How do the oxidation states of the elements vary in the transition series?

$(a)$ 

In the beginning of the series, very less number of $d$-electrons are available for the chemical bonding. Hence, less number of oxidation states are shown by elements present at the beginning of series. Ex. : Scandium $\left(d^{1}\right)$ is known only in $(+3)$. Titanium $\left(d^{2}\right)$ exists in $(+2),(+3)$ and $(+4)$. However, $\mathrm{Ti}(\mathrm{IV})$ is more stable than $\mathrm{Ti}(\mathrm{III})$ or $\mathrm{Ti}(\mathrm{II})$.

At the end of the series, there are too many $d$-electrons and $d$-orbitals are completely occupied. Ex.: $\mathrm{Zn}$ and $\mathrm{Cu}$. Hence, there are very less number of orbitals available to share electrons with others for higher valance. Hence, these elements show very less number of oxidation states.Ex. : $\mathrm{Zn}$ $(II)$ is only known while copper is known to exist in $\mathrm{Cu}$ $(I)$ or $\mathrm{Cu}$ $(II)$.

The greatest number of oxidation states of the elements are known in the middle of the series. Ex. : Mn shows oxidation state from $(+2)$ to $(+7)$. The stability of the higher oxidation states corresponds in value to the sum of $s$ - and $d$-electrons upto manganese and then decreases abruptly.

Ex. : $\mathrm{Ti}^{\mathrm{IV}} \mathrm{O}_{2}, \mathrm{~V}^{\mathrm{V}} \mathrm{O}_{2+}, \mathrm{Cr}^{\mathrm{VI}} \mathrm{O}_{4}^{2-}, \mathrm{Mn}^{\mathrm{VII}} \mathrm{O}_{4}^{-}, \mathrm{Fe}^{\mathrm{II}, \mathrm{III}}, \mathrm{Co}^{\mathrm{II}, \mathrm{III}}, \mathrm{Ni}^{\mathrm{II}}, \mathrm{Cu}^{\mathrm{I}, \mathrm{II}}, \mathrm{Zn}^{\mathrm{II}}$

The elements showing more than one oxidation states have the difference on oxidation state of unity. This is opposite to the oxidation states shown by non-transition elements which normally differs by two. Ex.: $V^{I I}, V^{I I I}, V^{I V}, V^{V}$.

Down the group, the stability of elements in higher oxidation states increases because the removal of electrons from $d$-orbitals becomes easy. Ex. : Mo(VI) and W(VI) are found to be more stable than $\mathrm{Cr}(\mathrm{VI})$. As a result, $\mathrm{Cr}(\mathrm{VI})$ in the form of dichromate act as strong oxidizing agent in acidic medium where as $\mathrm{MoO}_{3}$ and $\mathrm{WO}_{3}$ are not. Thus in $p$-block elements, the lower oxidation state is favoured by higher members of the group due to innert pair effect while in transition elements from group-4 to group-10, higher members show high oxidation state.

Low oxidation states are found when a complex compound has ligands capable of $\pi$-acceptor character in addition to $\sigma$-bonding.

Ex. : In $\mathrm{Ni}(\mathrm{CO})_{4}$ and $\mathrm{Fe}(\mathrm{CO})_{5}$, the oxidation state of nickel and iron is zero.