Acid-Base Equilibria

Acid-Base Equilibria

Correlation with Textbook

• Textbook Reference: Chemistry, 5th Edition by Gilbert, Kirss, Foster, Bretz, Davies
• Topics Covered:
  • Weak Acid and Base Ionization (15.4)
  • Polyprotic Acids and Bases (15.6)
  • Acidic and Basic Salts (15.8)
  • Common Ion Effect (16.2)
  • Buffers (16.3)
  • pH Indicators (16.4)
  • Titrations (16.4)

Percent Ionization

• Definition: Degree of dissociation of a substance in a solution.
• Formula:
  $\text{Percent Ionization} = \left( \frac{\text{Amount Ionized}}{\text{Initial Concentration}} \right) \times 100$
• Observation: % ionization decreases as initial concentration of acid increases.

Weak Acids and Bases

• Weak Acid Reaction Example:
  $\text{HNO}<em>2(aq) + \text{H}</em>2\text{O}(l) \rightleftharpoons \text{H}<em>3\text{O}^+(aq) + \text{NO}</em>2^{-}(aq)$
• Acid Dissociation Constant (Ka):
  $K<em>a = \frac{[\text{H}</em>3\text{O}^+][A^-]}{[HA]}$
• Small Ka indicates partial dissociation (weaker acid).
• Weak Base Reaction Example:
  $\text{NH}<em>3(aq) + \text{H}</em>2\text{O}(l) \rightleftharpoons \text{NH}_4^+(aq) + \text{OH}^-(aq)$
• Base Dissociation Constant (Kb):
  $K_b = \frac{[BH^+][OH^-]}{[B]}$

Relationship between Ka and Kb

• Water Autoionization:

$K<em>w = K</em>a \times K_b$

where $K_w$ is the ion product of water.

Acid and Base Dissociation Constants at 25 °C

• Example of Weak Acids
  

• Definition: Can donate multiple protons (e.g., H2SO4, H2CO3).
• pH Calculation: Usually based on the first ionization step (dominates pH).
• Diprotic Example:
  • Sulfuric Acid:
    \text{H}2\text{SO}4(aq) \rightleftharpoons \text{H}^+(aq) + \text{HSO}4^{-}(aq), \; K{a1} >> 1

Acid-Base Properties of Salts

• Weak Acid Example:
  • Reaction:
    $\text{HF}(aq) + \text{H}<em>2\text{O}(l) \rightleftharpoons \text{F}^{-}(aq) + \text{H}</em>3\text{O}^+(aq)$
  • Salt:
    $\text{NaF} \rightarrow \text{Na}^+(aq) + \text{F}^{-}(aq)$
  • F- acts as a weak base.
• Predicting pH:
  • Identify if cations and anions are acidic, basic, or neutral.

Common Ion Effect

• Definition: Shift in equilibrium due to addition of an ion involved in the reaction.
• Example:
  $\text{H}<em>2\text{CO}</em>3(aq) \rightleftharpoons \text{H}^+(aq) + \text{HCO}<em>3^{-}(aq) \quad \text{Adding } ext{NaHCO}</em>3 \Rightarrow ext{shifts left}$

Buffers

• Composition: Mixture of weak acid and its conjugate base.
• Function: Resists pH changes upon addition of small amounts of acid or base.
• Buffer Range:
  • $\text{Buffer Range} = pK_a \pm 1$

Titrations

• Definition: Process of adding a solution of known concentration to determine the concentration of an unknown solution.
• Equivalence Point: Stoichiometric amounts of acid and base have reacted.
• Strong Acid-Strong Base Titration Example:
  $\text{HCl}(aq) + \text{NaOH}(aq) \rightleftharpoons ext{H}_2\text{O}(l) + \text{NaCl}(aq)$
• pH Calculation: At equivalence point, pH = 7.00.

Henderson-Hasselbalch Equation

• Equation:
  $\text{pH} = \text{pK}_a + \log \left( \frac{[A^-]}{[HA]} \right)$
• Application: Useful for buffers and estimating pH.

Practice Problems

• Calculate Ka and degree of ionization for various acids.
• Determine pH for given concentrations of salts and titrant solutions.
• Use Henderson-Hasselbalch equation to estimate pH for buffer solutions.