Oxidation state

i. The group 16 elements have the valence shell electronic configuration $\mathrm{ns^2np^4}$. They attain a noble gas configuration either by gaining two electrons, forming $\mathrm{E^{2-}}$ ions, or by sharing two electrons, forming two covalent bonds. These elements thus show −2 and +2 oxidation states in their compounds.

Oxygen, being highly electronegative, shows a common oxidation state of −2 except in two cases: in $\mathrm{OF_2}$ its oxidation state is +2, and in peroxides it shows oxidation state −1 ($\mathrm{H_2O_2}$, $\mathrm{Na_2O_2}$). The other elements of the group exhibit +2, +4 and +6 oxidation states, with +4 and +6 being more common. The stability of the higher (+6) oxidation state decreases down the group while the stability of the lower oxidation state (+4) increases down the group, due to the inert pair effect. Bonding in the +4 and +6 oxidation states is primarily covalent.

ii. The group 17 elements are represented by their valence shell electronic configuration $\mathrm{ns^2np^5}$. They attain a noble gas configuration either by gaining one electron, forming $\mathrm{E^{-}}$ ions, or by sharing one electron, forming one covalent bond. All halogens exhibit the −1 oxidation state. However Cl, Br and I exhibit +1, +3, +5 and +7 oxidation states as well. This is because they are less electronegative than F and possess empty d-orbitals in the valence shell and therefore can expand the octet. The oxidation states +4 and +6 occur in the oxides and oxoacids of Cl and Br. The fluorine atom has no d-orbitals in its valence shell and therefore cannot expand its octet; being the most electronegative element, fluorine exhibits mostly the −1 oxidation state. …