The first ionization enthalpy values (in $kJ\,mol^{-1}$) of group $13$ elements are :

$B$	$801$
$Al$	$577$
$Ga$	$579$
$In$	$558$
$Tl$	$589$

How would you explain this deviation from the general trend?

(N/A) On moving down a group,ionization enthalpy generally decreases due to an increase in the atomic size and shielding effect.
Thus,on moving down group $13$,ionization enthalpy decreases from $B$ to $Al$.
However,$Ga$ has a higher ionization enthalpy than $Al$ because $Ga$ follows the $d$-block elements.
The shielding provided by $d$-electrons is poor,which leads to a higher effective nuclear charge experienced by the valence electrons in $Ga$.
Moving from $Ga$ to $In$,the ionization enthalpy decreases due to the increase in atomic size.
Finally,from $In$ to $Tl$,the ionization enthalpy increases because $Tl$ follows the $4f$ and $5d$ orbitals.
The poor shielding effect of $f$ and $d$ electrons results in a higher effective nuclear charge,holding the valence electron more strongly.