Acid-Base Chemistry, Buffers, Titrations, and Solubility Equilibria: Study Notes

Acid-Base Chemistry

Introduction

Acid-base chemistry is a fundamental topic in general chemistry, focusing on the properties, calculations, and behaviors of acids, bases, and their solutions. Understanding how to calculate pH, use equilibrium constants, and distinguish between strong and weak acids and bases is essential for mastering this area.

• Solving for pH of Weak Base: Use the base dissociation constant ($K_b$) and an ICE table to determine the concentration of $OH^-$, then calculate pOH and pH.

• Solving for pH of Strong Base: Strong bases dissociate completely. Calculate $[OH^-]$ directly from the base concentration, then use $pOH = -\log[OH^-]$ and $pH = 14 - pOH$.

• Solving for pH of Salt: Determine if the salt produces acidic, basic, or neutral solutions by considering the hydrolysis of its ions.

• Qualitatively Assessing pH of Salt: Salts from strong acid and strong base are neutral; from strong acid and weak base are acidic; from weak acid and strong base are basic.

• Strength of Acids and Bases: Strong acids/bases dissociate completely; weak acids/bases only partially dissociate.

• Solving for $K_a$ if Given $K_b$ (and vice versa): Use the relationship $K_a \times K_b = K_w$, where $K_w = 1.0 \times 10^{-14}$ at 25°C.

• Knowing Which to Use: $K_a$ or $K_b$: Use $K_a$ for acids and $K_b$ for bases, depending on the species in question.

• pKa and pKb: $pK_a = -\log K_a$ and $pK_b = -\log K_b$; lower values indicate stronger acids/bases.

• Polyprotic Acids: Acids that can donate more than one proton (e.g., $H_2SO_4$); each dissociation step has its own $K_a$.

• pOH: $pOH = -\log[OH^-]$; $pH + pOH = 14$ at 25°C.

Example:

Calculate the pH of a 0.10 M solution of acetic acid ($K_a = 1.8 \times 10^{-5}$):

• Set up ICE table for $CH_3COOH \rightleftharpoons CH_3COO^- + H^+$

• Solve for $[H^+]$ using $K_a$ expression

• Calculate $pH = -\log[H^+]$

Buffers

Introduction

Buffers are solutions that resist changes in pH upon the addition of small amounts of acid or base. They are essential in many chemical and biological systems.

• Buffer Components: Typically consist of a weak acid and its conjugate base, or a weak base and its conjugate acid.

• Common Ion Effect: The suppression of ionization of a weak acid or base by the presence of a common ion from a strong electrolyte.

• Purpose of a Buffer: To maintain a relatively constant pH in a solution.

• Calculate the pH of a Buffer: Use the Henderson-Hasselbalch equation: $pH = pK_a + \log\left(\frac{[A^-]}{[HA]}\right)$

• Henderson-Hasselbalch Equation: Relates pH, $pK_a$, and the ratio of conjugate base to acid.

• Effect of Adding $H_3O^+$ or $OH^-$ to a Buffer: The buffer neutralizes added acid or base, minimizing pH change.

• Buffer Range: The pH range over which a buffer is effective, typically $pK_a \pm 1$.

• Buffer Capacity: The amount of acid or base a buffer can neutralize before a significant pH change occurs.

• Preparing a Buffer: Mix a weak acid with its salt (conjugate base) or a weak base with its salt (conjugate acid).

Example:

To prepare a buffer with pH 4.75 using acetic acid ($pK_a = 4.75$) and sodium acetate, mix equal concentrations of both.

Acid-Base Titrations

Introduction

Acid-base titrations are analytical techniques used to determine the concentration of an acid or base by reacting it with a standard solution of known concentration.

• Analyte: The solution of unknown concentration being analyzed.

• Titrant: The solution of known concentration added to react with the analyte.

• Indicators: Substances that change color at (or near) the equivalence point.

• Strong Acid-Strong Base Titration: The equivalence point occurs at pH 7.

• Weak Acid-Strong Base Titration: The equivalence point occurs at pH > 7 due to the formation of a weak base.

• Weak Base-Strong Acid Titration: The equivalence point occurs at pH < 7 due to the formation of a weak acid.

• Polyprotic Acid Titration: Multiple equivalence points corresponding to each proton donated.

• Amino Acids as Polyprotic Acids: Amino acids can act as acids and bases, with multiple $pK_a$ values.

• Four Regions of a Titration Curve:

  1. Initial

  2. Before equivalence point (buffer region)

  3. Equivalence point

  4. After equivalence point

Example:

Titrating 25.0 mL of 0.10 M acetic acid with 0.10 M NaOH: The pH at the equivalence point is greater than 7 due to the formation of acetate ion, a weak base.

Equilibria of Slightly Soluble Compounds

Introduction

The solubility of ionic compounds in water is governed by equilibrium principles. The solubility product constant ($K_{sp}$) quantifies the extent to which a compound dissolves.

• $K_{sp}$: The equilibrium constant for the dissolution of a sparingly soluble salt. For $AB \rightleftharpoons A^+ + B^-$, $K_{sp} = [A^+][B^-]$.

• Common Ion Effect (Le Châtelier's Principle): The solubility of a salt decreases in the presence of a common ion.

• Effect of pH on Solubility: The solubility of salts containing basic anions increases in acidic solutions.

• When Will the Solution Precipitate? ($Q_{sp}$ vs $K_{sp}$): If $Q_{sp} > K_{sp}$, precipitation occurs; if $Q_{sp} < K_{sp}$, no precipitation; if $Q_{sp} = K_{sp}$, the solution is saturated.

Example:

Will a precipitate form if $[Ag^+] = 1.0 \times 10^{-4}$ M and $[Cl^-] = 1.0 \times 10^{-4}$ M? $K_{sp}$ for $AgCl$ is $1.8 \times 10^{-10}$.

• $Q_{sp} = [Ag^+][Cl^-] = (1.0 \times 10^{-4})(1.0 \times 10^{-4}) = 1.0 \times 10^{-8}$

• Since $Q_{sp} > K_{sp}$, a precipitate will form.