Chapter 18: Aqueous Ionic Equilibrium

Introduction

This chapter explores the chemical equilibria that occur in aqueous solutions, focusing on buffer systems, acid-base titrations, solubility equilibria, and the formation of complex ions. Understanding these concepts is essential for predicting and controlling the behavior of ions in solution, which is fundamental in both laboratory and biological contexts.

Buffer Solutions

Buffer solutions are mixtures that resist changes in pH when small amounts of acid or base are added. They are crucial in maintaining stable environments in chemical and biological systems.

• Definition: A buffer contains significant amounts of a weak acid and its conjugate base, or a weak base and its conjugate acid.

• Function: Buffers neutralize added acids or bases, minimizing pH changes.

• Example: Blood acts as a buffer, maintaining pH for proper physiological function.

Making an Acidic Buffer Solution

• Requires both a weak acid and its conjugate base in significant amounts.

• Example: Acetic acid (CH3COOH) and sodium acetate (CH3COONa).

Buffer Action: Addition of Base or Acid

• Added base is neutralized by the weak acid; added acid is neutralized by the conjugate base.

• Le Châtelier’s principle explains how equilibrium shifts to maintain pH.

Common Ion Effect

• Adding a salt containing the conjugate ion shifts equilibrium, affecting pH.

• Example: Adding NaA to a weak acid solution decreases [H+], increasing pH.

Calculating Buffer pH: Henderson–Hasselbalch Equation

• The equation relates pH, pKa, and the ratio of concentrations of acid and conjugate base:

• Valid when the "x is small" approximation holds (initial concentrations much larger than Ka).

Buffer Effectiveness

• Most effective when [acid] ≈ [base].

• Buffer capacity: Amount of acid/base a buffer can neutralize.

• Buffer range: pH range where buffer is effective, typically pKa ± 1.

Titration and Titration Curves

Titration is a technique to determine the concentration of an unknown solution by reacting it with a solution of known concentration.

• Equivalence Point: Moles of acid equal moles of base.

• Indicator: Changes color at a specific pH, marking the endpoint.

Titration Curves

• Plot of pH vs. volume of titrant added.

• Inflection point indicates equivalence point.

• Equivalence point pH depends on the nature of the salt formed:

  • Neutral salt: pH = 7

  • Acidic salt: pH < 7

  • Basic salt: pH > 7

Titration of Weak Acids and Bases

• Initial pH determined by weak acid/base equilibrium.

• Buffer region before equivalence point.

• At equivalence, solution contains only conjugate base/acid.

Acid–Base Indicators

• Indicators are weak acids or bases that change color depending on pH.

• Used to visually determine the endpoint of a titration.

Solubility Equilibria

Solubility equilibria describe the extent to which ionic compounds dissolve in water.

• Solubility Product (Ksp): Equilibrium constant for the dissolution of a solid salt.

• Molar Solubility: Number of moles of solute dissolved per liter of solution.

• Comparing Ksp values is only valid for salts with the same stoichiometry.

Common Ion Effect on Solubility

• Addition of a common ion decreases solubility by shifting equilibrium left.

• Example: Adding NaF to CaF2 solution decreases CaF2 solubility.

Effect of pH on Solubility

• Higher pH decreases solubility of ionic hydroxides.

• Lower pH increases solubility of salts containing anions of weak acids.

Precipitation and Selective Precipitation

• Precipitation occurs when ion concentrations exceed solubility.

• Selective precipitation separates ions by adding reagents that form insoluble salts with specific cations.

Qualitative Analysis

• Uses selective precipitation to identify ions in solution.

• Group separation based on solubility differences.

Complex Ion Formation

• Transition metals form complex ions with ligands (Lewis bases).

• Formation constant (Kf) quantifies stability of complex ions.

• Complex ion formation increases solubility of some salts.

Solubility of Amphoteric Metal Hydroxides

• Amphoteric hydroxides dissolve in both acidic and basic solutions.

• Example: Al(OH)3 becomes more soluble in acid and base.