Among the second period elements, the actual ionization enthalpies are in the order Li < B < Be < C < O < N < F < Ne. Explain the following: (i) Be has higher ΔiH than B. (ii) O has lower ΔiH than N and F.

(i) The electronic configuration of Be is 1s²2s². The electronic configuration of B is 1s²2s²2p¹. Be has higher ΔiH than B due to the following reasons: (a) The electronic configuration of Be has higher stability than the electronic configuration of B due to a completely filled 2s orbital. (b) During ionization of Be, an s electron is removed. During ionization of B, a p electron is removed. A 2s electron penetrates to the nucleus to a greater extent than a 2p electron. Thus, a 2p electron is more shielded than a 2s electron. The attraction of the nucleus for a 2s electron is higher than the attraction of the nucleus for a 2p electron. Thus, removal of a 2s electron requires higher energy than the removal of a 2p electron. Therefore, the ionization enthalpy of Be is higher than the ionization enthalpy of B. (ii) The electronic configuration of oxygen is 1s²2s²2p⁴. The 2p orbital contains 4 electrons, out of which 2 are present in the same 2p-orbital. Due to this, electron repulsion increases. N has a stable half-filled configuration. F has a greater nuclear charge. Hence, the ionization enthalpy of O is lower than that of N and F.