Which of the following is the
electron configuration for a magnesium(II) ion? A) 1s², B) 1s² 2s² 2p⁵, C) 1s² 2s²
2p⁶, D) 1s² 2s² 2p⁶ 3s¹, or E) 1s² 2s² 2p⁶ 3s¹ 3p².
The easiest way of working out the
electron configuration for a magnesium(II) ion is to start with the electron
configuration of a magnesium atom. The element magnesium is to be
found in group two of the periodic table. It’s also in the third period. Magnesium has atomic number 12. This means that if we have a
magnesium atom, there must be 12 protons in the nucleus and 12 electrons in the
surrounding space. This is because atoms are
electrically neutral. Protons and electrons have equal
and opposite charge. So if we have 12 protons in a
neutral particle, we must have 12 electrons. So for a magnesium atom, we have 12
electrons to work with.
If we lose two electrons to form
the magnesium two plus ion, we must have 10 electrons in total. We start by placing two of those
electrons in the 1s subshell, followed by two electrons in the 2s subshell and a
further six electrons in the 2p subshell. So we’ve used up all the electrons
we were given, exactly 10: two in the 1s subshell, two in the 2s subshell, and six
in the 2p subshell. There’s only over one orbital in an
s subshell. So we can only fit two electrons in
those subshells. The 2p subshell can fit six
electrons because it’s got three orbitals.
In the first electron shell, we
only have the 1s subshell. So we have a maximum of two
electrons there. And in the second electron shell,
we have the 2s and 2p subshells, giving us a maximum of eight electrons. So we filled the first and second
electron shells. And we don’t have to go further
because we’ve run out of electrons.
So our final answer for the
electron configuration for a magnesium(II) ion is 1s² 2s² 2p⁶.