Phosphorus – facts and properties
- Symbol
- P
- Atomic number
- 15
- Atomic mass
- 30.974 u
- Period
- 3
- Group
- 15
- Category
- polyatomic nonmetal
- Electron configuration
- [Ne] 3s² 3p³
- Electronegativity
- 2.19
- Atomic radius
- 98 pm
- Valence electrons
- 5
- State at standard conditions
- solid
- Melting point
- 44.2 °C
- Boiling point
- 280.5 °C
- Density
- 1.82 g/cm³
- Discovered
- 1669
Phosphorus is a non-metal with atomic number 15. It exists in several forms that differ so much they could pass for separate substances: white phosphorus is toxic and catches fire by itself in air, while red phosphorus is far tamer and sits on the striking strip of a matchbox. Life cannot do without it. As phosphate, phosphorus links the sugars in the backbone of DNA, hardens bones and teeth, and is one of the main nutrients in fertilisers.
Position in the periodic table
Occurrence and extraction
Elemental phosphorus is far too reactive to exist in nature, so it is found only in compounds, almost always as phosphate. The most important mineral is apatite, a calcium phosphate that forms large deposits of phosphate rock. By far the biggest reserves lie in Morocco and Western Sahara. Bones and teeth are themselves built largely from a calcium phosphate closely related to apatite.
Phosphorus was discovered in 1669 by Hennig Brand, a merchant and alchemist in Hamburg. Hoping to find the philosopher's stone, he boiled down large quantities of urine and heated the residue, obtaining a waxy substance that glowed in the dark. Its name comes from the Greek for light-bearer. Today white phosphorus is made in electric furnaces, where phosphate rock is heated with quartz sand and coke; phosphorus escapes as a vapour and is condensed under water. Most phosphate rock, however, never becomes the element at all. It is treated with sulfuric acid to make phosphoric acid and fertilisers. Because phosphate deposits are finite, recovering phosphorus from sewage sludge is becoming increasingly important.
Properties and reactions
Phosphorus has the electron configuration [Ne] 3s² 3p³ and five valence electrons. It belongs to period 3 and group 15, directly below nitrogen, and has an electronegativity of 2.19. Its most common oxidation states are minus three, plus three and plus five; in phosphates it is plus five.
The element has several allotropes. White phosphorus consists of tetrahedral molecules of four atoms. It is soft and waxy, with a density of 1.82 g/cm³, and it melts at 44.2 °C and boils at 280.5 °C. In air it reacts slowly with oxygen and glows faintly in the dark, and even slight warming makes it burst into flame. White phosphorus is highly toxic and causes deep, slow-healing burns, so it is stored under water. Heating it in the absence of air converts it into red phosphorus, which consists of linked chains, is stable in air, does not ignite spontaneously and is far less toxic. Black phosphorus forms under high pressure. It is the most stable form, has a layered structure like graphite and behaves as a semiconductor.
Phosphorus burns with a bright flame to give phosphorus pentoxide, a white powder that greedily absorbs water and reacts with it to form phosphoric acid. With metals phosphorus forms phosphides, and with chlorine it forms phosphorus chlorides.
Uses in everyday life and technology
Most mined phosphate ends up as fertiliser. Plants need phosphorus for growth, flowering and root development, and many fields must be supplied with it to keep yields stable. When too much phosphate washes into lakes and rivers, algae grow out of control, which is why phosphates have largely been removed from laundry detergents.
In living things, phosphate is everywhere. It links the sugar units in the backbone of DNA and RNA, carries energy in ATP and forms part of the phospholipids in cell membranes. Bones and teeth owe their hardness to calcium phosphate. We take in phosphate with almost every food, especially dairy products, meat and cereals.
In everyday life you meet red phosphorus on the striking surface of a safety match box. Only when the match is struck does it react with the oxidising agent in the match head and set it alight. Phosphoric acid gives cola drinks their tang and is used in rust converters. Phosphorus compounds act as flame retardants in plastics and textiles, as plasticisers and in crop protection products. White phosphorus is also used in incendiary munitions, whose use is internationally controversial because of the severe injuries they cause.
Molecules containing this element
- Phosphine PH3
- Aluminium phosphide AlP
- Phosphorus trifluoride PF3
Quick question
What is the atomic number of phosphorus (P)?
FAQ
Is phosphorus poisonous?
That depends on the form. White phosphorus is highly toxic and ignites spontaneously in air; even small amounts can be fatal. Red phosphorus is much less toxic and stable in air. Phosphates in food and in the body, on the other hand, are essential for life.
Why does white phosphorus glow in the dark?
White phosphorus reacts slowly with oxygen in the air even at room temperature. Part of the energy released escapes as a faint greenish light. This glow is chemiluminescence, not phosphorescence, even though the word phosphorescence was named after phosphorus.
Where is the phosphorus in a safety match?
In safety matches the red phosphorus is not in the match head but in the striking surface on the box. The head contains an oxidising agent. Only when the two meet as the match is struck does the match ignite.