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Doesn't ionization energy increase as you move from left to right on the periodic table?
Nitrogen is an exception because it really wants to have a half-full sub shell. But yes, ionization trend on the PT is increasing up and to the right. Just keep in mind Nitrogen has a larger ionization energy than Oxygen.
To add to this, Boron (and elements under) have higher ionization energies than Beryllium (and elements under) for a similar reason of electron shielding.
The reason lies in the electron configuration of the oxygen atom. Oxygen has 6 valence electrons in its n = 2 shell. 2 of these electrons will occupy the 2s subshell while the remaining 4 will occupy the 2p subshell. Because 2p subshell has 3 orbitals, the oxygen atom will have 1 pair of electrons in one orbital and the others will only have one unpaired electron. Because of the repulsion between the paired electrons, the ionization energy required to remove that electron that is being repelled by its pair and others is lower. This is because the repulsion between the parallel electrons in different orbitals isn't as strong as the repulsion between the electron pairs. So, if oxygen were to have 3 unpaired electrons in each of the orbitals, the ionization energy would have been greater.
Nitrogen's electron configuration is more stable than oxygens because it has more symmetry, making it more difficult to remove an electron. Because ionization energy is the energy required to remove an electron, nitrogen's is higher than oxygen's.
Draw the electronic configuration of the two atoms and you will see that Nitrogen has a half-full 2p shell. This makes it more stable than other unfull 2p shells.
nitrogen has all three of its sub shells filled while oxygen has unpaired electrons in its sub shells. this makes nitrogen more stable compared to oxygen. it is easier to remove these unpaired electrons from oxygen, giving it a lower ionization energy since it does not require as much energy as it would for nitrogen.
nitrogen has a half filled 2p orbital. this makes it more symmetrical therefore more stable. it is harder to get the electron form a highly stable atom. Even though application of periodic table trends indicates that Oxygen should have a higher first ionization energy, it is not the case with nitrogen.
Nitrogen has symmetry in it's 2p orbital (1 e- in each sub shell) and no electron-electron repulsion whereas Oxygen has one more electron than Nitrogen meaning its 2px orbital has electron-electron repulsion so its easier to ionize the fourth electron in the 2p orbital.
nitrogen is more stable, thus it takes more energy to remove an electron, compared to oxygen.
According to Hund's rule, an electron has lower energy when three subshells are filled within an orbital.
When you look at the electron configurations for nitrogen and oxygen you can see that oxygen has one group of paired electrons. Due to the electron electron repulsions that occur as a result of the paired electrons, it is pushed to a higher energy level which causes it to require less energy to remove compared to nitrogen where all the outer subshells are filled with a singular electron.
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