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Oxygen's ionization energy is lower than Nitrogen's despite the fact that ionization energy typically increases as you go across a period because of their electron configurations. Oxygen's electron configuration is 1s22s22p4. Nitrogen's electron configuration is 1s22s22p3. Because of Nitrogen's 2p3 orbital, Nitrogen has a half-filled orbital that is actually more stable than Oxygen's 2p4 orbital. Since Oxygen is less stable, it wants to become more stable by losing that extra electron it has to allow Oxygen to have a half-filled orbital like Nitrogen. Nitrogen, therefore, is more reluctant to have an electron taken away because it is more stable with a half filled orbital. As a result, the ionization energy, or energy needed to take away an electron, of Oxygen is lower than Nitrogen's ionization energy.
This is beacause an atom that has half-filled or full orbitals are more stable. For oxygen, its electron configuration is 1s2 2s2 2p4, and nitrogen's is 1s2 2s2 2p3. Nitrogen has half-filled 2p orbitals, so losing an electron would cause it to become more unstable. However, oxygen can become more stable if it loses on electron, so it has a lower ionization energy compared to Nitrogen.
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