Common Ion Effect: MDCAT Chemistry notes
Common Ion Effect MDCAT notes: definition, purification of NaCl by HCl gas, KClO3 with KCl, suppression of weak acid ionization and use in salt analysis.
Definition
The common ion effect is the decrease in solubility of a salt, or the suppression of ionization of a weak electrolyte, when a solution already contains an ion common to it. It is a direct application of Le Chatelier's principle: the added common ion shifts the equilibrium backward.
Suppression of ionization means decreasing the degree of ionization.
How it works
For a saturated solution $\mathrm{AB(s) \rightleftharpoons A^+(aq) + B^-(aq)}$ with $K_{sp} = [\mathrm{A^+}][\mathrm{B^-}]$, adding extra $\mathrm{B^-}$ from another source makes the ionic product exceed $K_{sp}$. The equilibrium shifts left and solid AB precipitates until the product falls back to $K_{sp}$. $K_{sp}$ itself does not change.
Important examples
| System | Common ion added | Result |
|---|---|---|
| Saturated brine (NaCl) + HCl gas | $\mathrm{Cl^-}$ | Pure NaCl crystallizes out |
| Saturated $\mathrm{KClO_3}$ + KCl | $\mathrm{K^+}$ | Solubility of $\mathrm{KClO_3}$ decreases |
| $\mathrm{CH_3COOH}$ + $\mathrm{CH_3COONa}$ | $\mathrm{CH_3COO^-}$ | Ionization of acetic acid suppressed |
| $\mathrm{NH_4OH}$ + $\mathrm{NH_4Cl}$ | $\mathrm{NH_4^+}$ | Ionization of $\mathrm{NH_4OH}$ suppressed, $[\mathrm{OH^-}]$ lowered |
| $\mathrm{H_2S}$ + HCl | $\mathrm{H^+}$ | $[\mathrm{S^{2-}}]$ lowered |
Purification of common salt
When HCl gas is passed through saturated brine, it increases the concentration of chloride ions. Since $[\mathrm{Na^+}][\mathrm{Cl^-}]$ now exceeds $K_{sp}$ of NaCl, pure NaCl precipitates while soluble impurities stay in solution.
Use in salt analysis
The common ion effect finds extensive application in qualitative salt analysis:
- Group II: $\mathrm{H_2S}$ is passed in the presence of dilute HCl. The $\mathrm{H^+}$ ions suppress ionization of $\mathrm{H_2S}$, keeping $[\mathrm{S^{2-}}]$ low. Only sulphides with very low $K_{sp}$ (group II) precipitate.
- Group III: $\mathrm{NH_4Cl}$ is added before $\mathrm{NH_4OH}$. The $\mathrm{NH_4^+}$ ions suppress ionization of $\mathrm{NH_4OH}$, so $[\mathrm{OH^-}]$ is just enough to precipitate hydroxides of low $K_{sp}$ such as $\mathrm{Fe(OH)_3}$ and $\mathrm{Al(OH)_3}$, but not those of later groups.
The same effect is the basis of buffer solutions.
Worked example
A salt MX has $K_{sp} = 1.0\times10^{-10}$. In pure water $s = \sqrt{1.0\times10^{-10}} = 1.0\times10^{-5}\ \mathrm{mol\,dm^{-3}}$.
In $0.10\ \mathrm{M}$ NaX, $[\mathrm{X^-}] \approx 0.10$, so $s = \dfrac{K_{sp}}{0.10} = 1.0\times10^{-9}\ \mathrm{mol\,dm^{-3}}$. The solubility has fallen by a factor of $10^4$.
Key formulas
- Solubility of MX in a solution containing $c$ mol/dm³ of $\mathrm{X^-}$: $s \approx \dfrac{K_{sp}}{c}$
- Precipitation when ionic product > $K_{sp}$
Common MDCAT traps
- HCl gas in brine increases the chloride ion concentration, not sodium or hydrogen as the key effect.
- The common ion lowers solubility; it does not change $K_{sp}$.
- Adding KCl to $\mathrm{KClO_3}$ solution decreases its solubility; it does not raise $[\mathrm{ClO_3^-}]$.
- Do not confuse the common ion effect with Henry's law or Hess's law.
Quick revision
- A common ion decreases solubility of a salt.
- A common ion suppresses ionization of a weak acid or base.
- NaCl is purified by passing HCl gas through saturated brine.
- Salt analysis groups II and III rely on the common ion effect.