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Oxygen has a lower ionization energy because when you add an electron, it is added to an already half full orbital while for nitrogen, only half the orbital is full. As a result, there is electron-electron repulsion which lowers the ionization energy for oxygen.
To add on, I like to visualize it by seeing the orbitals within the 2p sub shell. there's 3 orbitals in the 2p sub shell and N has 3 electrons in the 2p sub shell and O has 4. in N, the three electrons each get their own orbital so it's really stable and the electrons don't want to be removed. On the other hand, oxygen having four electrons means that one of the orbitals has two electrons spinning in opposite directions, so one of the electrons in this orbital will want to leave the orbital more than the stable nitrogen electrons.
It's kind of similar to why Cr and Cu have electron configurations different from what we expect - stability. It is more stable to have those 3 half filled orbitals in nitrogen than it is to have 1 pair of electrons, and then two unpaired electrons like in oxygen. Since it wants to be more stable, it is easier to remove an electron from oxygen, so the ionization energy is lower.
Is this the only exception to the general trend for ionization energy? ( That ionization energy increases across a period and decreases down a group?) Or do all group 16 elements have lower ionization energies than their preceding group 15 element?
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