Check Assertion A

Assertion: The first ionization enthalpy of oxygen is lower than that of nitrogen.

Electronic configurations: $\mathrm{N}: 1s^2\,2s^2\,2p^3$ $\mathrm{O}: 1s^2\,2s^2\,2p^4$

Nitrogen has a half-filled $2p^3$ configuration, which is especially stable.

Oxygen has $2p^4$, so one of the $2p$ orbitals contains a paired set of electrons, causing extra repulsion.

Therefore, removing one electron from oxygen is easier than from nitrogen.

So, Assertion A is correct.

Check Reason R

Reason: The four electrons in 2p orbitals of oxygen experience more electron-electron repulsion.

In oxygen, the $2p^4$ arrangement means one orbital has two electrons paired in it. This produces extra electron-electron repulsion compared with nitrogen's $2p^3$ arrangement, where all three $2p$ electrons remain unpaired in separate orbitals.

So, Reason R is also correct.

Does R correctly explain A?

The lowered first ionization enthalpy of oxygen relative to nitrogen is explained specifically by:

• the extra repulsion due to pairing in one $2p$ orbital of oxygen, and
• the special stability of half-filled $2p^3$ configuration of nitrogen.

The given reason mentions increased electron-electron repulsion in oxygen, which is true and contributes to the lower ionization enthalpy. Hence it does explain why oxygen loses an electron more easily than nitrogen.

Therefore, R is the correct explanation of A.

Correct option

$\boxed{\text{A}}$

Comparison with stored answer

Stored correct answer: B

My derived answer is A. The stored answer appears incorrect because the stated reason directly accounts for the anomalously lower first ionization enthalpy of oxygen compared with nitrogen.