Problem 88 - Chemistry The Science in Context

Question

Give an example of how the common-ion effect limits the dissolution of a sparingly soluble ionic compound.

Step-by-Step Solution

Answer

Answer: The common-ion effect limits the dissolution of a sparingly soluble ionic compound by decreasing its solubility when a soluble compound consisting of an ion in common with the sparingly soluble salt is added to the solution. This occurs due to the shift in equilibrium driven by Le Châtelier's principle. In the example of silver chloride (AgCl), adding sodium chloride (NaCl) to a saturated solution of AgCl increases the concentration of the common chloride ion (Cl^-) and shifts the equilibrium to the left, causing more AgCl to precipitate out of the solution and decrease its solubility.

Step 1: Understanding the common-ion effect

The common-ion effect refers to the decrease in solubility of a sparingly soluble salt when a soluble compound which consists of an ion in common with the sparingly soluble salt is added to the solution. This occurs due to the shift in equilibrium as an effect of Le Châtelier's principle.

Step 2: Choosing an example of a sparingly soluble salt

Let's choose the sparingly soluble ionic compound, silver chloride (AgCl), as our example of the common-ion effect. AgCl has limited solubility in water, and it dissolves according to the following equilibrium: AgCl(s) <=> Ag^+(aq) + Cl^-(aq)

Step 3: Introducing a soluble compound with a common ion

Let's introduce sodium chloride (NaCl) to the already saturated solution of silver chloride. NaCl is a soluble salt, and it shares a common ion (chloride, Cl^-) with AgCl. NaCl(s) -> Na^+(aq) + Cl^-(aq)

Step 4: Explaining the effect of the common-ion on the dissolution of the sparingly soluble salt

When NaCl is added to the saturated solution of AgCl, the concentration of the common ion (Cl^-) in solution increases. By Le Châtelier's principle, the equilibrium shifts to counteract the increase in Cl^- concentration, and this happens by shifting the equilibrium back to the left: AgCl(s) <= Ag^+(aq) + Cl^-(aq) As a result, more of the sparingly soluble silver chloride precipitates out of the solution and the solubility of AgCl decreases.

Step 5: Concluding the example

In the example of silver chloride, the addition of sodium chloride (a soluble compound with a common ion) to a saturated solution of AgCl decreases the solubility of AgCl. This is due to the common-ion effect, and it demonstrates how the dissolution of a sparingly soluble ionic compound can be limited.

Key Concepts

Le Châtelier's Principle

Le Châtelier's principle is a fundamental concept in chemistry, which states that if a dynamic equilibrium is disturbed by changing the conditions, the system responds to minimize the change and restore a new equilibrium.

In the context of solubility, when an ionic compound is dissolved in water, it reaches a point where the rate of dissolving equals the rate of crystallization, establishing a solubility equilibrium. If we add more of one of the ions present in the dissolved compound, say Cl⁻ from NaCl, the increased concentration of that ion pushes the equilibrium to reduce its effect.

Think of it as a see-saw: If one person on a see-saw suddenly adds weight, the see-saw will tilt; to balance it again, you'd need to adjust the weight on the other side. Similarly, the system adjusts by precipitating out the excess ions as the sparingly soluble solid, thus reducing their concentration in the solution.

Solubility Equilibrium

Solubility equilibrium refers to the state of a saturated solution, where the rate of dissolution of a substance equals the rate at which it precipitates. It's represented by a balanced chemical equation, often for a sparingly soluble salt in water.

For example, the equilibrium for silver chloride dissolving in water can be written as: AgCl(s) ⇌ Ag⁺(aq) + Cl⁻(aq)

At this equilibrium, there is a fixed ratio of silver ions to chloride ions in the solution. Any disturbance to this ratio by adding more of either ion causes a shift in the equilibrium following Le Châtelier's principle, resulting in a change in the solubility of the salt.

Sparingly Soluble Salts

Sparingly soluble salts are compounds that have a very low solubility in water. These salts do not completely dissolve, leaving behind a residue or precipitate.

Silver chloride (AgCl) is a classic example of a sparingly soluble salt. When AgCl is placed in water, only a small fraction dissolves to form silver and chloride ions. The solubility of these salts is often so low that they are used in chemical analysis and other applications where a restrained solubility is beneficial.

Silver Chloride (AgCl) Solubility

Silver chloride's solubility is pivotal in understanding the common-ion effect. AgCl is sparingly soluble in water, dissolving to a small extent to give silver ions (Ag⁺) and chloride ions (Cl⁻). Its solubility in water is further decreased in the presence of additional Cl⁻ ions.

This can be shown practically by adding a salt like sodium chloride (NaCl), which dissociates to yield more chloride ions into the solution. This common-ion presence shifts the equilibrium, resulting in more solid AgCl forming, and showcasing the common-ion effect. These concepts are not just isolated chemical curiosities but have broad applications, including controlling reactions in industrial processes, and even in medical treatments where precise ion concentrations are crucial.