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: …
The common ion effect is the suppression of the ionisation of a weak electrolyte when a strong electrolyte supplying an ion common to that equilibrium is added, e.g. adding sodium acetate suppresses acetic acid's ionisation via the shared acetate ion. …
By Le Chatelier's principle, adding an ion already present in a weak electrolyte's ionisation equilibrium shifts that equilibrium backward, suppressing further ionisation — the common ion effect, illustrated by adding sodium acetate to acetic acid.
Consider the ionisation equilibrium of a weak acid, acetic acid: CH3COOH⇌CH3COO−+H+ If sodium acetate (CH3COONa), a strong electrolyte, is added to this solution, it dissociates completely, greatly increasing the concentration of CH3COO− ions in solution. Because CH3COO− is common to (already appears in) the acetic acid equilibrium above, this added common ion shifts the equilibrium backward (toward the left, favouring un-ionised CH3COOH), by Le Chatelier's principle. As a result, the degree of ionisation of acetic acid decreases, and so does [H+] (and hence the solution's acidity). This suppression of a weak electrolyte's own ionisation, caused specifically by adding a strong electrolyte with a shared/common ion, is called the common ion effect. Another standard example: ad …