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. Concentration affects kinetics
Orders of reactions
Rate laws
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. Rates laws
Rate laws – mathematical equations that describe how reaction rate is proportional to concentration of reactants
Where:
k is the rate constant – determines how temp. affects rate
m & n are reaction orders – small whole numbers
aA + bB  cC + dD
Rate = k[A]m[B]n
NH NO2-1  N H2O
Rate = k[NH4+1][NO2-1]
So:
As the conc of NH4 doubles so does the reaction rate
As the conc of NO2 doubles so does the reaction rate
Exp’t
Initial
[NH4+1]
[NO2-1]
M/s (x10-7) 1
0.0100
0.200
5.4 2
0.0200
10.8 3
0.0400
21.5 4
0.0202 5
0.0404
21.6 6
0.0808
43.3
Given the data in the table,
calculate k. 5.4x10-7 M/s = k(0.0100)(0.200)
k = 2.7x10-4/M/s
…now any conc can be calc. p

4. Reaction orders
NH NO2-1  N H2O
Rate = k[NH4+1][NO2-1]
For this reaction the exponents m & n are both 1 and so not explicitly shown. We therefore say that the reaction is first order for both reactants. The overall order of reaction is the sum of the individual orders, so 2 in this case.
Reaction orders must be experimentally determined & are hard to predict.
CHCl3 + Cl2  CCl4 + HCl Rate = k[CHCl3][Cl2]1/2
2NO + 2H2  N H2O Rate = k[NO]2[H2]
So reaction orders can be either whole numbers or fractions.
Looking at the second reaction:
What are the reaction orders of each reactant?
What’s the overall reaction order?
How does doubling the concentration of NO change the rate?
How does doubling the concentration of H2 change the rate?
What are units of rate constants?
Rate = M/s
concentration = M
k = M/s = s-1 M p

5. Rate laws & initial data
A + B  C
Exp’t
[A]
[B]
Initial Rate (M/s) 1
0.100
4.0x10-5 2
0.200 3
16.0x10-5
Determine the rate law for the reaction.
Determine the rate constant.
Calculate rate when [A] = M and [B] = M
Rate = k[A]m[B]n Reactant A – the rate increases by 4X as concentration of A doubles, so m = 2. Reactant B – the rate doesn’t change when the conc of B doubles, so n = 0. Rate = k[A]2[B]0
b) k = rate/[A]2 = 4.0x10-5/(0.100)2 = 4.0x10-3 M-1 s-1
c) Rate = (4.0x10-3 M-1s-1)(0.050)2(0.100)0 = 1.0x10-5 M/s p