Reaction Mechanisms

Definition of Reaction Mechanisms

  • Reaction mechanisms consist of multiple steps that detail how a reaction progresses.

  • Each step is known as an elementary step, reflecting actual molecular collisions in a specific order.

  • The coefficients of reactants in elementary steps correspond to the exponents in rate expressions, as they statistically represent collisions.

Differences in Elementary Steps vs. Overall Reaction

  • The elementary steps often differ from the overall stoichiometry, but the sum of the steps equals the overall reaction.

  • If the sum does not equal, the proposed mechanism is incorrect.

Key Definitions

  • Molecularity: The number of molecules involved in an elementary step.

    • Unimolecular: One molecule (vibrates apart).

    • Bimolecular: Two molecules collide (most common).

    • Trimolecular: Rare, involves three colliding molecules simultaneously.

  • Intermediates: Species formed during the mechanism and not present at the end products, detectable through spectroscopy.

  • Catalysts: Increase the rate of reaction by lowering activation energy without being consumed.

    • Recognizable at both the start and end of a reaction, similar to intermediates but involved in the overall equation.

Rate Determining Step (RDS)

  • The RDS is the slowest elementary step, determining the overall reaction rate.

  • Comparative analogy: Traffic flow—slowest part of the journey influences total time.

    • Must match the experimental rate law for the overall reaction.

Experimentally Determining Rate Law

  1. Determine the rate law experimentally (isolation method).

  2. Propose a mechanism based on those findings.

  3. Test the mechanism—there must be a step that matches the experimental rate law.

Catalysis Details

  • Catalysis: All catalysts facilitate reactions by lowering activation energy.

    • Different from simply raising temperature (which increases energy but does not affect the barrier itself).

  • Types of Catalysis:

    • Heterogeneous Catalysis: Involves different phases, e.g., metal catalysts in gas reactions (e.g., ammonia production).

    • Homogeneous Catalysis: All reactants in the same phase, e.g., enzyme-substrate interactions in biochemistry.

Examples of Elementary Steps and Rate Laws

  1. Unimolecular Step A:

    • $O3 o O2 + O$

    • Rate: $rateA = kA [O_3]^1$ (unimolecular, molecularity = 1)

  2. Bimolecular Step B:

    • $NOCl + O o products$

    • Rate: $rateB = kB [NOCl]^1 [O]^1$ (bimolecular, molecularity = 2)

Example Mechanism

  • Isomerization of Cyclopropane:

    • Slow Step: Rate corresponds to the rate law.

    • Rate expression for the slow step equals the overall rate.

Final Mechanism Example

  • Hydrogen-Bromide Reaction:

  • Step 1: Breaks I2 into 2 I (fast)

  • Step 2: Combines 2 I with H2 (slow)

    • Rate law derived must match experimental parameters.

Transition to Equilibrium Concepts

  • Equilibrium represents a dynamic balance where the forward and reverse reactions occur at the same rate.

  • Formulas:

    • $rate{forward} = K1 [I_2]^1$

    • $rate_{reverse} = K^{-1} [I]^2$

  • Setting these equal leads to insight about reaction dynamics and solving for intermediates in the expressions.

End note: Prepare for upcoming chapters focusing on equilibrium concepts.