Sulphur and the rest of the elements of group 16 are less electronegative than oxygen. Therefore, their atoms cannot accept electrons easily. They can acquire the $n{s}^{2}n{p}^6$ configuration by sharing two electrons with the atoms of the other elements and thus, exhibit a $+2$ oxidation state in their compounds. In addition to this, their atoms have vacant $d$-orbitals in their valence shell. Electrons can be promoted here from $s$ and $p$-orbitals of the same shell, hence they can show $+4$ and $+6$ oxidation states.
Like sulphur, oxygen does not show $+4$ and $+6$ oxidation states. The reason is that oxygen: