Common ion effect (as applicable) - Chemistry AP Study Notes

The common ion effect describes the shift in equilibrium position when an ion that is already present in an equilibrium system is added from an external source. This phenomenon is a direct application of Le Chatelier's Principle, which states that a system at equilibrium will respond to counteract any imposed change.

Key Definition (Cambridge Standard): The common ion effect is the suppression of the ionization of a weak acid or weak base by the addition of a common ion, typically from a soluble salt.

Fundamental Equation: For a weak acid HA: HA(aq) ⇌ H⁺(aq) + A⁻(aq)

If we add a salt containing A⁻ (such as NaA), the equilibrium shifts left according to Le Chatelier's Principle, reducing the ionization of HA and decreasing [H⁺].

Solubility Equilibrium Application: For a sparingly soluble salt like AgCl: AgCl(s) ⇌ Ag⁺(aq) + Cl⁻(aq)

The solubility product constant Kₛₚ = [Ag⁺][Cl⁻] remains constant at a given temperature. Adding NaCl (common Cl⁻ ion) forces the equilibrium left, reducing [Ag⁺] and thus the solubility of AgCl.

Memory Aid - "CLEF":

• Common ion added
• Le Chatelier applies
• Equilibrium shifts away
• Forward reaction suppressed

Critical Understanding: The common ion effect always decreases the extent of ionization or dissolution of the original compound. The equilibrium constant (Ka, Kb, or Kₛₚ) remains unchanged; only the equilibrium position shifts.

Real-World Application 1: Blood pH Regulation Human blood maintains pH 7.35-7.45 using the carbonate buffer system: HCO₃⁻(aq) + H⁺(aq) ⇌ H₂CO₃(aq). When you hyperventilate, CO₂ is expelled, reducing H₂CO₃ and shifting equilibrium right, consuming H⁺ and raising pH (respiratory alkalosis). The kidneys compensate by adding HCO₃⁻ (common ion), which shifts equilibrium left, suppressing further H⁺ consumption—a biological example of the common ion effect stabilizing pH.

Real-World Application 2: Water Softening Hard water contains Ca²⁺ ions from dissolved CaCO₃. Adding sodium carbonate (Na₂CO₃) introduces CO₃²⁻ as a common ion to the equilibrium: CaCO₃(s) ⇌ Ca²⁺(aq) + CO₃²⁻(aq). The increased [CO₃²⁻] drives equilibrium left, precipitating CaCO₃ and removing Ca²⁺ from solution—softening the water.

Analogy - The Crowded Car Park: Imagine a reversible equilibrium like cars entering/leaving a car park at equal rates (equilibrium). If external cars (common ions) suddenly flood one section, fewer spaces exist for original cars to park there. The system adjusts by having more original cars stay outside (shift left) or move to other areas. The total parking capacity (equilibrium constant) hasn't changed, but the distribution (position) has.

Industrial Example: Soap Manufacture Soap production involves saponification followed by "salting out." Adding concentrated NaCl increases Na⁺ (common ion) concentration, reducing soap (sodium stearate) solubility via the common ion effect, causing soap to precipitate for easy collection.