Question

The +1 oxidation state in group 13 and +2 oxidation state in group 14 become more and more stable with increasing atomic number.Explain why?

Answer 1

The outer electronic configuration of group 13 elements is ${{ns}^{2}}$ ${{np}^{1}}$ and hence, their group oxidation state is + 3. Both
B and Al show an oxidation state of + 3. However, from Ga to Tl, the elements have either d or d and f-electrons besides s- and p-electrons. Since, d and f-electrons have poor shielding effect, therefore, they do not screen s-electrons of the valence shell effectively from the attraction of the nucleus.
Consequently, two s-electrons of the valence shell remain attracted by the nucleus and hence, do not participate in bond formation while only the remaining one p-electron of the valence shell participates in bond formation. As a result, elements from Ga to Tl show oxidation states of +1 and + 3.
This is called inert pair effect. Further, as the number of d or d and f-electrons- increases down the group the inert pair effect becomes more and more prominent. Consequently, the stability of + 3 oxidation state decreases while that of +1 oxidation state increases down the’group.
The outer electronic configuration of group 14 elements in ${{ns}^{2}}$ ${{np}^{2}}$ and hence, their group oxidation state is + 4. Both
C and Si show an oxidation state of 4. However, from Ge to Pb, the elements have either d or both d- and f-electrons in addition to s and p-electrons. Due to poor shielding of s-electrons of valence shell by these d and f-electrons, the inert pair effect comes into play.
As a result, the elements from Ge to Pb show oxidation states of + 2 and + 4 . Further, as the number of d or d and f-electrons increases down the group, the inert pair effect becomes more and more predominant. Consequendy, the stability of + 4 oxidation state decreases while that of + 2 oxidation state increases down the group from Ge to Pb.