How does the presence of a common ion affect the solubility of a sparingly soluble salt?

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Multiple Choice

How does the presence of a common ion affect the solubility of a sparingly soluble salt?

The presence of a common ion decreases the solubility of a sparingly soluble salt due to the common ion effect. When a salt that is only slightly soluble is dissolved in a solution that already contains a common ion, the equilibrium of the dissolution process shifts to counteract the increase in concentration of that ion. For example, consider a sparingly soluble salt like silver chloride (AgCl). The dissolution can be represented by: AgCl(s) ⇌ Ag⁺(aq) + Cl⁻(aq) If the solution already contains chloride ions (Cl⁻) from another source, the increased concentration of Cl⁻ shifts the equilibrium to the left, according to Le Chatelier's principle. This shift reduces the concentration of Ag⁺ ions that can remain in solution because the system attempts to decrease the added concentration of Cl⁻ by forming more solid AgCl. As a result, fewer moles of the salt dissolve, leading to a decrease in solubility. This concept is essential in understanding the behavior of ionic compounds in solutions and is widely applicable in various aqueous chemistry scenarios.

The presence of a common ion decreases the solubility of a sparingly soluble salt due to the common ion effect. When a salt that is only slightly soluble is dissolved in a solution that already contains a common ion, the equilibrium of the dissolution process shifts to counteract the increase in concentration of that ion.

For example, consider a sparingly soluble salt like silver chloride (AgCl). The dissolution can be represented by:

AgCl(s) ⇌ Ag⁺(aq) + Cl⁻(aq)

If the solution already contains chloride ions (Cl⁻) from another source, the increased concentration of Cl⁻ shifts the equilibrium to the left, according to Le Chatelier's principle. This shift reduces the concentration of Ag⁺ ions that can remain in solution because the system attempts to decrease the added concentration of Cl⁻ by forming more solid AgCl. As a result, fewer moles of the salt dissolve, leading to a decrease in solubility.

This concept is essential in understanding the behavior of ionic compounds in solutions and is widely applicable in various aqueous chemistry scenarios.

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