Problem 98 - General Chemistry: Principles and Modern Applications

Question

To increase the molar solubility of \(\mathrm{CaCO}_{3}(\mathrm{s})\) in a saturated aqueous solution, add (a) ammonium chloride; (b) sodium carbonate; (c) ammonia; (d) more water.

Step-by-Step Solution

Answer

(a) The molar solubilty won't change with \(NH_4Cl\). (b) The molar solubilty will decrease with \(Na_2CO_3\). (c) The molar solubilty will increase with \(NH_3\). (d) The molar solubilty will increase with more water.

Step 1: Analyzing the effect of ammonium chloride

Adding ammonium chloride (\(NH_4Cl\)) to the solution won't affect the solubility of \(CaCO_3\). This would only provide more chloride ions, which doesn't interact with \(CaCO_3\), and thus solubility stays the same.

Step 2: Analyzing the effect of sodium carbonate

Adding sodium carbonate (\(Na_2CO_3\)) would increase the carbonate ions (\(CO_3^{2-}\)) in the solution. According to Le Chatelier’s principle, this would shift the reaction to the left, reducing the solubility of the \(CaCO_3\). This is due to the common ion effect.

Step 3: Analyzing the effect of ammonia

Adding ammonia (\(NH_3\)) would shift the equilibrium towards the right as \(NH_3\) can react with \(Ca^{2+}\) to form soluble complexes. This would remove \(Ca^{2+}\) ions from the solution and, according to Le Chatelier’s principle, the reaction would shift right to compensate making more \(CaCO_3\) dissolve.

Step 4: Analyzing the effect of more water

Adding more water increases the volume of the solution, decreasing the concentration of the ions. According to Le Chatelier's principle, to counter the decrease in concentration, more \(CaCO_3\) will dissolve.

Key Concepts

Le Chatelier’s Principle

Le Chatelier’s principle is a fundamental concept in chemistry that predicts how a change in conditions can affect chemical equilibria. Specifically, it states that if a system at equilibrium experiences a change in pressure, temperature, or concentration of reactants or products, the system will adjust, or shift, in such a way as to counteract the change and re-establish equilibrium.

Common Ion Effect

The common ion effect refers to the decrease in solubility of an ionic compound when a common ion is added to the solution. It occurs because the addition of an ion that is already present in the saturated solution causes the equilibrium to shift according to Le Chatelier’s principle. With more of the common ion in solution, the reaction tends to move in the direction that consumes the common ions, decreasing the concentration of the dissolved ions and leading to precipitation if the product of the ionic concentrations exceeds the solubility product.

Solubility Equilibrium

Solubility equilibrium is the dynamic balance between the dissolution and precipitation of a solute in a solvent. It is defined by the solubility product constant (\(K_{sp}\)), which is the product of the molar concentrations of the ions produced when the solute dissolves, each raised to the power of its coefficient in the balanced equation. If the concentrations of dissolved ions in a saturated solution are such that the product exceeds the solubility product, the excess will precipitate out until equilibrium is re-established.

Saturated Aqueous Solution

A saturated aqueous solution contains the maximum amount of solute that can dissolve at a given temperature. At this point, the solution is in dynamic equilibrium with the undissolved solute, with the rate of dissolution equal to the rate of precipitation. Any undissolved solute settles at the bottom, and the solution's concentration remains constant unless disturbed by changes such as temperature fluctuations or the addition of more solute or solvent.