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

This tells me the it need to be easier to remove an electron native oxygen than it is for nitrogen as the electron in oxygen is slightly additional away indigenous the nucleus meaning lesser nuclear charge.

But why is that harder to remove an electron from oxygen, i.e. Why is the very first ionization energy of oxygen higher?

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



You watch 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 because nitrogen, in a stable fifty percent filled orbit state, is comparatively much more stable than oxygen. Oxygen, top top the other hand, would have tendency to lose an electron quickly to attain it"s an ext stable half filled orbital state.

Also, as a rule, half fill and completely filled orbital claims are more stable as compared to other configurations due to the fact that they attribute to maximum exchange energies.



Oxygen has actually a lower first ionization power as the electron the is gotten rid of is comes from a paired orbital.

Electrons within the same orbital experience maximum repulsion together the circulation of their wavefunctions is the same, so the probability thickness distribution is the same and the electrons have the right to be assumed of as occupying the exact same space. This maximizes your repulsion and increases the potential power of the electrons in that orbital, do the electrons much easier to remove. This is in spite of the increased efficient nuclear charge experienced by the electron in oxygen and the reduced radius the the orbital.

See: "Physical", Atkins, P.W. Ar 13.4, p.p.370 (4th edition) - sorry, I have an old one!


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