Magnesium is found to have a higher first ionization energy value than aluminum, how would you explain this exception to the general trend in terms of electron arrangements and attraction/repulsion?

For Mg, $Z$ $=$ $12$. Therefore its electronic configuration is $1 {s}^{2} 2 {s}^{2} 2 {p}^{6} 3 {s}^{2}$; versus Al, $Z$ $=$ $13$; $1 {s}^{2} 2 {s}^{2} 2 {p}^{6} 3 {s}^{2} 3 {p}^{1}$. While aluminum metal has extra nuclear charge, $p$ orbitals demonstrably have zero electron density at the nucleus, and should be easier to ionize than the $s$ electrons of the alkaline earth (why?), given that $s$ orbital electrons have some electron density at the nucleus.