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:

   • $O<em>3 o O</em>2 + O$

   • Rate: $rate<em>A = k</em>A [O_3]^1$ (unimolecular, molecularity = 1)

2. Bimolecular Step B:

   • $NOCl + O o products$

   • Rate: $rate<em>B = k</em>B [NOCl]^1 [O]^1$ (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:

  • $rate<em>{forward} = K</em>1 [I_2]^1$

  • $rate_{reverse} = K^{-1} [I]^2$

• 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.