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Because nitrogen has 3 parallel spin electrons in separate orbitals in the 2p subshell, while Oxygen has 3 parallel spin electrons and one antiparallel electron in the same orbital. Because of electron-electron repulsion and the fact that having a half-full 2p subshell is more stable, it is easier (requires less energy) to remove the electron in oxygen than in nitrogen.
Nitrogen has a half-full shell, which is more stable, so it would be harder to remove an electron, thus increasing the ionization energy. In addition, the 4th electron in the oxygen atom has an electron-electron repulsion with the other electron in the orbital, thus making it easier to remove the electron, which lowers the ionization energy.
by looking at the periodic table one would assume nitrogen to be less electronegative than oxygen following the trend on the periodic table, but that in fact due to the symmetry that the nitrogen atom has in the 2p subshell(compared to that of oxygen), more clearly seen in a arrow diagram, gives it a higher electronegativity than oxygen despite having one more electron!
From periodic trends, it can be incorrectly assumed that Oxygen has a higher ionization energy than Nitrogen. However, Oxygen wants a full shell so is more willing to give up its electrons (requiring less energy to its electrons and therefore giving Oxygen a lower ionization energy than Nitrogen).
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