Applications of Aqueous Equilibria

Overview

• Covers OpenStax Chapters 14.6 to 14.7 and Chapter 15.
• Focuses on buffers, common ion effect, solubility products.

Common Ion Effect

• Definition: Addition of a substance containing an ion in an equilibrium expression causes a shift according to Le Chatelier's Principle.
• Example 1: Solubility of Sodium Chloride
  • Dissolves into sodium ions and chloride ions.
  • Solubility limit: point at which ions re-precipitate as fast as they dissolve.
  • Adding potassium chloride (another chloride source) decreases solubility.
• Example 2: Aluminum Hydroxide Solubility
  • Dissolves poorly; addition of HCl affects solubility.
  • HCl adds hydrogen ions which neutralize hydroxide, increasing solubility.

Ionization of Weak Acids

• Example: Hydrofluoric Acid (HF) in presence of Sodium Fluoride (NaF)
  • NaF provides fluoride ions, shifts equilibrium, decreases HF ionization.
  • Results in higher pH (less acidic environment).
• Calculation: pH and percent ionization can be calculated using ICE tables and equilibrium expressions.

Buffered Solutions

• Definition: Solution with a weak acid and its conjugate base.
• Requirements:
  • Sufficient quantities and close to equal concentrations.
  • Acts to minimize pH changes.
• Buffer vs. Acid/Base:
  • Adding H+ shifts equilibrium, increasing HA and reducing A-.
  • Adding OH- shifts equilibrium, decreasing HA and increasing A-.

Henderson-Hasselbalch Equation

• Equation: pH = pKa + log([A-]/[HA])
• Used for buffered solutions; assumes no further ionization occurs.
• Provides a simplified method for calculating pH in buffered systems.

Solubility Product (Ksp)

• Solubility: Maximum amount that can dissolve, given as molar solubility (moles/liter).
• Solubility Product (Ksp): Equilibrium constant specific to dissolution of ionic compounds.
• Calculations: Relate Ksp to molar solubility; different ionic compounds may require different approaches.

Comparing Solubility

• Can compare Ksp directly if compounds dissociate into same number of ions.
• If different, must calculate molar solubility to compare directly.

Titration Curves

• Calculating pH at any point on a titration curve involves detailed steps (covered in Experiment 5 notes).
• Relates to identifying equivalence points and choosing appropriate indicators.