CH 12

Reaction Mechanisms

Introduction to Reaction Mechanisms

Reaction mechanisms describe the step-by-step sequence of elementary reactions by which an overall chemical change occurs. Understanding mechanisms helps explain how reactants are converted into products and how reaction rates are determined.

• Overall Reaction: The net chemical equation showing reactants and products.

• Elementary Steps: Simpler reactions that occur as part of the overall mechanism.

• Intermediates: Species produced in one step and consumed in another; do not appear in the overall reaction.

Breaking Down Reactions into Elementary Steps

Complex reactions can be broken down into a series of elementary steps. Each step represents a single molecular event.

• Example: Overall: NO2 + CO → NO + CO2 Mechanism:

  1. NO2 + NO2 → NO3 + NO (slow)

  2. NO3 + CO → NO2 + CO2 (fast)

• Intermediates: NO3 is produced in the first step and consumed in the second.

Elementary Steps and Rate Laws

Each elementary step has its own rate law, which is determined by the molecularity (number of reactant molecules involved).

Elementary Step	Rate Law
A → something	Rate = k[A]
2A → something	Rate = k[A]^2
A + B → something	Rate = k[A][B]
2A + B → something	Rate = k[A]^2[B]

• Key Point: The rate law for an elementary step is directly related to the stoichiometry of the reactants in that step.

• Third-order reactions: Steps involving three reactant molecules are rare but possible.

Rate-Determining Step

The slowest elementary step in a reaction mechanism is called the rate-determining step. The overall rate law is determined by this step.

• Example: NO2 + NO2 → NO3 + NO (slow) NO3 + CO → NO2 + CO2 (fast) Rate Law: $\text{Rate} = k[\text{NO}_2]^2$

• Important: Intermediates cannot appear in the rate law for the overall reaction.

Plausibility of Reaction Mechanisms

A proposed mechanism must satisfy certain criteria to be considered plausible:

• Elementary steps must add up to the overall balanced equation.

• No elementary step should involve more than three molecules (termolecular steps are rare).

• The rate law derived from the mechanism must match experimental evidence.

Examples of Mechanism Analysis

Consider the reaction: SO2 + SO3 → S2O + 2 O2

• Proposed mechanism:

  1. SO2 + SO2 → S2O + O3 (slow)

  2. SO3 + O3 → 2 O2 + SO2 (fast)

• Analysis:

  • Does not add up to the overall reaction.

  • The suggested rate law does not match the evidence.

  • An elementary step involves more than three molecules.

  • The mechanism is not plausible.

Identifying Intermediates

Intermediates are produced in one step and consumed in another. They do not appear in the overall reaction.

• Example: In the mechanism above, O3 is an intermediate.

Catalysts and Their Role in Reaction Mechanisms

Definition and Function of Catalysts

A catalyst is a substance that increases the rate of a chemical reaction without being consumed in the overall process. Catalysts work by lowering the activation energy, making it easier for reactants to convert to products.

• Effect on Reaction:

  • Speeds up reaction by increasing the rate constant (k).

  • Lowers activation energy ($E_a$).

  • Does not change the amount of product formed.

• Equation: $k = A e^{-E_a/RT}$

Types of Catalysts

• Homogeneous Catalyst: Catalyst is in the same phase as the reactants (e.g., all are in solution).

• Heterogeneous Catalyst: Catalyst is in a different phase (e.g., a solid catalyst with gaseous reactants).

Catalysts in Reaction Mechanisms

In a mechanism, a catalyst is consumed in one step and regenerated in another. It appears in the mechanism but not in the overall reaction.

• Example: Reaction: O2 + O → 2 O2 Mechanism:

  1. Cl + O2 → ClO + O

  2. ClO + O → Cl + O2

  Cl is the catalyst; ClO is an intermediate.

Distinguishing Intermediates and Catalysts

• Intermediates: Produced first, then consumed; appear on the right first in the mechanism.

• Catalysts: Consumed first, then produced; appear on the left first in the mechanism.

Practice and Review Questions

Effect of Temperature and Concentration on Reaction Rate

• Increasing temperature increases the rate constant (k) and thus the reaction rate.

• Increasing the concentration of reactants increases the reaction rate.

• Lowering the activation energy increases the rate.

Vapor Pressure and Solutions

• Adding an ionic compound to a solution lowers the vapor pressure.

• Lowering the temperature of the solution lowers the vapor pressure.

• Raising external pressure does not affect vapor pressure directly.

Summary Table: Elementary Steps and Rate Laws

Elementary Step	Rate Law
A → something	Rate = k[A]
2A → something	Rate = k[A]^2
A + B → something	Rate = k[A][B]
2A + B → something	Rate = k[A]^2[B]

Key Point: The order of the reaction in an elementary step matches the number of reactant molecules.

Additional info: These notes are suitable for introductory college-level chemistry and general biology students studying chemical kinetics and reaction mechanisms.