# White phosphorus

White phosphorus, yellow phosphorus, or simply tetraphosphorus (P$_4$) is an allotrope of phosphorus. It is a translucent waxy solid that quickly yellows in light (due to its photochemical conversion into red phosphorus), and impure white phosphorus is for this reason called yellow phosphorus. White phosphorus is the first allotrope of phosphorus, and in fact the first elementary substance to be discovered that was not known since ancient times. It glows greenish in the dark (when exposed to oxygen) and is highly flammable and pyrophoric (self-igniting) upon contact with air. It is toxic, causing severe liver damage on ingestion and phossy jaw from chronic ingestion or inhalation. The odour of combustion of this form has a characteristic garlic odor, and samples are commonly coated with white "diphosphorus pentoxide", which consists of P$_4$O$_{10}$ 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. P$_4$ is soluble in benzene, oils, carbon disulfide, and disulfur dichloride.

## Structure

White phosphorus exists as molecules of four phosphorus atoms in a tetrahedral structure, joined by six phosphorus—phosphorus single bonds. The tetrahedral arrangement results in ring strain and instability. Although both are called "white phosphorus", in fact two different crystal allotropes are known, interchanging reversibly at 195.2 K. The element's standard state is the body-centered cubic α form, which is actually metastable under standard conditions. The β form is believed to have a hexagonal crystal structure.
Molten and gaseous white phosphorus also retains the tetrahedral molecules, until 800 °C (1,500 °F; 1,100 K) when it starts decomposing to P $_2$ molecules. The P $_4$ molecule in the gas phase has a P-P bond length of r$_g$ = 2.1994(3) Å as was determined by gas electron diffraction. The β form of white phosphorus contains three slightly different P $_4$ molecules, i.e. 18 different P-P bond lengths — between 2.1768(5) and 2.1920(5) Å. The average P-P bond length is 2.183(5) Å.

## Chemical properties

Despite white phosphorus not being the most stable allotropes of phosphorus, its molecular nature allows it to be easily purified. Thus, it's defined to have a zero enthalpy of formation.
In base, 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 50 °C (122 °F), and at much lower temperatures if finely divided (due to melting-point depression). Phosphorus reacts with oxygen, usually forming two oxides depending on the amount of available oxygen: P$_4$O$_6$ (phosphorus trioxide) when reacted with a limited supply of oxygen, and P$_4$O$_{10}$ when reacted with excess oxygen. On rare occasions, P$_4$O$_7$, P$_4$O$_8$, and P$_4$O$_9$ are also formed, but in small amounts. This combustion gives phosphorus(V) oxide:
    P$_4$ + 5 O$_2$ → P$_4$O$_{10}$

## 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 (although phosphate rock contains substantial amounts of fluoroapatite, which would also form silicon tetrafluoride):
    2 Ca$_3$(PO$_4$)$_2$ + 6 SiO$_2$ + 10 C → 6 CaSiO$_3$ + 10 CO + P$_4$
In this way, an estimated 750,000 tons were produced in 1988.
Most (83% in 1988) 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:
    P$_4$ + 10Cl$_2$ → 4PCl$_5$
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 isolated polyhedra.
A cubane-type cluster, in particular, is unlikely to form, and the closest approach is the half-phosphorus compound P$_4$(CH)$_4$, 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 rather acutely toxic, with a lethal dose of 50-100 mg (1 mg/kg body weight). 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.