Now Ionisation energy(I.E) is the energy required to remove an electron from the valence shell of an atom.As seen in the configuration,N atom has stable half filled valence p-orbital ,therefore large amount of energy is required to remove the valence electron from N atom. Therefore I.E of N is higher than that of O.

helium

On the periodic table, first ionization energy generally increases as you move left to right across a period. This is due to increasing nuclear charge, which results in the outermost electron being more strongly bound to the nucleus.

On the periodic table, first ionization energy generally decreases as you move down a group. This is because the outermost electron is, on average, farther from the nucleus, meaning it is held less tightly and requires less energy to remove.

Boron

Re: Ionization Energies Yes, Helium has the highest ionization energy! This is because the electrons in helium are very close to the nucleus and so the electrostatic attraction is very high. This makes it difficult to remove an electron.

Higher ionization energy means that it takes more energy to remove one electron from an atom. Ionization energy increases across a period. Going across a period, Effective Nuclear Charge (Zeff) increases. Distance and shielding remain constant.

Ionization energy (IE) is the energy required to remove the highest-energy electron from a neutral atom. In general, ionization energy increases across a period and decreases down a group. Across a period, effective nuclear charge increases as electron shielding remains constant.

Conversely, as one moves right along a period, the number of core electrons remains constant, while the Zeff increases. Electron pairing causes some mutual electron-electron repulsion, making these electrons more energetic, resulting in a drop in ionization energy for group VI by comparison to group V.

The atomic size, however, is larger for chlorine than it is for fluorine because chlorine has three energy levels (chlorine is in period 3)….Group and Period Trends in Ionization Energy.

The first ionization of Beryllium is greater than that of Boron because Beryllium has a stable complete electronic configuration (1s22s2) so it, require more energy to remove the first electron from it, where as Boron has electronic configuration (1s22s23s1) which need lesser energy than that of Beryllium.

Oxygen also has an unexpectedly low ionisation energy, less than that of nitrogen. This is due to an electron being added to an already half full orbital in oxygen, which results in electron electron repulsion, which will lower the ionisation energy.

Boron has a smaller value than beryllium, and oxygen has a smaller value than nitrogen. The first break occurs when the first electron is added to a p subshell. The repulsion between these two electrons makes one of them easier to remove, and so the ionization energy of oxygen is lower than might be expected.

The ionisation energy of Boron is less than that of Beryllium because in Boron there is a complete 2s orbital. The increased shielding of the 2s orbital reduces the ionisation energy. Similarly, the I.E. of Oxygen is less than that of Nitrogen because the extra electron is shielded by the half-filled 2p orbital.