# What are the electronic configurations of common atomic ions?

Excluding the transition metals, the common ions of the 1st short rows are the same as that of the neutral inert gas. Consider the common ions of nitrogen, ${N}^{3 -}$, ${O}^{2 -}$, and ${F}^{-}$; all of these ions have 8 valence electrons, and a formal electronic configuration of $1 {s}^{2} 2 {s}^{2} 2 {p}^{6}$. While these ions have negative charges, these are isolelectronic with $\left[N e\right]$, the inert gas with 8 valence electrons.
$N {a}^{+}$, and $M {g}^{2 +}$ in the same way have the same formal $\left[N e\right]$ configuration. Why?
How do you propose to represent ${P}^{3 -}$, ${S}^{2 -}$, and $C {l}^{-}$?