1. Chapter 14
Homework pages 463 – 466
1, 2, 5 – 11, 13, 15, 19, ,
27, 30, 31, 39, 40, 41, 43 – 51,
56, 64.

2. Chapter 14 - Chemical Kinetics
Thermodynamics – does a reaction take place?
Kinetics – how fast does a reaction proceed?
Reaction rate is the change in the concentration of a reactant or a product with time (M/s).
Ave. rates
A B
rate = -
D[A] Dt
D[A] = change in concentration of A over
time period Dt
rate =
D[B] Dt
D[B] = change in concentration of B over
time period Dt
Because [A] decreases with time, D[A] is negative.
13.1

3. A B
time
rate = -
D[A] Dt
rate =
D[B] Dt
13.1

4. Br2 (aq) + HCOOH (aq) 2Br- (aq) + 2H+ (aq) + CO2 (g)
time
393 nm
light
Detector
393 nm
Br2 (aq)
[Br2] a Absorbance
13.1

5. Br2 (aq) + HCOOH (aq) 2Br- (aq) + 2H+ (aq) + CO2 (g)
Calculate the average rate for:
The first 200 s
From 300 s to 350 s
average rate = -
D[Br2] Dt
= -
[Br2]final – [Br2]initial
tfinal - tinitial
13.1

6. Br2 (aq) + HCOOH (aq) 2Br- (aq) + 2H+ (aq) + CO2 (g)
slope of
tangent
slope of
tangent
slope of
tangent
average rate = -
D[Br2] Dt
= -
[Br2]final – [Br2]initial
tfinal - tinitial
instantaneous rate = rate for specific instance in time
13.1

7. How does the rate depend on concentration? rate = k [Br2] k = rate
= rate constant
= 3.50 x 10-3 s-1
13.1

8. Reaction Rates and Stoichiometry
2A B
Two moles of A disappear for each mole of B that is formed.
rate = -
D[A] Dt 1 2
rate =
D[B] Dt
aA + bB cC + dD
rate = -
D[A] Dt 1 a
= -
D[B] Dt 1 b =
D[C] Dt 1 c =
D[D] Dt 1 d
13.1

9. If concentration of O2 decreases at a
Write the rate expression for the following reaction:
CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (g)
rate = -
D[CH4] Dt
= -
D[O2] Dt 1 2 =
D[CO2] Dt =
D[H2O] Dt 1 2
If concentration of O2 decreases at a
rate of 0.10 M/s, what is the reaction rate?
What is the rate of appearance of CO2?
13.1

10. The Rate Law
The rate law expresses the relationship of the rate of a reaction to the rate constant and the concentrations of the reactants raised to some powers.
aA + bB cC + dD
Rate = k [A]x[B]y
reaction is xth order in A
reaction is yth order in B
reaction is (x +y)th order overall
13.2

11. Double [F2] with [ClO2] constant
F2 (g) + 2ClO2 (g) FClO2 (g)
rate = k [F2]x[ClO2]y
Double [F2] with [ClO2] constant
Rate doubles
x = 1
rate = k [F2][ClO2]
Quadruple [ClO2] with [F2] constant
Rate quadruples
y = 1
13.2

12. Rate Laws Rate laws are always determined experimentally.
Reaction order is always defined in terms of reactant (not product) concentrations.
The order of a reactant is not related to the stoichiometric coefficient of the reactant in the balanced chemical equation.
F2 (g) + 2ClO2 (g) FClO2 (g) 1
rate = k [F2][ClO2]
13.2

13. S2O82- (aq) + 3I- (aq) 2SO42- (aq) + I3- (aq)
Determine the rate law and calculate the rate constant for the following reaction from the following data:
S2O82- (aq) + 3I- (aq) SO42- (aq) + I3- (aq)
Experiment
[S2O82-]
[I-]
Initial Rate (M/s) 1
0.08
0.034
2.2 x 10-4 2
0.017
1.1 x 10-4 3
0.16
rate = k [S2O82-]x[I-]y
y = 1
x = 1
rate = k [S2O82-][I-]
Double [I-], rate doubles (experiment 1 & 2)
Double [S2O82-], rate doubles (experiment 2 & 3)
k =
rate
[S2O82-][I-] =
2.2 x 10-4 M/s
(0.08 M)(0.034 M)
= 0.08/M•s
13.2

14. Finding Reaction Orders by Experiment
Given the reaction: O2 (g) + 2 NO(g)  2 NO2 (g)
The rate law for this reaction in general terms is:
Rate = k [O2]m[NO]n
To find the reaction orders, we run a series of experiments, each with a
different set of reactant concentrations, and determine the initial
reaction rates.
Experiment Initial Reactant Concentrations (mol/l) Initial Rate
O NO (mol/L·s)
1 < x * x * >3.21 x 10-3
2 < x x >6.40 x 10-3
x * x * x 10-3
x x x 10-3
x x x 10-3

16. What are the units of k for a first order reaction?
Zero order? Second order?
The rate equation tells how the rate depends on the
reactant concentrations. Can we derive an equation
that tells how the concentration depends on time?
These are the integrated rate laws.
How can we determine the reaction order graphically?

17. First-Order Reactions
rate = -
D[A] Dt
A product
rate = k [A]
D[A] Dt
= k [A] -
rate
[A]
M/s M =
k =
= 1/s or s-1
[A] is the concentration of A at any time t
[A]0 is the concentration of A at time t=0
[A] = [A]0exp(-kt)
ln[A] = ln[A]0 - kt
13.3

18. The reaction 2A B is first order in A with a rate constant of 2
The reaction 2A B is first order in A with a rate constant of 2.8 x 10-2 s-1 at 800C. How long will it take for A to decrease from 0.88 M to 0.14 M ?
[A]0 = 0.88 M
[A] = 0.14 M
What is the concentration of A after 100 s?

19. The half-life of a reaction, t1/2, is the time required for the concentration of a reactant to decrease to 1/2 of its initial concentration.
For a first order reaction only, the half-life is independent of the initial concentration.
t1/2 = (1/k) ln 2

20. First-Order Reactions
The half-life, t½, is the time required for the concentration of a reactant to decrease to half of its initial concentration.
t½ = t when [A] = [A]0/2
ln2 k =
0.693 k = t½
What is the half-life of N2O5 if it decomposes with a rate constant of 5.7 x 10-4 s-1? t½
ln2 k =
0.693
5.7 x 10-4 s-1 =
= 1200 s = 20 minutes
How do you know decomposition is first order?
units of k (s-1)
13.3

21. First-order reaction A product # of half-lives [A] = [A]0/n 1 2 2 4 38 4 16
13.3

23. k = (1/200s)ln2
= s-1

24. Summary of the Kinetics of Zero-Order, First-Order
and Second-Order Reactions
Order
Rate Law
Concentration-Time Equation
Half-Life
t½ =
[A]0 2k
rate = k
[A] = [A]0 - kt t½
ln2 k = 1
rate = k [A]
ln[A] = ln[A]0 - kt 1
[A] =
[A]0
+ kt
t½ = 1
k[A]0 2
rate = k [A]2
13.3

25. Experiment 13
Rate = k[A]x
Take ln of both sides.
Ln(rate) = x ln[A] + ln k
Ln(1/time) = x ln[A] + constant

26. 2. In order to react the colliding molecules must have
COLLISION THEORY OF CHEMICAL KINETICS
1. Chemical reactions occur as a result of collisions between reacting molecules.
2. In order to react the colliding molecules must have
a. the right orientation
b. a total kinetic energy equal to or greater than the activation energy, Ea, which is the minimum energy required to initiate the chemical reaction.

28. Increasing the temperature usually has a dramatic effect on reaction rates.
Chemists use a “rule of thumb” that a 10° rise in temperature doubles the rate of a reaction.

31. TRANSITION STATE THEORY
The species temporarily formed by the reactant molecules before they form the product is called the transition state or activated complex.
In the transition state old bonds are in the process of breaking and new bonds are beginning to form. The activation energy is the energy required to reach the transition state.

32. A + B C + D Exothermic Reaction Endothermic Reaction
The activation energy (Ea) is the minimum amount of energy required to initiate a chemical reaction.
13.4

34. Sample Problem
For the reaction
O3(g) + O(g) O2(g) Ea(fwd) = 19 kJ
and ΔH = -392 kJ. Draw a reaction energy
diagram, show a possible transition state, and
calculate Ea(rev).

35. Temperature Dependence of the Rate Constant
k = A • exp( -Ea/RT )
(Arrhenius equation)
Ea is the activation energy (J/mol)
R is the gas constant (8.314 J/K•mol)
T is the absolute temperature
A is the frequency factor
lnk = - Ea R 1 T
+ lnA
13.4

36. lnk = - Ea R 1 T
+ lnA
13.4

37. Test on chapters 14 & 15– May 26 EH Assignment Due May 23
Unit 16 Chemical Kinetics
Mininum score
Sec 1 Reaction Rates from Chemical Data 90
Sec 2 Rate Law Equations
Sec 3 Experimental Kinetics
Sec 4 First Order Rate Problems
Sec 5 Reaction Mechanisms
Sec 6 Temperature and Rate
40 (first 4 questions
Test on chapters 14 & 15– May 26

38. An elementary reaction is a simple reaction
that proceeds in one step.
Most reactions are complex and require more
than one step.
A reaction mechanism is the sequence of
elementary steps that leads to product formation.
An intermediate is a species that appears in the
mechanism, but not in the overall balanced rxn.

39. Reaction Mechanisms
The sequence of elementary steps that leads to product formation is the reaction mechanism.
2NO (g) + O2 (g) NO2 (g)
Elementary step:
NO + NO N2O2
Elementary step:
N2O2 + O NO2
Overall reaction:
2NO + O NO2
13.5

40. Unimolecular reaction – elementary step with 1 molecule
Intermediates are species that appear in a reaction mechanism but not in the overall balanced equation.
An intermediate is always formed in an early elementary step and consumed in a later elementary step.
Elementary step:
NO + NO N2O2
N2O2 + O NO2
Overall reaction:
2NO + O NO2 +
The molecularity of a reaction is the number of molecules reacting in an elementary step.
Unimolecular reaction – elementary step with 1 molecule
Bimolecular reaction – elementary step with 2 molecules
Termolecular reaction – elementary step with 3 molecules
13.5

41. Rate Laws and Elementary Steps
Unimolecular reaction
A products
rate = k [A]
Bimolecular reaction
A + B products
rate = k [A][B]
Bimolecular reaction
A + A products
rate = k [A]2
Writing plausible reaction mechanisms:
The sum of the elementary steps must give the overall balanced equation for the reaction.
The rate-determining step should predict the same rate law that is determined experimentally.
The rate-determining step is the slowest step in the sequence of steps leading to product formation.
13.5

42. For elementary reactions only, the exponents in the
rate law must correspond to the coefficients in the
balanced equation.
The elementary steps that make up a reaction
mechanism must satisfy two requirements.
1. The sum of the elementary steps must give the
overall balanced equation for the reaction.
2. The rate-determining step (the slow step) must
predict the same rate law as is determined by
experiment.

43. What is the equation for the overall reaction?
The experimental rate law for the reaction between NO2 and CO to produce NO and CO2 is rate = k[NO2]2. The reaction is believed to occur via two steps:
Step 1:
NO2 + NO NO + NO3
Step 2:
NO3 + CO NO2 + CO2
What is the equation for the overall reaction?
NO2+ CO NO + CO2
What is the intermediate?
NO3
What can you say about the relative rates of steps 1 and 2?
rate = k[NO2]2 is the rate law for step 1 so
step 1 must be slower than step 2
13.5

44. For the reaction 2 X + Y Z
rate = k[X][Y]
Can this be an elementary reaction?
For the reaction 2 A + B C
rate = k[A]2 [B]
Is this an elementary reaction?

45. Is this likely to be an elementary process?
Sample Problem
Problem: The following two reactions are proposed as elementary steps
in the mechanism for an overall reaction:
(1) NO2Cl (g) NO2(g) + Cl(g)
(2) NO2Cl (g + Cl(g) NO2(g) + Cl2(g)
(a) Write the overall balanced equation.
(b) Determine the molecularity of each step.
(c) Write the rate law for each step.
2 C2H6 + 7 O2  4 CO2 + 6 H2O
Is this likely to be an elementary process?

46. Mechanisms with a Slow Initial Step
2 NO2 + F NO2F
The experimental rate law for this reaction is first order in
both reactants. A proposed mechanism is:
NO2 + F NO2F + F slow
NO2 + F NO2F fast
Show that this is a reasonable mechanism. What is the
intermediate species?

47. A Catalyst is a substance that increases the rate
of reaction without itself being consumed.
A catalyst works by changing the mechanism. The
new mechanism has a lower activation energy.
It speeds up both forward and reverse reactions and
cannot increase the final equilibrium yield, but it
gets to the final equilibrium state faster.

48. A catalyst is a substance that increases the rate of a chemical reaction without itself being consumed.
k = A • exp( -Ea/RT ) Ea k
uncatalyzed
catalyzed
ratecatalyzed > rateuncatalyzed
Ea < Ea ‘
13.6

49. Depletion of the Earth’s Ozone Layer - I
The stratospheric ozone is vital to life on earth because ozone absorbs
short-wavelength (~ 3x10-7m) ultraviolet (UV) radiation from the sun.
UV photon
O3 (g) O2 (g) + O(g)
UVB
Normally two reactions maintain the ozone in the stratosphere, and
keep a level of ozone sufficient to absorb all of the UVB radiation that
reaches the earth’s atmosphere:
O2 (g) + O(g)  O3 (g) [formation]
O3 (g) + O(g)  O2 (g) [breakdown]
Under normal conditions, the atomic oxygen is produced by the
breakdown of molecular oxygen by it absorbing ultraviolet radiation
UV photon
O2 (g) O(g)
UVA

50. . . . . . . Depletion of the Earth’s Ozone Layer - II
Chlorofluorocarbons (CFC’s) were introduced in the 1930’s as working
gases to be used in refrigerators and air conditioners. They were found to
breakdown in the stratosphere by UVAradiation and release atomic
chlorine atoms to the atmosphere, which had a “catalytic” reaction with
ozone that has been causing it’s concentration in the stratosphere to drop. . .
Uv photon
Freon CF2Cl2 (g) CF2Cl(g) + Cl(g)
UVA
The • represents an unpaired electron, resulting from bond breaking, and
the resultant molecule is called a “free radical” and is very reactive. . .
O3 (g) + Cl(g) ClO(g) + O2 (g) . .
ClO(g) + O(g) Cl(g) + O2 (g)
O3 (g) + O(g) O2 (g)

51. Haber synthesis of ammonia
In heterogeneous catalysis, the reactants and the catalysts are in different phases.
Haber synthesis of ammonia
Ostwald process for the production of nitric acid
Catalytic converters
In homogeneous catalysis, the reactants and the catalysts are dispersed in a single phase, usually liquid.
Acid catalysis
Base catalysis
13.6

52. Haber Process
Fe/Al2O3/K2O
catalyst
N2 (g) + 3H2 (g) NH3 (g)
13.6

53. Enzyme Catalysis
13.6

54. uncatalyzed
enzyme
catalyzed
13.6