Chemical Equilibria: Acid-Base & Precipitation Reactions

Chemical Equilibria: Acid-Base & Precipitation Reactions

Chapter Objectives

  • Common Ion Effect (15.1): Understanding the shift in equilibrium due to the addition of a common ion.
  • Buffers (15.2-15.3): Definition and role of buffers in pH control and equilibrium.
  • Titration Curves (15.4): Evaluating the pH changes during titration.
  • Indicators (15.5): Using indicators to determine end points in titrations.

Acid-Base Reactions

  • Types of Reactions:
    • Strong Acid + Strong Base: e.g., HCl + NaOH → NaCl + H₂O.
    • Strong Acid + Weak Base: e.g., HCl + NH₃ → NH₄Cl.
    • Weak Acid + Strong Base: e.g., HOAc + NaOH → NaOAc + H₂O.
    • Weak Acid + Weak Base: e.g., HOAc + NH₃ →
pH Changes During Reactions
  • Analyze initial, during, and final pH for given reactions.
  • Important concepts include Common Ion Effect and Buffer Strength.

Common Ion Effect (Section 15.1)

  • Example: Addition of NH₄Cl to NH₃ solution.
    • The reaction: NH₃ + H₂O ⇄ NH₄⁺ + OH⁻. Adding NH₄⁺ shifts the equilibrium to the left, decreasing [OH⁻] and thus lowering pH.
  • Formula and calculations for pH in equilibrium scenarios.

Buffers (Section 15.2)

  • Function: Buffers resist changes in pH when acids or bases are added.
  • Composition: A weak acid and its conjugate base (e.g., HOAc/OAc⁻).
  • Example: The effect of adding OH⁻ to a buffer of acetic acid and acetate, demonstrating buffer action and reaction with added OH⁻ to maintain pH.
Henderson-Hasselbalch Equation
  • Relates pH, pKa, and concentrations of acid (HA) and conjugate base (A⁻):
    • pH = pKa + log([A⁻]/[HA]).

Titration and pH Curves (Chapter 15.4)

  • Understanding Titration Curves: Analyze how pH changes during titration of weak and strong acids/bases.
  • Equivalence Point: The point at which the amount of titrant added is enough to completely neutralize the analyte in the solution.
  • Halfway Point:
    • For weak acid titrations, pH = pKa.

Acid-Base Indicators (Section 15.5)

  • Role: Weak acids that change color depending on the pH of the solution, used to determine the equivalence point in titrations.
  • Key Principle: Choose indicators with pKa values close to the pH at the equivalence point.

Chapter 16: Solubility and Complex Ion Equilibria

  • Ksp: The solubility product constant, used to describe the saturation of slightly soluble salts.
  • Common Ion Effect: Adds a common ion to the solution, affecting the solubility of salts.
  • Example: Ksp calculations for salts like AgCl and their reactions with common ions.
Complex Ions
  • Complex ions are formed from the reaction of transition metals with ligands (Lewis bases) which increase solubility.
Solubility in Acidic Solutions
  • pH can greatly affect the solubility of salts; for instance, compounds with basic anions are more soluble in acidic conditions.
Practice Problems
  • Examples on calculating pH of buffered solutions, calculating Ksp values, and predicting changes in solubility due to common ions are included in practice sections.
Summary
  • Understanding the principles of acid-base equilibria, the behavior of buffers, titration functions, and the impact of solubility product constants is crucial for predicting and manipulating chemical behaviors in solutions.