Question

Question #8926e

Answer 1

Here is a plot of the ionization energies of the Period 2 elements.

For the elements between 2 and 10 (i.e. Li to F), the ionization energy generally increases, with two exceptions: B < Be, and O < N.

QUESTION I.

Why is the ionization energy of oxygen less than that of nitrogen?

Answer I.

The electron configuration of nitrogen is [He]`2s^2 2p_"x"2p_"y"2p_"z"`.

The electron configuration of oxygen is [He]`2s^2 2p_"x"^2 2p_"y"2p_"z"`.

In oxygen, two electrons must occupy the same orbital.

These electrons repel each other, so an oxygen atom is at a higher energy level than a nitrogen atom.

It takes less energy to remove an electron from an oxygen atom, so oxygen has a lower ionization energy.

Question II.

Why is the ionization energy of boron less than that of beryllium?

Answer II.

According to Slater's rules, electrons in `s` and `p` orbitals with the same value of `n` are equally effective in screening the nucleus. A `2s` electron screens another `2s` electron or `2p` electron by the same amount (0.35 charges).

So Option II is incorrect.

The correct answer must be Option III. Filled subshells are more stable than unfilled subshells.

CONCLUSION

This is a poorly worded question.

The correct answer is "I and III only", and that wasn't one of the options.