In nitrogen: $ce 2s^2 2p^3$ In oxygen: $ce 2s^2 2p^4$

This tells me that it need to be simpler to rerelocate an electron from oxygen than it is for nitrogen as the electron in oxygen is slightly even more ameans from the nucleus interpretation lesser nuclear charge.

But why is it harder to remove an electron from oxygen, i.e. why is the initially ionization energy of oxygen higher?




You are watching: Why does nitrogen have a higher ionization energy than oxygen

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You view from the electronic configurations:

nitrogen: $ce 2s^2 2p^3$ oxygen: $ce 2s^2 2p^4$

In reality, the first ionisation energy of nitrogen is better than the first ionisation energy of oxygen bereason nitrogen, in a steady fifty percent filled orbital state, is comparatively more stable than oxygen. Oxygen, on the various other hand also, would tend to lose an electron easily to accomplish it"s more secure fifty percent filled orbital state.

Also, as a dominance, half filled and also completely filled orbital claims are more secure as compared to other configurations because they attribute to maximum exreadjust energies.


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Oxygen has a lower first ionization power as the electron that is removed is coming from a paired orbital.
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Electrons within the very same orbital endure maximum repulsion as the distribution of their waveattributes is the same, so the probability density distribution is the same and the electrons can be believed of as occupying the same room. This maximizes their repulsion and also boosts the potential power of the electrons in that orbital, making the electrons much easier to rerelocate. This is despite the boosted efficient nuclear charge proficient by the electron in oxygen and also the reduced radius of the orbital.

See: "Physical yellowcomic.com", Atkins, P.W. Section 13.4, p.p.370 (fourth edition) - sorry, I have an old one!


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