The Effects of Temperature and Catalyst on Reaction Rate

1
The Effects of Temperature and Catalyst on
Reaction Rate
15.1 Activation Energy and Arrhenius
Equation 15.2 Interpretation of Rates of Gaseous
Reactions at Molecular
Level 15.3 Energy Profile 15.4 Effect of
Catalysts on Rates of Reactions
2
Activation Energy and Arrhenius Equation
3
15.1 Activation Energy and Arrhenius Equation
(SB p.49)
Activation Energy
Exothermic reaction
Activation energy energy required to break
bonds
(related to the rate of reaction)
4
15.1 Activation Energy and Arrhenius Equation
(SB p.49)
Activation Energy
Endothermic reaction
5
15.1 Activation Energy and Arrhenius Equation
(SB p.50)
Arrhenius Equation
where k is the reaction constant of the
reaction, A is a constant which is independent
of temperature, e is the base of the natural
logarithm, Ea is the activation energy of the
reaction in J mol-1, R is the ideal gas constant
(i.e. 8.314 J K-1 mol-1). and T is the
temperature in Kelvin.
6
15.1 Activation Energy and Arrhenius Equation
(SB p.50)
Arrhenius Equation
7
15.1 Activation Energy and Arrhenius Equation
(SB p.50)
Arrhenius Equation
8
15.1 Activation Energy and Arrhenius Equation
(SB p.50)
Arrhenius Equation
A graph showing the relationship between rate
constant and temperature
9
15.1 Activation Energy and Arrhenius Equation
(SB p.51)
Determination of Activation Energy Using
Arrhenius Equation
10
15.1 Activation Energy and Arrhenius Equation
(SB p.52)
Determination of Activation Energy Using Two Rate
Constants
By subtracting equation (2) from equation (1), we
obtain
11
Interpretation of Rates of Gaseous Reactions at
Molecular Level
12
15.2 Interpretation of Rates of Gaseous
Reactions at Molecular Level (SB p.54)
Distribution of Molecular Speeds in a Gas
Consider a sample of gas
Do all gas molecules move at the same speed?
13
15.2 Interpretation of Rates of Gaseous
Reactions at Molecular Level (SB p.54)
Distribution of Molecular Speeds in a Gas
Why is there a distribution of molecular speeds
even at a fixed temperature?
14
15.2 Interpretation of Rates of Gaseous
Reactions at Molecular Level (SB p.55)
Change in the Most Probable Speed

• Increase in most probable speed
• Curve becomes flattened
• Wider distribution of molecular speeds at a
  higher temp

Change in the distribution of molecular speeds as
temperature increases
15
15.2 Interpretation of Rates of Gaseous
Reactions at Molecular Level (SB p.55)
Increased proportion of fast-moving molecules
The proportion of fast-moving molecules increases
as temperature rises
16
15.2 Interpretation of Rates of Gaseous
Reactions at Molecular Level (SB p.56)
Simple Collision Theory

• Collision theory states that
• Reactant molecules must collide with each other
  to react
• The collisions that result in a reaction are
  called effective collisions
• The rate of a reaction is closely related to
  the frequency of collisions of molecules

17
15.2 Interpretation of Rates of Gaseous
Reactions at Molecular Level (SB p.56)
In Terms of Activation Energy

• The colliding molecules must have enough kinetic
  energy (i.e. activation energy) to break the
  bonds in the reactants
• Only collisions of molecules with kinetic energy
  greater than or equal to the activation energy
  lead to the formation of products

18
15.2 Interpretation of Rates of Gaseous
Reactions at Molecular Level (SB p.56)
In Terms of Activation Energy
Fraction of molecules with kinetic energy greater
than or equal to the activation energy
19
15.2 Interpretation of Rates of Gaseous
Reactions at Molecular Level (SB p.56)
In Terms of Collision Orientation
e.g. HCl(g) NH3(g) ?? NH4Cl(s)
Proper Orientation
20
15.2 Interpretation of Rates of Gaseous
Reactions at Molecular Level (SB p.57)
In Terms of Collision Orientation
Improper Orientation
21
15.2 Interpretation of Rates of Gaseous
Reactions at Molecular Level (SB p.57)
In Terms of Collision Orientation
Improper Orientation
22
15.2 Interpretation of Rates of Gaseous
Reactions at Molecular Level (SB p.57)
Explanation of Effect of Temperature on Rates of
Reactions in Terms of Collision Theory

• Higher temp
• ? higher K.E. of molecules
• ? Greater fraction of molecules can overcome
  the Ea
• ? no. of effective collisions increases
• ? reaction rate increases

23
15.2 Interpretation of Rates of Gaseous
Reactions at Molecular Level (SB p.57)
Explanation of Effect of Temperature on Rates of
Reactions in Terms of Collision Theory
The proportion of molecules having kinetic energy
greater than the activation energy increases as
temperature increases
24
15.2 Interpretation of Rates of Gaseous
Reactions at Molecular Level (SB p.58)
Explanation of Effect of Concentration on Rates
of Reactions in Terms of Collision Theory

• Higher concentrations of reactants
• ? collision frequency increases
• ? no. of effective collisions increases
• ? reaction rate increases

25
Energy Profile
26
15.3 Energy Profile (SB p.58)
Energy Profile
The energy profile is a graph showing the changes
in potential energy during a reaction.
27
15.3 Energy Profile (SB p.58)
Energy Profile
Energy profile of a single-stage exothermic
reaction
28
15.3 Energy Profile (SB p.59)
Single-Stage Reaction

• Chemical reactions take place in one step
• A-B C ? ABC ? A B - C
• In the transition state,
• Bond between A and B is partially broken
• Bond between B and C is partially formed

29
15.3 Energy Profile (SB p.59)
Single-stage Reaction
Energy profile of a single-stage exothermic
reaction
30
15.3 Energy Profile (SB p.59)
Example of Single-stage Reaction

• Substitution reaction of 1-bromobutane and
  hydroxide ions

31
15.3 Energy Profile (SB p.60)
Multi-stage Reaction

• Chemical reactions take place in two or more
  steps
• Formation of an intermediate

32
15.3 Energy Profile (SB p.60)
Multi-stage Reaction

• The slowest stage in the reaction mechanism is
  the rate determining step
• Involves the greatest amount of Ea

33
15.3 Energy Profile (SB p.60)
Multi-stage Reaction
Energy profile of a two-stage reaction
34
15.3 Energy Profile (SB p.61)
Example of Multi-stage Reaction

• Hydrolysis of 2-bromo-2-methylpropane

35
15.3 Energy Profile (SB p.61)
36
Effect of Catalysts on Rates of Reactions
37
15.4 Effect of Catalysts on Rates of Reactions
(SB p.63)
15.4 Effect of Catalysts on Rates of Reactions
(SB p.63)
Positive Catalysts and Negative Catalysts

• Catalysis is the action of the catalyst on the
  reaction
• A positive catalyst is one that speeds up a
  reaction
• A negative catalyst is one that slows down a
  reaction

38
15.4 Effect of Catalysts on Rates of Reactions
(SB p.63)
Working Principle of Catalysts and their Effects
on Reaction Rates

• By providing an alternative pathway for the
  reaction to take place

39
15.4 Effect of Catalysts on Rates of Reactions
(SB p.63)
Working Principle of Catalysts and their Effects
on Reaction Rates

• Positive catalyst
• Provide an alternative pathway with a lower
  activation energy
• Smaller Ea
• ? Greater fraction of molecules with K.E.
  greater than or equal to Ea
• ? Reaction proceeds faster

40
15.4 Effect of Catalysts on Rates of Reactions
(SB p.63)
Working Principle of Catalysts and their Effects
on Reaction Rates

• Negative catalyst
• Provide an alternative pathway with a higher
  activation energy
• Greater Ea
• ? Smaller fraction of molecules with K.E.
  greater than or equal to Ea
• ? Reaction proceeds slower

41
15.4 Effect of Catalysts on Rates of Reactions
(SB p.63)
Working Principle of Catalysts and their Effects
on Reaction Rates
Effect of catalysts on the fraction of molecules
possessing kinetic energy greater than or equal
to the activation energy
42
15.4 Effect of Catalysts on Rates of Reactions
(SB p.64)
Working Principle of Catalysts and their Effects
on Reaction Rates
Energy profiles of the uncatalyzed and catalyzed
pathways of a reaction
43
15.4 Effect of Catalysts on Rates of Reactions
(SB p.64)
Working Principle of Catalysts and their Effects
on Reaction Rates

• A catalyst alters the rate of a reaction
• Remains chemically unchanged at the end of the
  reaction

44
15.4 Effect of Catalysts on Rates of Reactions
(SB p.64)
Homogeneous Catalysis and Heterogeneous Catalysis
Catalyst
HeterogenousCatalyst
HomogenousCatalyst
(Reactants catalyst are NOT in the same phase)
(Reactants catalyst are in the same phase)
45
15.4 Effect of Catalysts on Rates of Reactions
(SB p.65)
Homogeneous Catalysis Intermediate Formation
46
15.4 Effect of Catalysts on Rates of Reactions
(SB p.65)
Homogeneous Catalysis Intermediate Formation
Uncatalyzed esterification of methanoic acid and
ethanol
47
15.4 Effect of Catalysts on Rates of Reactions
(SB p.65)
Homogeneous Catalysis Intermediate Formation
Catalyzed esterification of methanoic acid and
ethanol
48
15.4 Effect of Catalysts on Rates of Reactions
(SB p.66)
Homogeneous Catalysis Intermediate Formation

• H ions act as homogeneous catalyst
• Protonated carboxylic acid is formed as the
  intermediate
• Carbonyl carbon atom becomes more
  electron-deficient
• More easily attacked by the O atom from the
  alcohol
• H ions regenerated at the end of reaction

49
15.4 Effect of Catalysts on Rates of Reactions
(SB p.65)
Homogeneous Catalysis Intermediate Formation
Energy profiles of the uncatalyzed and
acid-catalyzed esterification of methanoic acid
and ethanol
50
15.4 Effect of Catalysts on Rates of Reactions
(SB p.66)
Heterogeneous Catalysis Adsorption
51
15.4 Effect of Catalysts on Rates of Reactions
(SB p.67)
Heterogeneous Catalysis Adsorption
Energy profiles of the uncatalyzed and catalyzed
decomposition of hydrogen peroxide solution
52
Applications of Catalysts
15.4 Effect of Catalysts on Rates of Reactions
(SB p.67)
Industrial Catalysts
53
Applications of Catalysts
15.4 Effect of Catalysts on Rates of Reactions
(SB p.68)
3. Nickel, platinium or palladium is used in the
hydrogenation of unsaturated oils to make
margarine
54
Applications of Catalysts
15.4 Effect of Catalysts on Rates of Reactions
(SB p.68)
55
Applications of Catalysts
15.4 Effect of Catalysts on Rates of Reactions
(SB p.68)
Catalytic Converters in Car Exhaust Systems
56
Applications of Catalysts
15.4 Effect of Catalysts on Rates of Reactions
(SB p.69)
57
Applications of Catalysts
15.4 Effect of Catalysts on Rates of Reactions
(SB p.69)
Enzymes in the Production of Alcoholic Drinks
58
The END
59
15.1 Activation Energy and Arrhenius Equation
(SB p.51)
Example 15-1A
For the following reaction C6H5N2 Cl(aq)
H2O(l) ?? C6H5OH(aq) N2(g) H(aq)
Cl(aq) the rate constants of the reaction at
different temperatures were measured and recorded
in the following table
60
15.1 Activation Energy and Arrhenius Equation
(SB p.51)
Example 15-1A
Determine the activation energy
graphically. (Given R 8.314 J K1 mol1)
Answer
61
15.1 Activation Energy and Arrhenius Equation
(SB p.52)
Example 15-1A
62
15.1 Activation Energy and Arrhenius Equation
(SB p.52)
Example 15-1A
63
15.1 Activation Energy and Arrhenius Equation
(SB p.52)
Back
Example 15-1A
64
The rate constant for a reaction at 110C is
found to be twice the value of that at 100C.
Calculate the activation of the reaction.(Given
R 8.314 J K-1 mol-1)
15.1 Activation Energy and Arrhenius Equation
(SB p.53)
Back
Example 15-1B
Answer
65
15.1 Activation Energy and Arrhenius Equation
(SB p.53)
Check Point 15-1
(a) The reaction 2A(g) B(g) ?? C(g) was
studied at a number of temperatures, and the
following results were obtained Determine
the activation energy of the reaction
graphically. (Given R 8.314 J K1 mol1)
Answer
66
15.1 Activation Energy and Arrhenius Equation
(SB p.53)
Check Point 15-1
67
15.1 Activation Energy and Arrhenius Equation
(SB p.53)
Check Point 15-1
68
15.1 Activation Energy and Arrhenius Equation
(SB p.53)
Check Point 15-1
69
15.1 Activation Energy and Arrhenius Equation
(SB p.53)
Check Point 15-1
Answer
70
15.1 Activation Energy and Arrhenius Equation
(SB p.53)
Check Point 15-1
Back
71
15.2 Interpretation of Rates of Gaseous
Reactions at Molecular Level (SB p.58)
Check Point 15-2
(a) Explain why not all collisions between
reactant molecules lead to the formation of
products.
Answer
(a) For a reaction to occur, colliding molecules
must have kinetic energy equal to or greater than
the activation energy to break the bonds in the
reactants, so that new bonds can form in the
products. Moreover, the collision must be in the
right geometrical orientation, and the atoms to
be transferred or shared do not come into direct
contact with each other, so that the atoms can
rearrange to form products. Products cannot be
formed if the kinetic energy of the reactant
molecules cannot overcome the activation energy,
or the collision orientation is not appropriate.
72
15.2 Interpretation of Rates of Gaseous
Reactions at Molecular Level (SB p.58)
Check Point 15-2
(b) Describe the effect of temperature on the
distribution of molecular speeds in a gaseous
system.
Answer
(b) An increase in temperature will lead to an
increase in the most probable speed of the
molecules. The peak of the curve of
Maxwell-Boltzmann distribution of molecular
speeds shifts to the right and the curve becomes
flattened. This indicates that the distribution
of molecular speed becomes wider and the number
of molecules having the most probable speed
decreases.
73
15.2 Interpretation of Rates of Gaseous
Reactions at Molecular Level (SB p.58)
Check Point 15-2
Back
(c) Explain why the rates of chemical reactions
increase with temperature.
Answer
(c) As temperature rises, the proportion of
fast-moving molecules increases. The kinetic
energy of the molecules also increases. A greater
fraction of molecules can overcome the activation
energy required for a reaction to occur.
Therefore, the number of effective collisions
increases and hence the rates of chemical
reactions increase.
74
15.3 Energy Profile (SB p.60)
Back
Check Point 15-3A
Draw an energy profile of a typical single-stage
endothermic reaction.
Answer
75
The energy profile of a multi-stage reaction is
shown below
15.3 Energy Profile (SB p.61)
Example 15-3
76
15.3 Energy Profile (SB p.62)
Back
Example 15-3
(a) Which stage is the rate determining step?
Explain your answer. (b) Is the reaction
exothermic or endothermic? Explain your answer.
Answer
(a) Stage 2 is the rate determining step. It is
because stage 2 has the greatest amount of
activation energy. (b) The reaction is
exothermic. It is because the potential energy
of the products is lower than that of the
reactants.
77
15.3 Energy Profile (SB p.62)
Check Point 15-3B
Referring to the energy profiles below, answer
the questions that follow. A
B
78
15.3 Energy Profile (SB p.62)
Check Point 15-3B
Referring to the energy profiles below, answer
the questions that follow. C
D
79
15.3 Energy Profile (SB p.62)
Check Point 15-3B
(a) Which reaction(s) is/are exothermic? (b) Which
reaction is the fastest? (c) Which reaction has
the greatest amount of activation energy?
Answer

• A, B and C
• B
• D

Back
80
15.4 Effect of Catalysts on Rates of Reactions
(SB p.69)
Check Point 15-4
(a) Explain what a negative homogeneous catalyst
is.
Answer

• A negative homogeneous catalyst is a catalyst
  that slows down a reaction. It exists in the same
  phase as the reactants and products in the
  reaction, and involves in the formation of an
  intermediate in the reaction.

81
15.4 Effect of Catalysts on Rates of Reactions
(SB p.69)
Check Point 15-4
(b) Explain what a positive heterogeneous
catalyst is.
Answer
(b) A positive heterogeneous catalyst is a
catalyst that speeds up a reaction but it is not
in the same phase as the reactant and products.
It provides an active surface for the reactant
particles to adsorb in a reaction.
82
15.4 Effect of Catalysts on Rates of Reactions
(SB p.69)
Back
Check Point 15-4
(c) Give three applications of catalysts.
Answer
(c) Iron used in the Haber process Platinum or
vanadium(V) oxide used in the Contact
process Nickel, platinum or palladium used in
the hydrogenation of unsaturated oils to make
margarine Nickel and nickel(II) oxide used in
the production of town gas Platinum (or
palladium) and rhodium used in catalytic
converters Enzymes used in fermentation of
glucose to produce ethanol Enzymes used in the
manufacture of biological washing powders. (any
3)