John Emsley looks at the element that's the devil in disguise
When bonded to four oxygen atoms, phosphorus is a key part of living things. But when the oxygen atoms are stripped away we uncover an element with terrible powers.
In the body
A normal diet provides around 5 g of phosphate per day and there is around 800 g of phosphorus in the human body. Most of it is in our bones as calcium phosphate, but more important are the organophosphates like DNA, adenosine triphosphate (ATP), guanosine metaphosphate (GMP, a messenger molecule), and the phospholipids that make up cell membranes.
The body produces and uses ATP on a truly remarkable scale - more than 1 kg per hour - and is the key to releasing energy from glucose.
Isolation and manufacture
Phosphorus can be obtained as white phosphorus (P4) or red phosphorus (Pn): the latter made by heating white phosphorus in a closed vessel. This converts a form that is flammable, deadly and glows in the dark, into one that has none of these features. There is even a graphite-like black phosphorus that is made under extreme pressure.
Phosphorus was first produced by Hennig Brandt in 1669 in Hamburg. He evaporated urine and heated the residues until they were red hot, and distilled off the phosphorus vapour. However, the most successful phosphorus producer was Ambrose Godfrey, laboratory assistant of the scientist Robert Boyle, the first person to investigate the chemistry of this element. Godfrey saw it as a means of becoming rich and for 50 years he supplied most of Europe with phosphorus from his premises in Covent Garden, London, selling it at £3 per ounce (£1000 in today's money).
Phosphorus became more widely available when it was discovered that bone was calcium phosphate, and it was manufactured by heating phosphoric acid (obtained by dissolving bone in sulfuric acid) with charcoal. Eventually phosphorus was produced from mineral phosphates like fluoroapatite, Ca5(PO4)3F, by heating it with coke in an electric furnace.
Despite its toxicity, phosphorus was a widely used pharmaceutical for 250 years, and even given to treat conditions such as tuberculosis and cholera. Yet it was known that breathing in phosphorus vapour led to a chronic condition known as phossy jaw, which slowly ate away the victim's jaw bone. It particularly afflicted those who made phosphorus matches in the 1800s, the so-called match girls, now seen as heroines of women's liberation.
In the middle of the last century, sodium tripolyphosphate (Na5P3O10) was added to detergents as a way of tying up calcium ions that cause hard water, but it caused environmental problems for aquatic systems and has been replaced by zeolites. There is now a legal requirement to remove phosphate from sewage in many parts of the world and new technology allows it to be recovered and recycled effectively.
White phosphorus was widely used in the wars of the last century in tracer bullets, incendiary bombs, and smoke grenades. It is still used as such.
Dishwasher powder and flares
Today, phosphoric acid (H3PO4) is the main feedstock from which other compounds are made. The acid is used to make phosphate fertilisers, animal feed supplements, dishwasher powders and food additives. Some is used industrially to give metals a protective layer and as a rust-remover and preventer.
A little elemental phosphorus is still produced to make compounds like magnesium phosphide (Mg3P2) for self-igniting flares. The flares also contain calcium carbide, and when the mixture gets wet the phosphide forms the spontaneously flammable gas diphosphine. This ignites the acetylene gas given off by the calcium carbide as it reacts with water.
Ultra-refined phosphorus is used to make metal phosphides, such as those of gallium and indium for light emitting diodes and for these the phosphorus has to be 99.9999% pure.
Atomic number 15; atomic weight 30.973761; melting point 44°C (white phosphorus); boiling point 280°C (white); density: 1.8 g cm-3 (white), 2.2 g cm-3 (red), 2.7 g cm-3 (black). Phosphorus is in group 15 of the periodic table. It has two principal oxidation states: +3, as in the oxide P4O6 and +5, as in P4O10.