CHEM111 Chapter 7 Chemical Reactions and Equilibria (1)

Introduction to Chemical Kinetics and Equilibria

  • Definition of Chemical Kinetics: The study of the rates of chemical reactions.

Reaction Rates

  • Measured by the concentration change of products/reactants over time.

  • Example reaction: Chloromethane and sodium iodide in acetone.

Determining Reaction Rate

  • Reaction rate = Increase in concentration of product (e.g., iodomethane) / Time interval.

  • Example: Concentration increases from 0 to 0.12 mol/L in 30 minutes.

Influencing Factors on Reaction Rates

  • Molecular Collisions: Reactions require collisions between reactants; many do not result in a reaction.

  • Activation Energy: Minimum energy required for a reaction to occur; higher energies lead to reaction.

  • Orientation of Collisions: Must properly orient for effective reactions.

Energy Diagrams in Reactions

  • Exothermic Reactions: Energy released is greater than energy needed to break bonds.

  • Endothermic Reactions: Energy required exceeds the energy released when bonds are formed.

Evaluating Reaction Factors

  1. Nature of Reactants: Fast reactions for ions, slower for covalent compounds.

  2. Concentration: Increased concentration typically increases reaction rates.

  3. Temperature: Increasing temperature raises rate, often doubling for every 10°C increase.

  4. Catalyst Presence: Increases rate without being consumed, helps reach equilibrium faster.

Reversible Reactions and Equilibrium

  • Reversible Reaction: Can proceed in either direction.

  • Equilibrium State: Forward and reverse reaction rates are equal, concentrations remain constant.

Equilibrium Constant (K)

  • Defined by the concentrations of products/reactants at equilibrium raised to their respective coefficients in the balanced equation.

Interpretation of K Values

  • K << 1: Favor reactants.

  • K ~ 1: Significant amounts of both reactants and products.

  • K >> 1: Favor products.

Le Chatelier's Principle

  • States that if an external stress is applied, the equilibrium will shift to counteract the change.

  1. Adding/removing reactants/products: Changes concentration.

  2. Changes in temperature: Affects forward/reverse reactions based on being exothermic/endothermic.

  3. Changes in pressure: Affects moles of gas present.

  4. Catalysts: Speed up reactions without changing equilibrium positions.

Practice Problems

  1. Calculate reaction rates given specific conditions (e.g., volume of O2 produced).

  2. Writing equilibrium constant expressions based on given reactions.