Ch.12.6 - Chemical Reaction Mechanisms Overview

Overview

• Defines reaction mechanisms as stepwise sequences of elementary reactions describing how overall reactions occur.
• Emphasizes that balanced overall equations show what reacts, not how; mechanisms reveal step-by-step processes.
• Connects elementary reaction molecularity to rate laws; rate laws for elementary steps can be written directly from their equations.

Learning Objectives

• Distinguish net reactions from elementary reactions (steps).
• Identify molecularity of elementary reactions.
• Write balanced overall equations from mechanisms.
• Derive rate laws consistent with a proposed mechanism.

Key Concepts

• Elementary Reaction: A single-step event that occurs exactly as written; its stoichiometry gives its rate law.
• Intermediate: Species produced in one step and consumed in a later step; not present in the overall equation.
• Molecularity: Number of reactant entities in an elementary step (unimolecular, bimolecular, termolecular).
• Rate-Determining Step: Slowest step of a multi-step mechanism that limits the overall reaction rate.
• Reversible Elementary Step: Can reach equilibrium; useful for expressing intermediate concentrations in terms of reactants.

Types Of Elementary Reactions

• Unimolecular
  • Involves one reactant entity decomposing or isomerizing.
  • Rate law: rate = k[A] (first order).
  • Example: O3 → O2 + O (part of ozone decomposition); cyclobutane decomposition to ethylene (can be unimolecular overall).
• Bimolecular
  • Involves collision of two entities (A + B or A + A).
  • Rate laws:
    • A + B → products: rate = k[A][B] (second order overall).
    • A + A → products: rate = k[A]^2 (second order).
  • Example: NO2 + CO → NO + CO2 can be a single-step bimolecular mechanism at high T.
• Termolecular
  • Simultaneous collision of three entities; rare due to low probability.
  • Rate laws include third-order dependences, e.g., rate = k[NO]^2[O2].
  • Examples: certain steps in NO + O2 and NO + Cl2 reactions.

Relating Mechanisms To Rate Laws

• If the slow (rate-determining) step is the first step, overall rate law equals that step’s rate law.
• If a slow step is preceded by a fast equilibrium, express intermediate concentration via the equilibrium expression, then substitute into rate law for slow step.
• Rate laws for overall reactions usually require experimental determination; mechanisms are then proposed to match observed rate law.

Example: Ozone Decomposition (Mechanistic Illustration)

• Two-step mechanism:
  1. O3 → O2 + O (elementary)
  2. O + O3 → 2 O2 (elementary)
• Intermediate: atomic O (produced in step 1, consumed in step 2).
• Overall: 2 O3 → 3 O2 (sum of steps).
• Demonstrates that overall stoichiometry does not imply direct collision of two O3 molecules.

Example: Temperature-Dependent Mechanisms (NO2 + CO)

• Above 225 °C: observed rate law rate = k[NO2][CO]
  • Consistent with a single-step bimolecular mechanism.
• Below 225 °C: observed rate law rate = k[NO2]^2
  • Not consistent with single-step mechanism; consistent with two-step mechanism:
    1. NO2 + NO2 ⇌ NO3 + NO (slow)
    2. NO3 + CO → NO2 + CO2 (fast)
  • Slow step controls rate (second-order in NO2).

Worked Example: Deriving Rate Law From Mechanism (NO + Cl2 → NOCl)

• Proposed two-step mechanism:
  1. NO + Cl2 ⇌ NOCl2 (fast equilibrium)
  2. NOCl2 + NO → 2 NOCl (slow, rate-determining)
• Overall equation: 2 NO + Cl2 → 2 NOCl
• Rate expressions:
  • Step 1 forward: rate1f = k1[NO][Cl2]
  • Step 1 reverse: rate1r = k-1[NOCl2]
  • Step 2: rate2 = k2[NOCl2][NO]
• Assume step 1 at equilibrium: k1[NO][Cl2] = k-1[NOCl2]
  • Therefore [NOCl2] = (k1/k-1)[NO][Cl2]
• Substitute into step 2 rate:
  • rate = k2[NO][NOCl2] = (k2·k1/k-1)[NO]^2[Cl2]
• Predicted overall rate law: rate = k'[NO]^2[Cl2] where k' = k2·k1/k-1

Check-Your-Learning Example (Fluorine Dimerization Equilibrium)

• Fast equilibrium step: F2 ⇌ 2 F
• At equilibrium: k1[F2] = k-1[F]^2
• Solve for [F]: [F] = sqrt((k1/k-1)[F2])

Key Terms And Definitions

• Elementary Reaction: Single-step reaction with molecularity equal to number of reactant entities.
• Intermediate: Transient species formed and consumed within mechanism steps.
• Molecularity: Number of molecules/atoms/ions colliding in an elementary step.
• Rate-Determining Step: The slowest step controlling overall reaction rate.
• Fast Equilibrium: A reversible step that rapidly reaches equilibrium relative to other steps.

Action Items / Study Tips

• Practice deriving rate laws from proposed mechanisms, especially when fast pre-equilibria exist.
• Distinguish when observed rate laws support single-step elementary mechanisms versus multistep mechanisms.
• Identify intermediates in mechanisms and eliminate them using equilibrium relations when needed.
• Memorize typical rate law forms for unimolecular, bimolecular, and termolecular elementary steps.

Summary Table: Elementary Reaction Types