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

This speak me that it should be easier to remove an electron from oxygen 보다 it is because that nitrogen together the electron in oxygen is slightly further away indigenous the nucleus meaning lesser nuclear charge.

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




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

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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 higher than the first ionisation energy of oxygen due to the fact that nitrogen, in a stable fifty percent filled orbit state, is comparatively much more stable 보다 oxygen. Oxygen, ~ above the other hand, would tend to lose an electron conveniently to attain it\"s an ext stable fifty percent filled orbital state.

Also, together a rule, half to fill and totally filled orbital claims are much more stable as contrasted to other configurations because they attribute to maximum exchange energies.


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Oxygen has actually a lower first ionization energy as the electron the is gotten rid of is comes from a combine orbital.
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Electrons within the very same orbital experience maximum repulsion as the circulation of their wavefunctions is the same, therefore the probability density distribution is the same and also the electrons have the right to be assumed of as occupying the exact same space. This maximizes their repulsion and increases the potential energy of the electron in the orbital, make the electrons simpler to remove. This is in spite of the increased reliable nuclear charge knowledgeable by the electron in oxygen and also the lessened radius of the orbital.

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


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