Download Lecture 1 - Introduction to catalysis

Industrial Catalysis
Lecture for Makrokierunek
Nikodem Kuźnik
Lecture 1
Scope of the lecture
• Role of catalysis in industrial chemistry, chemical lab and
life
• History of catalysis – the biggest inventors, discoveries,
improvements
• Characteristic concepts for catalysis: activity, selectivity
and lifetime
• Definition of catalyst, thermodynamic and kinetic
approach
• How does a catalyst work?
• Types of catalysis: homogeneous, heterogeneous,
autocatalysis, inhibition
• Few example of catalysis
Role of catalysis
Catalysts are big business.
The chemical industry depends upon catalysts.
We depend upon the chemical industry for our 21st
century life style.
We depend on catalysts.
Moreover, most of the chemical reactions in nature
are biocatalytic reactions.
History of catalysis
-
-
Over 2000 BC – enzymatic (biocatalytic) production of alcohol
1831 Catalytic (Pt) oxidation of SO2 to SO3 to produce H2SO4
1899 Catalytic (Ni) hydrogenation of olefins (P. Sabatier, J. B.
Sendérens)
1918 Catalytic synthesis of ammonia (F. Haber & C.Bosch)
1949: First commercial operation o f Platforming Process reforming
Early 1950s: commercial synthesis for zeolites
1953: Catalyst system for polymerizing ethylene at low temperature
and pressure - Nobel Prize awarded to Ziegler in 1963.
1954: "Beginning" of catalyst characterizations using instruments
1964: Olefin metathesis announced
1964: Startup by Monsanto of the world's first biodegradable detergents
1974: Auto Exhaust Treatment system accepted by Chrysler
1980s: beginning of enantioselective catalysis
Role of catalysis
• 80% of processes in the chemical industry use
catalysts.
• Wolrd catalyst demant in 2010 was 15 milliard
(15∙109) dollars.
• Growth in catalyst sales is increasing at between
5% and 10% per year.
• Turnover in industries using catalysts was about
14 trillion (14∙1012) dollars which is equal to the
gross domestic product of USA (2010).
• The conversion of coal or natural gas to anything
uses a catalyst.
What do they do for the chemical
industry?
Catalysts have an enormous impact on the chemical industry
because they:
• enable reactions to take place
• make processes more efficient
• a 0.5% to 1% increase in selectivity can lead to a up to
1 million dollar increase in operating profit
• make processes environmentally friendly.
Definition of a catalyst
A catalyst is a substance that accelerates the rate
of a chemical reaction, at some temperature, but
without itself being transformed or consumed by
the reaction.
A catalyst participates in the reaction but is neither
a chemical reactant nor a chemical product.
Catalysts provide an alternative pathway with a
lower activation energy, for a reaction to proceed.
Characteristic concepts for catalysis
• activity
• selectivity
• lifetime
Activity
• How fast reaction proceeds in the presence
of the catalyst
• Theory
– Reaction rate
– Rate constant k
– Activation energy Ea
Impact of catalyst on chemical reaction
– thermodynamic approach
E
E*
E*
cat
Reactants
Products
Reaction course
k  Ae
 E* /RT
Impact of catalyst on chemical reaction
– kinetic model and thermodynamic
approach
Activity
• Practice in industry
– Conversion under const. reaction conditions
– Space velocity for given constant conditions
𝑉0
𝑚3 𝑘𝑔−1 𝑠 −1
𝑚𝑐𝑎𝑡
V0 – volume flow rate, mcat – catalyst mass
𝑆𝑝𝑎𝑐𝑒 𝑣𝑒𝑙𝑜𝑐𝑖𝑡𝑦 =
– Space-time-yield (SPY)
𝐷𝑒𝑠𝑖𝑟𝑒𝑑 𝑝𝑟𝑜𝑑𝑢𝑐𝑡 𝑞𝑢𝑎𝑛𝑡𝑖𝑡𝑦
𝑆𝑇𝑌 =
(𝑚𝑜𝑙 𝐿−1 𝑕−1 )
𝐶𝑎𝑡𝑎𝑙𝑦𝑠𝑡 𝑣𝑜𝑙𝑢𝑚𝑒 ∙ 𝑡𝑖𝑚𝑒
– Temperature required for given conversion
Activity, lifetime
• Turnover frequency (TOF)
𝑣𝑜𝑙𝑢𝑚𝑒𝑡𝑟𝑖 𝑟𝑎𝑡𝑒 𝑜𝑓 𝑟𝑒𝑎𝑐𝑡𝑖𝑜𝑛
𝑇𝑂𝐹 =
(𝑠 −1 )
𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑐𝑒𝑛𝑡𝑒𝑟𝑠 ∙ 𝑣𝑜𝑙𝑢𝑚𝑒
𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑐𝑜𝑛𝑣𝑒𝑟𝑡𝑒𝑡 𝑠𝑢𝑏𝑠𝑡𝑟𝑎𝑡𝑒 −1
𝑇𝑂𝐹 =
(𝑠 )
𝑡𝑖𝑚𝑒 ∙ 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑐𝑎𝑡𝑎𝑙𝑦𝑠𝑡
Typical catalyst 10-2-102 s-1, enzymes 103-107 s-1
• Turnover number (TON)
TON = TOF [time-1]∙litefime of the catalyst [time]
Typical industrial catalyst 106-107
Autocatalysis
A single chemical reaction is said to have
undergone autocatalysis, or be
autocatalytic, if the reaction product is
itself the catalyst for that reaction.
Autocatalysis
Oxidation of oxalic acid with potassium permanganate:
2MnO4-(aq ) + 5H2C2O4(aq ) + 6H3O+(aq ) 
 2Mn2+(aq ) + 10CO2(aq ) + 14 H2O
KMnO4 is added.
MnSO4 is added to It catalyzes the
the solution on the reduction of
right.
KMnO4 to
colorless Mn2+.
Source: J.Chem.Edu.
The other
solution's reaction
rate eventually
increases as it
forms Mn2+, which
subsequently
autocatalyzes its
own formation.
Inhibition
An Inhibitor is a material (substance) which
is added to some system in order to retard
certain chemical reaction(s).
For example, 2,6-di-tert-butyl-p-cresol may be added to
tetrahydrofuran to prevent potentially dangerous
polymerization. Inhibitors are often added to chemicals
which tend to undergo self-induced free-radical
polymerization.
Types of catalysis
Heterogeneous,
when catalyst is in
separate phase
Homogeneous,
when all reactants
and catalyst are in
the same phase
Classification of Catalysts
Catalysts
Heterogenized
Homogeneous
catalysts
Homogeneous
Catalyst
Acid/base
catalysts
Transition
metal
compounds
Biocatalysts
(enzymes)
Heterogeneous
catalysts
Bulk
catalysts
Supported
catalysts
How does a catalyst work?
• Intermediate
compounds
H3C CH3
• Adsorption
H H
[Rh]
H3C CH2 [Rh] H
H [Rh] H
H
H
H
[Rh] H
H
H
H
H
H
H
Examples of catalytic reactions
H3O+/OH-
The experimentally determined rate law is of first
order with respect to the concentration of the
reagent, but the rate constant also depends upon
the pH :
The pH-dependence can be quantified as follows :
kobs = k0 + kH·[H+] + kOH·[OH-]
In this example, catalysis is performed either by
H3O+, or by OH-, as is often the case in such
hydrolysis reactions.
Examples of catalytic reactions
Hydrogenation of C=C