5.1.5 Enthalpy of Solution & Hydration

Enthalpy of Solution ($ΔsolH⁰$)

The enthalpy of solution is the enthalpy change when 1 mole of an ionic substance dissolves in enough solvent, typically water, to form an infinitely dilute solution.

For example, the dissolution of sodium chloride can be represented as:

$$\text{NaCl(s)} \rightarrow \text{Na}^+(\text{aq}) + \text{Cl}^-(\text{aq})$$

Ionic compounds dissolve when their lattice structure breaks down and their ions become surrounded by solvent molecules.

Enthalpy of Solution in Terms of Lattice and Hydration Enthalpies

The enthalpy of solution can be calculated using:

$$\Delta H_{sol} = \Delta H_{lattice} + \Delta H_{hydration}$$

where:

• Enthalpy of Lattice Dissociation ($ΔLH⁰$): The energy needed to separate 1 mole of an ionic compound into its gaseous ions. This process is endothermic.
• Enthalpy of Hydration ($ΔhydH⁰$): The enthalpy change when 1 mole of gaseous ions dissolves in water to form hydrated ions. This process is exothermic as energy is released when water molecules surround and bond with the ions. Using these values, we can create a cycle to calculate the enthalpy of solution.

Why Ionic Compounds Dissolve in Water

Water is a polar molecule with partial positive and negative charges on the hydrogen and oxygen atoms, respectively. This polarity allows water to interact with ions in ionic compounds:

• The negative end of water molecules is attracted to positive ions.
• The positive end of water molecules is attracted to negative ions. These interactions, known as hydration, release energy and enable ionic compounds to dissolve by breaking up the ionic lattice and surrounding ions with water molecules.

By understanding the enthalpies of lattice dissociation and hydration, we can determine whether a compound will dissolve in water and how much energy is involved in the process.