7th Edition #1.13

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Linyu Zeng 1H
Posts: 28
Joined: Fri Sep 28, 2018 12:19 am

7th Edition #1.13

Postby Linyu Zeng 1H » Wed Oct 31, 2018 6:01 pm

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 anomaly.

I thought in the review session yesterday 3pm the TA explained that this is because O is closer to reach the half-filled state by losing only one e-. (Or I remember it wrong) But the answer in the textbook is like this: O is the first element in which p-electron must be paired, which leads to the electron-electron repulsion. Does anyone know which explanation is correct or both factors play a role in this phenomenon?

Camille Marangi 2E
Posts: 60
Joined: Fri Sep 28, 2018 12:26 am

Re: 7th Edition #1.13

Postby Camille Marangi 2E » Wed Oct 31, 2018 6:14 pm

I think both explanations revolve around the same concept in regards to oxygen's ionization energy. Oxygen will have lower ionization energy than nitrogen and fluorine because it will be more stable by losing one electron to reach the half-filled state much like your TA said and the electron repulsions between the single pair of electrons in the 2px orbital also lead to a smaller ionization energy. So basically both factors lead to the lower ionization energy.

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