We have already known that the solubility of ionic solids in water varies a great deal. Some of these (like calcium chloride) are so soluble that they are hygroscopic in nature and even absorb water vapour from atmosphere.
Others (such as lithium fluoride) have so little solubility that they are commonly termed as insoluble. The solubility depends on a number of factors important amongst which are the lattice enthalpy of the salt and the solvation enthalpy of the ions in a solution.
For a salt to dissolve in a solvent the strong forces of attraction between its ions (lattice enthalpy) must be overcome by the ion-solvent interactions.
The solvation enthalpy of ions is referred to in terms of solvation which is always negative i.e. energy is released in the process of solvation.
The amount of solvation enthalpy depends on the nature of the solvent. In case of a non-polar (covalent) solvent, solvation enthalpy is small and hence, not sufficient to overcome lattice enthalpy of the salt.
Consequently, the salt does not dissolve in non-polar solvent. As a general rule , for a salt to be able to
dissolve in a particular solvent its solvation enthalpy must be greater than its lattice enthalpy so that the latter may be overcome by former.
Each salt has its characteristic solubility which depends on temperature. We classify salts on the basis of their solubility in the following three categories.
Category I – Soluble - Solubility > 0.1M
Category II - Slightly Soluble - 0.01M<Solubility< 0.1M
Category III- Sparingly Soluble - Solubility < 0.01M
We shall now consider the equilibrium between the sparingly soluble ionic salt and its saturated aqueous solution.