Enthalpies of solution and hydration
Enthalpies of solution and hydration
- Dissolving an ionic solid involves two enthalpy steps.
- Hydration releases energy; pulling the lattice apart costs it.
- A cycle links them to the enthalpy of solution.
The two changes
- enthalpy of hydration $\Delta H_{\text{hyd}}$: change when one mole of gaseous ions is surrounded by water (exothermic).
- enthalpy of solution $\Delta H_{\text{sol}}$: change when one mole of solute fully dissolves.
Practice
The enthalpy change of hydration is:
Water molecules are attracted to the ions, releasing energy, so hydration is exothermic.
The dissolving cycle
To dissolve, you first pull the lattice apart, then hydrate the ions:
$$\Delta H_{\text{sol}} = -\Delta H_{\text{latt}} + \Delta H_{\text{hyd}}$$
- Like lattice energy, $\Delta H_{\text{hyd}}$ is more exothermic for ions of higher charge and smaller radius.
Practice
The enthalpy of solution is given by:
You reverse the lattice energy (to separate the ions) then add the hydration enthalpy.
Practice
The enthalpy of hydration is more exothermic for ions with:
Higher charge and smaller radius pull the water molecules in more strongly, releasing more energy.
You've got it
Key idea
- hydration $\Delta H_{\text{hyd}}$ (exothermic): gaseous ions surrounded by water
- solution $\Delta H_{\text{sol}}$: one mole of solute dissolves
- the cycle: $\Delta H_{\text{sol}} = -\Delta H_{\text{latt}} + \Delta H_{\text{hyd}}$
- $\Delta H_{\text{hyd}}$ is more exothermic for higher charge and smaller radius