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
Activation Energy
Exothermic reaction
related to the rate of reaction
Activation energy, Ea energy required to start
the reaction
4
Activation Energy
Endothermic reaction
related to the rate of reaction
5
Most reactions have positive Ea since energy is
absorbed to break bonds in reactant particles.
6
Arrhenius Equation
Since rate kAaBb...
At fixed concentrations, rate depends on k which
in turn depends on
temperature (T) and the nature of the reaction (A
and Ea)
7
Q.23
R 8.31 J K?1 mol?1
Assume A is a constant
1.92 ? 2
A 10 K ? in T doubles the rate
8
Arrhenius Equation
Rate of reaction ? exponentially with temperature
9
Arrhenius Equation
T ?, ? A ?
(Minor effect)
? Rate ?
10
Arrhenius Equation
T ?,
(less negative)
(more positive)
(Major effect)
? Rate ?
11
Arrhenius Equation
T ?, ? A ?
(Minor effect)
? Rate ?
12
Arrhenius Equation
T ?,
(more negative)
(less positive)
(Major effect)
? Rate ?
13
Determination of Activation Energy
14
Determination of Ea by Graphical Method
logeA
15
Q.24
logek 1/T (K?1)
?14.9 1.80?10?3
?9.4 1.55?10?3
?6.8 1.43?10?3
?3.2 1.28?10?3
16
Ea -slope ? R ? 182 kJ mol?1
lt E(H I)
H H and I I bonds are formed
before the H I bond is completely broken (refer
to p.32)
17
If logerate is plotted against 1/T, since rate
kAxBy logerate logek logeAxBy
logek constant
18
Determination of Activation Energy Using Two Rate
Constants
19
Interpretation of Rates of Gaseous Reactions at
Molecular Level
20
The Kinetic Theory of Gases
- Developed by Maxwell and Boltzman
1. Gas particles are in a state of constant and
random motion in all directions, undergoing
frequent collisions with one another and with
the walls of the container.
2. The pressure exerted on the container is due
to the collisions between gas particles and the
walls of the containers.
21
The Kinetic Theory of Gases
- Developed by Maxwell and Boltzman
3. Gas particles are treated as point masses
because their volumes are negligible when
compared with the volume of the container.
4. There is no interaction among gas particles
except collisions.
22
The Kinetic Theory of Gases
- Developed by Maxwell and Boltzman

• Collisions between gas particles are perfectly
  elastic,
• i.e. the total kinetic energy is conserved.

23
Distribution of Molecular Speeds in a Gas
Consider a sample of gas
Transfer of K.E. among molecules ? Distribution
of molecular speeds
24
Distribution of Molecular Speeds in a Gas
25
The Kinetic Theory of Gases
- Developed by Maxwell and Boltzman
The mean kinetic energy of a sample of gas
particles is proportional to its absolute
temperature (T).
26
The Kinetic Theory of Gases
- Developed by Maxwell and Boltzman
n1 no. of molecules with velocity c1 and n1
n2 n3 ... n (total no. of molecules)
27
The distribution of velocity is not symmetrical
28
The Kinetic Theory of Gases
- Developed by Maxwell and Boltzman
For a sample of gas containing n molecules,
where m is the absolute mass of the gas molecule
29
Variation in the distribution of molecular Speeds
with T

• As T ?,
• Molecular speeds ?
• Curve becomes flattened
• Wider distribution of molecular speeds at a
  higher temp
• Area under the curve remains unchanged.

30
If n 1
M molar mass of gas in Kg M ? ? r.m.s velocity ?
31
The areas underneath the curves are the same The
lighter molecules are more spread out in
molecular speeds.
32
Q.25
H2
1845 ms?1
CO2
393 ms?1
The lightest gases (H2, He) can escape from the
gravitational pull of small planets ? Very rare
in the Earths atmosphere
33
Simple Collision Theory

• For a reaction to occur, the reactant particles
  must collide with
• kinetic energy ? Ea
• proper orientation.

34
HCl(g) NH3(g) ?? NH4Cl(s)
Proper Orientation
35
Improper Orientation
36
Improper Orientation
37
Simple Collision Theory

• For a reaction to occur, the reactant particles
  must collide with
• kinetic energy ? Ea
• proper orientation.

Z collision frequency
38
Simple Collision Theory

• For a reaction to occur, the reactant particles
  must collide with
• kinetic energy ? Ea
• proper orientation.

39
Simple Collision Theory

• For a reaction to occur, the reactant particles
  must collide with
• kinetic energy ? Ea
• proper orientation.

p fraction of collisions with proper
orientation
40
Theoretically, (from collision theory and kinetic
theory)
Experimentally,
41
If molarities of X, Y, are fixed
no. of collisions with proper orientation
42
Interpretation of the Effect of Temperature
Change on Rate of Reaction

• T ?
• speed of reactant particles ?
• collision frequency (Z) ?
• A ?
• ? rate ? (minor effect)

43
Interpretation of the Effect of Temperature
Change on Rate of Reaction

• T ?
• K.E. of reactant particles ?
• fraction of collisions with K.E. ? Ea ?

? rate ? exponentially (major effect)
44
No. of molecules
45
The shaded area no. of particles with K.E. gt E
46
If E Ea
47
As T ?, the fraction of particles with K.E. gt
Ea increases exponentially.
? Rate increases exponentially with T
48
Limitations of Collision Theory
Collision theory is based on the calculations
from kinetic theory of ideal gases. Thus, it is
ONLY applicable to reactions in gas phase.
In aqueous phase, the interactions between the
reactant particles and the solvent molecules have
to be considered.
The fraction of collisions with proper
orientation (the steric factor, p) cannot be
predicted. It can only be determined
experimentally.
49
Q.26
Consider the 2nd order single-step gas phase rx
R(g) R(g) ? products
Given k 1.00?10?2 mol?1 dm3 s?1 at 473 K,
R(g)initial 1?10?2 mol dm?3, L 6.02?1023
mol?1 Gas constant 8.31 J K?1 mol?1, Ea 100
kJ mol?1 Initial collision frequency(Z)
7.77?1032 s?1

• Estimate
• (i) The no. of effective collisions per m3 per
  second.
• (ii) The no. of collisions with K.E. ? Ea per m3
  per second.
• (b) Hence, deduce the steric factor, p of the
  reaction.

50
(a)(i)
(1.00?10-2 mol?1 dm3 s?1)(1.00?10-2 mol
dm?3)2 1.00?10-6 mol dm?3 s-1 1.00?10-6 mol ?
6.02?1023 mol-1 dm-3 s-1 6.02?1017 molecules
dm-3 s-1 6.02?1020 molecules m-3 s-1
6.02?1020 molecules of R are decomposed per cubic
meter per second
51
(a)(i)
Consider the 2nd order single-step gas phase rx
R(g) R(g) ? products
Rate 6.02?1020 molecules m-3 s-1
One effective collision leads to decomposition of
Two molecules of R. Thus, no. of effective
collisions per cubic meter per second 3.01?1020
52
(a)(ii)
3.01?1020 m-3 s-1
No. of collisions with K.E. ? Ea
7.77?1032 m-3 s-1 ? (8.93?10-12) 6.94?1021
m-3 s-1
53
(b)
3.01?1020 m-3 s-1
No. of collisions with K.E. ? Ea
6.94?1021 m-3 s-1
4.34
54
Q.27
If Ea ? 0
Rate is independent of T (A-level)
55
No. of molecules
56
No. of molecules
? Z ? p
57
Energy Profile Transition State Theory
58
Transition State Theory - focuses on what
happens after the collisions have started.
59
Energy profile - shows the variation of the
potential energy of the reaction mixture as the
reaction proceeds.
P.E.
reaction coordinate
60
Consider the one-step reaction, AB X
? A BX
P.E. of the reaction mixture are calculated for
any AB and BX distances, and the results are
plotted on a contour diagram
61
Consider the one-step reaction, AB X
? A BX
At R, A-B distance is short B-X distance is
long ? before reaction
62
Consider the one-step reaction, AB X
? A BX
The valley at R represents the potential energy
for the initial state of the system, i.e. AB and
X
63
Consider the one-step reaction, AB X
? A BX
At P, A-B distance is long B-X distance is
short ? after reaction
64
Consider the one-step reaction, AB X
? A BX
The valley at P represents the potential energy
for the final state of the system, i.e. A and BX
65
Consider the one-step reaction, AB X
? A BX
The energy contours rise in all directions from
the valleys at R and P,
but the easiest path is shown by the bold line
RTP
66
The transition state is like a col (??) in a
mountain region
67
(No Transcript)
68
A-B X ? A?????B?????X ? A B-X

• In the transition state,
• Bond between A and B is partially broken
• Bond between B and X is partially formed

Thus, Ea is lower than E(A-B)
69
Transition state (Activated complex) is the least
stable arrangement of the system in the most
probable reaction pathway.
70
Advantages of Transition State Theory
1. Ea and A can be calculated A ? Zp ? the
steric factor p can be predicted 2. It explains
why the reaction pathway is specific. 3. It is
applicable to gaseous and aqueous reactions.
71
Energy Profile One-step Mechanism
A-B X ? A?????B?????X ? A B-X
72
Example of One-step Mechanism
Rate kCH3ClOH?
Bimolecular One-step 2nd Order Nucleophilic
Substitution Reaction
SN2
73
CH3Cl OH? ? CH3OH Cl?
74
Energy Profile Multi-step Mechanism
E1 gt E2 ? Step 1 is the rate determining
step
75
Energy Profile Multi-step Mechanism
76
Multi-step Mechanism

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

77
Example of Multi-step Mechanism

• Hydrolysis of 2-chloro-2-methylpropane

(1)
(2)
78
Rate kC(CH3)3Cl
Unimolecular Two-step 1st Order Nucleophilic
Substitution Reaction
SN1
79
E1 gt E2 Step 1 is r.d.s.
Rate kC(CH3)3Cl
80
Reaction Mechanism and Rate Law
Reaction mechanisms are theoretical proposals
used to explain the experimentally determined
rate laws.
Reaction mechanism is the detailed sequence of
steps that occur in a reaction .
81
Reaction Mechanism and Rate Law
Each of the steps in a mechanism is called an
elementary step.
The number of reactant particles that takes part
in each elementary step is called the
molecularity of that step.
82
Reaction Mechanism and Rate Law
The number of reactant particles that takes part
in each elementary step is called the
molecularity of that step.
Unimolecular one particle collides with the
wall of the vessel or the excess solvent
83
Pseudo-1st order reaction
CH3COOCH3 H2O ? CH3COOH CH3OH
Rate kCH3COOCH3H2O If H2O is used as
solvent (in large excess) H2O ? constant
throughout the reaction Rate kCH3COOCH3 Unimo
lecular reaction
84
Reaction Mechanism and Rate Law
The number of reactant particles that takes part
in each elementary step is called the
molecularity of that step.
Unimolecular one particle collides with the
wall of the vessel or the excess solvent
Bimolecular two particles collide together
Termolecular three particles collide together
simultaneously (very rare)
85
The slowest step in a particular mechanism is
called the rate-determining step
Requirements for writing reaction mechanisms
1. The sum of elementary steps must give the
overall balanced equation for the reaction.
2. The mechanism must agree with the
experimentally determined rate law.
86
Consider the reaction
Rate kAB Only one intermediate
Proposed mechanism -
87
Q.28
88
E1 gt E2 ? E3 Step one is the r.d.s
Q.28
Rate kAB
P.E.
Reaction coordinate
89
Q.28
k
Rate kX
At equilibrium, k1AB k2X
90
Effect of Catalysts on Rates of Reactions
91
Working Principle of Catalysts and their Effects
on Reaction Rates
Catalysis ? Catalytic action
Catalysts alter the rates of reaction, 1. but
remain chemically unchanged at the end of the
reaction 2. by providing new, alternative
reaction pathways with different activation
energies.
92
Working Principle of Catalysts and their Effects
on Reaction Rates

• Positive catalyst
• Provides an alternative reaction pathway with a
  lower activation energy

93

• Lower Ea

? Greater fraction of molecules with K.E.
greater than or equal to Ea ? Reaction proceeds
faster
94
Working Principle of Catalysts and their Effects
on Reaction Rates

• Negative catalyst
• Provides an alternative reaction pathway with a
  higher activation energy

95

• Higher Ea

? Smaller fraction of molecules with K.E.
greater than or equal to Ea ? Reaction proceeds
slower
96
Working Principle of Catalysts and their Effects
on Reaction Rates
With catalysts, the contour diagrams and thus the
energy profiles are totally different from those
without
97
Catalyst
Reactants catalyst are in the same phase
Reactants catalyst are NOT in the same phase
98
Characteristics of Catalysts

1. For a given reversible reaction,

catalysts affect the rates of forward reaction
and backward reaction to the same extent.
99
Q.29
Without catalyst
With catalyst
100

101
Given T 298 K, R 8.31 J K?1 mol?1
5.9?108
102
Characteristics of Catalysts
2. Catalysts are chemically unchanged at the end
of reactions, but may undergo physical changes.
E.g. Lumps of MnO2 used in the decomposition
of H2O2 become powdered at the end of the
reaction.
103
Characteristics of Catalysts
3. Only small quantity is sufficient to catalyze
a reaction because catalysts can be regenerated.
However, if the catalysts are involved in the
rate equation, higher concentrations may affect
the rate more.
104
Characteristics of Catalysts
4. The effect of heterogeneous catalysts depends
on the surface area available for the catalytic
action.
Surface area of solid catalyst ? ? number of
reaction sites ? ? catalytic activity?
E.g. Finely divided Fe powder is used as the
catalyst in Haber process.
105
Characteristics of Catalysts
5. Catalytic actions are specific especially in
biological systems.
E.g. Enzymatic actions are highly specific.
106
Characteristics of Catalysts
6. The efficiency of a catalyst is often
enhanced by adding promoters. Promoters have no
catalytic actions on their own.
E.g. Fe2O3, KOH, Al2O3 in Haber process
107
Characteristics of Catalysts
7. The efficiency of a catalyst can be lowered
by adding poisons or inhibitors. Catalyst
poisons are specific in action.
E.g. Arsenic impurities may poison Pt but
not V2O5 in Contact process
108
Characteristics of Catalysts
8. Transition metals or compounds/ions
containing transition metals show marked
catalytic activities.
E.g. Pt, Ni, Fe, V2O5, MnO2, Mn2 Fe3
The catalytic actions are due to the presence
of low-lying partially filled d-orbitals.
109
Heterogeneous Catalysis Adsorption

• Occur on the surface of the catalyst.
• Reactants are adsorbed on the surface,

forming new bonds with the catalyst
while weakening bonds in reactants
2. Products, once formed, are desorbed from the
surface,
110
Examples of heterogeneous catalysis
111
CH2CH2 HH
slow
fast
fast
112
CH2CH2 HH
slow
fast
fast
113
CH2CH2 HH
slow
fast
fast
114
Q.30
115
Homogeneous Catalysis Intermediate Formation
Homogeneous catalysts participate in certain
stages of reactions and are regenerated at the
end or later stages of reactions.
116
Homogeneous Catalysis Intermediate Formation
117
Homogeneous Catalysis Intermediate Formation
118
Q.31
nucleophilic attack
?
Rate kCH3COOHCH3OH
Uncatalyzed esterification
119
Q.31
?
Uncatalyzed esterification
120
Acid-catalyzed esterification
Protonation at carbonyl O rather than hydroxyl O
since the former is more electron sufficient due
to polarization of pi electron cloud
121
Acid-catalyzed esterification
122
Acid-catalyzed esterification
123
Acid-catalyzed esterification
124
Acid-catalyzed esterification
125
Acid-catalyzed esterification
Most probable resonance structure
Carbonyl C becomes more electron-deficient ? More
easily attacked by nucleophile
126
Acid-catalyzed esterification
127
Acid-catalyzed esterification
Rate kRCOOHH
128
Acid-catalyzed esterification
r.d.s.
step 2
129
Acid-catalyzed esterification
r.d.s.
step 2
step 3
step 4
step 6
step 5
130
Acid-catalyzed esterification
r.d.s.
step 2
step 3
step 4
step 6
131
Acid-catalyzed esterification
r.d.s.
step 2
step 3
step 4
step 5
132
Acid-catalyzed esterification
r.d.s.
step 2
step 3
H is regenerated
step 4
step 6
step 5
133
Acid-catalyzed esterification
r.d.s.
step 2
step 3
For simplicity, steps 3 to 6 are combined
step 4
step 6
step 5
134
Rate kCH3COCH3H
135
(No Transcript)
136
kCH3COCH3H
137
Homogeneous Catalysis Using Transition Metal Ions
Principle - Transition metals exhibit variable
oxidation states
138
2I?(aq) S2O82?(aq) I2(aq)
2SO42?(aq)
The reaction is slow because colliding particles
carry like charges
139
Fe3(aq)
2I?(aq) S2O82?(aq) I2(aq)
2SO42?(aq)
Mechanism of catalyzed reaction -
2I?(aq) 2Fe3(aq) ? I2(aq) 2Fe2(aq)
2Fe2(aq) S2O82?(aq) ? 2Fe3(aq) 2SO42?(aq)
Both steps are fast because colliding particles
carry opposite charges.
140
Fe3(aq)
2I?(aq) S2O82?(aq) I2(aq)
2SO42?(aq)
Mechanism of catalyzed reaction -
2I?(aq) 2Fe3(aq) ? I2(aq) 2Fe2(aq)
2Fe2(aq) S2O82?(aq) ? 2Fe3(aq) 2SO42?(aq)
The mechanism is made possible by the variable
oxidation states of Fe
141
Q.31
Fe2(aq)
2I?(aq) S2O82?(aq) I2(aq)
2SO42?(aq)
Mechanism of catalyzed reaction -
2Fe2(aq) S2O82?(aq) ? 2Fe3(aq) 2SO42?(aq)
2I?(aq) 2Fe3(aq) ? I2(aq) 2Fe2(aq)
142
Mechanism of catalyzed reaction -
Ce4(aq) Mn2(aq) ? Ce3(aq) Mn3(aq)
Ce4(aq) Mn3(aq) ? Ce3(aq) Mn4(aq)
Mn4(aq) Tl(aq) ? Mn2(aq) Tl3(aq)
The mechanism is made possible by the variable
oxidation states of Fe
143
Applications of Catalysts
Industrial Catalysts
144
Applications of Catalysts
3. Nickel, platinium or palladium is used in the
hydrogenation of unsaturated oils to make
margarine
145
Applications of Catalysts
146
Applications of Catalysts
Catalytic Converters in Car Exhaust Systems
147
Applications of Catalysts
148
Applications of Catalysts
Enzymes in the Production of Alcoholic Drinks
Fermentation
149
The END
150
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
151
15.1 Activation Energy and Arrhenius Equation
(SB p.51)
Example 15-1A
Temperature (K) Rate constant (10-5 s-1)
278.0 0.15
298.1 4.10
308.2 20.00
323.0 140.00
Determine the activation energy
graphically. (Given R 8.314 J K1 mol1)
Answer
152
15.1 Activation Energy and Arrhenius Equation
(SB p.52)
Example 15-1A
153
15.1 Activation Energy and Arrhenius Equation
(SB p.52)
Example 15-1A
154
15.1 Activation Energy and Arrhenius Equation
(SB p.52)
Back
Example 15-1A
155
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
156
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)
Temperature (oC) 12 60 112 203 292
Rate constant (dm6 mol-2 s-1) 2.34 13.2 52.5 316 1000
Answer
157
15.1 Activation Energy and Arrhenius Equation
(SB p.53)
Check Point 15-1
158
15.1 Activation Energy and Arrhenius Equation
(SB p.53)
Check Point 15-1
159
15.1 Activation Energy and Arrhenius Equation
(SB p.53)
Check Point 15-1
160
15.1 Activation Energy and Arrhenius Equation
(SB p.53)
Check Point 15-1
Answer
161
15.1 Activation Energy and Arrhenius Equation
(SB p.53)
Check Point 15-1
Back
162
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.
163
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.
164
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.
165
15.3 Energy Profile (SB p.60)
Back
Check Point 15-3A
Draw an energy profile of a typical single-stage
endothermic reaction.
Answer
166
The energy profile of a multi-stage reaction is
shown below
15.3 Energy Profile (SB p.61)
Example 15-3
167
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.
168
15.3 Energy Profile (SB p.62)
Check Point 15-3B
Referring to the energy profiles below, answer
the questions that follow. A
B
169
15.3 Energy Profile (SB p.62)
Check Point 15-3B
Referring to the energy profiles below, answer
the questions that follow. C
D
170
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
171
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

1. 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.

172
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.
173
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)