phosphorusPnonmetallic chemical element of the nitrogen family (Group 15 [Va] of the periodic table) .

A brief treatment of phosphorus follows. For full treatment, see nitrogen group element: Phosphorus.

Properties, occurrence, and uses.Ordinarily that at room temperature is a colourless, semitransparent, soft, waxy solid that glows in the dark, it takes fire spontaneously upon exposure to air and forms dense white fumes of the oxide. Phosphorus was first prepared in elemental form in 1669 by a German alchemist, Hennig Brand, from a residue of evaporated urine.



Arabian alchemists of the 12th century may have isolated elemental phosphorus by accident, but the records are unclear. Phosphorus appears to have been discovered in 1669 by Henning Brand, a German merchant whose hobby was alchemy. Brand allowed 50 buckets of urine to stand until they putrified and “bred worms.” He then boiled the urine down to a paste and heated it with sand, thereby distilling elemental phosphorus from the mixture. Brand reported his discovery in a letter to Gottfried Wilhelm Leibniz, and, thereafter, demonstrations of this element and its ability to glow in the dark, or “phosphoresce,” excited public interest. Phosphorus, however, remained a chemical curiosity until about a century later when it proved to be a component of bones. Digestion of bones with nitric or sulfuric acid formed phosphoric acid, from which phosphorus could be distilled by heating with charcoal. In the late 1800s, James Burgess Readman of Edinburgh developed an electric furnace method for producing the element from phosphate rock, which is essentially the method employed today.

Occurrence and distribution

Phosphorus is a very widely distributed element—12th most abundant in the Earth’s crust, to which it contributes about 0.10 weight percent. Its cosmic abundance is estimated to be about one atom per 100 atoms of silicon, the standard. Its high chemical reactivity assures that it does not occur in the free state (except in a few meteorites). Phosphorus always occurs as the phosphate ion. The principal combined forms in nature are the phosphate salts. Nearly 190 different minerals have been found to contain phosphorus, but, of these, the principal source of phosphorus is the apatite series in which calcium ions exist along with phosphate ions and variable amounts of fluoride, chloride, or hydroxide ions, according to the formula [Ca10(PO4)6(F, Cl, or OH)2]. Other important phosphorus-bearing minerals are wavellite and vivianite. Commonly such metal atoms as magnesium, manganese, strontium, and lead substitute for calcium in the mineral; and silicate, sulfate, vanadate, and similar anions substitute for phosphate ions. Very large sedimentary deposits of fluoroapatite are found in many parts of the Earth. The phosphate of bone and tooth enamel is hydroxyapatite. (The principle of lessening tooth decay by fluoridation depends upon the conversion of hydroxyapatite to the harder, more decay-resistant, fluoroapatite.)



The element has about 10 forms (allotropes) that Estimates of the total phosphate rock in the Earth’s crust average about 50,000,000,000 tons, of which North Africa contains two-thirds, and Russia and the United States most of the remaining third. This estimate includes only ore that is sufficiently rich in phosphate for conversion to useful products by present methods. Vast quantities of material lower in phosphorus content also exist.

The only naturally occurring isotope of phosphorus is that of mass 31. The other isotopes from mass 24 to mass 46 have been synthesized by appropriate nuclear reactions. All of these are radioactive with relatively short half-lives. The isotope of mass 32 has a half-life of 14.3 days and has proven extremely useful in tracer studies involving the absorption and movement of phosphorus in living organisms.

Commercial production and uses

Two principal techniques for converting phosphate rock to usable materials are practiced. One involves acidulation of the crushed rock—with either sulfuric or phosphoric acids—to form crude calcium hydrogen phosphates that, being water-soluble, are valuable additions to fertilizer. The other method is the reduction of the phosphate with carbon in an electric furnace to give elemental phosphorus. The latter reaction is extremely complex, and its precise details depend upon the composition of the mineral phosphate. A charge of sand, coke, and phosphate rock is melted at about 1,500° C in an electric furnace. The calcium and impurities are left in the form of carbon monoxide gas and a complex fluorosilicate slag, and elemental phosphorus vapour, at about 300° C, distills out and is collected, condensed, and stored underwater (to prevent spontaneous ignition) as the white allotropic form of the element. More than half a million tons of phosphorus are made annually in the United States in this way. Most of the output is burned to phosphoric anhydride and subsequently treated with water to form phosphoric acid, H3PO4.

Only about 5 percent of the 2,000,000 tons of phosphorus consumed per year in the United States is used in the elemental form. Pyrotechnic applications of the element include tracers, incendiaries, fireworks, and matches. Some is used as an alloying agent, some to kill rodents, and the rest is employed in chemical synthesis. A large amount is converted to sulfides used in matches and in the manufacture of insecticides and oil additives. Most of the remainder is converted to halides or oxides for subsequent use in synthesizing organic phosphorus compounds.

Properties and reactions

The electronic configuration of the phosphorus atom can be represented by 1s22s22p63s23p3. The outer shell arrangement therefore resembles that of nitrogen, with three half-filled orbitals each capable of forming a single covalent bond and an additional lone-pair of electrons. Depending on the electronegativity of the elements with which it combines, phosphorus can therefore exhibit oxidation states of +3 or −3, just as does nitrogen. The principal differences between nitrogen and phosphorus are that the latter is of considerably lower electronegativity and has larger atoms, with outer d orbitals available. For these reasons, the similarities between nitrogen and phosphorus chemistry are largely formal ones, tending to conceal the actual, wide differences. The outer d orbitals in phosphorus permit an expansion of the octet, which leads to the +5 state, with five actual covalent bonds being formed in compounds, a condition impossible for nitrogen to achieve.

The first striking difference in chemistry of the two elements is that elemental phosphorus exists under ordinary conditions in any of 10 modifications, or allotropic forms, all of which are solid and which occur within three major categories: white, red, and black. Phosphorus molecules of formula P2, structurally analogous to N2 molecules and evidently also triply bonded, exist only at very high temperatures. These P2 molecules do not persist at lower temperatures—below about 800°—because of the fact that three single bonds in phosphorus, in contrast to the situation with nitrogen, are energetically favoured over one triple bond. On cooling, the triply bonded P2 molecules condense to form tetrahedral P4 molecules, in which each atom is joined to three others by single bonds. White phosphorus has two allotropes: the alpha form, which is stable at ordinary temperatures, has a cubic crystal structure; the beta form, which is stable below −78° C (−108° F), has a hexagonal crystal structure. White phosphorus is poisonous. Exposure to sunlight or to heat converts it to red phosphorus, which neither phosphoresces nor spontaneously burns in air. Black phosphorus is flaky like graphite and is made by subjecting white phosphorus to high pressures. It is chemically the least reactive form; white is by far the most reactive. White phosphorus Because of the relatively weak intermolecular attractions (van der Waals forces) between the separate P4 molecules, the solid melts easily at 44.1° C and boils at about 280° C. Formation of tetrahedra requires bond angles of 60° instead of the preferred 90°–109° angles, so that white phosphorus is a relatively unstable, or metastable, form. It changes spontaneously, but slowly, at temperatures around 200° or higher, to a polymeric form called “red phosphorus.” This substance is amorphous when formed at lower temperatures, but it can become crystalline, with a melting point of about 590° C. At higher temperatures and pressures, or with the aid of a catalyst, at ordinary pressures and a temperature of about 200° C, phosphorus is converted to a flaky, black crystalline form, which somewhat resembles graphite. This may prove to be the most stable form of phosphorus, despite the relative difficulty in its preparation. In both the red and the black forms, each phosphorus atom forms three single bonds, which are spread apart sufficiently to be relatively strain free.

Consistent with the metastable condition of the white modification, and the crowding of its covalent bonds, this form is far more reactive chemically than the others. It is highly toxic, reacts vigorously with most reagents, and inflames in air at only 35° C, so it must be stored under water or other inert liquid. White phosphorus dissolves readily in solvents such as carbon disulfide, in which it maintains the composition P4. White phosphorus has been used for military purposes as a source of smoke and to fill incendiary shells and grenades. In contrast, red phosphorus is insoluble and relatively inert, although large quantities of the usual commercial form can ignite spontaneously in air and react with water to form phosphine and phosphorus oxyacids. Red phosphorus is used in preparing the striking surface for safety matches.


Principal compounds.

Black phosphorus is more inert and is capable of conducting electricity. Both these polymeric forms are insoluble and are very much less volatile than white phosphorus.

Principal compounds

Phosphorus is used almost entirely in the form of compounds, usually in the oxidation states of +3, +5, and −3. Unlike nitrogen and various other members of the family, phosphorus tends to exhibit a preference for the +5 state.


Among the most commercially important phosphorus compounds are the oxides and acids. Much of the industrially produced white phosphorus is burned to form phosphorus pentoxide, P4O10. Sometimes called phosphoric anhydride, or diphosphorus pentoxide, this compound can be obtained in the form of a soft white powder or colourless crystalline solid. It is widely used in chemical analysis as a dehydrating agent and in organic synthesis as a condensing agent. Large quantities are treated with water to make orthophosphoric acid (H3PO4), commonly called phosphoric acid (q.v.), which has diverse industrial applications, including the production of phosphates, salts that contain the phosphate ion (PO43−), the hydrogen phosphate ion (HPO42−), or the dihydrogen phosphate ion (H2PO4). Such salts are used as leavening agents in baking, as abrasives in toothpaste, and sometimes as additives to detergents. Another salt, prepared by the action of phosphoric acid on phosphate rock, is calcium dihydrogen phosphate, or superphosphate, Ca(H2PO4)2, the most widely used phosphate fertilizer.



Analytical chemistry

Elemental phosphorus can be detected by its phosphorescence. It can also be converted to phosphine with boiling sodium hydroxide solution or with zinc and sulfuric acid; the phosphine is identified by means of test paper containing either silver nitrate or mercuric chloride, both of which are reduced to the free metal by phosphine, thereby darkening the paper. Phosphorus vapour also readily darkens silver nitrate test paper. Phosphorus is determined quantitatively by oxidation to phosphate, followed by any of several standard procedures. Phosphate, for example, may be precipitated as the magnesium ammonium salt, MgNH4PO4, converted by ignition to magnesium pyrophosphate, Mg2P2O7, and weighed. Alternatively, phosphate may be precipitated as ammonium phosphomolybdate; this can be weighed as such, converted to magnesium pyrophosphate and weighed, or titrated directly with sodium hydroxide solution.

Biological and physiological significance

Phosphorus is an important constituent of bones and teeth, and it is essential to the growth of living organisms. In organisms the element usually appears as phosphate. In its other forms phosphorus is likely to prove very toxic. White phosphorus attacks the skin and, when ingested, causes a necrosis of the jawbone, called “phossy jaw.” Certain organic esters of phosphoric acid, used as lubricating-oil additives, have been found to cause permanent paralysis when accidentally ingested. Phosphine is extremely toxic, as are its organic derivatives. Some of the most toxic substances known to man, collectively termed nerve gas, are organic derivatives of phosphorus.


atomic number15atomic weight30.9738melting point (white)44.1° C (111.4° F)boiling point (white)280° C (536° F)density (white)1.82 g/cm3 at 20° C (68° F)oxidation states−3, +3, +5electron config.1s22s22p63s23p3