16.1:
Common Ion Effect
Acetic acid, a weak acid, partially dissociates in solution to produce hydronium and acetate ions, while its salt, sodium acetate, dissociates completely to produce sodium ions and acetate. Both acetic acid and sodium acetate have the acetate ion in common.
When sodium acetate is added to an acetic acid solution, it increases the total concentration of acetate ions and disturbs the equilibrium. To counterbalance that change, the equilibrium shifts to the left and causes production of acetic acid until the equilibrium is reestablished.
In this case, the presence of the common ion results in the decreased dissociation of a compound. This phenomenon is known as the common ion effect.
The common ion effect can be explained with the help of Le Châtelier’s principle, which states that a change in the concentration of the reactants or products at equilibrium will cause the system to shift in a direction that counterbalances the change.
The pH of a 0.050 molar ammonia solution is 10.97. If 0.040 molar ammonium chloride is added into the solution, the new pH can be determined using the base dissociation constant of ammonia and an ICE table.
Ammonium chloride ionizes completely to produce 0.040 molar of both ammonium and chloride ions. As chloride ions are pH-neutral, they can be ignored.
Ammonia partially dissociates to produce ammonium and hydroxide ions. The Kb for this reaction is 1.76 × 10−5 and is equal to the concentration of ammonium times the concentration of hydroxide divided by the concentration of the ammonia.
The values for the initial, change, and equilibrium concentrations are placed in the ICE table, with changes in the concentration denoted by x.
Due to the small value of x, 0.050 minus x is approximately equal to 0.050, and 0.040 plus x is approximately equal to 0.040, which can be verified later by the 5% rule.
Substituting these values into the expression for Kb, x equals 2.2 × 10−5 molar.
The approximation is valid as the hydroxide concentration is less than 5% of 0.040 molar.
The pOH and pH of the solution can be calculated using the standard equations and equal 4.66 and 9.34, respectively.
Therefore, the presence of the common ion, ammonium ion, causes decreased dissociation of ammonia and thereby reduces the pH of the solution from 10.97 to 9.34.
16.1:
Common Ion Effect
Compared with pure water, the solubility of an ionic compound is less in aqueous solutions containing a common ion (one also produced by dissolution of the ionic compound). This is an example of a phenomenon known as the common ion effect, which is a consequence of the law of mass action that may be explained using Le Châtelier’s principle. Consider the dissolution of silver iodide:
This solubility equilibrium may be shifted left by the addition of either silver(I) or iodide ions, resulting in the precipitation of AgI and lowered concentrations of dissolved Ag+ and I–. In solutions that already contain either of these ions, less AgI may be dissolved than in solutions without these ions.
This effect may also be explained in terms of mass action as represented in the solubility product expression:
The mathematical product of silver(I) and iodide ion molarities is constant in an equilibrium mixture regardless of the source of the ions, and so an increase in one ion’s concentration must be balanced by a proportional decrease in the other.
Common Ion Effect on Solubility
The common ion affects the solubility of the compound in a solution. For example, solid Mg(OH)2 dissociate into Mg2+ and OH− ions as follows;
If MgCl2 is added to a saturated solution of Mg(OH)2, the reaction shifts to the left to relieve the stress produced by the additional Mg2+ ion, in accordance with Le Châtelier’s principle. In quantitative terms, the added Mg2+ causes the reaction quotient to be larger than the solubility product (Q > Ksp), and Mg(OH)2 forms until the reaction quotient again equals Ksp. At the new equilibrium, [OH–] is less and [Mg2+] is greater than in the solution of Mg(OH)2 in pure water.
If KOH is added to a saturated solution of Mg(OH)2, the reaction shifts to the left to relieve the stress of the additional OH– ion. Mg(OH)2 forms until the reaction quotient again equals Ksp. At the new equilibrium, [OH–] is greater and [Mg2+] is less than in the solution of Mg(OH)2 in pure water.
This text is adapted from Openstax, Chemistry 2e, Section 15.1: Precipitation and Dissolution.