Ions and magnetism

Anions add electrons to the next available orbital in aufbau order. Cations lose electrons from the subshell with the highest $n$ (the outermost). Main-group ions usually end up isoelectronic with a noble gas:

• $\ce{O}$, $[\mathrm{He}]\,2s^2\,2p^4$, gains 2 → $\ce{O^2-}$, $[\mathrm{He}]\,2s^2\,2p^6$ (the same as Ne)
• $\ce{Cl}$, $[\mathrm{Ne}]\,3s^2\,3p^5$, gains 1 → $\ce{Cl-}$, $[\mathrm{Ne}]\,3s^2\,3p^6$ (the same as Ar)
• $\ce{Mg}$, $[\mathrm{Ne}]\,3s^2$, loses 2 → $\ce{Mg^2+}$, $[\mathrm{Ne}]$

Heavier p-block metals can lose only their p electrons: tin, $[\mathrm{Kr}]\,5s^2\,4d^{10}\,5p^2$, gives $\ce{Sn^2+}$ = $[\mathrm{Kr}]\,4d^{10}\,5s^2$ (the $5p$ electrons leave first) and $\ce{Sn^4+}$ = $[\mathrm{Kr}]\,4d^{10}$.

The $4s$ subshell fills before $3d$, but once $3d$ holds electrons the $4s$ electrons are the outermost (highest $n$) and they are removed first. Remove all the $ns$ electrons before any $(n-1)d$ electrons:

Always start from the true neutral configuration: Cr is $[\mathrm{Ar}]\,4s^1\,3d^5$, so $\ce{Cr^3+}$ is $[\mathrm{Ar}]\,3d^3$.

Species with the same number of electrons are isoelectronic and have the same configuration. Count electrons as $Z$ minus the charge.

• 10 electrons, $[\mathrm{He}]\,2s^2\,2p^6$: $\ce{N^3-}$, $\ce{O^2-}$, $\ce{F-}$, $\ce{Ne}$, $\ce{Na+}$, $\ce{Mg^2+}$, $\ce{Al^3+}$
• 18 electrons, $[\mathrm{Ne}]\,3s^2\,3p^6$: $\ce{S^2-}$, $\ce{Cl-}$, $\ce{Ar}$, $\ce{K+}$, $\ce{Ca^2+}$

Isoelectronic species differ only in nuclear charge, which is what sets their relative sizes (see the periodic-trends topic). All of these have filled subshells, so all are diamagnetic.