Common Ion Effect
The common ion effect describes the decrease in solubility of an ionic compound when a soluble salt containing a common ion is added to the solution. It’s essentially an application of Le Chatelier’s principle to equilibrium reactions involving slightly soluble salts.
Explanation
When a sparingly soluble salt (like AgCl) is added to water, it dissolves to a small extent, establishing an equilibrium:
$ AgCl(s) \rightleftharpoons Ag^+(aq) + Cl^-(aq) $
This equilibrium is governed by the solubility product constant, Ksp:
$ K_{sp} = [Ag^+][Cl^-] $
Now, if we introduce a common ion (like Cl-) by adding a soluble salt (like NaCl) to the solution, the concentration of Cl- ions increases. According to Le Chatelier’s principle, the system will shift to counteract this change, meaning the equilibrium will shift to the left, favoring the formation of solid AgCl. This results in a decrease in the solubility of AgCl; less AgCl will dissolve in the presence of the common ion.
Example
Consider the solubility of silver chloride (AgCl) in pure water and in a 0.10 M NaCl solution. The Ksp of AgCl is 1.8 x 10-10.
1. Solubility in pure water:
Let ’s’ be the molar solubility of AgCl. $ K_{sp} = [Ag^+][Cl^-] = s * s = s^2 $ $ s = \sqrt{K_{sp}} = \sqrt{1.8 \times 10^{-10}} \approx 1.3 \times 10^{-5} M $
2. Solubility in 0.10 M NaCl:
Now, the initial concentration of Cl- is 0.10 M (from NaCl). Let ’s’ be the molar solubility of AgCl in this solution.
$ K_{sp} = [Ag^+][Cl^-] = s * (s + 0.10) $
Since ’s’ is very small compared to 0.10 (due to the very small Ksp), we can approximate (s + 0.10) ≈ 0.10
$ K_{sp} = s * 0.10 $ $ s = \frac{K_{sp}}{0.10} = \frac{1.8 \times 10^{-10}}{0.10} = 1.8 \times 10^{-9} M $
As you can see, the solubility of AgCl in the NaCl solution is significantly lower than in pure water due to the common ion effect.
Applications
The common ion effect has several practical applications:
- Qualitative Analysis: Qualitative Analysis Used to selectively precipitate certain ions from a solution.
- pH Control: Buffers Plays a role in buffer solutions, where a weak acid and its conjugate base (or a weak base and its conjugate acid) are present. The common ion suppresses the ionization of the weak acid or base, helping to maintain a stable pH.
- Solubility Control: Used to control the precipitation of salts in various chemical processes.
Key Takeaways
- The common ion effect decreases the solubility of a sparingly soluble salt.
- It’s a direct consequence of Le Chatelier’s principle.
- The common ion must be the same as one of the ions comprising the sparingly soluble salt.