Question 1.13 from the Practice Problems

[Sydney Rohan 3G]

I do not understand Question 1.13 from the Practice Problems. The question states:

Ionization energies usually increase on going from left to right across the periodic table. The ionization energy for oxygen, however, is lower than that of either nitrogen or fluorine. Explain this anomaly.

Can someone please explain this question to me?

[Reza Hemmati 3L]

Hi!

The reason why fluorine's ionization energy is higher than oxygen's is because fluorine has a higher nuclear charge, and thus has a stronger pull on its electrons, making it harder to remove one(ionization energy = energy needed to remove 1 electron from the element in the gaseous state). This part of the question follows the trend that ionization increases as you go farther to the right.

The part where it might be confusing is that nitrogen, which is to the left of oxygen on the periodic table, has a higher ionization energy. This occurs because of how the electrons are arranged within the two's sublevels. In nitrogen, there is 1 electron per orbital in the 2p sublevel. In oxygen, the additional electron fills one of the three orbitals in the 2p sublevel, which ends up causing relatively more repulsion between the electrons. This repulsion causes oxygen's valence electrons to be pulled by the nucleus less strongly, thus making it easier to pull them from oxygen atoms(and making the ionization energy lower). Although the raw nuclear charge of oxygen is higher(bc it has 1 more proton), this extra repulsion is more significant.
(The same thing happens with phosphorus and sulfur, explained by their 3p sublevel)

Hope this helps! ^^