Phosphorus pentoxide is a chemical compound with molecular formula P4O10 (with its common name derived from its empirical formula, P2O5). This white crystalline solid is the anhydride of phosphoric acid. It is a powerful desiccant and dehydrating agent.
Phosphorus pentoxide crystallizes in at least four forms or polymorphs. The most familiar one, a metastable form, shown in the figure, comprises molecules of P4O10. Weak van der Waals forces hold these molecules together in a hexagonal lattice (However, in spite of the high symmetry of the molecules, the crystal packing is not a close packing). The structure of the P4O10 cage is reminiscent of adamantane with Td symmetry point group. It is closely related to the corresponding anhydride of phosphorous acid, P4O6. The latter lacks terminal oxo groups. Its density is 2.30 g/cm3. It boils at 423 °C under atmospheric pressure; if heated more rapidly it can sublimate. This form can be made by condensing the vapor of phosphorus pentoxide rapidly, the result is an extremely hygroscopic solid.
The other polymorphs are polymeric, but in each case the phosphorus atoms are bound by a tetrahedron of oxygen atoms, one of which forms a terminal P=O bond involving the donation of the terminal oxygen p-orbital electrons to the antibonding phosphorus-oxygen single bonds. The macromolecular form can be made by heating the compound in a sealed tube for several hours, and maintaining the melt at a high temperature before cooling the melt to the solid. The metastable orthorhombic, "O"-form (density 2.72 g/cm3, melting point 562 °C), adopts a layered structure consisting of interconnected P6O6 rings, not unlike the structure adopted by certain polysilicates. The stable form is a higher density phase, also orthorhombic, the so-called O' form. It consists of a 3-dimensional framework, density 3.5 g/cm3. The remaining polymorph is a glass or amorphous form; it can be made by fusing any of the others.
P4O10 is prepared by burning tetraphosphorus with sufficient supply of oxygen:
For most of the 20th century, phosphorus pentoxide was used to provide a supply of concentrated pure phosphoric acid. In the thermal process, the phosphorus pentoxide obtained by burning white phosphorus was dissolved in dilute phosphoric acid to produce concentrated acid. Improvements in filter technology is leading to the "wet phosphoric acid process" taking over from the thermal process, obviating the need to produce white phosphorus as a starting material. The dehydration of phosphoric acid to give phosphorus pentoxide is not possible as on heating metaphosphoric acid will boil without losing all its water.
Phosphorus pentoxide is a potent dehydrating agent as indicated by the exothermic nature of its hydrolysis:
However, its utility for drying is limited somewhat by its tendency to form a protective viscous coating that inhibits further dehydration by unspent material. A granular form of P4O10 is used in desiccators.
The indicated coproduct P4O9(OH)2 is an idealized formula for undefined products resulting from the hydration of P4O10.
The desiccating power of P4O10 is strong enough to convert many mineral acids to their anhydrides. Examples: HNO3 is converted to N2O5; H2SO4 is converted to SO3; HClO4 is converted to Cl2O7; CF3SO3H is converted to (CF3)2S2O5.
Related phosphorus oxides
Phosphorus pentoxide itself is not flammable. Just like sulfur trioxide, it reacts vigorously with water and water-containing substances like wood or cotton, liberates much heat and may even cause fire due to the highly exothermic nature of such reactions. It is corrosive to metal and is very irritating – it may cause severe burns to the eye, skin, mucous membrane, and respiratory tract even at concentrations as low as 1 mg/m3.