Imagine you have a glass of water with some salt dissolved in it — say, sodium chloride (NaCl). The salt has dissociated into Na⁺ and Cl⁻ ions floating around. Now, if you add more salt, some of it will dissolve, but eventually the water becomes saturated and no more salt dissolves.
Now imagine a different scenario. You have a solution of silver chloride (AgCl) — a sparingly soluble salt. Very little of it dissolves, giving you a tiny concentration of Ag⁺ and Cl⁻ ions. What happens if you now add some sodium chloride (NaCl) to this solution? The NaCl will dissociate completely, flooding the solution with extra Cl⁻ ions.
Here's the key: the system tries to maintain its equilibrium. The solubility equilibrium for AgCl is:
AgCl(s)⇌Ag+(aq)+Cl−(aq)
When you add extra Cl⁻ from NaCl, Le Chatelier's principle kicks in. The equilibrium shifts to the left — more AgCl precipitates out of solution. The presence of a common ion (Cl⁻) suppresses the solubility of AgCl.
That's the common ion effect in a nutshell: the solubility of a salt decreases when you add another salt that shares a common ion with it.
The Precise Statement
Important
Common Ion Effect: The suppression of the dissociation of a weak electrolyte (or the solubility of a sparingly soluble salt) by the addition of a strong electrolyte that provides an ion common to the equilibrium system.
In other words: when you have an equilibrium involving ions, adding more of one of those ions (from a different source) shifts the equilibrium away from the dissociated form.
Why It Matters (and Where You'll See It)
The common ion effect isn't just a textbook curiosity — it's used everywhere in chemistry:
Controlling pH of buffer solutions: Adding a common ion (like acetate ion to acetic acid) suppresses the dissociation of the weak acid, keeping the pH stable.
Qualitative analysis: In salt analysis, you selectively precipitate certain ions by adding a common ion. For example, to test for chloride, you add AgNO₃ — the Ag⁺ is common to AgCl, so even tiny amounts of Cl⁻ will precipitate.
Industrial processes: In the Solvay process for making sodium carbonate, the common ion effect is used to precipitate sodium bicarbonate.
A Concrete Example with Numbers
Consider the solubility of silver chloride in pure water. The Ksp of AgCl is 1.8×10−10.
In pure water:
Ksp=[Ag+][Cl−]=s2=1.8×10−10
s=1.8×10−10=1.34×10−5 M
Now, what if the solution already contains 0.10 M NaCl (from a separate source)? The Cl⁻ concentration is now 0.10 M (plus a tiny bit from AgCl). Let s be the new solubility of AgCl: …
Common-ion effect = suppression of a weak electrolyte's ionisation by adding a strong electrolyte with a common ion (Le Chatelier); e.g. CH3COONa lowers the ionisation of CH3COOH. (OR - the law of mass action relates reaction rate to the product of the active masses of the reactants.)
Common-ion effect: When a strong electrolyte having an ion in common with a weak electrolyte is added to the weak electrolyte's solution, the ionisation (dissociation) of the weak electrolyte is suppressed. By Le Chatelier's principle the added common ion pushes the ionisation equilibrium backward.
Example: Acetic acid ionises only slightly, CH3COOH <=> CH3COO- + H+. Adding sodium acetate (CH3COONa), which is fully ionised and supplies a large amount of the common CH3COO- ion, shifts the equilibrium to the left, decreasing the degree of ionisation of the acid and hence lowering [H+]. …
Same / Similar Concept — real previous-year questions on the same or a closely similar concept, not this exact question.
CBSE 2023Set ANNUAL1 markMCQ
Q.Which one of the following will cause common-ion-effect when added to the following dissociation equilibrium reaction ?
CH3COOH(aq)⇌CH3COO(aq)−+H(aq)+
(a) CH3COCl
(b) AgCl
(c) CH3Cl
(d) HCl
›Reveal solutionSolution
Adding HCl directly increases [H+], an ion common to the acetic-acid ionisation equilibrium, which by Le Chatelier's principle suppresses further dissociation of CH3COOH — the common-ion effect.
The equilibrium given is CH3COOH(aq)⇌CH3COO(aq)−+H(aq)+. A common-ion effect occurs only when the added species directly supplies an ion that already appears in this equilibrium, i.e. CH3COO− or H+. HCl is a strong acid that fully ionises to give H+ and Cl−; the extra H+ it contributes is common to the acetic-acid equilibrium, so by Le Chatelier's principle the equilibrium shifts left, suppressing ionisation of CH3COOH — a genuine common-ion effect. AgCl supplies Ag+ and Cl−, neither of which appears in this equilibrium. CH3Cl (chloromethane) and CH3COCl (acetyl chlori …
The common ion effect is the suppression of the ionisation (or solubility) of a weak electrolyte when a strong electrolyte containing an ion common to it is added to the solution.
When a strong electrolyte that shares an ion with a weak electrolyte is added to a solution of that weak electrolyte, the equilibrium of the weak electrolyte's ionisation is shifted backward (toward the un-ionised form), by Le Chatelier's principle, because the concentration of the common ion increases. This reduces the degree of ionisation (or, for a sparingly soluble salt, reduces its solubility).
NH4OH + NH4Cl show the common ion effect via the shared NH4+ ion.
The common ion effect is the decrease in the degree of ionisation of a weak electrolyte when a strong electrolyte providing a common ion is added, in line with Le Chatelier's principle. The textbook case is a weak base or acid plus its salt. Here NH4OH is a weak base: