Question

Why is the ionization energy of the oxygen atom LESS than that of the nitrogen atom? Should not the greater atomic number of oxygen increase ionization energy?

Answer 1

This is arguably a consequence of `"Hund's rule of maximum multiplicity."`

Explanation:

And thus for the `"2p orbitals"`.... we get for the `2p_x`, `2p_y`,

`2p_z`, orbitals of nitrogen, `Z=7`...

`stackrel(uarr)ul` `stackrel(uarr)ul` `stackrel(uarr)ul`

.....rather than spin pairing.... which must occur for the ground state of the oxygen atom, `Z=8`, which accommodates the extra electron.....

`stackrel(uarr darr)ul` `stackrel(uarr)ul` `stackrel(uarr)ul`

And thus the reaction....`"Atom(g)"+Deltararr"Atom(g)"^(+) +e^-` requires LESS energy for the oxygen ground state. This effect wins over atomic charge...

Answer 2

Because Nitrogen is more stable than Oxygen because of half-filled electronic configurations.

Explanation:

As we know that Hund's rule states half-filled and full-filled orbitals are more stable. The electronic configuration of Nitrogen and Oxygen are as below:

Nitrogen = `1s^2,2s^2 2p^3`
Oxygen = `1s^2,2s^2 2p^4`

Note: p-subshell orbital can carry 6 electrons.

When we look at the above electronic configuration of Nitrogen and Oxygen, the p-subshell orbital is half-filled for nitrogen whereas Oxygen has one extra electron than half-filled configuration. So, Nitrogen is more stable than Oxygen and it requires higher energies to ionizes. And also when oxygen is ionized, it goes to the stable electronic configuration (half-filled electronic configuration).