Chapter XIX: Phosphorus and the Other Elements of the Fifth Group (1)

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.

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]