Textbook Problem 1.13

[Bethany Yang 2E]

Can someone explain to me the answer for this:
"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."

Also, is oxygen just an exception that we would need to memorize?

[Faith St Amant 3D]

I'm not sure if we'll be tested on the exceptions of the periodicity of ionization energies, but, from what I understand, the electron configuration of oxygen is the first of the p-block elements to require an electron to be paired in a p-subshell (nitrogen, for example, has 3 electrons all in separate p-subshells). This adds some repulsion into the system and therefore lowers the ionization energy of the atom. Hope that makes sense!

[annabelchen2a]

Hi: the ionization energy for oxygen is lower than that of nitrogen because when you look at the ground state electron configurations of each atom, you will see that nitrogen has 1 electron in each of the 3 orbitals (1 electron in Px, Py, and Pz), giving nitrogen a special kind of stability; whereas oxygen has an extra antiparallel electron and in turn will have an easier time than nitrogen losing that electron to gain that special stability ground state nitrogen has.