Common-Ion Effect Definition

What Is the Common-Ion Effect?

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The common-ion effect impacts ionization of compounds that share a common cation or anion with each other. Mark Mawson, Getty Images

Common-ion effect describes the suppressing effect on ionization of an electrolyte ​when another electrolyte is added that shares a common ion.

How the Common-Ion Effect Works

A combination of salts in an aqueous solution will all ionize according to the solubility products, which are equilibrium constants describing a mixture of two phases. If the salts share a common cation or anion, both contribute to the concentration of the ion and need to be included in concentration calculations.

As one salt dissolves, it affects how well the other salt can dissolve, essentially making it less soluble. Le Chatelier's principle states equilibrium will shift to counter a change when more of a reactant is added.

Example of the Common-Ion Effect

For example, consider what happens when you dissolve lead(II) chloride in water and then add sodium chloride to the saturated solution.

Lead(II) chloride is slightly soluble in water, resulting in the following equilibrium:

PbCl2(s) ⇆ Pb2+(aq) + 2Cl-(aq)

The resulting solution contains twice as many chloride ions and lead ions. If you add sodium chloride to this solution, you have both lead(II) chloride and sodium chloride containing the chlorine anion. The sodium chloride ionizes into sodium and chloride ions:

NaCl(s) ⇆ Na+(aq) + Cl-(aq)

The additional chlorine anion from this reaction decreases the solubility of the lead(II) chloride (the common-ion effect), shifting the lead chloride reaction equilibrium to counteract the addition of chlorine.

The result is that some of the chloride is removed and made into lead(II) chloride.

The common-ion effect occurs whenever you have a sparingly soluble compound. The compound will become less soluble in any solution containing a common ion. While the lead chloride example featured a common anion, the same principle applies to a common cation.