The formation of the oxide ion O^2– (g), from the oxygen atom requires first an exothermic and then an endothermic step as shown below, 
$\mathrm{O}\left(\mathrm{g}\right)+{\mathrm{e}}^{-}\to {\mathrm{O}}^{-};$ ${∆}_{\mathrm{f}}{\mathrm{H}}^0=–141$ $\mathrm{kJ}$ ${\mathrm{mol}}^{-1}$
${\mathrm{O}}^{-}\left(\mathrm{g}\right)+{\mathrm{e}}^{-}\to {\mathrm{O}}^{2-}\left(\mathrm{g}\right);$ ${∆}_{\mathrm{f}}{\mathrm{H}}^0=+780$ $\mathrm{kJ}$ ${\mathrm{mol}}^{-1}$

Thus, the process of formation of O^2– in the gas phase is unfavorable even though O^2– is isoelectronic with neon. It is due to the fact that:

The species Ar, K⁺ and Ca²⁺ contain the same number of electrons. In which order do their radii increase?
1.  Ar < K⁺ < Ca²⁺
2. Ca²⁺ < Ar < K⁺
3. Ca²⁺ < K⁺ < Ar
4. K⁺ < Ar < Ca²⁺

The correct order of ionic radii is:

1. $H^{-}>H^{+}>H$ $$
2. $N{a}^{+}>F^{-}>O^{2-}$
3. $F^{-}>O^{2-}>N{a}^{+}$
4. $N^{3-}>M{g}^{2+}>A{l}^{3+}$

\(Be^{2+}\) is isoelectronic with which of the following ions?
1. $H^{+}$
2. $L{i}^{+}$
3. $N{a}^{+}$
4. $M{g}^{2+}$

The value of electron gain enthalpy of Na⁺, if IE₁ of Na = 5.1 eV, is:

1. +10.2 eV 

2. –5.1 eV

3. –10.2 eV 

4. +2.55 eV

The correct order of the decreasing ionic radii among the following isoelectronic species is:

1. $C{a}^{2+}>K^{+}>S^{2-}>C{l}^{-}$

2. $C{l}^{-}>S^{2-}>C{a}^{2+}>K^{+}$

3. $S^{2-}>C{l}^{-}>K^{+}>C{a}^{2+}$

4. $K^{+}>C{a}^{2+}>C{l}^{-}>S^{2-}$

The correct order of increasing electron affinity for the elements, O, S, F and Cl is:

Among the elements Ca, Mg, P and Cl, the correct order of increasing atomic radii is:
1. Cl < P < Mg < Ca
2. P < Cl < Ca < Mg
3. Ca < Mg < P < Cl
4. Mg < Ca < Cl < P

Amongst the elements with the following electronic configurations, which one of them may have the highest ionisation energy?

1. $\left[Ne\right]$ $3s^2$ $3p^3$

2. $\left[Ne\right]$ $3s^2$ $3p^2$

3. $\left[Ar\right]$ $3d^{10}$ $4s^2$ $4p^3$

4. $\left[Ne\right]$ $3s^2$ $3p^1$