#pH and Solubility: A Last-Minute Review 🚀

Hey, you've got this! Let's quickly review how pH affects solubility, focusing on the key ideas you'll need for the AP exam. Remember, it's all about how equilibrium shifts with changing conditions.

#The Big Picture: pH, Solubility, and Equilibrium

• pH tells us the concentration of H⁺ (and OH⁻) ions in a solution.
• Solubility is all about how much of a substance dissolves in a solution, and it's an equilibrium process. ⚖️
• Le Chatelier's Principle is our guide: If we change a condition (like pH), the equilibrium will shift to counteract that change.
• The common ion effect can also influence solubility by adding an ion that's already part of the equilibrium.

#### Effects of Acidic Solutions 🍋

• Key Idea: Acidic solutions (high [H⁺]) generally increase the solubility of compounds with basic conjugate bases.
• Why? The H⁺ ions react with the basic conjugate base (A⁻) to form HA, reducing [A⁻], and shifting the equilibrium towards dissolution.
  • $A^- + H^+ \rightleftharpoons HA$
• The weaker the acid, the more basic its conjugate base, and the more its solubility is affected by pH.
• Example: Fe(OH)₃ solubility increases in acidic solutions because H⁺ ions react with OH⁻, reducing [OH⁻] and pushing the dissolution equilibrium to the right.
  • $Fe(OH)_3 \rightleftharpoons Fe^{3+} + 3OH^-$
  • $H^+ + OH^- \rightleftharpoons H_2O$

#Image From Purdue University

#### Effects of Basic Solutions 🧪

• Key Idea: Basic solutions (high [OH⁻]) generally increase the solubility of compounds with acidic conjugate acids.
• Why? The OH⁻ ions react with the acidic conjugate acid (like NH₄⁺), reducing its concentration and shifting the equilibrium towards dissolution.
  • $NH_4^+ + OH^- \rightleftharpoons NH_3 + H_2O$
• However, basic compounds or salts with basic conjugate bases become less soluble due to the common ion effect (increased [OH⁻]).
• Example: The solubility of a salt with a conjugate base like CH₃COO⁻ decreases in a basic solution because the equilibrium $CH_3COO^- + H_2O \rightleftharpoons CH_3COOH + OH^-$ shifts left due to the added OH⁻.

#Image From Acids and Bases 101

#### pH Neutral Compounds 😐

• Key Idea: Compounds with neutral ions (conjugates of strong acids/bases) are not affected by pH.
• Why? Ions like Na⁺ and Cl⁻ don't react with H⁺ or OH⁻ because they are conjugates of strong acids/bases.
• Example: NaCl's solubility is not altered by pH changes because Na⁺ and Cl⁻ are neutral.
• Exception: The common ion effect can still affect solubility if a common ion is present, but this is independent of pH.

#

• You won't have to calculate solubility changes due to pH on the AP exam, but you must understand the qualitative effects.
• Acidic solutions: Increase solubility of compounds with basic conjugate bases.
• Basic solutions: Increase solubility of compounds with acidic conjugate acids, decrease solubility of compounds with basic conjugate bases.
• Neutral compounds: Not affected by pH changes.
• Focus on applying Le Chatelier's Principle to predict shifts in equilibrium.

#Final Exam Focus 🎯

• Highest Priority: Understanding how pH affects the solubility of salts with basic or acidic conjugate ions.
• Common Question Types:
  • Predicting how solubility changes in acidic or basic solutions.
  • Applying Le Chatelier's Principle to solubility equilibria.
  • Identifying conjugate acid-base pairs and their impact on solubility.
• Time Management: Quickly identify if the salt has a basic or acidic conjugate and use Le Chatelier's Principle.
• Common Pitfalls: Confusing the common ion effect with pH effects.

#Practice Questions 🧪

Practice Question

#Multiple Choice Questions

1. The solubility of which of the following salts is most affected by changes in pH? (A) NaCl (B) KNO₃ (C) CaF₂ (D) BaSO₄

2. Which of the following will increase the solubility of Mg(OH)₂ in water? (A) Adding NaOH (B) Adding HCl (C) Adding MgCl₂ (D) Adding NaCl

#Free Response Question

A 1.0 L solution contains 0.010 mol of solid PbCl₂. The $K_{sp}$ of PbCl₂ is $1.6 \times 10^{-5}$.

(a) Write the balanced chemical equation for the dissolution of PbCl₂.

(b) Calculate the concentration of $Pb^{2+}$ in the solution at equilibrium.

(c) If a small amount of HCl is added to the solution, will the solubility of PbCl₂ increase, decrease, or remain the same? Explain.

(d) If a small amount of NaOH is added to the solution, will the solubility of PbCl₂ increase, decrease, or remain the same? Explain.

#Answer Key

Multiple Choice:

1. (C) CaF₂ - Fluoride (F⁻) is the conjugate base of a weak acid (HF) and is therefore affected by pH changes.
2. (B) Adding HCl - HCl will react with OH⁻, decreasing its concentration and shifting the equilibrium towards dissolution.

Free Response Question:

(a) $PbCl_2(s) \rightleftharpoons Pb^{2+}(aq) + 2Cl^-(aq)$ (1 point)

(b) Let 's' be the molar solubility of PbCl₂. Then, $[Pb^{2+}] = s$ and $[Cl^-] = 2s$. $K_{sp} = [Pb^{2+}][Cl^-]^2 = s(2s)^2 = 4s^3 = 1.6 \times 10^{-5}$. Solving for s, $s = \sqrt[3]{\frac{1.6 \times 10^{-5}}{4}} = 0.0159 M$. Therefore, $[Pb^{2+}] = 0.0159 M$ (3 points: 1 for setup, 1 for calculation, 1 for correct units)

(c) The solubility of PbCl₂ will decrease. Adding HCl introduces a common ion (Cl⁻), which shifts the equilibrium to the left, reducing the solubility of PbCl₂. (2 points: 1 for correct prediction, 1 for explanation)

(d) The solubility of PbCl₂ will remain the same. Adding NaOH will increase the concentration of OH- ions, which do not react with either Pb2+ or Cl- ions. Therefore, the solubility of PbCl2 will not be affected. (2 points: 1 for correct prediction, 1 for explanation)

Remember, you've got this! Focus on understanding the concepts, apply Le Chatelier's Principle, and you'll do great. Good luck! 👍