PHOSPHORUS, a non-metallic element, first known as Phosphorus mirabilis or igneus. (Symbol P, atomic number 55.
atomic weight 30.98.) It is never found free, but is widely and abundantly distributed in combination as phosphates (q.v.). It is essential to animal and vegetable life, occurring in urine, blood and tissues, and, as calcium phosphate, forming 58% of bones. The element was first obtained in 1669 by Brand of Hamburg, who kept his process secret. In England, however, Kunckel (1678) and Boyle (168o) attempted its preparation and happened to use Brand's method—urine was evaporated to dryness and the residue distilled with sand. In 1775 Scheele prepared it from bones, and these still form one of the chief commercial sources of phos phorus. Degreased bones or mineral phosphates are treated with sufficient sulphuric acid to combine with all the calcium, this sulphate is filtered off, and the filtrate is mixed with charcoal, coke or sawdust, dried in a furnace, and then distilled from Stourbridge-clay retorts at a white heat. The vapours are led through malleable iron pipes into troughs containing water, where they condense. The reactions are: Modern Processes.—In more modern electrothermal processes, calcium phosphate is mixed with sand and carbon, and fed into a furnace heated by an alternating current. At the high temperature the silica of the sand displaces phosphoric oxide and forms calcium silicate, and the oxide is reduced by the carbon so that the phos phorus vapour and carbon monoxide pass over into a condenser. The silicate remains as a liquid slag and is run off periodically. Other processes seek to produce calcium carbide simultaneously with phosphorus by using only calcium phosphate and carbon, thus obtaining a more valuable by-product. In all cases the phosphorus is cast into sticks either under water or in water-cooled pipes.
chloride, benzene, oil of turpentine, and liquefied ammonia or sulphur dioxide. It is a non-conductor of electricity, has density 1.836 at o° C, melts at and boils at 287° giving a colour less vapour, the density of which corresponds to tetratomic mole cules, P4; above 1,500° C the vapour density indicates dissociation to P2. The elevation of boiling point of solutions of phosphorus in carbon disulphide, and the depression of freezing point of solu tions in benzene, both indicate P4.
The element is highly inflammable, taking fire in air at burning with a bright white flame and forming dense white clouds of phosphoric oxide ; H. B. Baker has shown, however, that in perfectly dry air or oxygen it can be distilled unchanged. (See DRYNESS, CHEMICAL.) When exposed to air which has not been excessively dried, a stick of phosphorus undergoes slow com bustion, as shown by the phosphorescence visible in the dark. In pure oxygen, however, this phosphorescence is not exhibited unless the gas pressure is reduced or the temperature raised; similarly, in compressed air the phosphorescence ceases when a certain pressure is reached. Moreover, the presence of certain substances (e.g., Cl, Br, I, NH3, N20, NO2, H2S, SO2, CH4, C2H4) entirely inhibits the phosphorescence. During the process, oxygen is slightly ionized and traces of ozone are formed. Some of the phenomena accompanying phosphorescence have been ascribed to the oxidation of the trioxide. A very complete bibliography of investigations is given by W. E. Downey (J. Chem. Soc., 125, p. 349), and H. J. Emeleus (ibid., 1926, p. 1336; P. 788; 1928, p. 628) gives a further discussion of the inhibition and of the ultra-violet spectrum of the glow. Phosphorus corn bines directly with the halogens, sulphur and selenium, and most metals burn in its vapour forming phosphides. When very finely divided it decomposes water, giving hydrogen phosphide, and when boiled with water it gives hypophosphorous acid in addition; slow oxidation of wet phosphorus yields hypophosphoric acid.