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Equilibria of Other Reaction Classes

Shifting the Solubility Equilibrium

Solubility can be shifted by the common ion effect, which is a method of forcing a solute to precipitation out of a solution by adding a salt that has an ion in common with the solute. Solubility can also be shifted by a change in the acidity of a solution.
Because a saturated solution is in equilibrium, the equilibrium can be shifted left or right according to Le Chatelier's principle, which states that a change in the temperature, pressure, or concentration of a component will cause the equilibrium condition of a chemical system to change in a way that reduces the change. This means that increasing the concentration of one or more of the dissolved ions on the right side of the equation shifts the equilibrium to the left and causes the solute to form a solid precipitate. For example, in a saturated solution of silver iodide, AgI:
AgI(s)Ag+(aq)+I(aq),Ksp=8.52×1017{\rm{AgI}}(s)\rightleftharpoons{\rm{Ag}^+}(aq)+{\rm I^-}(aq),\hspace{5pt}\;{{K}_{\rm{sp}}=8.52\times10^{-17}}
The amount of solid AgI can be increased by the addition of aqueous iodide ions to the solution. Initially, this would increase the overall concentration of I, making it greater than the molar solubility of AgI. Because the equilibrium requires that [I] be less than the molar solubility of AgI, the increased concentration of aqueous iodide ions would shift the equilibrium to the left. Ions would fall out of solution as AgI solid precipitate until [I] was again equal to molar solubility of AgI and the equation returned to equilibrium. One way to add [I] is via an aqueous solution made from a salt that shares an ion with AgI, such as potassium iodide, KI. Adding this to the original AgI solution would therefore cause the precipitation of solid AgI.

Common Ion Effect

A salt that shares an ion with a slightly soluble compound lowers the solubility of that compound.
This example illustrates the common ion effect, the effect of adding to a solution a salt that has an ion in common with a solute, causing the equilibrium to shift, lowering the solubility of the original solute so that it precipitates out of solution.

The pH of an aqueous solution can also change the solute's solubility. A lower pH increases the solubility of a soluble base and salts of bases. Inversely, a higher pH increases the solubility of an acid. Suppose, for example, Zn(OH)2 is added to an aqueous solution of neutral pH.
Zn(OH)2(s)Zn2+(aq)+2OH(aq){\rm{Zn(OH)}_2}(s)\rightleftharpoons {\rm{Zn}^{2+}}(aq)+2{\rm{OH}^-}(aq)
Adding an acid will increase the solubility of the Zn(OH)2 by removing the OH ions, shifting the equilibrium of the original reaction to the right.
OH(aq)+H3O+(aq)2H2O(l){\rm{OH}^-}(aq)+{\rm{H}_3\rm{O}^+}(aq) \leftrightharpoons{2\rm{H}_2\rm{O}}(l)