This question tests the effect of low pH on the solubility of fluoride salts through formation of weak acid HF. At pH 2, the high [H+] protonates F- ions from CaF2 dissolution to form HF, reducing [F-] in solution. This decrease in [F-] makes Q < Ksp, driving the equilibrium toward greater dissolution per Le Chatelier's principle. As a result, the solubility of CaF2 increases in the acidic solution compared to pure water. A tempting distractor is choice B, which asserts fluoride cannot react with H+ because it's from a strong acid, but this misconceives HF as weak, allowing protonation. To determine pH effects on solubility, identify if the anion is a conjugate base of a weak acid and evaluate how protonation shifts the dissolution equilibrium.

This question tests understanding of the common ion effect with metal hydroxides at high pH. Cu(OH)₂ dissolves according to: Cu(OH)₂(s) ⇌ Cu²⁺(aq) + 2OH⁻(aq). At pH 10, the solution has a relatively high [OH⁻] concentration (10⁻⁴ M) compared to pure water. These OH⁻ ions are a common ion with the dissolution products, so according to Le Chatelier's principle, the increased [OH⁻] shifts the equilibrium to the left, decreasing the solubility of Cu(OH)₂. Choice C incorrectly suggests that buffers prevent equilibrium shifts, but buffers only maintain constant pH—they don't prevent the common ion effect. For metal hydroxides, high pH (high [OH⁻]) always decreases solubility through the common ion effect, while low pH increases solubility by removing OH⁻ through neutralization.

This question tests understanding of how weak acids affect the solubility of salts containing their conjugate bases. Silver acetate dissolves as: AgC₂H₃O₂(s) ⇌ Ag⁺(aq) + C₂H₃O₂⁻(aq). Acetic acid (HC₂H₃O₂) is a weak acid that partially dissociates, providing H⁺ ions that can protonate the acetate ion (C₂H₃O₂⁻) to form more HC₂H₃O₂. This removes C₂H₃O₂⁻ from the solution, and by Le Chatelier's principle, the dissolution equilibrium shifts right to produce more ions, increasing solubility. Choice D incorrectly claims that acetic acid supplies acetate ions, but weak acids actually consume their conjugate bases through protonation rather than supplying them. When a weak acid is added to a solution containing its conjugate base as part of a sparingly soluble salt, the acid will increase the salt's solubility by removing the anion through protonation.

This question tests understanding of how pH affects the solubility of salts containing basic anions. When CaCO₃ dissolves, it produces Ca²⁺ and CO₃²⁻ ions, where CO₃²⁻ is a basic anion that can accept protons. In acidic solution (low pH), the high concentration of H⁺ ions reacts with CO₃²⁻ to form HCO₃⁻ and H₂CO₃, effectively removing CO₃²⁻ from the solution. According to Le Chatelier's principle, removing a product (CO₃²⁻) shifts the dissolution equilibrium CaCO₃(s) ⇌ Ca²⁺(aq) + CO₃²⁻(aq) to the right, increasing solubility. Choice B incorrectly suggests Cl⁻ is a common ion with CaCO₃, but CaCO₃ contains no chloride ions. To solve pH-solubility problems, identify whether the anion is basic (can accept H⁺) or neutral, then apply Le Chatelier's principle to predict how pH changes affect the dissolution equilibrium.

This question tests understanding of how pH affects the solubility of salts containing polyprotic basic anions. Ca₃(PO₄)₂ dissolves to produce Ca²⁺ and PO₄³⁻ ions, where PO₄³⁻ is a strongly basic anion that readily accepts protons. In acidic solution (HCl), H⁺ ions protonate PO₄³⁻ stepwise to form HPO₄²⁻, H₂PO₄⁻, and even H₃PO₄, effectively removing PO₄³⁻ from the solution. By Le Chatelier's principle, removing the product PO₄³⁻ shifts the dissolution equilibrium to the right, significantly increasing the solubility of Ca₃(PO₄)₂. Choice D incorrectly reverses the acid-base chemistry, claiming acid increases [PO₄³⁻] when it actually decreases it through protonation. For salts containing highly basic anions (especially polyprotic ones like PO₄³⁻, CO₃²⁻), acidic conditions dramatically increase solubility by removing the anion through multiple protonation steps.

This question tests understanding of how pH affects the solubility of metal hydroxides. Zn(OH)₂ dissolves according to: Zn(OH)₂(s) ⇌ Zn²⁺(aq) + 2OH⁻(aq). In acidic solution (HNO₃), H⁺ ions react with OH⁻ ions to form water through the neutralization reaction: H⁺ + OH⁻ → H₂O. This effectively removes OH⁻ from the solution, and by Le Chatelier's principle, the dissolution equilibrium shifts to the right to produce more ions, increasing solubility. Choice D incorrectly claims that adding acid increases [OH⁻], when acids actually decrease hydroxide concentration by neutralization. For metal hydroxides, acidic conditions always increase solubility because H⁺ removes OH⁻ through neutralization, while basic conditions decrease solubility through the common ion effect.

This question tests the impact of low pH on the solubility of carbonate salts via acid-base reactions with the carbonate ion. In Beaker 2 at pH 2, the added H+ from the strong acid converts $CO3^2$- from Ag2CO3 dissolution into HCO3- and H2CO3, lowering $[CO3^2$-]. This reduction in $[CO3^2$-] causes Q to fall below Ksp, promoting more dissolution to restore equilibrium. Consequently, the solubility of Ag2CO3 increases in the acidic beaker compared to pure water. A tempting distractor is choice A, which mistakenly applies the common-ion effect to H+, but H+ is not an ion in the Ksp expression and instead participates in an acid-base reaction. To evaluate pH-dependent solubility, check if the anion can accept protons and use Le Chatelier's principle to analyze how removing the anion affects the equilibrium.

This question tests the understanding of how pH affects the solubility of salts with anions that are conjugate bases of weak acids. In Beaker 2 with 0.10 M HCl, the high [H+] from the acid reacts with $CO3^2$- ions produced by the dissolution of CaCO3 to form HCO3- and H2CO3. This reaction lowers the $[CO3^2$-] in solution, causing the ion product Q to be less than Ksp and shifting the equilibrium toward more dissolution according to Le Chatelier's principle. Therefore, the solubility of CaCO3 increases in the acidic solution compared to pure water. A tempting distractor is choice B, which misapplies the common-ion effect by claiming acid increases $[CO3^2$-], but this ignores the acid-base reaction that actually decreases $[CO3^2$-]. To predict pH effects on solubility, determine if the anion can be protonated by checking if it is the conjugate base of a weak acid, and apply Le Chatelier's principle to see how it shifts the equilibrium.

The skill being tested is comparing solubility at different pH levels for carbonates. The correct answer is that Ag2CO3 is more soluble at pH 2 because H+ converts $CO3^2$- to HCO3-/H2CO3, reducing free $[CO3^2$-] and promoting dissolution per Le Châtelier's principle. At pH 12, the lack of H+ prevents this reaction, resulting in lower solubility similar to neutral conditions. $CO3^2$- acts as a base, making acidity enhance solubility. A tempting distractor is A, that it's more soluble at pH 12 because OH- removes $CO3^2$-, but this misconceives the acid-base behavior, as OH- does not effectively protonate or remove carbonate. Identify the acid-base nature of ions and use equilibrium shifts to assess pH impacts on solubility.

The skill being tested is understanding how pH affects the solubility of salts with basic anions through acid-base equilibria. The correct answer is that solubility is highest in 0.10 M HCl because H+ consumes $CO3^2$- to form HCO3-/H2CO3, shifting the dissolution equilibrium forward by Le Châtelier's principle. In pure water, there is no excess H+ to react with $CO3^2$-, resulting in lower solubility limited by Ksp. In 0.10 M NaOH, the high [OH-] does not remove $CO3^2$- and may slightly suppress solubility due to indirect effects, but the key is the acid's enhancement. A tempting distractor is C, that solubility is the same because Ksp fixes it regardless of pH, but this misconceives that while Ksp is constant, coupled acid-base reactions alter effective ion concentrations. Always consider if dissolved ions can participate in pH-dependent equilibria when predicting solubility changes.