Allotropes of phosphorus
Elemental phosphorus can exist in several allotropes; the most common of which are white and red solids. Solid violet and black allotropes are also known. Gaseous phosphorus exists as diphosphorus and atomic phosphorus.
White phosphorus, or yellow phosphorus, or simply tetraphosphorus (P4) exists as molecules made up of four atoms in a tetrahedral structure. Gaseous and yellow arsenic and antimony and solid arsenic phosphide (AsP3) each have a similar tetrahedral structure. The tetrahedral arrangement results in ring strain and instability. The molecule is described as consisting of six single P–P bonds. Two different crystalline forms are known. The α form, which is stable under standard conditions, has a body-centered cubic crystal structure. It transforms reversibly into the β form at 195.2 K. The β form is believed to have a hexagonal crystal structure.
White phosphorus is a translucent waxy solid that quickly becomes yellow when exposed to light. For this reason it is also called yellow phosphorus. It glows greenish in the dark (when exposed to oxygen), is highly flammable and pyrophoric (self-igniting) upon contact with air as well as 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 smell, and samples are commonly coated with white "(di)phosphorus pentoxide", which consists of P4O10 tetrahedra with oxygen inserted between the phosphorus atoms and at their vertices. White phosphorus is only slightly soluble in water and it can be stored under water. Indeed, white phosphorus is only safe from self-igniting when it is submerged in water. It is soluble in benzene, oils, carbon disulfide, and disulfur dichloride.
Production and applications
The white allotrope can be produced using several different 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):
- 2 Ca3(PO4)2 + 8 C → P4 + 8 CO2 + 6 Ca
White phosphorus has an appreciable vapour pressure at ordinary temperatures. The vapour density indicates that the vapour is composed of P4 molecules up to about 800 °C. Above that temperature, dissociation into P2 molecules occurs.
It ignites spontaneously in air at about 50 °C, and at much lower temperatures if finely divided. This combustion gives phosphorus (V) oxide:
4 + 5 O
2 → P
Because of this property, white phosphorus is used as a weapon.
Non-existence of cubic-P8
Although white phosphorus converts to the thermodynamically more stable red allotrope, the formation of the cubic P8 molecule is not observed in the condensed phase. Derivatives of this hypothetical molecule have been prepared from phosphaalkynes.
Red phosphorus may be formed by heating white phosphorus to 250 °C (482 °F) or by exposing white phosphorus to sunlight. Red phosphorus exists as an amorphous network. Upon further heating, the amorphous red phosphorus crystallizes. Red phosphorus does not ignite in air at temperatures below 240 °C, whereas pieces of white phosphorus ignite at about 30 °C. Ignition is spontaneous at room temperature with finely divided material. Heating red phosphorus in the presence of moisture creates phosphine gas, which is both highly flammable and toxic.
Hittorf's violet phosphorus
Monoclinic phosphorus, or violet phosphorus, is also known as Hittorf's Metallic Phosphorus. In 1865, Johann Wilhelm Hittorf heated red phosphorus in a sealed tube at 530 °C. The upper part of the tube was kept at 444 °C. Brilliant opaque monoclinic, or rhombohedral, crystals sublime. Violet phosphorus can also be prepared by dissolving white phosphorus in molten lead in a sealed tube at 500 °C for 18 hours. Upon slow cooling, Hittorf's allotrope crystallises out. The crystals can be revealed by dissolving the lead in dilute nitric acid followed by boiling in concentrated hydrochloric acid. In addition, a fibrous form exists with similar phosphorus cages. Below is shown a chain of phosphorus atoms which exhibits both the purple and fibrous forms.
Reactions of violet phosphorus
It does not ignite in air until heated to 300 °C, and it is insoluble in all solvents. It is not attacked by alkali and only slowly reacts with halogens. It can be oxidised by nitric acid to phosphoric acid.
If it is heated in an atmosphere of inert gas, for example nitrogen or carbon dioxide, it sublimes and the vapour condenses as white phosphorus. If it is heated in a vacuum and the vapour condensed rapidly, violet phosphorus is obtained. It would appear that violet phosphorus is a polymer of high relative molecular mass, which on heating breaks down into P2 molecules. On cooling, these would normally dimerize to give P4 molecules (i.e. white phosphorus) but, in vacuo, they link up again to form the polymeric violet allotrope.
Black phosphorus is the thermodynamically stable form of phosphorus at room temperature and pressure. It is obtained by heating white phosphorus under high pressures (12,000 atmospheres). In appearance, properties and structure it is very like graphite, being black and flaky, a conductor of electricity, and having puckered sheets of linked atoms.
Black phosphorus has an orthorhombic structure and is the least reactive allotrope: a result of its lattice of interlinked six-membered rings. Each atom is bonded to three other atoms. A recent synthesis of black phosphorus using metal salts as catalysts has been reported.
The diphosphorus allotrope (P2) can be obtained normally only under extreme conditions (for example, from P4 at 1100 kelvin). Nevertheless, some advancements were obtained in generating the diatomic molecule in homogenous solution under normal conditions with the use of transition metal complexes (based on, for example, tungsten and niobium).
Diphosphorus is the gaseous form of phosphorus, and the thermodynamically stable form above 1200 °C and until 2000 °C. The dissociation of tetraphosphorus (P
4) begins at lower temperature: the percentage of P
2 at 800 °C is ≈ 1%. At temperatures above about 2000 °C, the diphosphorus molecule begins to dissociate into atomic phosphorus.
The red-brown phase differs from red phosphorus and is also stable in air for weeks. Electron microscope showed the red-brown form as having long, parallel nanorods with a diameter between 0.34 nm and 0.47 nm.
|Space group||I43m||P1 No.2||P2/c No.13||Cmca No.64|
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