Among the second period elements the actual ionisation enthalpies are in the order `LiltBltBeltCltOltNltFltNe`.
Explain why (a) `Be` has higher `Delta_(i)H` than `B` and (b) `O` has lower `Delta_(i)H` than `N` and `F`?

Explanation

(i) Be has higer `Delta _(i)H` than B because Be has the elctronic configuration `1s ^(2)2s ^(2)` while B has the electronic configuration `1s ^(2)2s ^(2)2p ^(1).` The larger `Delta _(i)H` of Be in comparision to B is due to the fact that
The electronic configuration of Be is more stable (completely filled 2s orbitals) than that of B.
In Be, the electron to be removed during the ionization is an s-electron while the electron to be removed during ionization of B is a p-electron. The penetration of a 2s-electron to the nucleus is more than that of a 2p-electron and hence 2p electron of B is more shielded from the nucleus by the inner core of electrons than the 2s-electron of Be. As result, 2s-electron is attracted to the nucleus more than 2p-electron. Therefore, it is difficult to remove a 2s-electron from Be than to remove the 2p-electron from B. Thus, Be has higer ionization enthalpy than B.
Oxygen has four electrons in 2p orbitals and two of the four 2p electrons must occupy the same 2p-orbital resulting in increased electron-electron repulsion. On the other hand. N has stable half filled configuration, while F has greater nuclear charge. Therefore, O has ionization enthalpy less than N as well as F.