KINETICS
Kinetics
Factors that affect rates
State of reactants
Temperature
Concentration
Catalyst
Reaction rates
For the reaction A  B the concentration of the reactants and the products were measured at
various times.
At t = 0s there is 1.00mol A and 0mol B
At t = 20s there is 0.54mol A and 0.46mol B
At t = 40s there is 0.30mol A and 0.70mol B
Average rate =
C H Cl(aq) + H O(l)  C H OH(aq) + HCl(aq)
4
9
2
4
9
Reaction Rates and Stoichiometry
C H Cl(aq) + H O(l)  C H OH(aq) + HCl(aq)
4
9
2
4
9
1:1 ratio
But what if it is not 1:1?
2 HI(g)  H2(g) + I2(g)
To generalize, for reaction
aA + bB  cC + dD Rate =
1. How is the rate at which ozone disappears related to the rate at which oxygen appears
in the reaction: 2 O3(g) → 3 O2(g)?
2. If the rate at which O2 appears, Δ[O2]/ Δt, is 6.0 × 10–5 M/s at a particular instant, at
what rate is O3 disappearing at this same time, –Δ[O3]/Δ t?
Dependence of Rate on Concentration
Rate Law
NH4+ (aq) + NO2- (aq)  N2 (g) + 2H2O (l)
Rate = k[reactants]m
Rate Constant
Don’t forget units!
Order of reaction
Rate = k [NH4+] [NO2−]
Order with respect to NH4+ =
Order with respect to NO2− =
Overall order =
22 = what happens to rate
order
Double the concentration
3. What are the overall reaction orders for the reactions:
2N2O5  4NO2 + O2
rate = k [N2O5]
CHCl3 + Cl2  CCl4 + HCl
rate = k [CHCl3][Cl2]1/2
H2 + I2  2HI
rate = k [H2][I2]
NO2  NO + ½O2
rate = k [NO2]2
overall order =
overall order =
overall order =
overall order =
4. The initial rate of a reaction A + B → C was measured for several different starting
concentrations of A and B, and the results are:
Using these data, determine
(a) the rate law for the reaction,
(b) the rate constant,
(c) the rate of the reaction when [A] = 0.050 M and [B] = 0.100 M.
5. For the reaction: 2NO(g) + O2(g) 2NO2(g) The following data was collected:
Trial Initial [NO] Initial [O2] Rate of NO2formation
1
0.01
0.01
0.05
2
0.02
0.01
0.20
3
0.01
0.02
0.10
Determine the rate law and rate constant
6. The following data were measured for the reaction of nitric oxide with hydrogen:
(a) Determine the rate law for this reaction.
(b) Calculate the rate constant.
(c) Calculate the rate when [NO] = 0.050 M and [H2] = 0.150 M
The change of concentration with time
Integrated rate laws
1st order
7. The decomposition of a certain insecticide in water follows first-order kinetics with a rate
constant of 1.45 yr–1 at 12 °C. A quantity of this insecticide is washed into a lake on June 1,
leading to a concentration of 5.0 × 10–7 g/cm3. Assume that the average temperature of the
lake is 12 °C.
a) What is the concentration of the insecticide on June 1 of the following year?
b) How long will it take for the concentration of the insecticide to decrease to 3.0 × 10 –7
g/cm3?
2nd order
Half Life
For a 1st order process
First Order
Second Order
Second Order
Rate Laws
Rate = - k [A]
Rate = - k [A]2
Rate = - k [A] [B]
Integrated
Rate Laws
1 =kt + 1
[A]t
[A]0
Complicated!
ln [A]t - ln [A]0= - k t
Graph
ln [A] vs. t is linear
1 vs. t is linear
[A]
Half Life
0.693 = t1/2
k
1 = t1/2
k [A]0
Complicated !
Collision Theory
An effective collision must have particles that collide with
Activation energy
PE diagrams
Temperature and rate
8. Consider a series of reactions having the following energy profiles:
Rank the reactions from slowest to fastest
Reaction Mechanisms
Overall reaction:
A + 2B  E + F
A+BC
C+BD
DE+F
Molecularity -
Rate Law of an elementary step
9. It has been proposed that the conversion of ozone into O 2 proceeds by a two-step
mechanism:
a) Describe the molecularity of each elementary reaction in this mechanism.
b) Write the equation for the overall reaction.
c) Identify the intermediate(s).
Multistep Mechanisms
Rate Determining Step
NO2 (g) + CO (g)  NO (g) + CO2 (g)
The rate law for the above reaction is found experimentally to be
Rate = k [NO2]2
A proposed mechanism for this reaction is
Step 1: NO2 + NO2  NO3 + NO (slow)
Step 2: NO3 + CO  NO2 + CO2 (fast)
10. The decomposition of nitrous oxide, N2O, is believed to occur by a two-step mechanism:
a) Write the equation for the overall reaction.
b) Write the rate law for the overall reaction.
c) Identify any intermediates
Fast Initial Step
2 NO (g) + Br2 (g)  2 NOBr (g)
The rate law for the above reaction is found to be
Rate = k [NO]2 [Br2]
Because termolecular processes are rare, this rate law suggests a two-step mechanism.
A proposed
Step 1: NO + Br2
mechanism for this reaction is
NOBr2 (fast)
Step 2: NOBr2 + NO  2 NOBr
(slow)
Overall rate =
But how can we find the concentration of an intermediate??
Assume fast step is in equilibrium!
Now plug back into overall rate law
11. Ozone reacts with nitrogen dioxide to produce dinitrogen pentoxide and oxygen:
The reaction is believed to occur in two steps:
The
experimental rate law
is rate = k[O3][NO2]. What can you say about the relative rates of the two steps of the
mechanism?
Catalysts
Catalysts increase the rate of a reaction by