Question

The following reaction that has reached equilibrium:
$NaCl(s)\leftarrow \to NaCl(aq)$
What should be the nature of the solution of NaCl for the phase equilibrium to exist?
A. Concentrated
B. Saturated
C. Dilute
D. Heated
E. Unsaturated

Answer 1

Hint: Sodium chloride shows high solubility in water as it is an ionic compound. The solid sodium chloride dissolves in water to form an aqueous solution but there is a point where no more solute can be further added to the solvent. This shows that the phase equilibrium is attained by the solution.

Complete step by step answer:
Crystal Sodium chloride is an ionic compound where the sodium ion and chloride ion are closely packed when dissolved in water the ionic compound sodium chloride dissociates to generate its constituent ions forming an aqueous solution.
The dissociation reaction of sodium chloride in water is shown below.
\[NaCl(s)\to N{{a}^{+}}(aq)+C{{l}^{-}}(aq)\]
The sodium chloride is considered as a strong electrolyte as it rapidly dissociates on dissolving in water into its constituent sodium cation and chloride anion. When sodium chloride is mixed in water there comes a point where no more sodium chloride can be added in the solution as equilibrium is attained and the solute starts to form white precipitate which can be seen in the bottom of the container. This stage is known as saturation and the solution is called a saturated solution. The property of solutes to dissolve in a saturated solution is called solubility. The saturated solution contains the maximum amount of solute which it can dissolve.
Thus the nature of the solution of NaCl for the phase equilibrium to exist is saturated.
So, the correct answer is “Option B”.

Note:
At the saturation point the rate of dissolution of sodium chloride and rate of precipitation becomes equal. If more solvent is added then more solute can be added to form the solution.