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The common-ion effect states that in a chemical solution in which several species reversibly associate with each other by an equilibrium process, increasing the concentration of any one of its dissociated components by adding another chemical that al2so contains it will cause an increased amount of association. This result is a consequence of Le Chatelier's principle for the equilibrium reaction of the association/dissociation. The effect is commonly seen as an effect on the solubility of salts and other weak electrolytes. Adding an additional amount of one of the ions of the salt generally leads to increased precipitation of the salt, which reduces the concentration of both ions of the salt until the solubility equilibrium is reached. The effect is based on the fact that both the original salt and the other added chemical have one ion in common with each other.

Solubility effects

A practical example used very widely in areas drawing drinking water from chalk or limestone aquifers is the addition of sodium carbonate to the raw water to reduce the hardness of the water. In the water treatment process, highly soluble sodium carbonate salt is added to precipitate out sparingly soluble calcium carbonate. The very pure and finely divided precipitate of calcium carbonate that is generated is a valuable by-product used in the manufacture of toothpaste.

The salting-out process used in the manufacture of soaps benefits from the common-ion effect. Soaps are sodium salts of fatty acids. Addition of sodium chloride reduces the solubility of the soap salts. The soaps precipitate due to a combination of common-ion effect and increased ionic strength.

Sea, brackish and other waters that contain appreciable amount of sodium ions (Na) interfere with the normal behavior of soap because of common-ion effect. In the presence of excess Na, the solubility of soap salts is reduced, making the soap less effective.

Buffering effect

A buffer solution contains an acid and its conjugate base or a base and its conjugate acid. Addition of the conjugate ion will result in a change of pH of the buffer solution. For example, if both sodium acetate and acetic acid are dissolved in the same solution they both dissociate and ionize to produce acetate ions. Sodium acetate is a strong electrolyte, so it dissociates completely in solution. Acetic acid is a weak acid, so it only ionizes slightly. According to Le Chatelier's principle, the addition of acetate ions from sodium acetate will suppress the ionization of acetic acid and shift its equilibrium to the left. Thus the percent dissociation of the acetic acid will decrease, and the pH of the solution will increase. The ionization of an acid or a base is limited by the presence of its conjugate base or acid.

NaCH₃CO₂(s) → Na(aq) + CH₃CO₂(aq)

CH₃CO₂H(aq) ⇌ H(aq) + CH₃CO₂(aq)

This will decrease the hydronium concentration, and thus the common-ion solution will be less acidic than a solution containing only acetic acid.

Exceptions

Many transition-metal compounds violate this rule due to the formation of complex ions, a scenario not part of the equilibria that are involved in simple precipitation of salts from ionic solution. For example, copper(I) chloride is insoluble in water, but it dissolves when chloride ions are added, such as when hydrochloric acid is added. This is due to the formation of soluble CuCl₂ complex ions.

Uncommon-ion effect

Sometimes adding an ion other than the ones that are part of the precipitated salt itself can increase the solubility of the salt. This "salting in" is called the "uncommon-ion effect" (also "salt effect" or the "diverse-ion effect"). It occurs because as the total ion concentration increases, inter-ion attraction within the solution can become an important factor. This alternate equilibrium makes the ions less available for the precipitation reaction. This is also called odd ion effect.

See also

• Chelate effect

References

1. IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version: (2006–) "common-ion effect (on rates)". doi:10.1351/goldbook.Error in template * unknown parameter name (GoldBookRef): "page"
2. Mendham, J.; Denney, R. C.; Barnes, J. D.; Thomas, M. J. K. (2000), Vogel's Quantitative Chemical Analysis (6th ed.), New York: Prentice Hall, p. 28, ISBN 0-582-22628-7
3. Claude E. Boyd (14 July 2015). Water Quality: An Introduction. Springer. pp. 56–. ISBN 978-3-319-17446-4.

• Equilibrium chemistry
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