Chemistry_ The Central Science, 14th Edition

Introduction to Aqueous Equilibria

  • Understanding the chemistry of coral reef formation and ocean processes requires knowledge of aqueous equilibria.

  • Focus includes acid-base equilibria, slightly soluble salts, and metal complexes in solution.

The Common-Ion Effect

  • In solutions containing both a weak acid (e.g., acetic acid) and its soluble salt (e.g., sodium acetate), they share a common ion (CH₃COO⁻).

  • Sodium acetate dissociates completely in solution, while acetic acid is a weak electrolyte with partial ionization.

  • Le Châtelier’s Principle: Adding sodium acetate shifts equilibrium in the ionization of acetic acid to the left, reducing [H⁺] concentration.

  • This phenomenon is known as the common-ion effect, where the presence of a common ion decreases the ionization of the weak electrolyte.

Calculating pH with Common Ions

  • Sample Exercise 17.1

    • Mixture: 0.30 M acetic acid and 0.30 M sodium acetate in 1.0 L of solution.

    • Identify strong and weak electrolytes:

      • Strong: Na+ (spectator), CH₃COO⁻ (conjugate base)

      • Weak: CH₃COOH

    • Key Equilibrium Reaction: CH₃COOH ⇌ H⁺ + CH₃COO⁻

    • Impact of Common Ion on Equilibrium:

    • Initial Concentrations:

      • CH₃COOH: 0.30 M

      • H⁺: 0

      • CH₃COO⁻: 0.30 M

    • Changes in Concentrations:

      • CH₃COOH decreases by x

      • H⁺ increases by x

      • CH₃COO⁻ increases by x

    • Final Concentrations:

      • CH₃COOH: 0.30 - x, H⁺: x, CH₃COO⁻: 0.30 + x

    • Using the equilibrium constant (Ka = 1.8 × 10⁻⁵):

      • 1.8 × 10⁻⁵ = (x)(0.30 + x)/(0.30 - x)

Weak Bases and the Common Ion

  • Weak bases also experience decreased ionization from the addition of common ions.

  • Example: The addition of NH₄⁺ from a strong electrolyte (NH₄Cl) shifts equilibrium of NH₃.

Buffers

  • Definition: Solutions with high concentrations of weak acid-base pairs that resist significant pH changes upon the addition of H⁺ or OH⁻.

  • Example: Blood maintains pH around 7.4 via buffers like HCO₃⁻/CO₃²⁻.

  • Composition and Action: Buffers must have roughly equal concentrations of acid and base.

  • Preparation Methods:

    • Mix weak acid with its salt (e.g., CH₃COOH with CH₃COONa).

    • Neutralize weak acid with strong base to form conjugate base (e.g., CH₃COOH with NaOH).

Buffer Calculations

  • Method to calculate pH:

    1. Use the Henderson-Hasselbalch equation: ( pH = pKa + ext{log} \left( \frac{[A^-]}{[HA]} \right) )

  • Example: Calculating pH for lactic acid buffer (0.12 M lactic acid and 0.10 M sodium lactate).

  • Sample Exercise 17.3: pH = pKa + log([C₃H₅O₃⁻]/[C₃H₅O₃H])

Buffer Capacity and pH Range

  • Buffer capacity relates to the amount of acid/base the buffer can neutralize before pH changes appreciably.

  • Effective over a range of pH, ideally within 1 pH unit of pKa.

Acid-Base Titrations

  • Titration is a technique to determine concentrations and equilibria.

  • Types:

    1. Strong Acid-Strong Base: Equivalence point at pH 7.

    2. Weak Acid-Strong Base: pH at equivalence point > 7.

    3. Strong Base-Strong Acid: pH changes rapidly near equivalence point.

  • Observations: pH titration curves represent changes in acidity/alkalinity.

Polyprotic Acid Titrations

  • Polyprotic acids undergo neutralization in steps, each with its own equivalence point.

  • Example: Phosphorous acid (H₃PO₃) neutralizes in two steps: H₃PO₃ + OH⁻ → H₂PO₃⁻ + H₂O.

  • Estimate pKa values from titration curves according to the procedure highlighted.

Conclusion

  • Understanding aqueous equilibria, common-ion effects, buffers, and titrations is key to comprehending chemical behavior in aquatic environments and biological systems.