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Understanding electron configuration helps explain why MgO consists of \(\mathrm{Mg}^{2+}\) and \(\mathrm{O}^{2-}\). Magnesium, in Group 2 of the periodic table, naturally loses two electrons to achieve a stable configuration similar to a noble gas. Its electron configuration is originally \(1s^2 2s^2 2p^6 3s^2\). By losing two electrons, magnesium achieves \(\mathrm{Mg}^{2+}\) with a configuration of \(1s^2 2s^2 2p^6\), resembling the electron arrangement of neon.

Oxygen, on the other hand, is in Group 16 and needs to gain two electrons to fill its valence shell, transitioning from \(1s^2 2s^2 2p^4\) to \(1s^2 2s^2 2p^6\), mimicking the electron configuration of neon as well. This results in the formation of \(\mathrm{O}^{2-}\).

Both ions achieve a stable and lower energy state by completing their outer electron shells. This stability, provided by achieving a noble gas configuration, strongly supports the existence of \(\mathrm{Mg}^{2+}\) and \(\mathrm{O}^{2-}\) ions in magnesium oxide.

Lattice energy plays a crucial role in explaining the stability of ionic compounds like MgO. Lattice energy is the energy released when ions bond to form a crystalline lattice. The larger the charge and the smaller the ion, the higher the lattice energy. In the case of MgO, \(\mathrm{Mg}^{2+}\) and \(\mathrm{O}^{2-}\) ions have high charge densities.

The electrostatic attraction between these oppositely charged ions is strong, resulting in a high lattice energy. This strong attraction and high lattice energy contribute significantly to the stability of MgO, which is more stable compared to a system of \(\mathrm{Mg}^{+}\) and \(\mathrm{O}^{-}\) ions. Lower charged ions would generate less electrostatic attraction and thus, lower lattice energy, resulting in less stability.

Therefore, considering lattice energy, \(\mathrm{Mg}^{2+}\) and \(\mathrm{O}^{2-}\) ions are the favored configuration, serving as evidence for this composition in magnesium oxide.

The conductivity of molten MgO provides experimental evidence for the presence of \(\mathrm{Mg}^{2+}\) and \(\mathrm{O}^{2-}\) ions. When MgO is in its solid form, ions are locked in place within a crystalline lattice structure and cannot move freely, so the compound is a poor conductor of electricity.

However, when melted, these ions gain mobility. In the molten state, \(\mathrm{Mg}^{2+}\) and \(\mathrm{O}^{2-}\) ions can move freely and carry electrical current. This movement of charged particles is necessary for conductivity. The ability to conduct electricity when molten strongly supports that MgO is composed of highly charged ions, confirming the \(\mathrm{Mg}^{2+}\) and \(\mathrm{O}^{2-}\) ions formation.

Thus, molten conductivity experiments validate the ionic nature of magnesium oxide, reaffirming the arrangement of ions with high charges.

X-ray diffraction offers another experimental method to support the \(\mathrm{Mg}^{2+}\) and \(\mathrm{O}^{2-}\) ionic model of MgO. This technique involves directing X-rays onto the crystal and studying the diffraction pattern produced.

The diffraction pattern reflects the internal structure and arrangement of atoms or ions within the crystal lattice. A tightly packed lattice with high charge ions like \(\mathrm{Mg}^{2+}\) and \(\mathrm{O}^{2-}\) would show a specific diffraction pattern consistent with the high ion charges and increased lattice energy.

Thus, analyzing these patterns provides insight into the ionic sizes and charges, indirectly confirming the presence of \(\mathrm{Mg}^{2+}\) and \(\mathrm{O}^{2-}\) ions rather than potential alternatives like \(\mathrm{Mg}^{+}\) and \(\mathrm{O}^{-}\). Therefore, X-ray diffraction further proves the accurate formulation of magnesium oxide.