Reaction mechanisms
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Reaction Mechanisms. Collision Theory predicts that the rate law of any reaction should match the balanced chemical equation 2 A + 3B → C rate = k [A] 2 [B] 3. However, this assumes that the reaction proceeds by a single

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Reaction Mechanisms

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Reaction mechanisms

Reaction Mechanisms

Collision Theory predicts that the rate law of any reaction should

match the balanced chemical equation

2 A + 3B → C rate = k [A]2 [B]3

However, this assumes that the reaction proceeds by a single

collision. It is almost impossible for 2 A’s and 3 B’s to all collide

at the same time.

The above reaction most likely occurs by a series of reaction steps, called elementary steps, which as a whole lead to the overall reaction. Each elementary step has its own equation, activation energy, and rate law.


Reaction mechanisms1

Reaction Mechanisms

The sum of all of the elementary steps in a proposed reaction

mechanism must match the balanced equation for the overall

reaction. This is rule # 1 for evaluating reaction mechanisms.

One of the elementary steps will be the slowest in the

process. This step will limit the rate of the entire reaction and is

called the rate determining step.

The rate law for this one elementary step determines the rate

law for the entire reaction. Therefore, the expected rate

law of the slow, rate-determining step must match the

experimentally determined rate law for the overall reaction.

This is rule # 2 for evaluating reaction mechanisms.


No 2 co no co 2 empirical rate law rate k no 2 2

NO2 + CO  NO + CO2 Empirical rate law: rate = k [NO2] 2

1. Reaction does not occur by the simple collision of NO2and CO since the experimentally determined rate law does not match the expected rate law from the balanced equation.

2. The reaction occurs in a series of reactions, each with its own rate constant and activation energy.


Reaction mechanisms

Overall: NO2 + CO  NO + CO2Possible Reaction MechanismStep 1: NO2 + NO2  NO3 + NO Step 2: NO3 + CO  NO2 + CO2

Each step is an elementary step with its own equation, activation energy, and expected rate law.

Elementary Step equations can include catalysts and intermediates


Reaction mechanisms

Overall: NO2 + CO  NO + CO2Possible Reaction MechanismStep 1: NO2 + NO2  NO3 + NO Step 2: NO3 + CO  NO2 + CO2

Does the proposed mechanism include an intermediate or a catalyst?

NO3 is an intermediate since it is produced and then used up in the total reaction.

There is not catalyst in the reaction since a catalyst would begin a reaction and not be used up in the reaction.


Reaction mechanisms

Rule # 1. The sum of all the steps in a proposed reaction

mechanism must equal the overall balanced

chemical equation

Overall: NO2 + CO  NO + CO2 Proposed Reaction MechanismStep 1: NO2 + NO2  NO3 + NOStep 2: NO3 + CO  NO2 + CO2

Combine the steps to produce an overall reaction

NO2 + NO2 + NO3 + CO  NO2 + CO2 + NO3 + NOxx x x

NO2 + CO → CO2 + NO

Combined steps produces a balanced chemical equation

equal to the overall balanced chemical equation


Reaction mechanisms

Rule # 2. The rate law for the rate determining step must match the experimental rate law for the overall equationOverall equation ; NO2 + CO → NO + CO2Experimental Rate law: rate = k [NO2]2Step 1: NO2 + NO2  NO3 + NOStep 2: NO3 + CO  NO2 + CO2

Step 1: NO2 + NO2 → NO3 + NO rate = k1 [NO2]2

Step 2: NO3 + CO→ NO2 + CO2 rate = k [NO3] [CO] rate = k [NO2]2[CO]

Since the rate law must be based on reactants, and NO3 is an intermediate, account for [NO3] using the rate law that produced I it in the first step - rate = k [NO2]2


Reaction mechanisms

Overall equation ; NO2 + CO → NO + CO2Experimental Rate law: rate = k [NO2]2Step 1: NO2 + NO2  NO3 + NOStep 2: NO3 + CO  NO2 + CO2

Step 1: NO2 + NO2 → NO3 + NO rate = k [NO2]2

Step 2: NO3 + CO→ NO2 + CO2 rate = k [NO2]2 [CO]

The rate law for step 1 matches the empirical rate law and

the combined steps produce an equation equal to the balanced

chemical equation for the reaction

Step 1 in the reaction must be the slower, rate-determine step in the reaction.


Overall equation 2 no 2 f 2 2 no 2 f experimental rate law rate k no 2 f 2

Overall equation: 2 NO2 + F2 2 NO2F Experimental rate law: rate =k [NO2] [F2]

Proposed mechanism Step 1: NO2 + F2  NO2F + F Step 2: F + NO2  NO2F

Is there an intermediate or catalyst in the proposed mechanim?

A lone F atom is produced in step 1 and used up in step 2.

F is therefore an intermediate.


Overall equation 2 no 2 f 2 2 no 2 f experimental rate law rate k no 2 f 21

Overall equation: 2 NO2 + F2 2 NO2F Experimental rate law: rate =k [NO2] [F2]

Proposed mechanism Step 1: NO2 + F2  NO2F + F Step 2: F + NO2  NO2F

Does the sum of the elementary steps equal to overall balanced

equation?

NO2 + F2 + F+ NO2  NO2F + F + NO2FXX NO2 + F2 + NO2  NO2F + NO2F 2 NO2 + F2  2 NO2F yes. Rule # 1 is fulfilled.


Overall equation 2 no 2 f 2 2 no 2 f experimental rate law rate k no 2 f 22

Overall equation: 2 NO2 + F2 2 NO2F Experimental rate law: rate = k [NO2] [F2]

Proposed mechanism Step 1: NO2 + F2  NO2F + F Step 2: F + NO2  NO2F

Determine the rate law for steps 1 and 2.

Step 1: NO2 + F2  NO2F + F rate = k [NO2 ] [F2]

Step 2: F + NO2  NO2Frate = k [NO2 ] [F]

rate = k [NO2] [NO2] [F2]

rate = k [NO2]2 [F2]

The rate law for step 1 matches the empirical rate law so step 1

is the slow, rate-determining step in the reaction.


Reaction mechanisms

Overall equation 2 NO2 + F2 2 NO2F rate=k[NO2]2 [F2]Proposed mechanismStep 1: NO2 + F2  NO2F + FStep 2: F + NO2  NO2F

The proposed mechanism satisfies both rule 1 and 2 for a possible mechanism for the reaction and predicts that the first step is the rate-determining step in the process


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