7.4 The Mechanism of a Chemical Reaction

This topic explains the mechanism of a chemical reaction — the sequence of steps by which reactants are converted into products. It describes how the rate-determining step controls the overall rate and how the rate law provides information about the mechanism.

Definition

Mechanism: The path that reactants take to form products in a chemical reaction. It refers to the sequence of all chemical steps leading from starting materials to products.

Rate-Determining Step: The slowest step in the reaction mechanism, which determines the overall rate of the reaction.

Reaction Intermediate: A species that is produced in one step of a chemical reaction and consumed in another step.

Main Idea

The path that reactants take to form products in a chemical reaction is called the mechanism. A reaction rate equation is very useful because it provides information about the reaction mechanism. The reaction can take place in one step or in several steps. If a reaction proceeds in two or more steps, one of the steps is the slowest. The rate of the slowest step determines the overall reaction rate. This is because it sets a limit to the rate at which the overall reaction can occur. No reaction can proceed more slowly than the rate-determining step. All other steps in the reaction mechanism are usually fast. The slowest step in the reaction mechanism, which determines the overall rate of the reaction, is called the rate-determining step.

Core Concept

• Mechanism: Sequence of steps from reactants to products.
• Rate-determining step: Slowest step; determines overall rate.
• Reaction intermediate: Produced in one step, consumed in another.
• Rate law: Provides information about the rate-determining step.
• Balanced equation: May not give information about the mechanism.

Detailed Explanation

What is a Reaction Mechanism?

• The path that reactants take to form products in a chemical reaction is called the mechanism.
• A reaction rate equation is very useful because it provides information about the reaction mechanism.
• The reaction can take place in one step or in several steps.
• If a reaction proceeds in two or more steps, one of the steps is the slowest.
• The rate of the slowest step determines the overall reaction rate.
• This is because it sets a limit to the rate at which the overall reaction can occur.
• No reaction can proceed more slowly than the rate-determining step.
• All other steps in the reaction mechanism are usually fast.
• The slowest step in the reaction mechanism, which determines the overall rate of the reaction, is called the rate-determining step.

Example 7.4: Reaction of NO₂ with CO

Reaction: NO_2(g) + CO_(g) → NO_(g) + CO_2(g)

Rate Law: Rate = k[NO₂]²

What information do you get from this about the rate-determining step?

Solution:

The rate equation gives us following information:

i) The reaction is second order with respect to NO₂ and zero with respect to CO. Therefore, it is independent of the concentration of CO.

ii) Two molecules of NO₂ are involved in the rate-determining step.

iii) Reaction must proceed in more than one step.

The proposed mechanism for the reaction is as follows:

NO_2(g) + NO_2(g) Slow→ NO_3(g) + NO_(g)

NO_3(g) + CO_(g) Fast→ NO_2(g) + CO_2(g)

The first step is the rate-determining step. Species NO₃ that does not appear in the overall reaction is called a reaction intermediate. This example also proves that a balanced chemical equation may not give any information about the reaction mechanism.

Example 7.5: Decomposition of Hypochlorite Ion

Reaction: 3ClO^-_(aq) → ClO₃^- + 2Cl^-

The rate of the reaction is second order in ClO⁻ ion.

Rate = k[ClO^-]²

The following two-step mechanism is consistent with the rate law for the reaction.

ClO^-_(aq) + ClO^-_(aq) → ClO₂^-_(aq) + Cl^-_(aq)

Overall reaction:

3ClO^-_(aq) → ClO₃^-_(aq) + 2Cl^-_(aq)

Select the rate-determining step.

Solution:

Rate law for this reaction indicates that two ClO⁻ ions must participate in the rate-determining step. Therefore, Step I is the rate-determining step in this mechanism.

Example 7.6: Reaction of NO₂ with O₃

Reaction: 2NO_2(g) + O₃ → N₂O_5(g) + O_2(g)

Rate: Rate = k[NO₂][O₃]

This indicates that one molecule of NO₂ and one molecule of O₃ participate in the determining step.

Thus rate law includes the concentration of each of the reactants raised to the power that equals the coefficient for the reactant in the equation for the rate-determining step.

Concept Assessment Exercise 7.2

For the following reaction: 2NO_2(g) + F_2(g) → 2NO₂F_(g)

Rate = k[NO₂][F₂]

Which of the following mechanism is consistent with the rate law?

(a) 2NO₂ + F₂ → 2NO₂F

(b) NO₂ + F₂ → NO₂F + F (fast step) NO₂ + F → NO₂F (slow step)

(c) NO₂ + F₂ → NO₂F + F (slow step) NO₂ + F → NO₂F (fast step)

(d) F₂ → 2F (slow step) 2NO₂ + 2F → 2NO₂F (fast step)

Example 7.7: Reaction of NO with H₂

Reaction: 2NO_(g) + 2H_2(g) → N_2(g) + 2H₂O_(g)

The mechanism for this reaction involves two steps:

2NO + H₂ → N₂ + H₂O₂ (slow)

Write the experimental rate law for this reaction?

Intermediate Level

Mechanism and Rate-Determining Step

Example 7.4

Reaction: NO₂ + CO → NO + CO₂ Rate: Rate = k[NO₂]²

• Second order in NO₂
• Zero order in CO
• Two NO₂ molecules in rate-determining step
• More than one step

• NO₂ + NO₂ → NO₃ + NO (slow)
• NO₃ + CO → NO₂ + CO₂ (fast)

Intermediate: NO₃

Example 7.5

Reaction: 3ClO⁻ → ClO₃⁻ + 2Cl⁻ Rate: Rate = k[ClO⁻]²

• ClO⁻ + ClO⁻ → ClO₂⁻ + Cl⁻ (slow)
• ClO₂⁻ + ClO⁻ → ClO₃⁻ + Cl⁻ (fast)

Rate-determining step: Step I

Example 7.6

Reaction: 2NO₂ + O₃ → N₂O₅ + O₂ Rate: Rate = k[NO₂][O₃]

Information: - One NO₂ and one O₃ in rate-determining step.

Concept Assessment Exercise 7.2

Reaction: 2NO₂ + F₂ → 2NO₂F Rate: Rate = k[NO₂][F₂]

• NO₂ + F₂ → NO₂F + F (slow)
• NO₂ + F → NO₂F (fast)

Example 7.7

Reaction: 2NO + 2H₂ → N₂ + 2H₂O Mechanism: - 2NO + H₂ → N₂ + H₂O₂ (slow)

**Rate law:** Rate = k[NO]²[H₂]

Important Points

• 1. Mechanism: Sequence of steps.
• 2. Rate-determining step: Slowest step.
• 3. Rate law: Shows rate-determining step.
• 4. Reaction intermediate: Produced and consumed.
• 5. Example 7.4: Rate = k[NO₂]², mechanism with NO₃ intermediate.
• 6. Example 7.5: Rate = k[ClO⁻]², Step I is RDS.
• 7. Example 7.6: Rate = k[NO₂][O₃].
• 8. Concept Assessment 7.2: Mechanism (c).
• 9. Example 7.7: Rate = k[NO]²[H₂].

Remember

• - Mechanism: Step-by-step path
• - Rate-determining step: Slowest step
• - Rate law: Shows RDS
• - Intermediate: Produced and consumed
• - **Example 7.4:** Rate = k[NO₂]²
• - **Example 7.5:** Rate = k[ClO⁻]²
• - **Example 7.6:** Rate = k[NO₂][O₃]
• - **Example 7.7:** Rate = k[NO]²[H₂]

Common Mistakes

• 1. Thinking the balanced equation gives the mechanism: It does not.
• 2. Confusing rate-determining step with overall reaction: RDS is the slowest step.
• 3. Forgetting the reaction intermediate: It is produced and consumed.
• 4. Mixing up rate laws: Use experimental data.

Quick Revision

• Mechanism: Sequence of steps.
• Rate-determining step: Slowest step.
• Rate law: Shows rate-determining step.
• Reaction intermediate: Produced and consumed.
• Example 7.4: Rate = k[NO₂]², mechanism with NO₃ intermediate.
• Example 7.5: Rate = k[ClO⁻]², Step I is RDS.
• Example 7.6: Rate = k[NO₂][O₃].
• Concept Assessment 7.2: Mechanism (c).
• Example 7.7: Rate = k[NO]²[H₂].

60-Second Revision

• Mechanism: Step-by-step path
• Rate-determining step: Slowest step
• Rate law: Shows RDS
• Intermediate: Produced and consumed
• **Example 7.4:** Rate = k[NO₂]²
• **Example 7.5:** Rate = k[ClO⁻]²
• **Example 7.6:** Rate = k[NO₂][O₃]
• **Example 7.7:** Rate = k[NO]²[H₂]

QUALITY VALIDATION

• Topic: 7.4 The Mechanism of a Chemical Reaction
• Chapter: 7 (Chemical Kinetics)

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