hypervalency (expanded octet)

If the octet rule says eight electrons is the magic number, why does sulfur in SF6 happily sit at the centre of six bonds, surrounded by twelve electrons? Molecules in which a central atom appears to hold more than eight valence electrons are called hypervalent, or expanded-octet, species. They are common from the third period downward — phosphorus, sulfur, chlorine, xenon — in compounds like PF5, SF6, ClF3 and XeF4.

The old textbook explanation was that these heavier atoms borrow their empty d orbitals to make room for the extra pairs (sp3d, sp3d2 hybrids). This is now regarded as largely incorrect: high-level calculations show the d orbitals are too high in energy to contribute much. The modern, honest picture is that the bonding is delocalised — three-centre four-electron bonds, described by molecular-orbital theory, let a central atom hold several highly electronegative neighbours (usually F, O, or Cl) using mainly its s and p orbitals, with the bonding pairs spread over more than two atoms. The Lewis dot count still says twelve around sulfur, but no twelve electrons are truly crammed into one little atom's valence shell using d orbitals.

Hypervalency matters because it is how the chemistry of the heavier p-block actually works: sulfuric acid, phosphates, perchlorates, the noble-gas fluorides and countless inorganic reagents all involve expanded-octet centres. It is also why hypervalency is rare for the small second-period atoms (nitrogen never makes NF5): they are simply too small to fit five or six bulky neighbours, regardless of any d-orbital argument.