Reaction Mechanisms Chapter 12, Section 6. Reaction Mechanisms The sequence of events that describes...
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Transcript of Reaction Mechanisms Chapter 12, Section 6. Reaction Mechanisms The sequence of events that describes...
![Page 1: Reaction Mechanisms Chapter 12, Section 6. Reaction Mechanisms The sequence of events that describes the actual process by which reactants become products.](https://reader036.fdocuments.us/reader036/viewer/2022070401/56649f205503460f94c3879e/html5/thumbnails/1.jpg)
Reaction Mechanisms
Chapter 12, Section 6
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Reaction Mechanisms
The sequence of events that describes the actual process by which reactants become products is called the reaction mechanism.
Proposed Mechanism:Step 1: NO2 + NO2 NO3 + NO Step 2: NO3 + CO NO2 + CO2
NO2 + CO NO + CO2
Rate = k[NO2]2
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Reaction Mechanisms
• Reactions may occur all at once or through several discrete steps.
• Each of these processes is known as an elementary reaction or elementary process.
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Reaction Mechanisms
The molecularity of a process tells how many molecules are involved in the process.
Termolecular elementary reactions seldom occur.
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Multistep Mechanisms
• In a multistep process, one of the steps will be slower than all others.
• The overall reaction cannot occur faster than this slowest, rate-determining step.
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Reaction Mechanisms must satisfy 2 requirements:
1. The sum of the elementary steps must give the overall balanced equation.
2. The mechanism must agree with the experimentally determined rate law.
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Slow Initial Step
• The rate law for this reaction is found experimentally to be
Rate = k [NO2]2
• CO is necessary for this reaction to occur, but the rate of the reaction does not depend on its concentration.
• This suggests the reaction occurs in two steps.
NO2 (g) + CO (g) NO (g) + CO2 (g)
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Slow Initial Step• A proposed mechanism for this reaction is
Step 1: NO2 + NO2 NO3 + NO (slow)
Step 2: NO3 + CO NO2 + CO2 (fast)
• Does it make sense?
• What is the intermediate in this reaction?
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Fast Initial Step
• The rate law for this reaction is found to be
Rate = k [NO]2 [Br2]
• Because termolecular processes are rare, this rate law suggests a two-step mechanism.
2 NO (g) + Br2 (g) 2 NOBr (g)
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Fast Initial Step
• A proposed mechanism is
Step 2: NOBr2 + NO 2 NOBr (slow)
Step 1 includes the forward and reverse reactions.
Step 1: NO + Br2 NOBr2 (fast)
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Fast Initial Step
• The rate of the overall reaction depends upon the rate of the slow step.
• The rate law for that step would be
Rate = k2 [NOBr2] [NO]
• But how can we find [NOBr2]?
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Fast Initial Step
• NOBr2 can react two ways:
– With NO to form NOBr
– By decomposition to reform NO and Br2
• The reactants and products of the first step are in equilibrium with each other.
• Therefore,
Ratef = Rater
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Fast Initial Step
• Because Ratef = Rater ,
k1 [NO] [Br2] = k−1 [NOBr2]
• Solving for [NOBr2] gives us
k1
k−1
[NO] [Br2] = [NOBr2]
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Fast Initial Step
Substituting this expression for [NOBr2] in the rate law for the rate-determining step gives
k2k1
k−1
Rate = [NO] [Br2] [NO]
= k [NO]2 [Br2]
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Catalysts
• A catalyst speeds up a reaction without being consumed.
• Identify the catalyst in the reaction mechanism below:
Cl(g) + O3(g) ClO(g) + O2(g)
O(g) + ClO(g) Cl(g) + O2(g)
How is it different than an intermediate?