phosphorus halides

React phosphorus with a halogen such as chlorine and you get phosphorus halides, of which the two famous ones are phosphorus trichloride, PCl3, and phosphorus pentachloride, PCl5. They are fuming, reactive liquids and solids that hiss and smoke in moist air, and they are the everyday chlorinating reagents that chemists reach for when they want to swap an oxygen for a chlorine.

PCl3 is a trigonal pyramid: phosphorus with three P-Cl bonds and a lone pair, the same shape family as ammonia and phosphine. PCl5 is more interesting. In the gas phase it is a trigonal bipyramid — five chlorines around the phosphorus, three in an equatorial plane and two pointing along the axis. Older textbooks explained the five bonds by saying phosphorus "expands its octet" using empty 3d orbitals, but that picture is now regarded as largely wrong: the d orbitals are too high in energy to contribute much. The honest modern description uses delocalised molecular bonding (a three-centre, four-electron picture for the axial bonds) that needs no d orbitals. Stranger still, solid PCl5 is not made of PCl5 molecules at all: it is an ionic solid, [PCl4]+ [PCl6]-, a tetrahedral cation and an octahedral anion sitting in a lattice. Both chlorides react violently with water: PCl3 hydrolyses to phosphorous acid and HCl, PCl5 to phosphoric acid and HCl.

These halides are genuinely useful. PCl3 and PCl5 convert alcohols into chloroalkanes and carboxylic acids into acyl chlorides, which is why they turn up constantly in organic synthesis, and PCl3 is the starting point for many organophosphorus compounds, from flame retardants to pesticides to the phosphine ligands of catalysis. They also make a clean teaching example: PCl5 exists (five bonds to phosphorus) while NCl5 does not, because phosphorus is large enough to fit five chlorines around it whereas the small nitrogen atom cannot, and nitrogen has no low energy way to accommodate more than four bonds.