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[Andrew Berard 1F]

The problem states: "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 abnormality."
Can someone please explain this to me? Also, should we know the exceptions to the periodic table trends or just the general rules? Thanks!:-)

[tmehrazar]

So although the trend is increasing from left to right and down to up for ionization energy, atoms that are stable, meaning either they have all unpaired electrons in their shells or all paired, have a higher ionization energy because when it is stable, it's harder to remove an electron meaning it has a higher ionization energy. Oxygen for example is unstable because it has four electrons in the 3p orbital, so their is one pair and two unpaired. Nitrogen on the other hand has three electrons in the 3p orbital, so it is stable with all unpaired electrons. Flourine has 5 electrons in its 3p orbital, so it has two pairs and one unpaired so in comparison to oxygen, the one extra electron in oxygen will be easier removed than the two extra electron in fluorine I believe, so fluorine has a higher ionization energy (this follows the trend anyway). I think we should know the trends and then check the electron configuration to see how stable the atoms are in order to catch these exceptions. Hope this helps!

[Madison Hacker 1L]

Yes, because the electrons in the 2p orbital are unstable, Oxygen will want to gain or receive electrons badly making the ionization lower.

[105012653 1F]

"Oxygen will want to gain or receive electrons badly making the ionization lower" im lost on this concept, can you further explain?