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Chemical Kinetics
Chapter 15
H2O2 decomposition in
an insect
H2O2 decomposition
catalyzed by MnO2
Chemical Kinetics
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• We can use thermodynamics to tell
if a reaction is product or reactant
favored.
• But this gives us no info on HOW FAST
reaction goes from reactants to
products.
•KINETICS — the study of REACTION
RATES and their relation to the way the
reaction proceeds, i.e., its MECHANISM.
Energy Diagram
KINETICS
thermodynamics
KINETICS
the study of REACTION RATES
and their relation to the way the
reaction proceeds, i.e., its
MECHANISM.
•The reaction mechanism
is our goal!
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Reaction Mechanisms
The sequence of events at the molecular
level that control the speed and
outcome of a reaction.
Br from biomass burning destroys
stratospheric ozone.
(See R.J. Cicerone, Science, volume 263, page 1243, 1994.)
Step 1:
Br + O3 ---> BrO + O2
Step 2:
Cl + O3 ---> ClO + O2
Step 3:
BrO + ClO + light ---> Br + Cl + O2
NET:
2 O3 ---> 3 O2
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REACTION RATES
-
RR = D [P ] = D [R ]
Dt
P =products
Dt
R = reactants
Determining a Reaction Rate
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Dye Concentration
Rate = the change
in [dye] divided
by time
The rate is
determined from
the plot.
Time
Relative Rates
Reactant
2A g
-
4B + C
D [A ] = D [B ] = D [C ]
2D t
4D t
Dt
Rate Calculations
Factors Affecting RXN Rates
• *Nature of Reactants
• Temperature
• Concentration
• Surface Area/ Physical
state
• Catalysts
Collision Theory
NO
Collisions
Collisions
NO
YES
Energy
Collisions
Energy
Orientation
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Concentrations and Rates
To postulate a reaction
mechanism, we study
• reaction rate and
• its concentration
dependence
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Concentration and rate
What is concentration of reactant as function of
time?
The rate law is
D[A]
Rate  = k [A]
Dtime
REACTION ORDER
In general, for
a A + b B --> x X
Rate = k
m
n
[A] [B]
The exponents m,, and n
• are the reaction order
• can be 0, 1, 2 or fractions
• must be determined by experiment!
Rate contant: Arrhenius equation
Rate constant is dependent on only the
activation energy and temperature
Temp (K)
Rate
constant
k  Ae
Frequency factor
-E a / RT
Activation
energy
8.31 x 10-3 kJ/K•mol
Frequency factor = frequency of collisions with correct geometry.
Simulation: RATE
MECHANISMS
A Microscopic View of Reactions
Mechanism: how reactants are converted
to products at the molecular level.
RATE LAW ---->
MECHANISM
experiment ----> theory
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More on Mechanisms
Reaction is
UNIMOLECULAR
if only one reactant is
involved.
BIMOLECULAR
if two different molecules
must collide to form a
products
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A bimolecular reaction
Collision Theory
Reactions require
(a) activation energy and
(b) correct geometry.
O3(g) + NO(g) ---> O2(g) + NO2(g)
1. Activation energy
2. Activation energy
and geometry
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Mechanisms
O3 + NO reaction occurs in a single
ELEMENTARY step. Most others involve a
sequence of elementary steps.
Adding elementary steps gives NET reaction.
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2 I- + 2 H+ + H2O2---> I2 + 2 H2O
1.
Rate law determined from experiment is:
Rate = k [I-] [H2O2]
Most rxns. have sequence of elementary steps.
NOTE
Order and stoichiometric
coefficients NOT the same!
3. Rate law reflects all chemistry
down to and including the slowest
step in multistep reaction.
2.
Mechanisms
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2 I- + H2O2 + 2 H+ ---> I2 + 2 H2O
Rate = k [I-] [H2O2]
Proposed Mechanism
Step 1 — slow HOOH + I-
--> HOI + OH-
Step 2 — fast HOI + I- -->
Step 3 — fast 2 OH- + 2 H+
I2 + OH--> 2 H2O
Rate is controlled by slow step —
RATE DETERMINING STEP, RDS.
Rate can be no faster than RDS!
Mechanisms
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2 I- + H2O2 + 2 H+ ---> I2 + 2 H2O
Rate = k [I-] [H2O2]
Step 1 — slow
Step 2 — fast
Step 3 — fast
HOOH + I- --> HOI + OH-
HOI + I- --> I2 + OH2 OH- + 2 H+ --> 2 H2O
Elementary Step 1 is bimolecular and involves Iand HOOH. Therefore, this predicts the rate law
should be
Rate  [I-] [H2O2] — as observed!!
The species HOI and OH- are reaction
intermediates.
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Simulation:” Mechanisms
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Sovled problems: pg 144
NO2 + CO
NO + CO2
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Single step
NO2 + CO reaction:
Rate = k[NO2]2
Two possible
mechanisms
Two steps: step 1
Two steps: step 2
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Ozone Decomposition Mechanism
2 O3 (g) ---> 3 O2 (g)
[O3 ]2
Rate = k
[O2 ]
Proposed mechanism
Step 1: fast, equilibrium
O3 (g) <--> O2 (g) + O (g)
Step 2: slow
O3 (g) + O (g) ---> 2 O2 (g)
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