Notes on Chemical Kinetics

Chemical Kinetics

Section 14.1 Rates of Reactions

  • Definition: The rate of a chemical reaction is defined as the change in concentration of a reactant or product per unit time.
  • Measurement: Reaction rates can be measured via:
    • Disappearance of a reactant over time.
    • Appearance of a product over time.
  • Units: Common units include:
    • M/h for fast reactions.
    • mM/h or μM/h for slower reactions.

Changing Reaction Rates

  • Reaction rate can be influenced by several factors including:
    • Particle Size: Smaller solid particles increase surface area, enhancing contact and thus the reaction rate.
    • Concentration: Higher concentration of reactants leads to more frequent collisions, increasing the reaction rate.
    • Temperature: Increased temperature provides more kinetic energy, leading to more effective collisions.
    • Nature of Reactants: Varies in terms of potential energy, affecting the reaction kinetics.
    • Catalysts: Lower the activation energy required for the reaction, speeding up the process.

Measuring Reaction Rates

  • Mathematical expression:
    • For a simple reaction A → B, the rate can be expressed as:
    • rate = -Δ[A]/Δt = Δ[B]/Δt
  • Commonly measured in M/s or the equivalent in other time units, depending on the reaction speed.

Stoichiometry and Reaction Rates

  • Stoichiometric Relationships: The rate of reactions is directly related to the stoichiometry of the balanced chemical equation:
    • For example, in A → B, the rate of disappearance of A is equal to the appearance of B (1:1 relationship).
    • In reactions like 2 A → B, the rate of disappearance of A is twice that of the appearance of B.
  • General form for a balanced reaction aA → bB:
    • rate = -1/a Δ[A]/Δt = 1/b Δ[B]/Δt

Section 14.2 Reaction Rates and Concentration: Rate Laws

  • Rate Law: A mathematical relationship that describes how the rate of a reaction depends on the concentrations of the reactants.
  • Form: rate = k[A]^m[B]^n, where k is the rate constant, and m and n are the reaction orders for reactants A and B, respectively.
  • Order of Reaction:
    • Zero Order: rate is constant, independent of the concentration of reactants.
    • First Order: rate is directly proportional to the concentration of one reactant.
    • Second Order: rate is proportional to the square of the concentration.

Graphical Determination of Reaction Order

  • To find the order of reaction graphically:
    • Zero Order: A plot of [A] vs. time is linear.
    • First Order: A plot of ln[A] vs. time is linear.
    • Second Order: A plot of 1/[A] vs. time is linear.

Section 14.4 Reaction Rates and Temperature: Activation Energy

  • Activation Energy (Ea): The minimum energy required for reactants to form products. It represents an energy barrier for the reaction.
  • Collision Theory: Suggests that for a reaction to occur, reactant molecules must collide with sufficient energy and the proper orientation.
  • Arrhenius Equation: Describes how the rate constant (k) is affected by temperature and activation energy.

Section 14.5 Reaction Mechanisms

  • Reaction Mechanism: Detailed stepwise pathway of a chemical reaction. Each step is called an elementary step, which may involve intermediate species.
  • Rate-Determining Step: The slowest step in the reaction mechanism that determines the overall rate of reaction.
    • The rate law for the overall reaction reflects the stoichiometry of this step.

Section 14.6 Catalysis

  • Catalysts: Substances that speed up reactions without undergoing permanent changes themselves.
    • Provide an alternate pathway for a reaction, often with lower activation energy.
  • Types of Catalysts:
    • Homogeneous Catalysts: Same phase as the reactants (e.g., gas-liquid systems).
    • Heterogeneous Catalysts: Different phase (e.g., solid catalysts in liquid reactions).
  • Importance: Catalysts are crucial in many biological processes and industrial applications.