Why can sulfur hexafluoride have more than eight electrons around its sulfur atom?

In the context of Gilbert N. Lewis’s model, sulfur hexafluoride (SF₆) is classified as a hypervalent compound because it possesses more electrons around the central sulfur atom than the octet rule typically allows.1

While traditional explanations often suggest that sulfur uses its empty d orbitals to accommodate these extra electrons—a concept known as octet expansion—modern quantum mechanical theories offer a different perspective. These theories indicate that hypervalence is not necessarily dependent on d-orbital availability. Instead, it is often viewed as a consequence of the larger atomic radius of period-3 elements compared to period-2 elements, which allows more atoms to pack around a central atom.12

Furthermore, molecular orbital theory explains that the bonding in SF₆ involves delocalization, where the bonding effect of electron pairs is distributed over all the atoms in the molecule. This approach eliminates the need to invoke octet expansion, treating it as a useful rule of thumb rather than a fundamental physical explanation.2