Phosphorus is a nonmetal that is a solid in its natural state. It can vary in colors from red, silvery white, and black, has a melting point of 111.6 °F, and has a boiling point of 531 °F. The different forms of allotropes that phosphorus forms cause the variety in colors and reactivity characteristics. Some common uses of phosphorus include: as a component in safety matches, smoke bombs, tracer bullets, and fireworks, as an ingredient in fertilizers and pesticides, and in the production of chinaware, steel, and bronze.
Phosphorus is extremely reactive, and so it is not found free in nature. Isolation procedures to reduce phosphorus are carried out with great care.
Phosphorus sources that are used for large scale commercial isolation purposes come from minerals which contain phosphate rock. Most of these phosphate rocks are composed of a tri-calcium phosphate mineral called apatite; this mineral does not contain pure phosphate, but it provides a ready supply for isolation. The main chemical process that is used to isolate phosphorus from phosphorus minerals uses high amounts of energy; the high amounts of energy make this process inefficient for small scale laboratory isolations.
In the phosphorus isolation process, phosphorus minerals are heated with carbon and sand in an electric furnace. The calcium and phosphorus from the minerals, silicon from the sand, and carbon are reacted in the furnace at a temperature of 2732 °F. The resulting chemicals from the reaction are carbon dioxide, phosphorus, and calcium silicon oxygen compound. Pure phosphorus, depending on the allotrope that is isolated, can be toxic and highly explosive.
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