Question

The first ionization enthalpy values (in kJmol-1 ) of group 13 elements are :B	Al	Ga	In	Tl
801	577	579	558	589
How would you explain this deviation from the general trend?

Answer 1

On moving down a group, ionization enthalpy generally decreases due to an increase in the atomic size and shielding. Thus, on moving down group $13$, ionization enthalpy decreases from $B$ to $Al$. But, $Ga$ has higher ionization enthalpy than $Al$. $Al$ follows immediately after $s$-block elements, whereas $Ga$ follows after $d$-block elements. The shielding provided by $d$-electrons is not very effective. These electrons do not shield the valence electrons very effectively. As a result, the valence electrons of $Ga$ experience a greater effective nuclear charge than those of $Al$. Further, moving from $Ga$ to $In$, the ionization enthalpy decreases due to an increase in the atomic size and shielding. But, on moving from $In$ to $Tl$, the ionization enthalpy again increases. In the periodic table, $Tl$ follows after $4f$ and $5d$ electrons. The shielding provided by the electrons in both these orbitals is not very effective. Therefore, the valence electron is held quite strongly by the nucleus.

Hence, the ionization energy of $Tl$ is on the higher side.