Reaction Mechanisms

• The chance of more than two particles colliding simultaneously with correct geometry and minimum energy required is very small.• If there are more than 2 reactants, the reaction

will most likely occur by a number of simpler steps• A reaction mechanism is the step or series of

steps that make up a reaction.

Reaction Mechanisms

• A reaction mechanism is a series of simple steps that ultimately lead from the initial reactants to the final products of a reaction.• A valid mechanism must satisfy the following 2

criteria:• be consistent with the stoichiometry of the overall

or net reaction.• account for the experimentally determined rate

law.

Reaction Mechanisms• Molecularity : refers to the number of reactants involved in

an elementary step. Usually limited to 1 or 2, rarely 3.• Consider the production of NO2 from NO and O2.

2 NO + O2 2 NO2

Is it likely for this reaction to proceed in one step? Explain

This reaction consists of 2 elementary steps.1) 2 NO N2O2 (fast)

2) N2O2 + O2 2 NO2 (slow, rate determining)

2 NO + O2 2 NO2

• N2O2 is a reaction intermediate which is short lived and difficult to isolate.

• Often it is difficult to determine the actual mechanism of an reaction, most are proposals supported by experimentation. (Rate Law)• The rate-determining step is the crucial step in

establishing the rate of the overall reaction.• The rate determining step is the slowest step in the

reaction mechanism (bottle neck). Must relate to the experimentally determined rate law equation.• The rate determining step will therefore have the

highest activation energy.

Rate Law & Elementary Steps

• Can use the equation coefficients as the reaction orders in the rate law for an elementary step

Rate Law Relationship• Lets revisit the the production of NO2 from NO and O2.

2 NO + O2 2 NO2

The proposed mechanism consists of 2 elementary steps.

1)2 NO N2O2 (fast)

2)N2O2 + O2 2 NO2 (slow, rate determining)

2 NO + O2 2 NO2

• If this is a valid mechanism what is the rate law equation for this reaction?

Rate = k[O2]

• The equation reflects the rate determining (slow) step.

Catalyzed Decomposition of Hydrogen Peroxide

Slow step: H2O2 + I- H2O + OI-

Fast step: H2O2 + OI- H2O + O2 + I-

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Net equation: 2 H2O2 2 H2O + O2

The slow step is the rate-determining step.Rate of the reaction = rate of slow step = k[H2O2][I-]

OI- = ?- Reaction intermediate

I- = ?- Catalyst

Catalyzed Reactions• Reaction Intermediate• A species that is created in one step and

consumed in the other• Catalyst• A species that is present originally then

reforms later on during the reaction• It is not written in the overall equation, but

you may see it noted above the reaction arrow.• Catalysts provide a lower activation energy

pathway by producing an intermediate.

Homework• Read section 6.6• Page 386 #1-3• Page 387 #3,5