Common Ion EffectJEE Main

Add the salt, silence the acid — interactive Chemistry simulation for IIT-JEE.

Concept

Add sodium acetate to acetic acid: the extra acetate (common ion) pushes HA ⇌ H⁺ + A⁻ backwards (Le Chatelier), crushing the acid's dissociation and raising the pH. The equilibrium constant never changes — only the position does. This suppressed, poised state is exactly what a buffer is.

Key formula

[H^+] = K_a\frac{[HA]}{[A^-]} \;\Rightarrow\; pH = pK_a + \log\frac{[salt]}{[acid]}

Derivation

With salt concentration S dominating A⁻: $K_a = \frac{[H^+]\,S}{C_{acid}}$ , so $[H^+] = K_aC/S$ — linear suppression by S.

Taking −log gives Henderson–Hasselbalch. Note α collapses from √(Ka/C) to Ka·C/(C·S) ≈ Ka/S — often a 100× drop for equal concentrations.

Scenarios to explore

Common Ion Effect — Add the salt, crush the dissociation — buffers explained.

Real-world applications

Buffers (blood pH 7.4 via H₂CO₃/HCO₃⁻).
Salting-out soap with NaCl.
Qualitative analysis: controlling S²⁻ with HCl in group separation.

JEE exam tips

With common ion: [H⁺] = Ka·(acid/salt) — skip the quadratic.
Equal acid & salt ⇒ pH = pKa — the buffer midpoint.
Solubility version: s' = Ksp/[common ion] for a 1:1 salt.

Common mistakes

Thinking Ka changes — it's constant; only α shifts.
Forgetting the effect works on solubility too (AgCl in NaCl solution).
Using √(KaC) with a common ion present.

Exam traps to avoid

Adding NaCl to HCl does nothing acidic (both strong) — common ion needs a WEAK equilibrium to suppress.
The common ion must actually be common — Na⁺ is a spectator.