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.
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.
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This section is about the chemistry of white phosphorus. For military applications, see white phosphorus munitions. |
 White phosphorus sample with a chunk removed from the corner to expose un-oxidized material
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 Tetraphosphorus molecule
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| Names | |
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| IUPAC names
White phosphorus
Tetraphosphorus | |
| Systematic IUPAC name
1,2,3,4-Tetraphosphatricyclo[1.1.0.02,4]butane | |
Other names
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| Identifiers | |
3D model (JSmol)
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| ChemSpider | |
PubChem CID
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| UN number | 1381 |
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| 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 | |
| NFPA 704 (fire diamond) | |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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White phosphorus, yellow phosphorus or simply tetraphosphorus (P4) exists as molecules of phosphorus made up of four atoms in a tetrahedral structure. The tetrahedral arrangement results in ring strain and instability. The molecule is described as consisting of six single phosphorus—phosphorus bonds (P–P bonds). Two crystalline forms are known. The α form is defined as the standard state of the element, but is actually metastable under standard conditions.[1] It has a body-centered cubic crystal structure, and transforms reversibly into the β form at 195.2 K. The β form is believed to have a hexagonal crystal structure.[2]
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) 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 smell, and samples are commonly coated with white "diphosphorus pentoxide", which consists of P4O10 tetrahedral 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. Indeed, white phosphorus is safe from self-igniting when it is submerged in water; due to this, unreacted white phosphorus can prove hazardous to beachcombers who may collect washed-up samples while unaware of their true nature.[3][4] P4 is soluble in benzene, oils, carbon disulfide, and disulfur dichloride.
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.[5] 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):
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 (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: P4O6 (phosphorus trioxide) when reacted with a limited supply of oxygen, and P4O10 when reacted with excess oxygen. On rare occasions, P4O7, P4O8, and P4O9 are also formed, but in small amounts. This combustion gives phosphorus(V) oxide:
Because of this property, white phosphorus is used as a weapon.
Phosphorus pentachloride is prepared by the reaction of white phosphorus with excess of dry chlorine.[6]
It can also be prepared by the action of sulfuryl chloride on white phosphorus.[6]
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. Analogs of this hypothetical molecule have been prepared from phosphaalkynes.[7] White phosphorus in the gaseous state and as waxy solid consists of reactive P4 molecules.
Red phosphorus may be formed by heating white phosphorus to 300 °C (570 °F) in the absence of air or by exposing white phosphorus to sunlight. Red phosphorus exists as an amorphous network. Upon further heating, the amorphous red phosphorus crystallizes. Bulk red phosphorus does not ignite in air at temperatures below 240 °C (460 °F), whereas pieces of white phosphorus ignite at about 30 °C (86 °F).
Under standard conditions it is more stable than white phosphorus, but less stable than the thermodynamically stable black phosphorus. The standard enthalpy of formation of red phosphorus is −17.6 kJ/mol.[1] Red phosphorus is kinetically most stable.
It was first presented by Anton von Schrötter before the Vienna Academy of Sciences on December 9, 1847, although others had doubtlessly had this substance in their hands before, such as Berzelius.[8]
Red phosphorus can be used as a very effective flame retardant, especially in thermoplastics (e.g. polyamide) and thermosets (e.g. epoxy resins or polyurethanes). The flame retarding effect is based on the formation of polyphosphoric acid. Together with the organic polymer material, these acids create a char that prevents the propagation of the flames. The safety risks associated with phosphine generation and friction sensitivity of red phosphorus can be effectively minimized by stabilization and micro-encapsulation. For easier handling, red phosphorus is often used in form of dispersions or masterbatches in various carrier systems. However, for electronic/electrical systems, red phosphorus flame retardant has been effectively banned by major OEMs due to its tendency to induce premature failures.[9] One persistent problem is that red phosphorus in epoxy molding compounds induces elevated leakage current in semiconductor devices.[10] Another problem was acceleration of hydrolysis reactions in PBT insulating material.[11]
Red phosphorus can also be used in the illicit production of methamphetamine and Krokodil.
Red phosphorus can be used as an elemental photocatalyst for hydrogen formation from the water.[12] They display a steady hydrogen evolution rates of 633 μmol/(h⋅g) by the formation of small-sized fibrous phosphorus.[13]
Monoclinic phosphorus, or violet phosphorus, is also known as Hittorf's metallic phosphorus.[14][15] 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 sublimed as a result. 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.[16] In addition, a fibrous form exists with similar phosphorus cages. The lattice structure of violet phosphorus was presented by Thurn and Krebs in 1969.[17] Imaginary frequencies, indicating the irrationalities or instabilities of the structure, were obtained for the reported violet structure from 1969.[18] The single crystal of violet phosphorus was also produced. The lattice structure of violet phosphorus has been obtained by single‐crystal x‐ray diffraction to be monoclinic with space group of P2/n (13) (a = 9.210, b = 9.128, c = 21.893 Å, β = 97.776°, CSD-1935087). The optical band gap of the violet phosphorus was measured by diffuse reflectance spectroscopy to be around 1.7 eV. The thermal decomposition temperature was 52 °C higher than its black phosphorus counterpart. The violet phosphorene was easily obtained from both mechanical and solution exfoliation.
Violet phosphorus does not ignite in air until heated to 300 °C and 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 a vacuum, they link up again to form the polymeric violet allotrope.
Black phosphorus is the thermodynamically stable form of phosphorus at room temperature and pressure, with a heat of formation of −39.3 kJ/mol (relative to white phosphorus which is defined as the standard state).[1] It was first synthesized by heating white phosphorus under high pressures (12,000 atmospheres) in 1914. As a 2D material, in appearance, properties, and structure, black phosphorus is very much like graphite with both being black and flaky, a conductor of electricity, and having puckered sheets of linked atoms.[19]
Black phosphorus has an orthorhombic pleated honeycomb structure and is the least reactive allotrope, a result of its lattice of interlinked six-membered rings where each atom is bonded to three other atoms.[20][21] In this structure, each phosphorus atom has five outer shell electrons.[22] Black and red phosphorus can also take a cubic crystal lattice structure.[23] The first high-pressure synthesis of black phosphorus crystals was made by the Nobel prize winner Percy Williams Bridgman in 1914.[24] Metal salts catalyze the synthesis of black phosphorus.[25]
Black phosphorus based sensors exhibit several superior qualities over traditional materials used in piezoelectric or resistive sensors. Characterized by its unique puckered honeycomb lattice structure, black phosphorus provides exceptional carrier mobility. This property ensures its high sensitivity and mechanical resilience, making it an intriguing candidate for sensor technology.[26][27]
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Main article: phosphorene |
The similarities to graphite also include the possibility of scotch-tape delamination (exfoliation), resulting in phosphorene, a graphene-like 2D material with excellent charge transport properties, thermal transport properties and optical properties. Distinguishing features of scientific interest include a thickness dependent band-gap, which is not found in graphene.[28] This, combined with a high on/off ratio of ~105 makes phosphorene a promising candidate for field-effect transistors (FETs).[29] The tunable bandgap also suggests promising applications in mid-infrared photodetectors and LEDs.[30][31] Exfoliated black phosphorus sublimes at 400 °C in vacuum.[32] It gradually oxidizes when exposed to water in the presence of oxygen, which is a concern when contemplating it as a material for the manufacture of transistors, for example.[33][34] Exfoliated black phosphorus is an emerging anode material in the battery community, showing high stability and lithium storage.[35]
Ring-shaped phosphorus was theoretically predicted in 2007.[36] The ring-shaped phosphorus was self-assembled inside evacuated multi-walled carbon nanotubes with inner diameters of 5–8 nm using a vapor encapsulation method. A ring with a diameter of 5.30 nm, consisting of 23 P8 and 23 P2 units with a total of 230 P atoms, was observed inside a multi-walled carbon nanotube with an inner diameter of 5.90 nm in atomic scale. The distance between neighboring rings is 6.4 Å.[37]
The P6 ring shaped molecule is not stable in isolation.
Single-layer blue phosphorus was first produced in 2016 by the method of molecular beam epitaxy from black phosphorus as precursor.[38]
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Main article: Diphosphorus |
The diphosphorus allotrope (P2) can normally be obtained only under extreme conditions (for example, from P4 at 1100 kelvin). In 2006, the diatomic molecule was generated in homogeneous solution under normal conditions with the use of transition metal complexes (for example, tungsten and niobium).[39]
Diphosphorus is the gaseous form of phosphorus, and the thermodynamically stable form between 1200 °C and 2000 °C. The dissociation of tetraphosphorus (P4) begins at lower temperature: the percentage of P2 at 800 °C is ≈ 1%. At temperatures above about 2000 °C, the diphosphorus molecule begins to dissociate into atomic phosphorus.
P12 nanorod polymers were isolated from CuI-P complexes using low temperature treatment.[40]
Red/brown phosphorus was shown to be stable in air for several weeks and have properties distinct from those of red phosphorus. Electron microscopy showed that red/brown phosphorus forms long, parallel nanorods with a diameter between 3.4 Å and 4.7 Å.[40]
| Form | white(α) | white(β) | violet | black |
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| Symmetry | Body-centred cubic | Triclinic | Monoclinic | Orthorhombic |
| Pearson symbol | aP24 | mP84 | oS8 | |
| Space group | I43m | P1 No.2 | P2/c No.13 | Cmca No.64 |
| Density (g/cm3) | 1.828 | 1.88 | 2.36 | 2.69 |
| Bandgap (eV) | 2.1 | 1.5 | 0.34 | |
| Refractive index | 1.8244 | 2.6 | 2.4 |
