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You're correct in that IE increases from left to right, but Nitrogen has a higher IE because it has a half-filled p subshell which is more stable than Oxygen's 4 e- p subshell, so it actually requires more energy to take away that electron from Nitrogen than Oxygen
To explain this break in the periodic trend, we need to take a look at the electron configuration of nitrogen and oxygen. The electron configuration of nitrogen is 1s2 2s2 2p3, which means it has a stable, half-filled 2p shell. Oxygen's electron configuration is 1s2 2s2 2p4, one electron beyond stable in the 2p shell. So taking an electron from nitrogen would make it less stable, that's why nitrogen has a higher ionization energy than oxygen.
Ionization energy does increase going left to right across a periodic table. However, comparing the electronic configuration of nitrogen and oxygen, you can see that nitrogen ends at 2p3 and oxygen is 2p4. Half or fully filled orbitals are stable, and nitrogen's 2p orbital is filled halfway so it is more stable than oxygen, meaning it would require more energy to remove an electron from nitrogen than oxygen.
Rita Chen 1C wrote:I was wondering whether this goes for the ones below as well. I was a little confused whether these exceptions are for all of the ones that follow the same atom trend.
This doesn't apply to the ones below, hence why it's an exception to the general IE trend.
Possible explanation: In the other atoms, there are more electrons, including within the d-orbital, that would help to stabilize the atom. In other words, since there's more electrons, the difference of one electron would be less likely to destabilize the atom and have a large ionization energy jump between electrons within the same orbital. I hope that makes sense.
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