Acid-Base Chemistry, Buffers, Acid-Base Titrations, and Equilibria of Slightly Soluble Compounds

Acid-Base Chemistry

Introduction

Acid-base chemistry is fundamental to understanding chemical equilibria, solution properties, and many biological and industrial processes. This section covers calculations and concepts related to acids, bases, and their strengths.

• Solving for pH of Weak Base: Use equilibrium expressions and the base dissociation constant ($K_b$) to determine the pH of weak bases.

• Solving for pH of Strong Base: Strong bases dissociate completely; calculate pOH and convert to pH.

• Solving for pH of Salt: Salts may hydrolyze in water, affecting pH depending on their constituent ions.

• Qualitatively Assessing pH of Salt: Predict whether a salt solution will be acidic, basic, or neutral based on its ions.

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

• Solving for $K_a$ if given $K_b$: Use the relationship $K_a \times K_b = K_w$ where $K_w$ is the ion-product constant for water ($1.0 \times 10^{-14}$ at 25°C).

• Solving for $K_b$ if given $K_a$: Rearranged from above: $K_b = \frac{K_w}{K_a}$.

• 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, such as $H_2SO_4$ or $H_3PO_4$; 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 ($K_a = 1.8 \times 10^{-5}$).

Additional info: For weak acids, set up an ICE table and solve for [H+].

Buffers

Introduction

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

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

• Common Ion Effect: The presence of a common ion suppresses the ionization of a weak acid or base.

• Purpose of a Buffer: To maintain a stable 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, pKa, and the ratio of conjugate base to acid.

• Effect of Adding H3O+ or OH- to a Buffer: The buffer neutralizes added acid or base, minimizing pH change.

• Buffer Range: The effective pH range is typically $pK_a \pm 1$.

• Buffer Capacity: The amount of acid or base a buffer can neutralize before pH changes significantly.

• Preparing a Buffer: Mix a weak acid/base with its conjugate salt in appropriate proportions.

Example: Prepare a buffer with pH 5 using acetic acid ($pK_a = 4.76$) and sodium acetate.

Additional info: Use the Henderson-Hasselbalch equation to determine the required ratio of acid to base.

Acid-Base Titrations

Introduction

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

• Analyte: The substance 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.

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

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

• Amino Acids as Polyprotic Acids: Amino acids can act as polyprotic acids due to their multiple ionizable groups.

• Four Regions of a Titration Curve:

  1. Initial

  2. Before equivalence point (buffer region)

  3. Equivalence point

  4. After equivalence point

Example: Titrate 25.0 mL of 0.100 M acetic acid with 0.100 M NaOH and plot the pH curve.

Additional info: Use stoichiometry and equilibrium calculations for each region.

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 solubility product constant; for $AB \rightleftharpoons A^+ + B^-$, $K_{sp} = [A^+][B^-]$.

• Common Ion Effect (Le Châtelier's Principle): Adding a common ion decreases solubility by shifting equilibrium.

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

• When Will the Solution Precipitate? ($Q$ vs $K_{sp}$):

  • If $Q < K_{sp}$: No precipitation; solution is unsaturated.

  • If $Q = K_{sp}$: Solution is saturated; equilibrium.

  • If $Q > K_{sp}$: Precipitation occurs; solution is supersaturated.

Example: Will a solution containing 0.01 M $Ag^+$ and 0.01 M $Cl^-$ precipitate if $K_{sp}$ for $AgCl$ is $1.8 \times 10^{-10}$?

Additional info: Calculate $Q = [Ag^+][Cl^-]$ and compare to $K_{sp}$.