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Appendix
6. Catalysis
theory and practice
A catalyst is a
substance that alters the rate of chemical reaction without itself
being permanently chemically changed.
It will chemically change
temporarily e.g. change in ligand or oxidation state or other bonding
arrangement, but will return to is original state often via a 2–3
stage 'catalytic cycle'.
A catalyst provides a
reaction pathway with a lower activation energy Ea,
compared to the uncatalysed reaction.
See the diagrams immediately below, simple
exothermic/endothermic reaction, and, more realistically, a complex two
stage cycle profile further on).
Two primary modes
of catalytic action – heterogeneous and homogeneous
In the case of transition
metals, both the 3d and 4s electrons can be involved and in many
cases where there the catalyst is a transition metal compound, the
ability to easily change oxidation state is really important in
facilitating an alternative lower activation energy pathway.
See also my kinetics
notes on
Heterogeneous
and homogeneous catalysis mechanisms
HETEROGENEOUS CATALYSIS: (e.g. nickel catalysing the hydrogenation of an alkene)
-
The catalyst and
reactants are in different phases (usually solid catalyst and
gaseous/liquid reactants).
-
The reaction occurs
on the catalyst surface which may be the transition metal or one
of its compounds e.g. an oxide.
-
The reactants must be
adsorbed onto the catalyst surface at the 'active sites'.
-
This can be physical
adsorption or chemically bonding to the catalyst surface. Either
way, it has the effect of concentrating the reactants close to
each other and weakening the original intra–molecular bonds of
the reactant molecules.
-
Below I've described examples of
heterogeneous catalysis to reduce the activation energy by providing
an alternative pathway via the active sites on the surface of the
catalyst.
-
The strength of
adsorption is crucial to having a 'fruitful' catalyst surface.
-
If
adsorption too strong,
the reactant/product molecules are too strongly 'chemisorped'
inhibiting reaction progress e.g. can happen with tungsten (W).
-
If
adsorption too weak, the
reactants are not chemisorped strongly enough to allow the
initial bond breaking processes to happen e.g. can happen with silver (Ag),
though silver is used in some industrial processes.
-
Just right:
Nickel (Ni), platinum (Pt), rhodium (Rh) etc. will adsorb the
reactants sufficiently to enable the bond breaking process
to be initiated but to not strong to retain the product
molecules. These three metals are used in many industrial
processes e.g. hydrogenating oils to make margarine (Ni) and
catalytic converters in vehicle exhausts (Pt, Rh).
-
It is usual
to use
the catalyst in a finely divided form to maximise surface area to
give the greatest and therefore most efficient rate of reaction.
-
This means the catalyst must be physically supported, since it
will have no bulk strength in its own right e.g.
-
Platinum–rhodium
metal is produced on a temperature resistant ceramic support
in catalytic converters of motor vehicle exhausts.
-
A support medium is essential to
maximise the surface area of a heterogeneous catalyst.
-
Catalyst poisoning
should be avoided. This inhibiting effect is caused by
impurity molecules being strongly chemisorbed on the most active sites
of the catalyst surface.
-
This blocking of the active
catalytic sites considerably reduces the efficiency of the
catalyst and increases production costs if the catalyst has to
be replaced or functions with less efficiency e.g.
-
sulfur poisons
the iron catalyst in the Haber Process for making ammonia,
-
and lead poisons the
platinum–rhodium surface in car exhaust catalytic converters.
-
In the case of transition
metals, the catalytic behaviour usually depends on the solid metal
forming weak temporary bonds with the substrate reactant molecules
(usually gaseous, but can be liquid).
-
The transition metals (1st
series) can use the 3d and 4s electrons on the metal surface to form
these weak temporary bonds.
-
When the chemical transformation
is complete, these bonds (now between metal and product) easily
break to release the product molecules. The hydrogenation
mechanism of alkenes using nickel is illustrated below.
-
Formation of intermediate involving the catalyst - which temporarily
changes.
-
A two stage reaction profile for a
catalytic cycle (Ea = activation energy)
-
This sort of diagram is most
applicable to homogeneous catalysis where definite intermediates are
formed, but in general principle it applies to heterogeneous catalysis
too where the adsorption (particularly chemical) is equivalent to
forming a transition state or complex.
-
Ea1 is the activation energy leading to the formation of an
intermediate complex.
-
Ea2 is the activation energy for the change of the
intermediate complex into products.
-
Ea3 is the activation energy of the uncatalysed reaction.
-
-
An example of heterogeneous catalysis is
illustrated above i.e. the hydrogenation of alkenes (e.g. ethene + hydrogen
===> ethane). Nickel is the solid phase catalyst and the reactant gases in
the different gaseous phase.
-
In terms of activation energies, with
reference to the reaction profile before the nickel diagram:
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(1) ==> (2) is represented by
Ea1, the
absorption of the reactant molecules onto the catalyst surface.
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(3) represents the minimum potential
energy trough where the molecules are absorbed onto the catalyst surface.
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(2) ==> (3-5) is represented by
Ea2 the
formation of the products form the intermediate absorbed states of the
molecules.
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Vanadium(V) oxide, V2O5,
is used as a heterogeneous catalyst in the 'Contact Process'
for the production of sulfur trioxide in the manufacture
of sulfuric acid.
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The catalysing of the
conversion of sulfur dioxide into sulfur trioxide is explained via change in oxidation state
changes i.e. some classic transition metal chemistry.
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(i)
SO2(g)
+ 1/2O2(g) ===> SO3(g)
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The mechanism,
somewhat simplified, goes
via the catalytic cycle ...
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(ii) SO2 + V2O5 ==> SO3
+ V2O4, then (iii)
V2O4
+ 1/2O2 ==> V2O5
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The vanadium
changes oxidation state from +5 to +4 and back to +5 in the catalytic cycle, a classic combination of two characteristics of
transition metals – variable oxidation state and catalytic
properties.
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The intermediate here is vanadium(IV) oxide V2O4.
-
This is a good example of
the chemical involvement of a catalyst, which temporarily
changes from V2O5 to V2O4
and then is converted back to V2O5
hence the catalyst is unchanged overall but continues its
function in the catalytic cycle.
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An iron/iron(III)
oxide mixture is used as
a heterogeneous catalyst in the Haber synthesis of ammonia from hydrogen and
nitrogen.
HOMOGENEOUS CATALYSIS:
-
All the catalyst and
reactants are in the same phase (usually a solution), and so the
catalysed reaction can happen throughout the bulk of the reaction
medium.
-
Most homogeneously catalysed
reactions involve several reaction steps to complete the process
from reactants to products.
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Even those described here are
simplifications, but it is the concepts involved that are important,
and can usually be illustrated with a few simplified equations to
illustrate homogeneous catalysis.
-
Below I've described examples of
homogeneous catalysis to reduce the activation energy by providing
an alternative pathway.
-
The catalysis
is usually due to temporary changes in oxidation state of a
transition metal ion and results in a 'catalytic
cycle'. In other words, the homogeneous catalysed reactions occur via some
intermediate species.
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e.g. (i) Either iron(II) Fe2+ ions or iron(III)
Fe3+ ions catalyse the oxidation of iodide ions by
peroxodisulfate
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uncatalysed (Ea3
in diagram above) the overall
reaction is:
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(i)
S2O82–
(aq) + 2I– (aq) ==> 2SO42–
(aq) + I2 (aq)
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However, this
'direct' uncatalysed reaction involves the collision of two
repelling negative ions and so has a high activation energy.
Activation energies arise from outer electron shell
repulsions and bond energies.
-
BUT, the
collision of an Fe3+ ion and an I– ion
involves positive–negative attraction which helps overcome the
repulsion component activation energy due to two negative
ions colliding.
-
-
so initially
for catalysed (Ea1 in diagram above)
-
followed by (Ea2
in diagram above)
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The iron(III) ion
is regenerated in the cycle whether you start with Fe2+
or Fe3+ , showing the iron ions act in a
genuine catalytic way and the iron ions are not consumed
overall in the process.
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If
you added up the two equations of the cycle you get the
equation of overall reaction change.
-
All the reactants,
products and catalytic ions are all in the same phase i.e.
aqueous solution and the activation energy is considerable
reduced by enabling a different pathway for the reaction.
-
This
is an excellent example of why transition element
compounds can act as catalysts in specific redox
reactions i.e. they can exist in, and interchange
between, two (or more) oxidation states that facilitate
the overall reaction.
-
Note 3: Eø arguments can be used to check on the
feasibility of the reaction or mechanism steps.
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(ii)
The
autocatalysis by Mn2+ ions when the oxidising
agent potassium manganate(VII), KMnO4, is used to
titrate the ethanedioate ion, C2O42–,
(from acid/salt, old name 'oxalic/oxalate').
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2MnO4–(aq)
+ 16H+(aq) + 5C2O42–(aq)
==> 2Mn2+(aq) + 8H2O(l) +
10CO2(g)
-
Initially,
the reactant collisions are between two anions which will
have a high activation energy, hence the slow start to the
reaction.
-
In the titration you see it gradually gets
faster and faster because there is catalytic cycle
involving the hexa–aqua Mn(II) ion and ethanedioate
complexes of Mn(II) and Mn(III).
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[MnII(H2O)6]2+(aq) ==> [MnII(C2O4)3]4–(aq) ==>
[MnIII(C2O4)3]3–(aq) ==> [Mn(H2O)6]2+(aq) + CO2(aq/g)
-
Note
that the catalytic cycle involves changes in ligand and
oxidation state in the manganese metal ions and two
intermediate complexes.
-
Without the formation of
the intermediate ethanedioate complex ion of manganese(II)
and manganese(III) the manganate(VII) ion cannot readily
oxidise the ethanedioate ion.
-
-
In the diagram above, the
trough between the peaks for Ea1 and Ea2
represents the formation of the intermediate complex ions of
manganese(II) and manganese(III), which considerably reduce
the activation energy Ea3.
-
(iii) Cobalt(II)
ions catalyse the oxidation of the 2,3–dihydroxybutandioate
ion (acid/salt, old name 'tartaric/tartrate') to water, methanoate
ion and carbon dioxide with hydrogen peroxide solution.
-
The likely scheme of events is outlined below,
showing the ease of conversion of cobalt between a +2 oxidation
state complex and a +3 oxidation state complex.
-
The overall equation for this
oxidation of 2,3-dihydroxybutanoate acid is:
-
2,3-dihydroxybutanoate
ion + hydrogen peroxide ===> carbon
dioxide + methanoate ion + water
-
-OOC-CH(OH)-CH(OH)-COO-
+ 3H2O2 ===> 2CO2
+ 2HCOO- + 4H2O
-
The rest of the
equations are NOT
meant to be balanced.
-
Starting
with the
pink hexa–aqa Co2+ ion, which is a Co(II)
complex
-
[Co(H2O)6]2+(aq)
==> [Co(OOCCH(OH)CH(OH)COO)3]4–(aq)
-
the
pink Co(II) complex changes ligand
from water to the organic acid, but no change in oxidation
state or co–ordination number, and I
don't know its colour?, but it perhaps it doesn't
exist long enough to be seen?
-
[Co(OOCCH(OH)CH(OH)COO)3]4–(aq) == via
H2O2 ==>
[Co(OOCCH(OH)CH(OH)COO)3]3–(aq)
-
the
Co(II)–acid complex is oxidised by the
hydrogen peroxide to a Co(III) –acid complex
which is green,
-
and this green complex is
the intermediate in the reaction profile diagram below.
-
[Co(OOCCH(OH)CH(OH)COO)3]3–(aq) ==> [Co(H2O)6]2+(aq),H2O(l),HCOO–(aq),CO2 (aq/g)
-
the
green Co(III) complex then breaks down to
give the products,
-
and
you see the bubbles of carbon dioxide and the 'return'
of the
pink hexa–aqa Co2+ complex ion.
-
In
the above sequence, the change in ligand affects the
relative stability of the oxidation states. The CoII–acid
complex is stable as regards 'breakdown', but is readily
oxidised to the CoIII–acid complex, which is
NOT stable to breakdown.
-
Again, the middle trough represents the formation of the
intermediate complexes which overall reduces the activation energy of
the uncatalysed reaction.
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(iv) -
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