Chemistry Chapter 13: Fundamental Equilibrium Concepts

Chapter 13: Fundamental Equilibrium Concepts

13.1 Chemical Equilibrium

  • Definition: Chemical equilibrium is the state in which the concentrations of reactants and products remain constant over time in a closed system.
    • Example: For the reaction N2O4 (g) ⇌ 2NO2 (g)
    • N2O4 is colorless, while NO2 is brown.
    • At equilibrium, the concentrations of N2O4 and NO2 no longer change.
  • Dynamic Nature: Equilibrium is not static. Reactions continue to occur in both the forward and reverse directions at equal rates.

13.2 Equilibrium Constants

  • Equilibrium Constant (K): Represents the ratio of the concentrations of products to reactants at equilibrium.
    • Notation: K = [Products]^(coefficients) / [Reactants]^(coefficients)
    • Units: K is dimensionless; depends on the reaction.
  • Reaction Quotient (Q): A measure of concentrations of reactants and products at any point in time during the reaction.
    • Can be calculated before reaching equilibrium.
  • Comparison of Q and K: Determines the direction of the reaction to reach equilibrium:
    • If Q < K, the reaction shifts to the right (toward products).
    • If Q > K, the reaction shifts to the left (toward reactants).
    • If Q = K, the system is at equilibrium.

13.3 Shifting Equilibria: Le Châtelier’s Principle

  • Principle Overview: If a system at equilibrium is subjected to change (stress), the system adjusts to counteract that change and restore balance.
  • Types of Disturbances:
    • Adding/Removing a Reactant/Product: A system will shift to consume the added species or to replenish a removed species.
    • Changing Temperature: An increase in temperature shifts toward the endothermic direction (absorbing heat), while a decrease shifts toward the exothermic direction (releasing heat).
    • Adding/Removing a Pure Solid or Liquid: Generally has no effect unless the substance is completely removed from the mixture.
  • Catalysts: They speed up both the forward and reverse reactions but do not change the position of equilibrium or the equilibrium constant.

13.4 Equilibrium Calculations

  • Types of Calculations:
    • Equilibrium Constant: Calculate K using equilibrium concentrations.
    • Finding Missing Concentrations: Use stoichiometry and ICE (Initial, Change, Equilibrium) tables.
    • Initial Conditions: Calculating equilibrium from initial concentrations often requires setting up an algebraic equation.
  • Example Calculations:
    • Set up ICE tables to track concentration changes from initial values to equilibrium states for a variety of reactions.

Examples:

  • Example 13.6: Calculate changes in concentrations using stoichiometric coefficients.
  • Example 13.8 & 13.9: Further examples illustrating varying initial concentrations and their impacts on equilibrium calculations.

Summary of Key Concepts

  • An equilibrium system is dynamic and maintains constant concentrations of reactants and products.
  • Equilibrium calculations involve understanding how to use K and Q to predict reaction behavior under various conditions.
  • Le Châtelier’s Principle provides a framework for predicting how equilibria respond to disturbances in a system.