Question

With suitable explanation, give a brief account of the oxidation states exhibited by the elements of group 16.

Answer 1

1. All Group 16 elements commonly show an oxidation state of −2 due to their valence electron configuration, ns² np⁴, needing two electrons to complete their octet.
2. Oxygen mostly exhibits −2 state but can also show +2 in OF₂ and −1 in peroxides \[\ce{(O^{2-}_2)}\]. Its electronegativity is highest, so it prefers to gain two electrons.
3. Down the group, the tendency to have −2 oxidation state decreases because electronegativity decreases from oxygen to polonium. Polonium does not show −2 state.
4. Besides −2, sulphur and heavier elements show positive oxidation states of +2, +4, and +6, with +4 and +6 being more stable.
5. These higher oxidation states arise because heavier elements have vacant d-orbitals. They can excite electrons to d-orbitals and form more bonds (e.g., in +4 and +6 states).
6. Oxygen does not have d-orbitals and hence does not show positive oxidation states like other elements of the group.
7. The +4 state involves sp³d hybridisation, and the +6 state involves sp³d² hybridisation, allowing the formation of 4 or 6 bonds, respectively.