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Rh other oxides, when the element is burnt in a limited supply of air or in pure oxygen under reduced pressure (E. Jungfleisch, Abs. Jour. Chem. Soc., 1907, ii. 761), and also when a solution of phosphorus in the trichloride or tribromide is exposed to light. It is a yellow or red powder which becomes dark red on heating; it is stable in air and can be heated to 300° without decomposition. Its existence, however, has been denied by A. Stock (Abs. Jour. Chem. Soc., 1910, ii. 121). The oxide P2O was obtained by Besson (Comptes rendus, 1897, 124, p. 763; 1901, pp. 132, 1556) by heating a mixture of phosphonium bromide and phosphorus oxychloride in sealed tubes to 50°.

Phosphorus oxide, P$2$O$6$, discovered by Sage in 1777, is a product of the limited combustion of phosphorus in air. It may be conveniently prepared by passing a rapid current of air over burning phosphorus contained in a combustion tube, and condensing the product in a metal condenser, from which it may be removed by heating the condenser to 50°–60° (Thorpe and Tutton, Jour. Chem. Soc., 1890, pp. 545, 632; 1891, p. 1019). Jungfleisch has obtained it by carrying out the combustion with oxygen under reduced pressure, or diluted with an inert gas. It forms crystals, apparently monoclinic, which melt at 22.5° to a clear, colourless, mobile liquid of boiling-point 173.1°. Its specific gravity is 2.135 at 21°. Vapour density and cryoscopic determinations point to the double formula, P$4$O$6$. It is comparatively stable up to 200°, but when heated in a sealed tube to 430° it gives phosphorus and the tetroxide P$2$O$5$. It is unaffected by light when pure, but if phosphorus be present, even in minute quantity, it turns yellow and ultimately dark red. It oxidizes on exposure to air to the pentoxide, and with a brilliant inflammation when thrown into oxygen at 50°–60°. It slowly reacts with cold water to form phosphorous acid; but with hot water it is energetically decomposed, giving much red phosphorus or the suboxide being formed with an explosive evolution of spontaneously inflammable phosphoretted hydrogen; phosphoric acid is also formed. With dilute alkalis phosphates are slowly formed, but with concentrated solutions the decomposition follows the same course as with hot water. With chlorine it gives phosphoryl and “ metaphosphoryl ” chlorides, the action being accompanied with a greenish flame; bromine gives phosphorus pentabromide and pentoxide which interact to give phosphoryl and “ metaphosphoryl ” bromides; iodine gives phosphorus di-iodide, P$2$I$4$, and pentoxide, P$2$O$5$; whilst hydrochloric acid gives phosphorus trichloride and phosphorous acid, which interact to form free phosphorus, phosphoric acid and hydrochloric acid. It combines violently with sulphur at 160° to form phosphorus sulphoxide, P$4$O$6$S$4$, which forms highly lustrous tetragonal plates (after sublimation), melting at 102° and boiling at 295°, it is decomposed by water into sulphuretted hydrogen and meta phosphoric acid, the latter changing on standing into or tho phosphoric acid. Sulphur trioxide and sulphuric acid oxidize phosphorus oxide, giving the pentoxide and sulphur dioxide, whilst sulphur chloride, S$2$Cl$2$, gives phosphoryl and thiophosphoryl chlorides, free sulphur and sulphur dioxide. Ammonia also reacts immediately. giving phosphorus diamide, P(OH)(NH$2$)$2$, and the corresponding ammonium salt. Phosphorous oxide is very poisonous, and is responsible for the caries set up in the jaws of those employed in the phosphorus industries (see below). It is probable, however, that pure phosphorous oxide vapour is odourless, and the odour of phosphorus as ordinarily perceived is that of a mixture of the oxide with ozone.

Phosphorus tetroxide, P$2$O$4$, was obtained by Thorpe and Tutton by heating the product of the limited combustion of phosphorus in vacuo as a sublimate of transparent, highly lustrous, orthorhombic crystals. They are highly deliquescent, and form with water a mixture of phosphorous and phosphoric acids: P$2$O$4$+3H$2$O=H$3$PO$3$+H$2$PO$4$. The vapour density at about 1400° is 230, i.e. slightly less than that required by P$8$O$16$ (West, Jour. Chem. Soc., 1902, p. 923).

Phosphorus oxide, or phosphorus pentoxide, P$4$O$10$, formed when phosphorus is burned in an excess of air or oxygen, or from dry phosphorus and oxygen at atmospheric pressure (Jungfleisch, loc. cit.), was examined by Boyle and named “ flowers of phosphorus ” by Marggraf in 1740. It is a soft, flocculent powder, which on sublimation forms transparent, monoclinic crystals. It is extremely deliquescent, hissing when thrown into water, with which it combines to form phosphoric acid. It is reduced when heated with carbon to phosphorus, carbon monoxide being formed simultaneously. Its vapour density at 1400° points to the double formula (West, Jour. Chem. Soc., 1896, p. 154).

Oxyacids.—Phosphorus forms several oxyacids: hypophosphorous acid, H$3$PO$2$, and hypophosphoric acid, H$4$P$2$O$6$ or H$2$PO$3$, of which the anhydrides are unknown; phosphorous acid, H$3$PO$3$, derived from P$4$O$6$; monoperphosphoric acid, H$3$PO$5$; perphosphoric acid, H$2$P$4$O$8$; and meta-, pyro-, and ortho-phosphoric acidls, derived from P$4$O$10$, for which see.

Hypophosphorous acid, HP(OH)$2$, discovered by Dulong in 1816, and obtained crystalline by Thomson in 1874 (Ber., 7, p. 994), is prepared in the form of its barium salt by warming phosphorus with baryta water, removing the excess of baryta by carbon dioxide, and crystallizing the nitrate. The acid may be prepared by evaporating in a vacuum the solution obtained by decomposing the barium salt with the equivalent amount of sulphuric acid. The acid forms a white crystalline mass, melting at 17.4° and having a strong acid reaction. Exposure to air gives phosphorous and phosphoric acids, and on heating it gives phosphine and phosphoric acid. A characteristic reaction is the formation of a red precipitate of cuprous hydride, Cu$2$H$2$, when heated with copper sulphate solution to 60°. It is a monobasic acid forming salts which are permanent in air, but which are gradually oxidized in aqueous solution. On heating they yield phosphine and leave a residue of pyrophosphate, or a mixture of meta- and pyrophosphates, with a little phosphorus. They react as reducing agents. On boiling with caustic potash they evolve hydrogen, yielding a phosphate.

Phosphorous acid, P(OH)$3$, discovered by Davy in 1812, may be obtained by dissolving its anhydride, P$4$O$6$, in cold water; by immersing sticks of phosphorus in a solution of copper sulphate contained in a well-closed flask, filtering from the copper sulphide and precipitating the sulphuric acid simultaneously formed by baryta water, and concentrating the solution in vacuo; or by passing chlorine into melted phosphorus covered with water, the first formed phosphorus trichloride being decomposed by the water into phosphorous and hydrochloric acids. It may also be reparcd by leading a current of dry air into phosphorus trichloride at 60° and passing the vapours into water at 0°, the crystals thus formed being drained, washed with ice-cold water and dried in a vacuum. The crystals melt at 70°. The acid is very deliquescent, and oxidizes on exposure to air to phosphoric acid. It decomposes on heating into phosphine and phosphoric acid. It is an energetic reducing agent; for example, when boiled with copper sulphate metallic copper is precipitated and hydrogen evolved. Although nominally tribasic the commonest metallic salts are dibasic. Organic ethers, however, are known in which one, two and three of the hydrogen atoms are substituted (Michaelis and Becker, Ber., 1897, 30, p. 1003). The metallic phosphates are stable both dry and in solution; when strongly heated they evolve hydrogen and yield a pyrophosphate, or, especially with the heavy metals, they give hydrogen and a mixture of phosphide and pyrophosphate.

Hypophosphoric acid, H$4$P$2$O$6$; or H$2$PO$3$, discovered by Salzer in 1877 among the oxidation products of phosphorus by moist air, may be prepared by oxidizing phosphorus in an aqueous solution of copper nitrate, or by oxidizing sticks of phosphorus under water, neutralizing with sodium carbonate, forming the lead salt and decomposing this with sulphuretted hydrogen (J. Cavalier and E. Cornee, Abs. Jour. Chem. Soc., 1910, ii. 31). The aqueous solution may be boiled without decomposition, but on concentration it yields phosphorous and phosphoric acids. Deliquescent, rectangular tablets of H$4$P$2$O$6$·2H$2$O separate out on concentrating a solution in a vacuum, which on drying further give the acid, which melts at 55°, and decomposes suddenly when heated to 70° into phosphorous and metaphosphoric acids with a certain amount of hydrogen phosphide. The solution is stable to oxidizing agents such as dilute hydrogen peroxide and chlorine, but is oxidized by potassium permanganate to phosphoric acid; it does not reduce salts of the eavy metals. With silver nitrate it gives a white precipitate, Ag$4$P$2$O$6$. The sodium salt, Na$4$P$2$O$6$·10H$2$O, forms monoclinic prisms and in solution is strongly alkaline; the acid salt, Na$3$HP$2$O$6$·9H$2$O, forms monoclinic tablets. The formula of the acid is not quite definite. Cryoscopic measurements on the sodium salt points to the double formula, but the organic esters appear to be derived from H$2$PO$3$, (see A. Rosenheim and M. Pritze, Ber., 1908, 41, 2708; E. Cornee, Abs. Jour. Chem. Soc., 1910, ii. 121).

Monoperphosphoric and perphosphoric acids, H$2$PO$5$ and H$4$P$2$O$8$, were obtained by J. Schmidlin and P. Massini (Ber., 1910, 43, 1162). The first is formed when 30% hydrogen peroxide reacts with phosphorus pentoxide or meta- or pyrophosphoric acids at low temperatures and the mixture diluted with ice-cold water. The solution is strongly oxidizing, even converting manganous salts to permanganates in the cold, a property not possessed by monopersulphuric acid. Perphosphoric acid is formed when pyrophosphoric acid is treated with a large excess of hydrogen peroxide.

Halogen Compounds.—Phosphorus trifluoride, PF$3$, discovered by Davy, may be obtained mixed with the pentafluoride; by direct combination of its elements; from the tribromide and arsenic trifluoride (MacIvor); from the tribromide and zinc fluoride, and from dried copper phosphide and lead fluoride (H. Moissan). It is a colourless, non-fuming gas, which gives a colourless, mobile liquid at –10° and 20 atmospheres; the liquid boils at –95° and solidifies at 160° (Moissan, Comptes rendus, 1904, 138, p. 789). It does not burn in air, but explodes, under the action of a flame or the electric spark, when mixed with half its volume of oxygen, giving the oxyfluoride, POF$3$. It is slowly decomposed by water giving hydrofluoric and phosphorous acids, or, in addition, fluorphosphorous acid, HPF$4$. It has no action on glass in the cold, but when heated it gives phosphorus and silicon tetrafluoride. Phosphorus pentafluoride, PF$5$, discovered by Thorpe (Proc. Roy. Soc., 1877, 25, p. 122), may be obtained by burning the trifluoride in fluorine, from the pentachloride and arsenic trifluoride and from the trifluoride and bromine, the first formed fluorobromide, PF$3$Br$2$, decomposing into the pentabromide and pentafluoride: 5PF$3$Br$2$=3PF$5$+2PBr$5$. It is a colourless gas 4½ times heavier than air, and liquefies at 15° under 40 atmospheres, solidifying when the pressure is diminished. It is incombustible and extinguishes flame. It fumes in moist air and is quickly decomposed by water giving hydrofluoric and phosphoric