aluminum chloride dimer (Al2Cl6)

Write down aluminum chloride and you naturally write AlCl3. But heat the solid until it sublimes, catch the vapor, and weigh a molecule of it, and you find it is twice as heavy as AlCl3: nature pairs two units together into Al2Cl6. Why would a salt-looking compound travel around as a double molecule? The answer is the same electron deficiency that haunts all of group 13.

In a single AlCl3 unit the aluminum has only six valence electrons, two short of an octet, and an empty orbital, just like boron in BF3. To satisfy itself, each aluminum borrows a lone pair from a chlorine on a neighboring AlCl3 unit. Two units lock together through two bridging chlorines, giving Al2Cl6: each aluminum is now four-coordinate and tetrahedral, with two terminal Cl and two shared bridging Cl that each donate a lone pair across to the other aluminum (Al-Cl-Al bridges, here ordinary dative bonds, not the 3c-2e bonds of boranes). The result is a neat edge-sharing pair of tetrahedra. On strong heating the dimer cracks back into monomeric AlCl3, and in water it hydrolyzes violently, fuming as it forms aqua ions and HCl.

This dimerization matters because it is the cleanest demonstration that group-13 trihalides crave a fourth electron pair, the same drive that makes them strong Lewis acids. AlCl3 is the classic Lewis-acid catalyst of Friedel-Crafts reactions, where it pulls a halide off an alkyl or acyl halide to make a reactive cation. The dimer story also shows that a formula is not a structure: knowing the empirical formula AlCl3 tells you nothing about whether the species is a giant ionic lattice (the solid), a bridged dimer (the vapor and nonpolar solvents), or a solvated monomer.