Lecture Notes on Chemical Equilibrium

Introduction to Equilibrium

• Equilibrium is a state with no observable changes, but reactions still occur.
• Chemical Equilibrium: Achieved when the rate of the forward reaction equals the rate of the reverse reaction.
• Physical Equilibrium: Phase change equilibrium, e.g., evaporation equal to condensation.

Real-World Examples

• Hemoglobin Adaptation: Mountain climbers produce more hemoglobin at high altitudes for better oxygen transport.

Definitions

• Equilibrium Constant (K): Expresses the ratio of product to reactant concentrations at equilibrium.
  • Kc: Uses concentrations in molarity.
  • Kp: Uses pressures.
  • Different constants for different scenarios: e.g., Kc, Kp, KW, KA, KB, KEQ.

Equilibrium Expressions

• Expressions are written based on the balanced chemical equation.
• Products over Reactants: Raised to the power of their stoichiometric coefficients.
• Exclusions: Solids and liquids typically not included due to their large concentrations.

Reaction Examples

• N2O4 and NO2: Common textbook example for equilibrium understanding due to its clear color change and response to temperature.
• Equilibria in Different Conditions: Starting with only NO2, only N2O4, or a mixture yields different equilibrium states.

Understanding Equilibrium Constants

• Magnitude of K: Indicates direction of reaction favorability.
  • K >> 1: Favor products (right side dominance)
  • K << 1: Favor reactants (left side dominance)

Calculations and Applications

• ICE Tables: Initial, Change, Equilibrium - a method to track concentrations throughout a reaction.
• Determining K: Use equilibrium concentrations to compute K using the equilibrium expression.

Homogeneous vs. Heterogeneous Equilibrium

• Homogeneous: All reactants/products in the same phase.
• Heterogeneous: Different phases present, solids/liquids often excluded from expressions.

Example Reactions

• Decomposition Reactions: One reactant forming multiple products, e.g., NH4HS decomposing into NH3 and H2S.

Special Considerations

• Delta N: Change in moles of gases affects the relationship between Kp and Kc.
  • When ΔN = 0, Kp = Kc.
  • ΔN = moles of gaseous products - moles of gaseous reactants.

Practical Calculations

• Equilibrium Constants: Calculated by inserting equilibrium concentrations or pressures into the expression.
• Relation of Kc to Kp: Kp = Kc(RT)^Δn, where Δn is the change in moles of gas.

Important Constants

• Equilibrium Constant: A key concept in understanding chemical systems, related to reaction conditions and concentration.
• Constants in Science: Discovery of constants (e.g., equilibrium constants, Avogadro's number) is crucial for understanding scientific phenomena.