Which one of the following orders is correct for the first ionisation energies of the elements?

How would you explain the fact that the first ionisation enthalpy of sodium is lower than that of magnesium,but its second ionisation enthalpy is higher than that of magnesium?

Assertion $(A)$: $16$th group elements have higher ionisation enthalpy values than $15$th group elements in the corresponding periods.
Reason $(R)$: $15$th group elements have half-filled stable electronic configurations.

The elements of Group $13$ with highest and lowest first ionisation enthalpies are respectively$:$

Match List-$I$ with List-$II$:

List-$I$ (Electronic configuration)	List-$II$ ($\Delta_{i}H$ in $kJ\ mol^{-1}$)
$(a)$ $1s^{2} 2s^{2}$	$(i)$ $801$
$(b)$ $1s^{2} 2s^{2} 2p^{4}$	$(ii)$ $899$
$(c)$ $1s^{2} 2s^{2} 2p^{3}$	$(iii)$ $1314$
$(d)$ $1s^{2} 2s^{2} 2p^{1}$	$(iv)$ $1402$

Choose the most appropriate answer from the options given below:

What are the various factors due to which the ionization enthalpy of the main group elements tends to decrease down a group?