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Lattice energy is the energy released when gaseous ions come together to form a solid ionic compound. It is a measure of the strength of the ionic bonds in a compound. Ionic compounds form due to the strong electrostatic attraction between positively and negatively charged ions.
The lattice energy of \( \mathrm{MgO}(s) \) is so different from that of \( \mathrm{NaF}(s) \) because it depends on the charges of the ions and the sizes of the ions. The lattice energy is directly proportional to the charges of the ions and inversely proportional to the sizes of the ions. In the case of \( \mathrm{MgO}(s) \), both the magnesium and oxygen ions have a 2- charge, whereas in \( \mathrm{NaF}(s) \), the sodium ion has a 1+ charge and the fluoride ion has a 1- charge. The higher charges of the ions in \( \mathrm{MgO}(s) \) result in a much higher lattice energy compared to \( \mathrm{NaF}(s) \).
The formation of \( \mathrm{Mg}^{2+} \mathrm{O}^{2-} \) instead of \( \mathrm{Mg}^{+} \mathrm{O}^{-} \) can be explained by the need to achieve a stable octet configuration for both magnesium and oxygen. Magnesium typically forms ionic compounds by losing two electrons to achieve a stable octet configuration, resulting in the formation of \( \mathrm{Mg}^{2+} \) ions. Oxygen, on the other hand, gains two electrons to achieve a stable octet configuration, leading to the formation of \( \mathrm{O}^{2-} \) ions. This results in the formation of a stable ionic compound with a 1:1 ratio of magnesium and oxygen ions, represented as \( \mathrm{Mg}^{2+} \mathrm{O}^{2-} \).
As for why \( \mathrm{Mg}^{3+} \mathrm{O}^{3-} \) does not form, it is because the formation of such a compound would require an excessive transfer of electrons, leading to highly unfavorable energetics. The high charge on both the magnesium and oxygen ions would result in strong electrostatic repulsion, making the formation of \( \mathrm{Mg}^{3+} \mathrm{O}^{3-} \) energetically unfavorable and unlikely to occur under normal conditions.
