Phosphorus Trioxide, P₂O₃

The formation of an oxide by the slow combustion of phosphorus had already been noticed in the eighteenth century, and the existence of such a compound was known to Lavoisier and Davy, while the conditions of formation and the composition were established by Dulong.

The method of preparation by the slow oxidation of white phosphorus has already been described. When a current of dry air, or of oxygen, the pressure of which lies between certain limits, is passed over solid white phosphorus, this oxide is formed at a slow rate. At higher temperatures, at or above the melting-point of white phosphorus, the rate of formation is sufficiently great, but the product may contain less than 10 per cent, of phosphorous oxide, with nearly 80 per cent, of phosphoric oxide, and variable quantities of phosphorus remain unburnt in the form of the white mixed with the red element. It is possible to separate the phosphorus pentoxide, as is done in the standard method due to Thorpe and Tutton, according to which a rapid stream of dry air is passed over liquid phosphorus in a hard glass tube, which is connected with a brass tube containing a glass wool filter and maintained at about 60° C. by an external current of water. This tube retains the phosphorus pentoxide, while the trioxide passes on and is condensed in a U-tube which is immersed in a freezing mixture. The phosphorous oxide is melted and filtered into another U-tube. The further purification is described below.

Phosphorous oxide has also been prepared by the action of the trichloride on phosphorous acid, thus

H₃PO₃ + PCl₃ = P₂O₃ + 3HCl

It is also produced when phosphorus trichloride is used to replace the hydroxyl group of anhydrous acetic acid, thus

3CH₃COOH + 2PCl₃ = P₂O₃ + 3HCl + 3CH₃COCl

The trioxide sublimes as a mist of solid particles which are difficult to condense completely. The condensate forms a snow-like mass of minute crystals, or large feathery crystals, or a wax-like mass which is very fusible, deliquescent and inflammable. The vapour smells of garlic and is poisonous. The material contains small quantities, up to about 1 per cent., of phosphorus, droplets of which may appear in the later stages of the distillation. The complete removal is difficult, but the amount may be greatly reduced by the following method. The trioxide is crystallised from dry carbon disulphide in an atmosphere of carbon dioxide at -18° C. The crystals are filtered off on a perforated porcelain disc, washed with petroleum ether and recrystallised from carbon disulphide. These operations are carried out in an atmosphere of dry carbon dioxide. The purified oxide is analysed by melting under water at 40° to 50° C. It is completely absorbed, whereas the usual preparation, which is saturated with phosphorus at 25° C., shows a small residue of undissolved phosphorus under the same conditions. The solubility of phosphorus in the purified oxide at 25° C. is 1.7 grams in 100 grams. The solution on solidifying assumes the opaque waxy appearance characteristic of the ordinary preparation.

The density (D_4°^21°) is 2.135. The melting-point is 22.5° C. The crystalline system is monoclinic. When large crystals are examined in polarised light they may show pinacoid, prism and complementary pairs of pyramid faces. When the oxide is freed from all except traces of the dissolved phosphorus as just described, it appears as a transparent crystalline solid and the physical properties are slightly modified; in particular, the melting-point was 23.8° C. instead of 22.4°.

The density of the slightly supercooled liquid at 21° C. (D_4°^21°) is 1.9431 and that of the liquid at its boiling-point, 173.1° C., 1.6897. The smoothed values of the specific volume as derived from measurements in a dilatometer from 27.10° to 140.30° C. are as follows, and give the coefficient of expansion of the liquid up to its boiling-point:—

Relative specific volumes of phosphorus trioxide

t° C.	25	30	50	75	100	125	150	170	173.1
v	1.000	1.0046	1.0228	1.0459	1.0695	1.0941	1.1200	1.1419	1.1454

The vapour pressures of the liquid are as follows:—

Vapour pressures of phosphorus trioxide

t° C	22.4	30.8	40.8	50.0	64.4	72.7	91.2
p (mm.)	2.7	4.1	6.0	9.5	18.4	50.8	297.9

A straight line is obtained by plotting log p against 1/T up to T = 336°, and if this is produced it leads to a boiling-point under normal pressure of T = 458° C. abs. (cf. 446° above). The corresponding graph above 336° is also a straight line which, however, leads to a boiling-point of T = 374°. It is considered that the measurements at higher temperatures are defective, probably owing to interaction of the trioxide and moisture, with production of phosphine.

The refractive index of the liquid has been determined for several of the standard wavelengths in the visible spectrum:—

λ, mμ	670.5	589.2	434.0
n	1.5350	1.5403	1.5616

These and other results are represented by the dispersion formula

n = 1.5171 + 817670λ^-2 - 31659070₇λ^-4

The dielectric constant at 22° C. is 3.2.

The molecular weight and structure have been deduced in the usual manner from physical constants. The vapour density, as determined by Hofmann's method, varied between 7.67 and 7.83 (air = l) at temperatures between 132° and 184° C., which corresponds to a molecular weight which is represented by the formula P₄O₆. This agrees with the molecular weight calculated from the lowering of the freezing-point of benzene.

The molar volume at the boiling-point, 130.2, minus the atomic volumes of six singly-linked oxygen atoms (—O—) leaves 83.4, which is very nearly equal to four times the atomic volume of elementary phosphorus at its boiling-point, i.e. 4×20.9. It is concluded that the molecule _P4 in P₄O₆ occupies the same volume as the molecule _P4 of the liquid element. Now the atomic volume of phosphorus in its trivalent combination, as in PCl₃, etc., is certainly greater than that of elementary phosphorus at the same vapour pressure. Whence it follows that elementary phosphorus, as well as phosphorus in P₄O₆, is exerting its highest valency, i.e., is tercovalent with a mixed bond. The respective structural formulae for the element and the trioxide would therefore be

and

A somewhat similar formula involving the transfer of electrons has been suggested by Henstock.

Phosphorus trioxide turns yellow and then red on exposure to sunlight, and slightly yellow in ordinary diffused light. After months of exposure red phosphorus is formed. The change may be represented by the equation

5P₂O₃ = 4P + 3P₂O₅

When the oxide was heated to about 200° C. in a sealed tube it became turbid, then yellow and finally red. These changes proceeded still further at higher temperatures, and at 445° C. the whole was converted into solid products according to the equation

2P₄O₆ = 3P₂O₄ + 2P

The oxide is unaffected by molecular hydrogen.

The liquid ignites in air or oxygen at about 50° C. and burns with a vivid flame to the pentoxide. Like phosphorus itself phosphorous oxide undergoes slow combustion with the emission of a glow, which may be due in part to the small quantities of the element already mentioned. This combustion, however, differs from that of phosphorus in several respects. Ozone is produced only in small quantities, if at all, and may be due to the action of light of wavelength λ = 120 to 180 mμ on the oxygen. Hydrogen peroxide is not produced by the oxidation of phosphorous oxide. Dry oxygen combines with the oxide with increasing speed from about 10° C., and over a certain range the speed varies as the square root of the oxygen pressure. Glowing oxidation in the presence of water vapour has been connected with the intermediate formation of phosphine. Ozonised air gives a strong glow, close to the surface of the oxide.

Products of oxidation vary according to the conditions. Dry oxygen at ordinary temperatures, and especially under reduced pressure, gave phosphorus pentoxide, while moist oxidation at atmospheric pressure gave tetroxide. Further experiments have shown that (a) the pure trioxide, (b) the ordinary trioxide, when oxidised either by air or oxygen or ozonised oxygen, gives a mixture of P₂O₄ and P₂O₅ in the constant proportions which correspond to an empirical formula PO_2.19. The equation suggested for this oxidation is

8P₄O₆ + 4O₃ + 5O₂ = 5P₄O₈ + 3P₄O₁₀

Phosphorous oxide ignites spontaneously in chlorine, burning with a greenish flame and forming a clear liquid which on distillation yielded phosphoryl trichloride and a residue of metaphosphoryl chloride:—

P₄O₆ + 4Cl₂ = 2POCl₃ + 2PO₂Cl₂

The oxide also inflames in bromine. Slow combination yields the pentabromide and pentoxide, thus

5P₄O₆ + 20Br₂ = 8PBr₅ + 6P₂O₅

Combination with iodine is less vigorous, but when the oxide and iodine in carbon disulphide are heated in a sealed tube, orange-red P₂I₄ is deposited.

The oxide does not mix with sulphur, but when heated together in an inert atmosphere to about 160° C. the two combine with the production of diphosphorus dithiotrioxide, P₄O₆S₄.

The liquid oxide slightly above the melting-point does not mix with water, but combines with it slowly to give phosphorous acid:—

P₄O₆ + 6H₂O = 4H₃PO₃

Hot water reacts explosively giving phosphine and red phosphorus.

Hydrogen chloride dissolves in phosphorous oxide and reacts with it in a manner which is the reverse of that by which the oxide can be formed, this reverse equation being

P₄O₆ + 6HCl = 2PCl₃ + 2H₃PO₃

Such a mixture may also give phosphoric acid and phosphorus, according to the equation

4H₃PO₃ + PCl₃ = 3H₃PO₄ + 2P + 3HCl

Ammonia reacts violently with the oxide, reducing it to red phosphorus. When, however, the ammonia is passed into a solution of the oxide in ether, it gives a diamide of phosphorous acid, HOP(NH₂)₂.

The oxide acts vigorously upon ethyl alcohol, and by the regulation of this reaction diethyl phosphite has been obtained:—

P₄O₆ + 8C₂H₅OH = 4HOP(OC₂H₅)₂ +2H₂O

The chemical properties of the purified oxide differ in some respects from those of the ordinary preparation. It has a pungent acid smell which does not recall the smell of phosphorus. When exposed to light, either in an evacuated bulb or in one filled with carbon dioxide, it does not turn red. It does not absorb oxygen at pressures of about 100 to 80 mm. even at 25° C. and in the presence of water vapour. When heated in a sealed tube containing dry oxygen at 300 mm. it shows no glow below 200° C., and the oxygen is only slowly absorbed at 220° C.