Question

Is each compound more soluble in acidic solution or in pure water? Explain. a. \(\mathrm{Hg}_{2} \mathrm{Br}_{2}\) b. \(\mathrm{Mg}(\mathrm{OH})_{2}\) c. \(\mathrm{CaCO}_{3}\) d. \(\mathrm{AgI}\)

Short answer

\(\mathrm{Hg}_{2} \mathrm{Br}_{2}\), \(\mathrm{Mg}(\mathrm{OH})_{2}\), and \(\mathrm{CaCO}_{3}\) are more soluble in acidic solution because the reaction with H+ ions reduces the concentration of their respective anions. \(\mathrm{AgI}\) does not react with H+ to form weakly ionized species, so its solubility is not significantly increased in acidic solution compared to pure water.

Step-by-step solution

Understanding Solubility in Acidic Solution vs. Pure Water

To determine if a compound is more soluble in an acidic solution compared to pure water, we must consider the reaction of the compound with H+ ions. If the resulting species after reaction with H+ ions are less abundant than the original anion, the compound's solubility should increase in an acidic solution. This is because the reaction with H+ reduces the ion concentration, driving the dissolution reaction forward according to Le Chatelier's principle.

Analyze Solubility of \(\mathrm{Hg}_{2} \mathrm{Br}_{2}\)

For \(\mathrm{Hg}_{2} \mathrm{Br}_{2}\), the dissociation in water is as follows: \[ \mathrm{Hg}_{2} \mathrm{Br}_{2} (s) \rightleftharpoons \mathrm{Hg}_{2}^{2+} (aq) + 2\mathrm{Br}^{-} (aq) \] In an acidic solution, the bromide ions (\mathrm{Br}^{-}) would react with the H+ ions to form HBr, which is a weak acid and would exist mainly as molecules rather than ions. This decreases the concentration of bromide ions, thus shifting the equilibrium to the right and increasing the solubility of \(\mathrm{Hg}_{2} \mathrm{Br}_{2}\).

Analyze Solubility of \(\mathrm{Mg}(\mathrm{OH})_{2}\)

For \(\mathrm{Mg}(\mathrm{OH})_{2}\), the dissociation in water is: \[ \mathrm{Mg}(\mathrm{OH})_{2} (s) \rightleftharpoons \mathrm{Mg}^{2+} (aq) + 2\mathrm{OH}^{-} (aq) \] The hydroxide ions (\mathrm{OH}^{-}) will react with H+ ions in an acidic solution to form water: \[ \mathrm{OH}^{-} (aq) + \mathrm{H}^{+} (aq) \rightarrow \mathrm{H}_{2}\mathrm{O} (l) \] This reaction removes hydroxide ions, driving the dissolution reaction forward and increasing the solubility of \(\mathrm{Mg}(\mathrm{OH})_{2}\) in acidic solutions.

Analyze Solubility of \(\mathrm{CaCO}_{3}\)

For \(\mathrm{CaCO}_{3}\), the dissociation in water is: \[ \mathrm{CaCO}_{3} (s) \rightleftharpoons \mathrm{Ca}^{2+} (aq) + \mathrm{CO}_{3}^{2-} (aq) \] In an acidic solution, the carbonate ions (\mathrm{CO}_{3}^{2-}) react with H+ ions to form \(\mathrm{HCO}_{3}^{-}\) and ultimately \(\mathrm{CO}_{2}\) gas and water, thus decreasing the carbonate ion concentration and increasing solubility of \(\mathrm{CaCO}_{3}\) in acidic solution.

Analyze Solubility of \(\mathrm{AgI}\)

For \(\mathrm{AgI}\), the dissociation in water is: \[ \mathrm{AgI} (s) \rightleftharpoons \mathrm{Ag}^{+} (aq) + \mathrm{I}^{-} (aq) \] However, since neither the silver nor the iodide ions react significantly with H+ ions to form weakly ionized species or gases, the solubility of \(\mathrm{AgI}\) should not increase when placed in an acidic solution. So it is more soluble in pure water than in an acidic solution.

Key concepts

Le Chatelier's Principle

Understanding how a system at equilibrium reacts to disturbances is crucial for solving many chemistry problems. Le Chatelier's principle states that if an external change is applied to a system at equilibrium, the system will adjust itself to partially counteract the change and re-establish equilibrium. For example, when a solute's ions in a saturated solution react with added reagents, like H+ ions in an acidic solution, the equilibrium shifts to replace the ions that are removed by this reaction.

This principle can explain why certain salts are more soluble in an acidic solution than in pure water. When an acid reacts with one of the ions produced by the salt's dissociation, the equilibrium shifts to produce more of those ions, effectively increasing the solubility of the salt. The steps provided in the exercise solution apply Le Chatelier's Principle to predict the effect of an acidic environment on the solubility of different compounds.

Solubility Product

The solubility product, or K_sp, is a constant that provides a measure of the solubility of a compound under a set of conditions, typically at a fixed temperature. It is the product of the molar concentrations of the ions in a saturated solution, each raised to the power of its stoichiometric coefficient in the dissociation equation.

Using the solubility product, we can assess how the solubility of a compound changes in response to different conditions. It is important to remember that the K_sp remains constant for a given compound at a certain temperature. However, as the exercise shows, adding an acid can change the concentration of the ions in the solution, pushing the dissolution reaction forward and effectively increasing solubility even though K_sp is unchanged. Understanding the interaction between the K_sp and Le Chatelier's principle enhances comprehension of solubility behavior in various environments.

Common Ion Effect

The common ion effect refers to the decrease in the solubility of an ionic compound when a common ion is added to the solution. The presence of a common ion from an added salt or acid will shift the equilibrium, as described by Le Chatelier's principle, reducing the solubility of the compound.

For instance, if we add HCl to a solution of AgCl, the increased concentration of Cl⁻ ions will shift the equilibrium toward the solid AgCl, making the solution less saturated with AgCl. However, this effect can work in reverse when dealing with insoluble bases or acidic salts, where removing the common ion (e.g., OH⁻ or CO₃²⁻) via an acid-base reaction increases the solubility of the compound. This is the rationale applied in the exercise to analyze how acidic conditions affect solubility.

Acid-Base Reactions

Acid-base reactions are a type of chemical reaction that typically involves the transfer of a proton (H+) from an acid to a base. In the context of solubility, these reactions can greatly influence the saturation level of a solution.

For compounds that produce hydroxide or other anions that react with acids, the addition of an acid to the solution can reduce the concentration of these anions through neutralization or formation of less soluble species. In the steps provided, compounds such as Mg(OH)₂ and CaCO₃ display increased solubility in an acidic solution due to the acid-base reactions that consume their anions, thus demonstrating how these reactions play a pivotal role in determining solubility in various environments.