The incorrect relation among the following is

Which of the following electrons should have the highest value of ionisation energy (for the same value of the principal quantum number)?

Consider the following ionization enthalpies of two elements $A$ and $B$.

Element	Ionization enthalpy $(kJ \ mol^{-1})$ $(1^{st}, 2^{nd}, 3^{rd})$
$A$	$899, 1757, 14847$
$B$	$737, 1450, 7731$

Which of the following statements is correct?

The first ionization energy of magnesium is smaller as compared to that of elements $X$ and $Y$,but higher than that of $Z$. The elements $X$,$Y$ and $Z$,respectively,are :

$B$ has a smaller first ionization enthalpy than $Be$. Consider the following statements:
$(I)$ It is easier to remove a $2p$ electron than a $2s$ electron.
$(II)$ The $2p$ electron of $B$ is more shielded from the nucleus by the inner core of electrons than the $2s$ electrons of $Be$.
$(III)$ The $2s$ electron has more penetration power than the $2p$ electron.
$(IV)$ The atomic radius of $B$ is more than $Be$.
(Atomic number: $B=5, Be=4$)
The correct statements are:

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