Question

Knowing the electron gain enthalpy values for $ O \to {O^ - } $ and $ O \to {O^{2 - }} $ as $ - 141 $ and $ 702\,kJmo{l^{ - 1}} $ respectively, how can you account for the formation of a large number of oxides having $ {O^{2 - }} $ and not $ {O^ - } $ ?

Answer 1

Hint: A concept that comes into play in this question is lattice energy. The explanation for formation of oxides having $ {O^{2 - }} $ rather than $ {O^ - } $ can be done by comparing the lattice energy of compounds formed by both these ions.

Complete answer:
Electron gain enthalpy is the amount of energy released when a neutral, isolated gaseous atom accepts or gains an electron and forms the respective gaseous negative ion. In terms of oxygen, we are given that the electron gain enthalpy value for $ O \to {O^ - } $ is $ - 141\,kJ $ . This is the first electron gain enthalpy, which means it is the electron gain enthalpy when the first electron was accepted by the gaseous atom. It is negative. The electron gain enthalpy value of $ O \to {O^{2 - }} $ is $ 702\,kJmo{l^{ - 1}} $ , which is the second electron gain enthalpy and is positive. Knowing this, we understand that the formation of $ {O^{2 - }} $ is more difficult than the formation of $ {O^ - }. $ However, most oxides have $ {O^{2 - }} $ instead of $ {O^ - } $ and this can be explained in terms of lattice energy.
Lattice energy is the amount of energy that is required to dissociate one mole of an ionic compound into its constituent gaseous ions. Taking an example of a metal M, when M forms oxides with both $ {O^{2 - }} $ and $ {O^ - }, $ the lattice energy of $ {M^{2 + }}{O^{2 - }}\; $ is higher than those of $ {M^ + }{O^ - }. $ Therefore, we can account for the formation of a large number of oxides having $ {O^{2 - }} $ and not $ {O^ - }. $

Note:
The lattice energy of a compound relies on two factors: the charge carried by the constituent ions and the distance between the ions. The greater the charges carried by the constituent ions, the greater is the force of attraction between them. This implies that lattice energy is directly proportional to the charge carried by the constituent ions. On the other hand, lattice energy is inversely proportional to the distance between the ions.