10 Fun and Interesting Phosphorus Facts

Phosphorus is element 15 on the periodic table, with the element symbol P. Because it is so chemically reactive, phosphorus is never found free in nature, yet you encounter this element in compounds and in your body. Here are 10 interesting facts about phosphorus:

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Phosphorus Facts (Atomic Number 15 or Element Symbol P)

By Anne Marie Helmenstine, Ph.D.

Interesting Phosphorus Facts

1. Phosphorus was discovered in 1669 by Hennig Brand in Germany. Brand isolated phosphorus from urine. The discovery made Brand the first person to discover a new element. Other elements such as gold and iron were known before that, but no specific person found them.
2. Brand called the new element "cold fire" because it glowed in the dark. The name of the element comes from the Greek word phosphoros, which means "bringer of light." The form of phosphorus Brand discovered was white phosphorus, which reacts with oxygen in air to produce a green-white light. Although you might think the glow would be phosphorescence, phosphorus is chemiluminescent and not phosphorescent. Only the white allotrope or form of phosphorus glows in the dark.
3. Some texts refer to phosphorus as the "Devil's Element" because of its eerie glow, tendency to burst into flame, and because it was the 13th known element.
4. Like other nonmetals, pure phosphorus assumes markedly different forms. There are at least five phosphorus allotropes. In addition to white phosphorus, there is red, violet, and black phosphorus. Under ordinary conditions, red and white phosphorus are the most common forms.
5. While the properties of phosphorus depend on the allotrope, they share common nonmetallic characteristics. Phosphorus is a poor conductor of heat and electricity, except black phosphorus. All types of phosphorus are solid at room temperature. The white form (sometimes called yellow phosphorus) resembles wax, the red and violet forms are noncrystalline solids, while the black allotrope resembles graphite in pencil lead. The pure element is reactive, so much so that the white form will ignite spontaneously in air. Phosphorus typically has an oxidation state of +3 or +5.
6. Phosphorus is essential to living organisms. There are about 750 grams of phosphorus in the average adult. In the human body, it's found in DNA, bones, and as an ion used for muscle contraction and nerve conduction. Pure phosphorus, however, can be deadly. White phosphorus, in particular, is associated with negative health effects. Matches made using white phosphorus are associated with a disease known as phossy jaw which causes disfiguration and death. Contact with white phosphorus can cause chemical burns. Red phosphorus is a safer alternative and is considered non-toxic.
7. Natural phosphorus consists of one stable isotope, phosphorus-31. At least 23 isotopes of the element are known.
8. The primary use of phosphorus is for fertilizer production. The element is also used in flares, safety matches, light-emitting diodes, and steel production. Phosphates are used in some detergents. Red phosphorus is also one of the chemicals used in illegal production of methamphetamines.
9. According to a study published in the Proceedings of the National Academies of Sciences, phosphorus may have been brought to Earth by meteorites. The release of phosphorus compounds seen early in Earth's history (yet not today) contributed to the conditions needed for the origin of life. Phosphorus is abundant in the Earth's crust at a concentration of about 1,050 parts per million, by weight.
10. While it's certainly possible to isolate phosphorus from urine or bone, today the element is isolated from phosphate-bearing minerals. Phosphorus is obtained from calcium phosphate by heating the rock in a furnace to yield tetraphosphorus vapor. The vapor is condensed into phosphorus underwater to prevent ignition.

Sources

• Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd Ed.), Oxford:Butterworth-Heinemann.
• Hammond, C. R. (2000). The Elements, in Handbook of Chemistry and Physics (81st ed.). CRC press.​
• Meija, J.; et al. (2016). "Atomic weights of the elements 2013 (IUPAC Technical Report)". Pure and Applied Chemistry. 88 (3): 265–91.
• Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110.