Common-ion effect

The common ion effect is responsible for the reduction in the solubility of an ionic precipitate when a soluble compound containing one of the ions of the precipitate is added to the solution in equilibrium with the precipitate. It states that if the concentration of any one of the ions is increased, then, according to Le Chatelier's principle, some of the ions in excess should be removed from solution, by combining with the oppositely charged ions. Some of the salt will be precipitated until the ion product is equal to the solubility product. In short, the common ion effect is the suppression of the degree of dissociation of a weak electrolyte containing a common ion.[1]

Solubility effects

The solubility of a sparingly soluble salt is reduced in a solution that contains an ion in common with that salt. For instance, the solubility of silver chloride in water is reduced if a solution of sodium chloride is added to a suspension of silver chloride in water.[2]

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 Na+ interfere with the normal behavior of soap because of common ion effect. In the presence of excess sodium ions 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.[3] 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 hydrogen ion 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 equilibra involved in simple precipitation of salt 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 effect 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.[4] This alternate equilibrium makes the ions less available for the precipitation reaction.

See also

References

  1. ↑ Mannam Krishnamurthy; Subba Rao Naidu (2012). "8". In Lokeswara Gupta. Chemistry for ISEET - Volume 1, Part A (2012 ed.). Hyderabad, India: Varsity Education Management Limited. p. 298.
  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. 39, ISBN 0-582-22628-7
  3. ↑ 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
  4. ↑ Claude E. Boyd (14 July 2015). Water Quality: An Introduction. Springer. pp. 56–. ISBN 978-3-319-17446-4.


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