Catalyst Fundamentals
朱信
Hsin Chu
Professor
Dept. of Environmental Eng.
National Cheng Kung University
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1. The Basics: Activity and Selectivity
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Catalyst
A material that increases the rate of a chemical reaction while itself
not undergoing any permanent change
Complete oxidation example (nonselective)
Pt
C2 H 4  3O2 
 2CO2  2H 2O
The adsorption of C2H4 and O2 onto the catalyst Pt provides a
chemical shortcut → reaction at lower temperatures
Partial oxidation example (selective)
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VO
C2 H 4  O2 
 CH 3CHO (aldehyde)
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Next slide (Fig. 1.1)
Activation energy
2 5
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2
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Both the net enthalpy (△H) and net free energy (△G) are
unaffected by the presence of the catalyst. (and consequently the
equilibrium constant (Ke))
The catalyst influences selectivity by preferentially lowering the
activation energy for a particular step in the reaction sequence and
increases the rate at which this step proceeds.
e.g., V2O5: The reaction path leading to the formation of the
aldehyde product is favored because it has a lower
activation energy than the complete combustion to CO2
and H2O.
Pt: The opposite is true.
In catalytic air pollution control:
All the processes utilize solid heterogeneous catalysts through which
gaseous reactants pass.
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2. Dispersed Catalyst Model
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The number of reactant molecules converted to products in a given
time is directly related to the number of catalytic sites available to
the reactants. → Maximize the number of active sites by dispersing
the catalytic components (Pt, Fe, Ni, Rh, Pd, CuO, PdO, CoO, etc.)
onto a surface
High surface area, porous oxide particles (carriers):
Al2O3, SiO2, TiO2, SiO2-Al2O3 combinations, activated carbon, etc.
Next slide (Fig. 1.2)
Pt catalyst dispersed on a high surface area Al2O3 carrier
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The internal surface of the Al2O3 in rich in surface OH– groups,
depending on the type of Al2O3 and its thermal history.
These OH species, represent sites upon which one can chemically or
physically bond a catalytic substance.
The higher the catalytic surface area, the higher the rate of reaction
for a process controlled by kinetics rather than mass transfer.
The rate is influenced by the hydrodynamics of the fluid flow, the
pore size and structure of the carrier, and the molecular dimensions
of the diffusing molecule.
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3. The Steps in Heterogeneous Catalysis
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catalyst
A 
B
(1) Reactant A → outer surface of the carrier containing the catalytic
sites Bulk molecular diffusion rates vary approximately with T3/2
Typical activation energy, E1=2-4 kcal/mol
(2) Reactant A → internal catalytic sites
Pore diffusion process varies approximately with the square root of k,
the rate constant for the chemical conversion reaction.
Activation energy, E2, is approximately one-half that of a chemical
reaction or about 6~10 kcal/mol.
(3) Reactant A chemisorb onto the catalytic site
Chemisorption rates generally follow exp (-E3/RT).
E3, activatin energy for chemisorptions , is typically greater than 10
kcal/mol.
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(4) Catalytic surface reactions involve chemisorbed A → B Reaction
rates follow exp (-E4/RT).
Activation energy, E4, is typically grater than 10 kcal/mol.
(5) Product B must desorb from the site
Desorption rates follow exp (-E5/RT).
Activation energy, E5, is typically greater than 10 kcal/mol.
(6) Product B diffuse through the porous network
Kinetics are similar to step 2.
(7) Product B diffuse into the bulk gas
Kinetics are similar to step 1.
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Next slide (Fig. 1.3)
Rate-controlling steps versus temperature profile
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4. The Arrhenius Equation
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A simple first order reaction
rate = kCA
k: rate constant
CA: concentration of the reactant A
Arrhenius Equation
 Ea 
k  k0 exp  
 RT 
ln k  ln k0 
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Ea  1 
R  T 
k0: proportional to the number of active sites
Next slide (Fig. 1.4)
Arrhenius plot
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