Phosphorus: An Overview
Phosphorus is in Group V, period 3, with atomic number 15. It exists in several allotropic forms, the two most common being white and red phosphorus.
White Phosphorus
Extracted by heating a mixture of sand (SiO₂), coke and calcium phosphate in an electric furnace at about 1500°C: 2Ca₃(PO₄)₂(s) + 6SiO₂(s) → 6CaSiO₃(s) + P₄O₁₀(g); the phosphorus(V) oxide is then reduced by the coke: P₄O₁₀(g) + 10C(s) → P₄(s) + 10CO(g). The phosphorus vapour is cooled in cold water to solidify it, then compressed.
Properties: four phosphorus atoms arranged in a strained tetrahedron, making it unstable; denser than water; highly toxic, causing liver damage if inhaled; melts at 44°C, boils at 277°C; a white, waxy, translucent solid that glows green in the dark on exposure to oxygen (phosphorescence); ignites spontaneously in air, so it’s stored under water.
Red Phosphorus
Made by heating white phosphorus at 250°C, away from air, using a small amount of iodine or sulphur as a catalyst.
Properties: a red, opaque solid; denser than white phosphorus; a less strained structure makes it more stable; non-toxic when pure; sublimes at 290°C in the absence of air.
Chemical Properties
- With air: burns vigorously to form dense white fumes of phosphorus(V) oxide: P₄(s) + 5O₂(g) → P₄O₁₀(g), which dissolves in water to give phosphoric acid: P₄O₁₀(g) + 6H₂O(l) → 4H₃PO₄(aq). Burned in limited air instead, it forms phosphorus(III) oxide: P₄(s) + 3O₂(g) → P₄O₆(g), which dissolves in water to give phosphorous acid: P₄O₆(g) + 6H₂O(l) → 4H₃PO₃(aq).
- With chlorine: forms phosphorus(III) chloride or phosphorus(V) chloride, depending how much chlorine is available: P₄(s) + 6Cl₂(g) → 4PCl₃(l); P₄(s) + 10Cl₂(g) → 4PCl₅(l).
Uses of Phosphorus
Making phosphorus and phosphoric acid; manufacturing fertilisers such as superphosphate, ammonium phosphate and triple phosphate; making special glass for sodium lamps; manufacturing insecticides, pesticides, matches, explosives and fireworks; and making alloys such as phosphor bronze.
