Q.Magnesium oxide ($\text{MgO}$) has a much higher lattice enthalpy than sodium chloride ($\text{NaCl}$), even though both crystallise in the same rock-salt structure. Explain this difference in terms of ionic charge and ionic size.

An ionic (or electrovalent) bond forms when one atom transfers one or more valence electrons completely to another atom, generating a positively charged cation and a negatively charged anion that are then held together purely by electrostatic (Coulombic) attraction — unlike a covalent bond, no electron pair is shared between the two.

Ionic bonding is generally favoured between an electropositive element (typically a metal, on the left of the periodic table) and an electronegative element (typically a non-metal, on the right), because this arrangement lets both atoms move toward a stable, low-energy electron configuration — usually the eight-electron octet of the nearest noble gas. Three factors together determine whether ionic bond formation is energetically favourable: a low ionisation enthalpy for the atom losing electrons (so removing the electron costs little energy), a high (strongly negative) electron gain enthalpy for the atom gaining electrons (so the process releases a large amount of energy), and a high lattice enthalpy for the resulting solid (which supplies the energy needed to overcome ionisation and ultimately makes the overall process favourable).

The classic example is sodium chloride: a sodium atom, configuration $2,8,1$, loses its single outer electron to become $\text{Na}^+$ (configuration $2,8$, matching neon), while a chlorine atom, configuration $2,8,7$, gains that electron to become $\text{Cl}^-$ (configuration $2,8,8$, matching argon). Both resulting ions have achieved a complete, stable octet. In the solid state, this attraction is not confined to a single isolated ion pair; instead, each cation surrounds itself with several neighbouring anions and vice versa, extending indefinitely throughout the crystal to build a three-dimensional ionic lattice, such as the well-known rock-salt structure of $\text{NaCl}$, in which every $\text{Na}^+$ ion is surrounded by six $\text{Cl}^-$ neighbours and every $\text{Cl}^-$ ion by six $\text{Na}^+$ neighbours.

Lattice enthalpy quantifies the overall strength of this ionic attraction: it is defined as the energy released when one mole of the ionic solid forms from its constituent ions in the gaseous state (equivalently, the energy needed to completely separate one mole of the solid back into gaseous ions). Because the attraction between ions is fundamentally governed by Coulomb's law, lattice enthalpy depends directly on the product of the ionic charges and inversely on the sum of the ionic radii: $U \propto \dfrac{Q_+Q_-}{r_++r_-}$. This single relationship explains a great deal of periodic behaviour in ionic compounds — for instance, why doubly-charged ion pairs such as $\text{Mg}^{2+}/\text{O}^{2-}$ produce dramatically higher lattice enthalpies than singly-charged pairs such as $\text{Na}^+/\text{Cl}^-$, and why lattice enthalpy generally falls as ionic radius increases down a group. The complete energetics of forming an ionic solid — combining atomisation, ionisation, electron gain and lattice formation — is summarised in the Born–Haber cycle, a practical tool for checking these individual enthalpy terms against the experimentally measurable enthalpy of formation.