What is the lattice energy trend across the periodic table?

Lattice energy increases as you move across a period in the periodic table, particularly when comparing ionic compounds of metals from successive groups. Lattice energy is defined as the energy required to completely separate one mole of an ionic solid into gaseous ions, essentially the energy needed to break apart the ionic lattice:

:::center $\ce{M_aX_b}\textrm{(s)} \longrightarrow a \ce{\,M^{b+}}\textrm{(g)} + b\ce{\,X^{a-}}\textrm{(g)}$ :::

This value is always positive because energy must be supplied to overcome the electrostatic attractions holding the ions together. The magnitude of lattice energy directly reflects the strength of the ionic bonding in the compound; stronger ionic bonds require more energy to break, resulting in higher lattice energy values.

This increasing trend across a period occurs because the charge on the cations increases as you move from left to right across the metal groups. For example, comparing oxides across Period 3: $\ce{Na2O}$ contains $\ce{Na+}$ cations, MgO contains $\ce{Mg^{2+}}$ cations, and $\ce{Al2O3}$ contains $\ce{Al^{3+}}$ cations. According to Coulomb's law, the electrostatic attraction between ions is directly proportional to the product of their charges and inversely proportional to the distance between them. As the cation charge increases from +1 to +2 to +3, the electrostatic attraction to the oxide anion $(\ce{O^{2-}})$ becomes progressively stronger. Additionally, moving across a period, the cations generally become smaller due to increasing nuclear charge pulling the electrons closer, which further increases the lattice energy since the ions can pack more closely together. This combination of higher charge and smaller ionic radius results in much stronger coulombic attractions and therefore significantly higher lattice energies as you progress across the periodic table.