1. CHE1031 Lecture 10: Reaction kinetics
Lecture 10 topics Brown chapter 14
1. Reaction rates
Factors that effect reaction rates
Visualizing rates & units
Average reaction rates
Instantaneous reaction rates
Stoichiometry & reaction rates
2. Concentration & reaction rates
Rate laws
Reaction orders
3. Change in concentration with time
First- & second-order reactions
Half-life
4. Temperature & reaction rate
Collision, orientation & Ea
5. Reaction mechanisms
Elementary
Multistep
6. Catalysis `

2. Elementary vs. multistep reactions
Reaction mechanisms
Elementary vs. multistep reactions
Rate-limiting steps
The general process of advancing scientific knowledge by making experimental observations and by formulating hypotheses, theories, and laws.
It’s a systematic problems solving process AND it’s hands-on….. Experiments must be done, data generated, conclusions made.
This method is “iterative”; it requires looping back and starting over if needed. [Why do you think they call it REsearch?] Often years, decades or more of experiments are required to prove a theory.
While it’s possible to prove a hypothesis wrong, it’s actually NOT possible to absolutely prove a hypothesis correct as the outcome may have had a cause that the scientist hasn’t considered.

3. Reaction mechanisms
Reaction mechanism – describe the details by which a chemical rxn takes place. There are a series of basic & common mechanisms.
Elementary mechanisms: single-step reactions
Multistep reactions: require several steps, or reactions to reach completion
Molecularity? Uni-, bi-, or termolecular describe reactions of 1, 2 or 3 molecules.
Unimolecular: H3C – N = C  H3C – C = N = =
Bimolecular: NO + O3  NO2 + O2
Two-step reaction:
(1) NO2 + NO2  NO3 + NO
(2) NO3 + CO  NO2 + CO2
sum NO2 + CO  NO + CO2
Molecules that don’t appear in the summed (overall) reaction are intermediates. p

4. Elementary reactions & rate laws
The relationship is quite simple as you can see here: p

5. Rate-limiting steps
When chemical reactions require more than one step, their overall rate is often limited by the slowest of the steps.
So this slowest stop is called the rate-limiting step because it limits the overall rate of reaction.
Step 1: NO2 + NO2  NO3 + NO (slow)
Step 2: NO3 + CO  NO2 + CO2 (fast)
Overall: NO2 + CO  NO + CO2 k1
k2 >> k1 k2
What is the rate law of the overall reaction?
Because step 1 is much slower than step 2, it is rate-limiting.
The rate of the overall reaction is equal to the rate of the slow step (1).
Step 1 is bimolecular, so
rate = k1[NO2]2 p

6. Rate-limiting examples
Nitrous oxide decomposes by a two-step mechanism.
N2O  N2 + O (slow)
N2O + O  N2 + O2 (fast)
Write the equation for the overall reaction.
Write the rate law for the overall reaction.
a) 2N2O  2N2 + O2
b) Rate = k[N2O]2
Ozone reacts with nitrogen dioxide by a two-step mechanism:
Step 1: O3 + NO2  NO3 + O2
Step 2: NO3 + NO2  N2O5
Overall: O NO2  N2O5 + O2
Overall experimental rate law is: rate = k[O3][NO2].
Which step is slower?
Step 1 is the slow step, since it is used in the rate law for the overall reaction. p