Question

Which of these statements about the common-ion effect is most correct? (a) The solubility of a salt MA is decreased in a solution that already contains either \(\mathrm{M}^{+}\) or \(\mathrm{A}^{-} .\) (b) Common ions alter the equilibrium constant for the reaction of an ionic solid with water. \((\mathbf{c})\) The common-ion effect does not apply to unusual ions like \(\mathrm{SO}_{3}^{2-} .(\mathbf{d})\) The solubility of a salt MA is affected equally by the addition of either \(\mathrm{A}^{-}\) or a noncommon ion.

Short answer

The most correct statement about the common-ion effect is (a): "The solubility of a salt MA is decreased in a solution that already contains either M+ or A-."

Step-by-step solution

Understanding the Common-Ion Effect

The common-ion effect refers to the decrease in the solubility of a sparingly soluble salt when an ionic compound containing a common ion is added to a solution of the salt. This occurs due to the equilibrium principle which states that when a system at equilibrium is subjected to a change in concentration, temperature, or pressure, the system will adjust itself to counteract the effect of the change and maintain the new equilibrium. Therefore, adding the common ion to the solution will make the system compensate and shift its equilibrium position, resulting in the decreased solubility of the salt. Now let's analyze each statement against this understanding.

Analyzing Statement (a)

Statement (a) says that "The solubility of a salt MA is decreased in a solution that already contains either M+ or A-." This statement is in line with the concept of the common-ion effect, as adding an ion that is already part of the salt MA will shift the equilibrium to reduce the solubility of the salt.

Analyzing Statement (b)

Statement (b) says that "Common ions alter the equilibrium constant for the reaction of an ionic solid with water." This statement is incorrect because the common-ion effect only affects solubility and causes a shift in the equilibrium position, but the equilibrium constant itself remains unchanged.

Analyzing Statement (c)

Statement (c) says that "The common-ion effect does not apply to unusual ions like SO3^2-." This statement is not correct because the common-ion effect applies to all ionic compounds, whether the ions are "unusual" or not, as long as there is a common ion in the solution.

Analyzing Statement (d)

Statement (d) says that "The solubility of a salt MA is affected equally by the addition of either A- or a noncommon ion." This statement is also incorrect because the common-ion effect only applies when a common ion is added to the solution. Adding a noncommon ion will not affect the solubility through a common-ion effect.

Conclusion

Based on the analysis of each statement, the most correct statement about the common-ion effect is (a): "The solubility of a salt MA is decreased in a solution that already contains either M+ or A-."

Key concepts

Solubility Equilibrium

Solubility equilibrium is a fundamental concept in chemistry that describes the balance between the dissolved ions in a solution and the undissolved solid salt at a given temperature. When a salt is dissolved in water, it breaks into its constituent ions. This process continues until the solution becomes saturated, forming a dynamic equilibrium.
At this point, the rate at which the ions leave the solid to dissolve into the solution is equal to the rate at which the ions in solution precipitate back to form the solid. This balance can be described mathematically by the solubility product constant, denoted as Ksp.

• The solubility product constant (Ksp) is unique for each salt and depends on the temperature of the solution.
• It has a fixed value and signifies the maximum concentration of ions that a solution can hold before precipitation begins.

Solubility equilibrium is crucial in understanding phenomena like the common-ion effect. When a solution containing a common ion is added, the added ions shift the equilibrium position according to Le Chatelier's Principle. Consequently, this affects the solubility of the salt in the solution, often leading to decreased solubility.

Le Chatelier's Principle

Le Chatelier's Principle is a key principle in understanding how a system at equilibrium responds to disturbances or changes. It states that if a chemical system at equilibrium experiences a change in concentration, temperature, or pressure, the equilibrium will shift in a direction that opposes the change to achieve a new balance.
For instance, in the context of the common-ion effect, when a common ion is introduced to a solution, it results in increased concentration of that ion in the equilibrium mix. Applying Le Chatelier's Principle:

• The system shifts to counter this change by favoring the reaction that uses up the excess ion.
• This often pushes the equilibrium towards the formation of more solid salt, thus decreasing solubility.

By understanding Le Chatelier's Principle, students can predict how various factors impact equilibrium and solubility. It can also explain why the addition of a common ion causes decreased solubility, as the system strives to maintain equilibrium by precipitating fewer ions into the solution.

Equilibrium Constants

Equilibrium constants are vital in describing the extent to which a chemical reaction proceeds. They provide a quantitative measure of the position of equilibrium for reversible reactions, expressed in terms of concentration of products and reactants. For solubility equilibria, the equilibrium constant is specifically known as the solubility product constant (Ksp).
The value of Ksp is generally unaffected by the presence of a common ion, which is a crucial point to remember. While the common-ion effect can shift the equilibrium position, thus altering the solubility of a salt, the equilibrium constant itself remains the same.

• This invariability means Ksp is characteristic for a compound at a given temperature.
• It allows scientists and students to predict solubility under various conditions, provided the temperature remains constant.

Understanding equilibrium constants facilitates a deeper grasp of chemical reactions and solubility, highlighting how changes in concentration affect reaction directions without altering the inherent equilibrium properties of the system.