- An essential nutrient for plants and animals that is commonly applied to crops as a phosphate fertilizer. Phosphorus can contribute to the eutrophication of lakes and other water bodies. Sources of excess phosphorus include sewage and agricultural runoff.
- CRC Handbook of Chemistry and Physics and the American Chemical Society
Atomic Number: 15
Atomic Symbol: P
Atomic Weight: 30.97376
Electron Configuration: [Ne]3s23p3
(Gr. phosphoros, light bearing; ancient name for the planet Venus when appearing before sunrise) Brand discovered phosphorus in 1669 by preparing it from urine.
Phosphorus exists in four or more allotropic forms: white (or yellow), red, and black (or violet). Ordinary phosphorus is a waxy white solid; when pure it is colorless and transparent. White phosphorus has two modifications: alpha and beta with a transition temperature at -3.8° C.
It is insoluble in water, but soluble in carbon disulfide. It takes fire spontaneously in air, burning to the pentoxide.
Never found free in nature, it is widely distributed in combination with minerals. Phosphate rock, which contains the mineral apatite, an impure tri-calcium phosphate, is an important source of the element. Large deposits are found in Russia, in Morocco, and in Florida, Tennessee, Utah, Idaho, and elsewhere.
It is very poisonous, 50 mg constituting an approximate fatal dose. Exposure to white phosphorus should not exceed 0.1 mg/m3 (8-hour time-weighted average - 40-hour work week). White phosphorus should be kept under water, as it is dangerously reactive in air, and it should be handled with forceps, as contact with the skin may cause severe burns.
When exposed to sunlight or when heated in its own vapor to 250° C, it is converted to the red variety, which does not phosphoresce in air as does the white variety. This form does not ignite spontaneously and is not as dangerous as white phosphorus. It should, however, be handled with care as it does convert to the white form at some temperatures and it emits highly toxic fumes of the oxides of phosphorus when heated. The red modification is fairly stable, sublimes with a vapor pressure of 1 atm at 17° C, and is used in the manufacture of safety matches, pyrotechnics, pesticides, incendiary shells, smoke bombs, tracer bullets, etc.
White phosphorus may be made by several methods. By one process, tri-calcium phosphate, the essential ingredient of phosphate rock, is heated in the presence of carbon and silica in an electric furnace or fuel-fired furnace. Elementary phosphorus is liberated as vapor and may be collected under phosphoric acid, an important compound in making super-phosphate fertilizers.
In recent years, concentrated phosphoric acids, which may contain as much as 70% to 75% P2O5 content, have become of great importance to agriculture and farm production. World-wide demand for fertilizers has caused record phosphate production. Phosphates are used in the production of special glasses, such as those used for sodium lamps.
Bone-ash, calcium phosphate, is used to create fine chinaware and to produce mono-calcium phosphate, used in baking powder.
Phosphorus is also important in the production of steels, phosphor bronze, and many other products. Trisodium phosphate is important as a cleaning agent, as a water softener, and for preventing boiler scale and corrosion of pipes and boiler tubes.
Phosphorus is also an essential ingredient of all cell protoplasm, nervous tissue, and bones.
Sources: CRC Handbook of Chemistry and Physics and the American Chemical Society.