Ionization energy of O vs N

[Sofia Barker 2C]

Why does oxygen have a lower ionization energy than nitrogen? This question was in the homework but doesn’t oxygen having a lower ionization energy violate the periodic trend that ionization energy increases down a period?

[Ada Chung 1C]

My UA said that if you draw out the 2p orbital of O and N you will see that N has 3 half-filled orbitals which are supposedly more stable than O which has 2 half-filled orbitals.

[Christine Honda 2I]

Since there are two electrons in the first orbital of the O sublevel, there is greater electronic repulsion in the 2p sublevel for O than N. Therefore, it is easier to remove an electron from the O than the N, and the ionization energy of O is lower than N. Also nitrogen has a lower ionization energy than oxygen because nitrogen is half-filled which according to the Hund rule, half-filled and full filled orbitals are more stable.

[Rohan Kubba Dis 4B]

This is the main deviation of the periodic trend of ionization energy increasing to the right and up (fluorine). In this case, nitrogen has a a higher ionization energy than oxygen. The reason is because if you look at the 2p shells of each element, you will see that oxygen has 4 total electrons in this shell, with orientation: up+down, up, up while nitrogen has orientation: up,up,up. These are based off of Hund’s Rule and the Pauli Exclusion Principle. Essentially, it is easier to remove that down electron from oxygen’s shell than to remove one from nitrogen’s because nitrogen’s symmetrical orbital is more stable than oxygen’s.

[gabbymaraziti]

It's easier understand this if you look at it visually, by drawing out the p-orbitals for each atom. N has one electron in each p-orbital, and O has one electron in 2 of the orbitals, and two electrons in one of the orbitals. N's electron configuration is more stable because its orbitals are half-filled.

Half-filled and fully-filled orbitals have greater stability than other configurations, because of their symmetry.

Therefore, it would be easier to remove an electron from oxygen (i.e. lower ionization energy) and achieve a half-filled configuration, than it would be to violate the stability of nitrogen's half-filled configuration (i.e. higher ionization energy).

[Shimei_2F]

The ionization energy trend in the periodic table has a few exceptions with oxygen and nitrogen being one of them. The ionization energy of oxygen is lower than nitrogen because each p-orbital in a nitrogen atom is occupied by one electron but in oxygen, the first orbital is occupied by a pair of electrons. The energies of the paired electron rises due to the repulsion making one of them easier to remove than if two electrons had been in different orbitals.

[Noah Canio 3C]

This is based on the electron configuration of both oxygen and nitrogen. Oxygen has two unpaired electrons in its 2p orbitals while Nitrogen has three unpaired electrons. Based on my experience with one of the UA's, the symmetry of the electrons in nitrogen indicated by having a half complete valence shell makes nitrogen more willing to hold on to its electrons more tightly than expected.