Concept:The purification of table salt relies on disrupting the solubility equilibrium by adding an external source of a shared ion.
Formula:$$ \text{NaCl}_{(s)} \rightleftharpoons \text{Na}^+_{(aq)} + \text{Cl}^-_{(aq)} $$
Solution:- When \( \text{HCl} \) gas is passed through the solution, it heavily dissociates, providing a massive influx of \( \text{Cl}^- \) ions.
- Because \( \text{Cl}^- \) is "common" to both \( \text{NaCl} \) and \( \text{HCl} \), Le Chatelier's principle dictates the equilibrium shifts left to consume the excess \( \text{Cl}^- \).
- This forces pure \( \text{NaCl} \) to crystallize and precipitate, leaving impurities in the solution. This is a classic application of the Common Ion Effect.
Why other options are incorrect:Hess's law relates to enthalpy changes. Henry's law relates to gas solubility. The Law of mass action defines the equilibrium expression but doesn't specifically name this precipitation phenomenon.
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