 White phosphorus sample with a chunk removed from the corner to expose un-oxidized material | |
 Tetraphosphorus molecule | |
| Names | |
|---|---|
| IUPAC names White phosphorus tetrahedro-Tetraphosphorus | |
| Systematic IUPAC name 1,2,3,4-Tetraphosphatricyclo[1.1.0.02,4]butane | |
Other names
| |
| Identifiers | |
3D model (JSmol) | |
| ChEBI | |
| ChemSpider | |
| ECHA InfoCard | 100.107.967 |
| 1856 | |
| UN number | 1381 |
| |
| |
| Properties | |
| P4 | |
| Molar mass | 123.895 g·mol−1 |
| Density | 1.82 g/cm3 |
| Melting point | 44.1 °C; 111.4 °F; 317.3 K |
| Boiling point | 280 °C; 536 °F; 553 K |
| Hazards[1] | |
| GHS labelling: | |
    | |
| Danger | |
| H250,H300+H330,H314,H400 | |
| P210,P222,P260,P264,P270,P271,P273,P280,P284,P301+P310+P330,P301+P330+P331,P303+P361+P353,P304+P340+P310,P305+P351+P338+P310,P335+P334,P363,P370+P378,P391,P403+P233,P405,P422,P501 | |
| NFPA 704 (fire diamond) | |
Threshold limit value (TLV) | 0.1 mg/m3 |
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |
White phosphorus,yellow phosphorus, or simplytetraphosphorus (P4) is anallotrope of phosphorus. It is a translucentwaxy solid that quickly yellows in light (due to itsphotochemical conversion intored phosphorus),[2] and impure white phosphorus is for this reason called yellow phosphorus. White phosphorus is the first allotrope ofphosphorus that was discovered, isolated for the first time in 1669 byHenning Brand.[3]
When in an oxygen-containing atmosphere, it will exhibit a faint green glow in the absence of light. White phosphorus is also highlyflammable andpyrophoric (self-igniting) upon contact with air. It istoxic, causing severeliver damage upon ingestion andphossy 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 ofP4O10 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.P4 is soluble inbenzene,oils,carbon disulfide, anddisulfur dichloride.
White phosphorus exists asmolecules of four phosphorusatoms in a tetrahedral structure, with each phosphorus atom making three phosphorus—phosphorussingle bonds for a total of six P-P bonds per tetrahedron. Thetetrahedral arrangement results inring strain and instability.[4] White phosphorus can take on one of two crystal allotropes that interechange reversibly above 195.2 K (−78.0 °C; −108.3 °F). The element'sstandard state is thebody-centered cubic α form, which ismetastable understandard conditions.[4] The β form is believed to have ahexagonal crystal structure.[5]
Molten and gaseous white phosphorus are also composed of these tetrahedra until 800 °C (1,500 °F; 1,100 K) when they start decomposing intoP
2 molecules.[6] TheP
4 molecule in the gas phase has a P-P bond length ofrg = 2.1994(3) Å as was determined bygas electron diffraction.[7] The β form of white phosphorus contains three slightly differentP
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.[6]
Despite white phosphorus not being the most stable allotrope of phosphorus (see:black phosphorus), it is still used as the reference state for solid phosphorus and defined to have astandard enthalpy of formation of zero. This is because it is much easier to handle and purify for the purposes of collecting reference thermodynamic data.
Inbasic media, white phosphorus spontaneouslydisproportionates tophosphine and various phosphorusoxyacid salts.[8]
Many reactions of white phosphorus involve insertion into the P-P bonds, such as the reaction with oxygen, sulfur,phosphorus tribromide and theNO+ ion.
It ignites spontaneously in air at about 50 °C (122 °F), and at much lower temperatures if finely divided (due tomelting-point depression). Phosphorus reacts with oxygen, usually formingtwo oxides depending on the amount of available oxygen:P4O6 (phosphorus trioxide) when reacted with a limited supply of oxygen, andP4O10 when reacted with excess oxygen. On rare occasions,P4O7,P4O8, andP4O9 are also formed, but in small amounts. This combustion gives phosphorus(V) oxide:
The white allotrope can be produced using several methods. In the industrial process,phosphate rock is heated in an electric or fuel-firedfurnace in the presence ofcarbon andsilica.[9] Elemental phosphorus is then liberated as a vapour and can be collected underphosphoric acid. An idealized equation for thiscarbothermal reaction is shown forcalcium phosphate (although phosphate rock contains substantial amounts offluoroapatite, which would also formsilicon tetrafluoride):
In this way, an estimated 750,000 tons were produced in 1988.[10]
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.[needs update] Much of the remaining 17% is mainly used for the production of chlorinated compoundsphosphorus trichloride,phosphorus oxychloride, andphosphorus pentachloride:[11]
Other products derived from white phosphorus includephosphorus pentasulfide and various metal phosphides.[10]
Although white phosphorus forms thetetrahedron, the simplest possiblePlatonic solid, no other polyhedral phosphorus clusters are known.[12] White phosphorus converts to the thermodynamically-stabler red allotrope, but that allotrope is not composed of isolated polyhedra.
Acubane-type cluster, in particular, is unlikely to form,[12] and the closest approach is the half-phosphorus compoundP4(CH)4, produced fromphosphaalkynes.[13] Other clusters are more thermodynamically favorable, and some have been partially formed as components of larger polyelemental compounds.[12]
White phosphorus is 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 andphossy jaw which appears to be caused by the formation ofamino bisphosphonates.[10][14][15]
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 forbeachcombers who may collect washed-up samples while unaware of their true nature.[16][17]
