White phosphorus

White phosphorus, yellow phosphorus, or simply tetraphosphorus is an allotrope of phosphorus. It is a translucent waxy solid that quickly yellows in light, and impure white phosphorus is for this reason called yellow phosphorus. White phosphorus is the first allotrope of phosphorus, and was discovered in 1669 by Henning Brand.
When in an oxygen-containing atmosphere, it will exhibit a faint green glow in the absence of light. White phosphorus is also highly flammable and pyrophoric upon contact with air. It is toxic, causing severe liver damage upon ingestion and phossy jaw from chronic ingestion or inhalation. The combustion of this form has a characteristic garlic odor, and samples are commonly coated with white "diphosphorus pentoxide", which consists of tetrahedra with oxygen inserted between the phosphorus atoms and at their vertices. White phosphorus is only slightly soluble in water and can be stored under water. is soluble in benzene, oils, carbon disulfide, and disulfur dichloride.

Structure

White phosphorus exists as molecules of four phosphorus atoms in a tetrahedral structure, with each phosphorus atom making three phosphorus—phosphorus single bonds for a total of six P-P bonds per tetrahedron. The tetrahedral arrangement results in ring strain and instability. White phosphorus can take on one of two crystal allotropes that interechange reversibly above. The element's standard state is the body-centered cubic α form, which is metastable under standard conditions. The β form is believed to have a hexagonal crystal structure.
Molten and gaseous white phosphorus are also composed of these tetrahedra until when they start decomposing into molecules. The molecule in the gas phase has a P-P bond length of r_g = 2.1994 Å as was determined by gas electron diffraction. The β form of white phosphorus contains three slightly different molecules, i.e. 18 different P-P bond lengths — between 2.1768 and 2.1920 Å. The average P-P bond length is 2.183 Å.

Chemical properties

Despite white phosphorus not being the most stable allotrope of phosphorus, it is still used as the reference state for solid phosphorus and defined to have a standard enthalpy of formation of zero. This is because it is much easier to handle and purify for the purposes of collecting reference thermodynamic data.
In basic media, white phosphorus spontaneously disproportionates to phosphine and various phosphorus oxyacid salts.
Many reactions of white phosphorus involve insertion into the P-P bonds, such as the reaction with oxygen, sulfur, phosphorus tribromide and the NO⁺ ion.
It ignites spontaneously in air at about, and at much lower temperatures if finely divided. Phosphorus reacts with oxygen, usually forming two oxides depending on the amount of available oxygen: when reacted with a limited supply of oxygen, and when reacted with excess oxygen. On rare occasions,,, and are also formed, but in small amounts. This combustion gives phosphorus oxide:

Production and applications

The white allotrope can be produced using several methods. In the industrial process, phosphate rock is heated in an electric or fuel-fired furnace in the presence of carbon and silica. Elemental phosphorus is then liberated as a vapour and can be collected under phosphoric acid. An idealized equation for this carbothermal reaction is shown for calcium phosphate :
In this way, an estimated 750,000 tons were produced in 1988.
Most white phosphorus is used as a precursor to phosphoric acid, half of which is used for food or medical products where purity is important. The other half is used for detergents. Much of the remaining 17% is mainly used for the production of chlorinated compounds phosphorus trichloride, phosphorus oxychloride, and phosphorus pentachloride:
Other products derived from white phosphorus include phosphorus pentasulfide and various metal phosphides.

Other polyhedrane analogues

Although white phosphorus forms the tetrahedron, the simplest possible Platonic solid, no other polyhedral phosphorus clusters are known. White phosphorus converts to the thermodynamically-stabler red allotrope, but that allotrope is not composed of isolated polyhedra.
A cubane-type cluster, in particular, is unlikely to form, and the closest approach is the half-phosphorus compound, produced from phosphaalkynes. Other clusters are more thermodynamically favorable, and some have been partially formed as components of larger polyelemental compounds.

Safety

White phosphorus is acutely toxic, with a lethal dose of 50-100 mg. Its mode of action is not known but is thought to involve its reducing properties, possibly forming intermediate reducing compounds such as hypophosphite, phosphite, and phosphine. It damages the liver, kidneys, and other organs before eventually being metabolized to non-toxic phosphate. Chronic low-level exposure leads to tooth loss and phossy jaw which appears to be caused by the formation of amino bisphosphonates.
White phosphorus is used as a weapon because it is pyrophoric. For the same reasons, it is dangerous to handle. Measures are taken to protect samples from air since it will react with oxygen at ambient temperatures, and even in small samples this can lead to self-heating and eventual combustion. There are anecdotal reports of problems for beachcombers who may collect washed-up samples while unaware of their true nature.