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: …
NH4Cl supplies a high concentration of the common ion NH4+, which by Le Chatelier's principle pushes NH4OH's own dissociation equilibrium back to the left, suppressing its ionization. …
Step 1. NH4OH is a weak base and ionizes only partially: NH4OH(aq)⇌NH4+(aq)+OH−(aq).
Step 2. NH4Cl is a salt, and salts are strong electrolytes (section 3.2.1) -- so NH4Cl dissociates almost completely: NH4Cl(aq)→NH4+(aq)+Cl−(aq), adding a large extra amount of NH4+ from a source entirely separate from NH4OH's own ionization.
Step 3. NH4+ is a PRODUCT of NH4OH's own dissociation equilibrium. Increasing the concentration of a product (from any source) is a 'stress' on that equilibrium, and by Le Chatelier's principle the equilibrium shifts to counteract it -- here, by shifting toward the reactant (un-ionized NH4OH) side. …
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: