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Lecture 60 Kinetics III

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Title: Lecture 60 Kinetics III


1
Lecture 60 - Kinetics III
2
A Model for Chemical Kinetics
A B C
3 things affect the reaction rate
1. Rate of collisions between A and B 2. Fraction
of collisions having correct orientation 3.
Fraction of collisions having sufficient
energy to cause reaction
3
1. Rate of Collisions
  • Collision rate Z AB

related to molecular speeds
higher concentrations more collisions
more reactions
4
2. Orientation of Collisions
e.g. 2 BrNO Br2 2 NO
Br



no reaction
5
2. Orientation of Collisions
  • p steric factor
  • fraction of collisions with correct
  • orientation
  • 0 lt p lt 1

6
3. Energy of Collisions
Transition State
O-N---Br---Br---N-O
Ea
Potential Energy
Reactants
BrNO BrNO
DE
Products
Br2 2 NO
Extent of Reaction
7
3. Energy of Collisions
  • Ea is the activation energy
  • If energy of collision gt Ea,
  • products may be formed
  • fraction of collisions with energy gt Ea is

8
Reaction Rate
  • rate p f collision rate
  • p f Z AB
  • but rate k AB

thus, k p Z f
-Ea/RT
pZe
(Arrhenius)
9
Arrhenius Equation
-Ea/RT
k Ae
if Ea gt 0, k increases as T increases (more
energetic collisions)
Also, A µ T
10
Arrhenius Equation
1. Knowing k vs. T, can find A, Ea 2. Knowing A,
Ea, can find k at any T
11
Finding A and Ea
  • re-cast the equation
  • ln(k) ln(A) - Ea/RT

-Ea
1
ln(k)
ln(A)
R
T
y m x b
slope intercept
12
Finding A and Ea
intercept ln(A)
slope -Ea/R
ln(k)
1/T
13
for example,
  • 2 N2O5(g) 4 NO2(g) O2(g)
  • T, oC k, s-1 T, K ln(k) 1/T
  • 20 2.0 x 10-5 293 -10.82 .00341
  • 60 2.9 x 10-3 333 -5.84 .00300

14
for example,
-5.84
x
slope -Ea/R, intercept ln(A)
x
ln k
x
-10.82
x
.00300
.00341
1/T
-5.84 - (-10.82)
slope
-1.2 x 104 K
(.00300 - .00341) K-1
15
for example,
  • but, slope -Ea/R
  • thus, Ea - R (slope)
  • -8.314 J K-1 mol-1 (1.2 x 104 K)
  • 1.0 x 105 J/mol

16
another example...
  • O3 Cl OCl O2
  • A 2.8 x 10-11 cm3 molecule-1 s-1
  • Ea 2.1 kJ mol-1
  • find k at -50oC

17
O3 Cl OCl O2
-Ea/RT
k Ae
2.8 x 10-11 cm3 molecule-1 s-1 x
-2100 J/mol
8.314 J K-1mol-1 x 223 K)
e

9.02 x 10-12 cm3 molecule-1 s-1
18
Catalysis
  • A catalyst speeds up a reaction,
  • but is not consumed itself
  • The catalyst provides
  • a different (and faster) mechanism
  • for the reaction to proceed

19
Catalysis
Cl
e.g. O3(g) O(g) ? 2 O2(g)
20
Catalysis
O3 O ? 2 O2
O3 Cl ? ClO O2
Potential Energy
with a catalyst Ea is smaller
no catalyst
Extent of Reaction
21
Heterogeneous Catalysis
  • e.g. hydrogenation reactions

H
H
H
H
C C
H2(g)
H - C - C - H(g)
H
H(g)
H
H
ethylene
ethane
slow due to strong H-H bond!
22
Heterogeneous Catalysis
Metal
23
Heterogeneous Catalysis
Metal
1. Adsorption
24
Heterogeneous Catalysis
Metal
1. Adsorption 2. Migration across surface
25
Heterogeneous Catalysis
Metal
1. Adsorption 2. Migration across surface 3.
Reaction
4. Desorption
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