Question

Consider the following ionization energy data $\left(kJ/mol\right)$:
$N=1402$
$O=1314$
$F=1681$
$Ne=2081$
What is the BEST explanation for why the ionization energy for oxygen is the lowest for these elemens?

• A. O has additional electron-electron repulsions.
• B. Shielding is higher in O than the other elements.
• C. Penetration of electrons to the nucleus is high in O.
• D. Effective nuclear energy is low in O.

Answer 1

Answer: (A)

O has additional electron-electron repulsions.

Electronic configuration of Oxygen is

$1s^2{2s}^{2}2p^4$

Electronic configuration of Nitrogen is

$1s^2{2s}^{2}2p^3$

You must be knowing that an atom containing half filled or fully filled orbitals are very stable. Nitrogen has a half filled $2$p orbital. So when it is already stable, why would it want to lose an electron upon supplying energy and become unstable? This is why one requires large amounts of energy to ionised the nitrogen atom.

On the other hand, since Oxygen is already unstable relative to Nitrogen, by losing one electron it attains a stable half filled $2$p orbital. So oxygen undergoes ionisation at a relatively lower energy. This is why nitrogen has a higher ionisation energy than oxygen.