieces of phosphorus
through a wide tube, fixed into the upper neck of a large glass globe, on
to a cup suspended in the centre of the globe. These lumps are set alight
by touching them with a hot wire, and the phosphorus burns into
P_{2}O_{5}. The dry air necessary for its combustion is forced into the
globe through a lateral neck, and the white flakes of phosphoric
anhydride formed are carried by the current of air through a second
lateral neck into a series of Woulfe's bottles, where they settle as
friable white flakes. Phosphoric anhydride may also be formed by passing
dry air through a solution of phosphorus in carbon bisulphide. All the
materials for the preparation of this substance must be carefully dried,
because it _combines_ with great eagerness _with water_, at the same time
developing a large amount of heat and forming metaphosphoric acid,
HPO_{3}, from which the water cannot be separated by heat. Phosphoric
anhydride is a colourless snow-like substance, which attracts moisture
from the air with the utmost avidity. It fuses at a red heat, and then
_volatilises_. Its affinity for water is so great that it takes it up
from many substances. Thus it converts sulphuric acid into sulphuric
anhydride, and carbohydrates (wood, paper) are carbonised, and give up
the elements of water when brought into contact with it.
[10] The air must first be expelled from the flask by hydrogen, or some
other gas which will not support combustion, as otherwise an
explosion might take place owing to the spontaneous inflammability
of the phosphuretted hydrogen.
The combustion of phosphuretted hydrogen in oxygen also takes
place under water when the bubbles of both gases meet, and it is
very brilliant. The phosphuretted hydrogen obtained by the action
of phosphorus on caustic potash always contains free hydrogen, and
often even the greater part of the gas evolved consists of
hydrogen.
_Pure phosphuretted hydrogen_ (not containing hydrogen or liquid
or solid phosphides) is obtained by the action of a solution of
potash on phosphonium iodide: PH_{4}I + KHO = PH_{3} + KI + H_{2}O
(in just the same way as ammonia is liberated from ammonium
chloride). The reaction proceeds easily, and the purity of the gas
is seen from the fact that it is entirely absorbed by bleaching
powder and is not spontaneously inflammable. Its mixture with
oxygen explodes when the pressure is diminished (Chapter XVIII.,
Note 8). The vapours of bromine, nitric acid, &c., cause it to
again acquire the property of inflaming in the air; that is, they
partially decompose it, forming the liquid hydride, P_{2}H_{4}.
Oppenheim showed that when red phosphorus is heated at 200° with
hydrochloric acid in a closed tube it forms the compound
PCl_{3}(H_{3}PO_{3}), together with phosphine.
[10 bis] If there be a deficiency of oxygen, _phosphorous anhydride_
P_{2}O_{3} is formed. It was obtained by Thorpe and Tutton (1890)
and is easily volatilised, melts at 22°·5, boils without change
(in an atmosphere of N_{2} or CO_{2}) at 173°, and is therefore
easily separated from P_{2}O_{3}, which volatilises with
difficulty. The vapour density shows that the molecular weight is
double, _i.e._ P_{4}O_{6} (like As_{2}O_{3}). Although colourless,
phosphorous anhydride (its density in a state of fusion at 24° =
1·936) turns yellow and reddens in sun-light (possibly red
phosphorus separates out ?), and decomposes at 400° forming
hypophosphorous anhydride P_{2}O_{4} (Note 11) and phosphorus. It
passes into P_{2}O_{5} in air and oxygen, and when slightly heated
in oxygen becomes luminous, and ultimately takes fire. Cold water
slowly transforms P_{2}O_{3} into phosphoric acid, but hot water
gives an explosion and leads to the formation of PH_{3},
(P_{4}O_{6} + 6H_{2}O = PH_{3} + 3PH_{3}O_{4}). Alkalis act in the
same manner. It takes fire in chlorine and forms POCl_{3} and
PO_{2}Cl, and combines with sulphur at 160°, forming
P_{2}S_{2}O_{3} (the molecular formula is double this) a substance
which volatilises in vacuo and is decomposed by water into H_{2}S
and phosphoric acid, _i.e._ it may be regarded as P_{2}O_{5}, in
which O_{2} has been replaced by two atoms of sulphur. Judging
from the above, the mixture of P_{2}O_{3} and P_{2}O_{5} formed in
the combustion of phosphorus in air is transformed into P_{2}O_{5}
in an excess of oxygen.
When moist phosphorus slowly oxidises in the air, it not only forms
phosphorous and phosphoric acids, but also _hypophosphoric acid_,
H_{4}P_{2}O_{6}, which when in a dry state easily splits up at 60° into
phosphorous and metaphosphoric acids (H_{4}P_{2}O_{6} = H_{3}PO_{3} +
HPO_{3}), but differs from a mixture of these acids in that it forms
well-characterised salts, of which the sodium salt,
H_{2}Na_{2}P_{2}O_{6}, is but slightly soluble in water (the sodium salts
of phosphoric and phosphorous acids are easily soluble), and that it does
not act as a reducing agent, like mixtures containing phosphorous
acid.[11]
[11] Salzer proved the existence of hypophosphoric acid (it is also
called subphosphoric acid), in which many chemists did not
believe. Drawe (1888) and Rammelsberg (1892) investigated its
salts. It may be obtained in a free state by the following method.
The solution of acid produced by the slow oxidation of moist
phosphorus is mixed with a solution (25 p.c.) of sodium acetate. A
salt, Na_{2}H_{2}P_{2}O_{6},6H_{2}O, crystallises out on cooling;
it is soluble in 45 parts of water, and gives a precipitate of
Pb_{2}P_{2}O_{6} with lead salts (Ag_{4}P_{2}O_{6} with salts of
silver). The lead salt is decomposed by a current of hydrogen
sulphide, when lead sulphide is precipitated, while the solution,
evaporated under the receiver of an air-pump, gives crystals of
H_{4}P_{2}O_{6},2H_{2}O, which easily lose water and give
H_{4}P_{2}O_{6}. The salts in which the H_{4} is replaced by
Ni_{2}, or NiNa_{2}, or CdNa_{2}, &c., are insoluble in water.
In order to see the relation between phosphoric acid and
hypophosphoric acid which does not contain the elements of
phosphorous acid (because it does not reduce either gold or
mercury from their solutions), but which nevertheless is capable
of being oxidised (for example, by potassium permanganate) into
phosphoric acid, it is simplest to apply the law of substitution.
This clearly indicates the relation between oxalic acid,
(COOH)_{2}, and carbonic acid, OH(COOH). The relation between the
above acids is exactly the same if we express phosphoric acid as
OH(POO_{2}H_{2}), because in this case P_{2}H_{4}O_{6}, or
(POO_{2}H_{2})_{3}, will correspond with it just as oxalic does
with carbonic acid. A similar relationship exists between
hyposulphuric or dithionic acid, (SO_{2}OH)_{2}, and sulphuric
acid, OH(SO_{2}OH), as we shall find in the following chapter.
Dithionic acid corresponds with the anhydride S_{2}O_{5},
intermediate between SO_{2} and SO_{3}; oxalic acid with
C_{2}O_{3}, intermediate between CO and CO_{2}; hypophosphoric
acid corresponds with the anhydride P_{2}O_{4}, intermediate
between P_{2}O_{3} and P_{2}O_{5}, and the analogue of N_{2}O_{4}.
Judging by the general law of the formation of acids (Chapter XV.), the
series of phosphorus compounds should include the following _ortho-acids_
and their corresponding anhydrides, answering to phosphuretted hydrogen,
H_{3}P:--
H_{3}PO_{4}, phosphoric acid, and P_{2}O_{5}, anhydride,
H_{3}PO_{3}, phosphorous acid, and P_{2}O_{3}, anhydride,
H_{3}PO_{2}, hypophosphorous acid, and P_{2}O, anhydride.[12]
The last of these (the analogue of N_{2}O) is almost unknown. Phosphoric
anhydride (P_{2}O_{5}) with a small quantity of water does not at first
give orthophosphoric acid, PH_{3}O_{4}, but a compound P_{2}O_{5},H_{2}O,
or PHO_{3}, whose composition corresponds with that of nitric acid; this
is _metaphosphoric acid_. Even with an excess of water, combining with
phosphoric anhydride, this metaphosphoric acid, and not the ortho-,
passes at first into solution. Metaphosphoric acid in solution only
passes into orthophosphoric acid when the solution is heated or after a
lapse of time.
[12] Besides the hydrates enumerated, a compound, PH_{3}O, should
correspond with PH_{3}. This hydrate, which is analogous to
hydroxylamine, is not known in a free state, but it is known as
triethylphosphine oxide, P(C_{2}H_{5})_{3}O, which is obtained by
the oxidation of triethylphosphine, P(C_{2}H_{5})_{3}. It must be
observed that there may also be lower oxides of phosphorus
corresponding with PH_{3}, like N_{2}O and NO, and there are even
indications of the formation of such compounds, but the data
concerning them cannot be considered as firmly established.
_Orthophosphoric acid_[13] is obtained by oxidising phosphorus with
nitric acid until the phosphorus entirely passes into solution and the
lower oxides of nitrogen cease to be evolved. The reaction takes place
best with dilute nitric acid, and when aided by heat. The resultant
solution is evaporated to a syrup. If a weighed quantity of phosphorus
(dried in a current of dry carbonic anhydride) be taken, a crystalline
mass of the acid can be obtained by evaporating the solution until it
consists only of the quantity[14] of phosphoric acid corresponding with
the amount of phosphorus taken (from 31 parts of P, 98 parts of
solution). The acid fuses at +39°; specific gravity of the liquid 1·88.
Phosphorus pentachloride, PCl_{5}, and oxychloride, POCl_{3} (see further
on), give orthophosphoric acid and hydrochloric acid with water. The two
other varieties of phosphoric acid, with which we shall presently become
acquainted, give the same ortho-acid when under the influence of acids,
with particular ease when boiled and more slowly in the cold. By itself
orthophosphoric acid (either in solution or when dry) does not pass into
the other varieties; it does not oxidise, and therefore presents the
limiting and stable form. When heated to 300°, it loses water and passes
into pyrophosphoric acid, 2H_{3}PO_{4} = H_{2}O + H_{4}P_{2}O_{7}, whilst
at a red heat it loses twice as much water and is converted into
metaphosphoric acid, H_{3}PO_{4} = H_{2}O + HPO_{3}. In aqueous solution
orthophosphoric acid differs clearly from pyro- or metaphosphoric acids,
because the solutions of these latter acids give different reactions:
thus orthophosphoric acid does not precipitate albumin, does not give a
precipitate with barium chloride, and forms a yellow precipitate of
silver orthophosphate, Ag_{3}PO_{4}, with silver nitrate (in the presence
of alkalis, but not otherwise); whilst a solution of pyrophosphoric acid,
H_{4}P_{2}O_{7}, although it does not precipitate albumin or barium
chloride, gives a white precipitate of silver pyrophosphate,
Ag_{4}P_{2}O_{7}, with silver nitrate; and a solution of metaphosphoric
acid, HPO_{3}, precipitates both albumin and barium chloride, and gives a
white precipitate of silver metaphosphate, AgPO_{3}, with silver nitrate.
These points of distinction were studied by Graham, and are exceedingly
instructive. They show that the solution of a substance does not
determine the maximum of chemical combination with water, that solutions
may contain various degrees of combination with water, and that there is
a clear difference between the water serving for solution and that
entering into chemical combination. Graham's experiments also showed that
the water whose removal or combination determines the conversion of
ortho- into meta- and pyrophosphoric acids differs distinctly from water
of crystallisation, for he obtained the salts of ortho-, meta-, and
pyrophosphoric acids with water of crystallisation, and they differed in
their reactions, like the acids themselves. This water of crystallisation
was expelled with greater ease than the water of constitution of the
hydrates in question.[14 bis]
[13] Phosphoric acid, being a soluble and almost non-volatile
substance, cannot be prepared like hydrochloric and nitric acids
by the action of sulphuric acid on the alkali phosphates, although
it is partially liberated in the process. For this purpose the
salts of barium or lead may be taken, because they give insoluble
salts, thus Ba_{3}(PO_{4})_{2} + 3H_{2}SO_{4} = 3BaSO_{4} +
2H_{3}PO_{4}. Bone ash contains, besides calcium phosphate, sodium
and magnesium phosphates, and fluorides and other salts, so that
it cannot give directly a pure phosphoric acid.
[14] If this is not done the orthophosphoric acid, PH_{3}O_{4}, loses a
portion of its water, and then, as with an excess of water, it
does not crystallise.
[14 bis] The difference between the reactions of ortho-, meta- and
pyrophosphoric acids, established by Graham (_see_ p. 163), is of
such importance for the theory of hydrates and for explaining the
nature of solutions, that in my opinion its influence upon
chemical thought has been far from exhausted. At the present time
many such instances are known both in organic (for instance, the
difference between the reactions of the solutions of certain
anhydrides and hydrates of acids), and inorganic chemistry (for
example, the difference between the rose and purple cobalt
compounds, Chapter XXII. &c.) They essentially recall the long
known and generalised difference between C_{2}H_{4} (ethylene),
C_{2}H_{6}O (ethyl alcohol = ethylene + water), and C_{4}H_{10}O
(ethyl ether = 2 ethylene + water = 2 alcohol - water); but to the
present day the numerous analogous phenomena existing among
inorganic substances are only considered as a simple difference in
degrees of affinity, distinguishing the water of constitution
(hydration), crystallisation, and solution without penetrating
into the difference of the structure or distribution of the
elements, which exists here and gives rise to a distinct isomerism
of solutions. In my opinion the progress of chemistry, especially
with regard to solutions, should make rapid strides when the cause
of the isomerism of solutions, for instance, of ortho- and
pyrophosphoric acids, has become as clear to us as the cause of
many well-studied instances of the isomerism, polymerism, and
metamerism of organic compounds. Here it forms one of those many
important problems which remain for the chemistry of the future in
a state of only indistinct presentiments and in the form of facts
empirically known but insufficiently comprehended.
Orthophosphoric acid has a pleasant acid taste and a distinctly acid
reaction; it is used as a medicine, and is not poisonous (phosphorous
acid is poisonous). Alkalis, like sodium, potassium, and ammonium
hydroxides, saturate the acid properties of phosphoric acid when taken in
the ratio 2NaHO : H_{3}PO_{4}--that is, when salts of the composition
HNa_{2}PO_{4} are formed. When taken in the ratio NaHO : H_{3}PO_{4}, a
solution having an acid reaction is obtained, and when 3NaHO :
H_{3}PO_{4}--that is, when the salt Na_{3}PO_{4} is formed--an alkaline
reaction is obtained. Hence many chemists (Berzelius) even regarded the
salts of composition R_{2}HPO_{4} as normal, and considered phosphoric
acid to be bibasic. But the salt Na_{2}HPO_{4} also shows a feeble
alkaline reaction, so that it is impossible to judge the characteristic
peculiarities of acids by the reactions on litmus paper, as we already
know from many examples. Orthophosphoric acid is tribasic, because it
contains three equivalents of hydrogen replaceable by metals, forming
salts, such as NaH_{2}PO_{4}, Na_{2}HPO_{4}, and Na_{3}PO_{4}. It is also
tribasic, because with silver nitrate its soluble salts always give
Ag_{3}PO_{4},[15] a salt with three equivalents of silver, and because by
double decomposition with barium chloride it forms a salt of the
composition Ba_{3}(PO_{4})_{2}, and silver and barium hardly ever give
basic salts. With the metals of the alkalis, phosphoric acid forms
soluble salts, but the normal salts of the metals of the alkaline earths,
R_{3}(PO_{4})_{2} and even R_{2}H_{2}(PO_{4}), are insoluble in water,
but dissolve in feeble acids, such as phosphoric and acetic, because they
then form soluble acid salts, especially RH_{4}(PO_{4})_{2}.[16]
[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_