Short Answer

The common ion effect occurs when a solution already contains an ion that is part of a sparingly soluble salt. This reduces the solubility of the salt because the equilibrium shifts according to Le Chatelier’s principle to counter the increased ion concentration.

For example, adding NaCl to a solution of AgCl reduces the solubility of AgCl because of the common Cl⁻ ion. The effect is important in predicting precipitation, solubility, and industrial chemical processes.

Detailed Explanation :

Definition of Common Ion Effect

The common ion effect is the decrease in solubility of a salt when a solution already contains one of the ions present in the salt. It is based on Le Chatelier’s principle, which states that a system at equilibrium responds to changes in concentration, pressure, or temperature by shifting the equilibrium to counteract the change.

1. How the Common Ion Effect Works

Consider a sparingly soluble salt, AgCl:

• The solubility product (Ksp) is given by:

• If the solution already contains Cl⁻ ions (e.g., from NaCl), the [Cl⁻] increases.
• To maintain Ksp constant, the concentration of Ag⁺ must decrease, meaning less AgCl dissolves.

Result: Solubility of AgCl decreases.

2. Factors Influencing the Effect

1. Concentration of Common Ion
   • Higher concentration of the common ion leads to a larger decrease in solubility.
2. Temperature
   • Solubility changes with temperature; the common ion effect is more noticeable at lower temperatures for many salts.
3. Nature of Salt
   • Strongly ionized salts with low Ksp show greater effect.

3. Examples of Common Ion Effect

1. Silver Chloride (AgCl)
   • AgCl solubility decreases in NaCl solution due to Cl⁻ ions.
2. Calcium Carbonate (CaCO₃)
   • Solubility decreases in the presence of Na₂CO₃, a source of CO₃²⁻ ions.
3. Lead Sulfate (PbSO₄)
   • Solubility decreases when Na₂SO₄ is added due to SO₄²⁻ ions.

4. Importance of the Common Ion Effect

1. Controlling Precipitation
   • Useful in qualitative analysis to selectively precipitate one salt.
2. Solubility Prediction
   • Helps in predicting how much of a salt will dissolve in a solution containing other ions.
3. Industrial Applications
   • Used in water treatment, purification of chemicals, and controlling unwanted precipitation.
4. Biological Systems
   • Maintains ion balance in biological fluids, e.g., calcium and phosphate solubility in blood.

5. Relation to Solubility Product (Ksp)

• Solubility product (Ksp) remains constant at a given temperature.
• Addition of a common ion increases one ion concentration.
• To keep Ksp unchanged, solubility of the other ion decreases.

Example: AgCl in water:

• If [Cl⁻] increases from 0.01 M, [Ag⁺] must decrease.
• Solubility is lower than in pure water.

6. Summary

• Common ion effect reduces solubility of salts in solutions containing one of the ions.
• It follows Le Chatelier’s principle.
• Important in predicting precipitation, controlling solubility, and industrial or laboratory applications.
• Examples: AgCl with NaCl, CaCO₃ with Na₂CO₃, PbSO₄ with Na₂SO₄.

Conclusion

The common ion effect is a key principle in chemistry that decreases the solubility of a salt in the presence of a solution containing a common ion. It is explained by Le Chatelier’s principle, where equilibrium shifts to reduce additional ion concentration. Understanding this effect is crucial in laboratory analysis, industrial processes, water treatment, and even biological systems, helping chemists control precipitation, solubility, and ion concentrations effectively.