Solubility Equilibria


  • The solubility of ionic solids in water varies greatly depending on factors such as 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, which is determined by the solvation enthalpy of the ions in the solvent.
  • The solvation enthalpy of ions depends on the nature of the solvent. In 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 solvents.
  • For a salt to dissolve in a particular solvent, its solvation enthalpy must be greater than its lattice enthalpy so that the latter may be overcome by the former.
  • Salts can be classified into three categories based on their solubility in water:

·         Soluble salts: these salts dissolve readily in water to form a clear and homogeneous solution. Examples include sodium chloride, potassium nitrate, and calcium chloride.

·         Insoluble salts: these salts have very low solubility in water and do not form clear solutions. Examples include lead (II) chloride, silver chloride, and calcium sulfate.

·         Sparingly soluble salts: these salts have solubility between that of soluble and insoluble salts. Examples include calcium carbonate, barium sulfate, and lead (II) sulfate.


Solubility Product Constant

When we try to dissolve a solid into water, if it dissolves, there are three possibilities:

1. The solid is a non-electrolyte and it dissolves as neutral molecules.

2. The solid is a highly soluble electrolyte; it dissolves almost completely.

3. The solid is a sparingly soluble electrolyte; it dissolves to a limited extent.

It is the third possibility that interests us here. Let us take the example of dissolution of AgCl to understand the equilibria in such cases. When silver chloride is added to water, the following equilibrium is established.

This is an example of a heterogeneous equilibrium because it involves both a solid and a solution. This equilibrium is known as the solubility equilibrium for which the equilibrium constant expression is

As a matter of convention the concentration of the undissolved solid is taken as one. We can rewrite the equilibrium as

The equilibrium constant now is the product of the concentrations of the ions. It is called solubility product constant or simply solubility product. A new symbol, Ksp, has been assigned to this constant. The mass expression on the right , is called, ion product or ionic product. The solubility product constant of a given salt is constant at a given temperature.

 Relationship between Solubility and Solubility Product Constant

The solubility product constant for a substance is related to its solubility. The nature of relationship depends on the nature of the salt.

Common Ion Effect on Solubility of ionic salts

  • Le Chatelier's principle predicts that if the concentration of any ion in a solution of an ionic salt is increased, some of the salt will precipitate until the solubility product constant (Ksp) equals the ion product (Qsp).
  • Conversely, if the concentration of any ion in a solution of an ionic salt is decreased, more of the salt will dissolve until Ksp equals Qsp.
  • Le Chatelier's principle applies to soluble salts such as sodium chloride, and in such cases, the activities of ions are used instead of their molarities in the expression for Qsp.
  • Passing HCl gas through a saturated solution of sodium chloride can cause the precipitation of sodium chloride due to an increased concentration (activity) of chloride ion available from the dissociation of HCl.
  • The common ion effect is used for the almost complete precipitation of a particular ion as its sparingly soluble salt with a very low value of solubility product for gravimetric estimation.
  • Common ion effect can be used to precipitate silver ion as silver chloride, ferric ion as its hydroxide (or hydrated ferric oxide), and barium ion as its sulfate for quantitative estimations.


 The solubility of salts of weak acids like phosphates increases at lower pH. This is because at lower pH the concentration of the anion decreases due to its protonation. This in turn increase the solubility of the salt so that Ksp = Qsp. We have to satisfy two equilibria simultaneously i.e.,

Now, again taking inverse, we get

[X] / {[X] + [HX]} = f = Ka / (Ka + [H+]) and it can be seen that ‘f’ decreases as pH decreases. If S is the solubility of the salt at a given pH then

Ksp = [S] [f S] = S2 {Ka / (Ka + [H+])} and

S = {Ksp ([H+] + Ka ) / Ka }1/2 (7.46)

Thus solubility S increases with increase in [H+] or decrease in pH.