Solubility Product

the product of the ion concentrations in a saturated solution of a slightly soluble strong electrolyte. The exponents of the ion concentrations in the solubility product are equal to the corresponding ion coefficients in the equation of electrolyte dissociation. For nonideal solutions, the concentrations are replaced by activities, and the product thus obtained is called the activity product. There is a characteristic value of the solubility product for each electrolyte at a given temperature in a given solvent.

The constancy of the solubility product is derived from the law of mass action. It represents a specific application of this law to the case of the equilibrium between an electrolyte in the solid phase and a saturated solution of the same electrolyte. It is assumed that the electrolyte in the solution is fully dissociated. The solubility product is most accurately measured by the electromotive force (emf) method; another common method of measurement derives from the determination of solubility by the electrical conductivity of saturated solutions. The solubility product for many compounds is established with sufficient accuracy for all practical purposes. In tables, the solubility product is usually given for a temperature of 25°C (sometimes for 18°C).

Since the solubility product is a constant, it follows that if the product of the ion concentrations in a solution exceeds the solubility product’s value, precipitation will occur. If the opposite is true, no precipitate is formed. This consequence permits a control of the ion content in a solution by means of precipitation, dissolving, and salting-out—all key processes in analytical chemistry and chemical technology. When, for example, the concentration of one ion is increased by introducing a new electrolyte with the same type of cation or anión into the solution, the concentration of the other ion is decreased by the precipitation of part of the poorly soluble electrolyte. The solubility decreases only to a certain minimum value, after which an increase in solubility may be observed owing to either the formation of complex ions or the increase in the ionic strength of the solution.

An increase in solubility may be achieved by binding one ion in the solution so as to form another ion that does not yield a slightly soluble compound. For example, in order to dissolve a CaCO₃ precipitate, the ion is bound with the aid of a H⁺ ion into a weakly dissociated ion: . The concentration of ions is thereby decreased, and the precipitate will dissolve until the solubility product is achieved.