Revise Enthalpies of solution and hydration for Cambridge 9701 (23.2)

1. Defining the Key Enthalpy Changes

• Standard Enthalpy Change of Solution ($\Delta H_{\text{sol}}^{\ominus}$): This is the enthalpy change when one mole of a substance dissolves in a large excess of solvent to form a solution that is 'infinitely dilute'. For sodium chloride, the equation is: $\text{NaCl(s)} + \text{aq} \rightarrow \text{Na}^{+}(\text{aq}) + \text{Cl}^{-}(\text{aq})$. This value can be exothermic (negative) or endothermic (positive).

• Standard Enthalpy Change of Hydration ($\Delta H_{\text{hyd}}^{\ominus}$): This is the enthalpy change when one mole of isolated gaseous ions is dissolved in water to form hydrated aqueous ions. For a sodium ion, the equation is: $\text{Na}^{+}(\text{g}) + \text{aq} \rightarrow \text{Na}^{+}(\text{aq})$. This process is always exothermic because new, stable ion-dipole forces are formed between the ion and polar water molecules.

2. The Dissolving Process: An Enthalpy Cycle

3. Factors Affecting Enthalpy of Hydration

• Ionic Charge: The greater the charge on an ion, the stronger its attraction to the partial charges on water molecules ($\delta-$ on oxygen, $\delta+$ on hydrogen). Therefore, $\Delta H_{\text{hyd}}^{\ominus}$ becomes more exothermic (more negative). For example, $\Delta H_{\text{hyd}}^{\ominus}$ for Mg²⁺(g) is significantly more exothermic than for Na⁺(g).

• Ionic Radius: For ions with the same charge, a smaller ionic radius means the charge is concentrated in a smaller volume. This higher charge density leads to a stronger electric field and a more powerful attraction to water molecules, making $\Delta H_{\text{hyd}}^{\ominus}$ more exothermic. For example, $\Delta H_{\text{hyd}}^{\ominus}$ for Li⁺(g) is more exothermic than for K⁺(g).