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title: "A student investigates the solubility of calcium fluoride, \\(\\text{CaF}_2(s)\\), in different aqueous environments at \\(25^\\circ\\text{C}\\). The relevant equilibrium processes and their equilibrium constants at this temperature are given below.  \\[\\text{CaF}_2(s) \\rightleftharpoons \\text{Ca}^{2+}(aq) + 2\\,\\text{F}^-(aq) \\quad K_{sp} = 3.9 \\times 10^{-11}\\] \\[\\text{HF}(aq) \\rightleftharpoons \\text{H}^+(aq) + \\text{F}^-(aq) \\quad K_a = 6.3 \\times 10^{-4}\\]  The student compares the molar solubility of \\(\\text{CaF}_2(s)\\) in pure water at \\(\\text{pH} = 7.0\\) to its molar solubility in a buffer solution maintained at \\(\\text{pH} = 3.0\\). Which of the following correctly predicts and justifies the relative solubility of \\(\\text{CaF}_2(s)\\) in the buffer solution at \\(\\text{pH} = 3.0\\) compared to pure water?"
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url: "https://nerd-notes.com/ubq/121474/"
date_modified: "2026-08-23T05:04:46+00:00"
---

# A student investigates the solubility of calcium fluoride, \(\text{CaF}_2(s)\), in different aqueous environments at \(25^\circ\text{C}\). The relevant equilibrium processes and their equilibrium constants at this temperature are given below.

\[\text{CaF}_2(s) \rightleftharpoons \text{Ca}^{2+}(aq) + 2\,\text{F}^-(aq) \quad K_{sp} = 3.9 \times 10^{-11}\]
\[\text{HF}(aq) \rightleftharpoons \text{H}^+(aq) + \text{F}^-(aq) \quad K_a = 6.3 \times 10^{-4}\]

The student compares the molar solubility of \(\text{CaF}_2(s)\) in pure water at \(\text{pH} = 7.0\) to its molar solubility in a buffer solution maintained at \(\text{pH} = 3.0\). Which of the following correctly predicts and justifies the relative solubility of \(\text{CaF}_2(s)\) in the buffer solution at \(\text{pH} = 3.0\) compared to pure water?

A student investigates the solubility of calcium fluoride, \(\text{CaF}_2(s)\), in different aqueous environments at \(25^\circ\text{C}\). The relevant equilibrium processes and their equilibrium constants at this temperature are given below.

\[\text{CaF}_2(s) \rightleftharpoons \text{Ca}^{2+}(aq) + 2\,\text{F}^-(aq) \quad K_{sp} = 3.9 \times 10^{-11}\]
\[\text{HF}(aq) \rightleftharpoons \text{H}^+(aq) + \text{F}^-(aq) \quad K_a = 6.3 \times 10^{-4}\]

The student compares the molar solubility of \(\text{CaF}_2(s)\) in pure water at \(\text{pH} = 7.0\) to its molar solubility in a buffer solution maintained at \(\text{pH} = 3.0\). Which of the following correctly predicts and justifies the relative solubility of \(\text{CaF}_2(s)\) in the buffer solution at \(\text{pH} = 3.0\) compared to pure water?

- **A.** The solubility is greater at \(\text{pH} = 3.0\) because \(\text{H}^+(aq)\) reacts with \(\text{F}^-(aq)\) to form the weak acid \(\text{HF}(aq)\), decreasing \([\text{F}^-]\) and shifting the dissolution equilibrium toward products.
- **B.** The solubility is greater at \(\text{pH} = 3.0\) because the added \(\text{H}^+(aq)\) ions increase the numerical value of the solubility product constant, \(K_{sp}\), for \(\text{CaF}_2\).
- **C.** The solubility is less at \(\text{pH} = 3.0\) because the excess \(\text{H}^+(aq)\) ions compete with \(\text{Ca}^{2+}(aq)\) for water molecules, exerting a common-ion effect that shifts the dissolution equilibrium toward reactants.
- **D.** The solubility is unchanged at \(\text{pH} = 3.0\) because \(\text{F}^-(aq)\) is the conjugate base of a strong acid and therefore does not accept protons from \(\text{H}^+(aq)\) to alter the equilibrium.

*The answer key and step-by-step explanation are available to logged-in users at https://nerd-notes.com/ubq/121474/*