CHAPTER XIX. Phosphorus and the Other Elements of the Fifth Group

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.

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]

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 (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 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.

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.).

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.

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. 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]

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.

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. 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; 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).

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.

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. 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. 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.

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.

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.

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. 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.

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.

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.

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.

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.

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 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 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]

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.

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.

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}, 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}.

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.

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 nNaHO + HX = 14 and NaHO + nHX = 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; nNaHO + 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 + nH{3}PO{4} = 14 (approximately) and nNaHO + 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. When ignited they lose the last molecule of water of constitution, and give sodium pyrophosphate, Na{4}P{2}O{7}.

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.

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.

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.

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.

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.

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, 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.

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. 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.

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, 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.

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°, 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.

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