Question

Give reasons for the following:

(a) Oxygen has lower ionisation energy than that of nitrogen.

(b) Electron gain enthalpy of chlorine is more negative than that of fluorine.

(c) Arrange $O^{2-}$, $A{l}^{3+}$, $N{a}^{+},F^{-}$ and $M{g}^{2+}$ in the increasing order of their sizes.

Answer 1

$\left(a\right)$ Oxygen has lower ionisation energy than that of nitrogen because the electronic configuration of nitrogen has half-filled stability where for oxygen it is partially filled, as shown above, It is easier to remove the electron from partially filled oxygen atom as in doing so it attains half-filled stability. In the case of nitrogen, it required more energy to remove an electron from half-filled or full filled orbitals.

$\left(b\right)$ The negative electron gain enthalpy of an electron is more than that of fluorine because of the small size of the fluorine atom. As a result, there is strong interelectronic repulsion in the relatively small $2p$- orbital of fluorine & thus, the incoming electron does not experience much net attract force from the nucleus to hold the incoming electron.

$\left(c\right)$ For isoelectronic species - High the positive charge smaller is the size and high the negative charge larger is the size.

$A{l}^{3+}<M{g}^{2+}<N{a}^{+}<F^{-}<O^{2-}$

Since they are isoelectronic species having $10$ electrons in each case.