Describe, in terms of bond breaking and bond making, what occurs to the solid ionic lattice when an ionic compound dissolves in water.
What does the term enthalpy change of solution, $\Delta H_{sol}$, mean?
Use the data below to calculate the standard enthalpy change of hydration, $\Delta H^{\circ}_{hyd}$, of chloride ions, $\text{Cl}^-(g)$. It may help to draw an energy cycle. Data given: $\Delta H^{\circ}_{hyd}(\text{Mg}^{2+}(g)) = -1925\ \text{kJ mol}^{-1}$; lattice energy of $\text{MgCl}_2(s) = -2524\ \text{kJ mol}^{-1}$; enthalpy change of solution for $\text{MgCl}_2(s) = -155\ \text{kJ mol}^{-1}$. Hence determine $\Delta H^{\circ}_{hyd}(\text{Cl}^-(g))$.
The enthalpy change of hydration for $\text{Na}^+$, $\Delta H^{\circ}_{hyd}(\text{Na}^+(g))$, is $-410\ \text{kJ mol}^{-1}$. Suggest an explanation for why the $\Delta H^{\circ}_{hyd}$ of the $\text{Na}^+$ ion is less exothermic than the $\Delta H^{\circ}_{hyd}$ of the $\text{Mg}^{2+}$ ion.
Describe and explain how the solubility of the Group II sulfates changes down the group.