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Doc Brown's
Advanced A Level Chemistry Advanced A Level Chemistry - Kinetics-Rates
revision notes Part 6
5.2 The theory of catalytic
mechanisms is discussed using heterogeneous examples and homogeneous examples
5.2.
Catalytic mechanisms
5.2a
Introduction
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REMINDER OF WHAT A CATALYST IS AND HOW
DO CATALYSTS WORK?
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ARE THERE DIFFERENT KINDS OF CATALYSTS?
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A catalyst is a
substance that alters the rate of chemical reaction without itself
being permanently chemically changed. Never
state things like "it doesn't react, just speeds it up".
It must take part in the reaction and it must change chemically,
albeit on a temporary basis. This may involve chemical intermediates
or even superficially on the surface of a solid catalyst.
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A catalyst lowers the activation
energy (Ea) by providing a different 'pathway'
or mechanism that makes the bond breaking processes (or other
electronic changes in the reactants) occur more readily.
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Catalysis
also involves the formation of intermediates, not just a matter of
an 'activated complex' or 'transition state'.
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e.g. for a
transition metal the reactant molecules may be adsorbed and their
bonds weakened, or, for a transition metal compound, it may involve a change in ligand or oxidation state or other
bonding re–arrangement, but will return to is original state often via a
2–3 stage 'catalytic cycle'.
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In organic chemistry, e.g. in acid
catalysis, the reactant molecule may be protonated to form some 'active'
protonated intermediate e.g. a carbocation in the case of adding water
to alkenes to form alcohols.
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By providing an alternative reaction pathway
of lower activation energy Ea,
compared to the uncatalysed reaction, e.g. see the diagram above
for a simple exothermic reaction, although more realistically, it is
usually a more complex cycle profile of at least two stages (see the 2nd
reaction profile further down).
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The argument based on Maxwell
- Boltzmann distribution described in detail in section 5.1
shows how reducing the activation energy considerably increases the
proportion of particles with sufficient kinetic energy to overcome the
'hump' of the activation energy in a given time.
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We are talking about the increased
frequency of a fruitful collision - more chance of a successful
collision in a given time (despite the fact that most collisions do
NOT lead to product formation - the particles just bounce of each other!
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A catalyst can be
changed physically e.g. the granules can end up more powdery or the
surface become roughened. This may be due to a heat effect from
exothermic reactions or just side effect of regeneration in the
catalytic cycle.
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Some examples of comparing the activation
energies of uncatalysed and catalysed reactions.
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There are two primary modes of
catalytic action – heterogeneous and homogeneous,
and both are described with examples below
5.2b
Heterogeneous catalysts and theory
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HETEROGENEOUS
CATALYST THEORY:
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A HETEROGENEOUS CATALYST IS IN A DIFFERENT PHASE
(often solid state') THAN THE REACTANTS (often gaseous or liquid/solution)
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The catalyst
and reactants are in different phases (usually solid catalyst
and liquid/gaseous reactant)
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The reaction
occurs on the catalyst surface which may be the transition metal
or one of its compounds, examples quoted above. The reactants
must be adsorbed onto the catalyst surface at the 'active sites'.
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This can be
physical adsorbed or 'weakly' chemically bonded 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 within the reactant molecules and so allows a greater
chance of 'fruitful' collision.
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The diagram
above illustrates a typical heterogeneous catalysis e.g.
hydrogenation of alkenes with hydrogen and a nickel catalyst.
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The strength
of adsorption is crucial to having a 'fruitful' catalyst
surface,
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The
bonding to the catalyst surface (chemisorption/adsorption)
must be strong enough to weaken reactant molecule bonds but
weak enough to allow new bonds to form and the products to
'escape' from the catalyst surface (desorption).
Typical examples to illustrate this idea ....
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If the catalyst–reactant
bonding is too
strong, most reactant/product molecules will be too strongly 'chemisorped'
inhibiting reaction progress e.g. tungsten (W),
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If the catalyst–reactant
bonding is too
weak, many reactants are not chemisorped strongly enough to
allow the initial bond breaking processes to happen e.g. gold
(Au) and silver (Ag) tend to be more limited catalysts,
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just
right: copper (Cu), nickel (Ni), platinum (Pt), rhodium (Rh), palladium*
(Pd) catalyse many reactions
such as hydrogenation, redox reactions involving CO and NO
etc.
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*Palladium
can catalyse the spontaneous combustion/combination of
hydrogen and oxygen at room temperature!
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In
catalytic converters the very expensive metals Pt, Rh and
Pd are used.
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Cu and Ni are cheaper alternatives but they
are more vulnerable to catalytic poisoning of the active sites
by traces of sulphur dioxide in the exhaust gases.
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Unfortunately it would
appear that the cheaper alternative metals chemisorb the 'catalytic
poisons' more strongly the more expensive Pt and Rh.
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Once
poisoned, the catalyst in a converter cannot be
regenerated and its a new costly converter!
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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.
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Catalyst
poisoning should be avoided if at all possible and in industry
catalysts have to be replaced or 're–furbished' via suitable
physical/chemical treatment. The poisoning inhibiting effect is caused by impurity molecules being strongly
chemisorbed (chemically bonded) to the most
active sites*
of the catalyst surface. It
considerably reduces the efficiency of the catalyst, especially as
the most effective catalyst sites bind impurities the strongest,
competing with the reactant molecules e.g.
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Sulphur
poisons the iron catalyst in the Haber Process for making
ammonia (probably chemisorbed to form iron sulphide).
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Lead
poisons the platinum–rhodium surface in car exhaust catalytic
converters, hence the need for 'un–leaded' fuel. Lead is
strongly absorbed preferentially onto active sites.
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Active sites*
Not all the surface of a catalyst is effective due to
minute imperfections in crystal structure at the atomic
level. The real crystal lattices are far from the
geometrical perfection we present them as in our diagrams
for teaching purposes. For catalytic surfaces, this
appears to be desirable up to a point! It has been shown
that an irregular catalytic surface is more effective than
a 'perfect' plane one. These irregularities in the first
few layers of atoms can be an atomic holes (vacancies),
steps and terraces, and planes of atoms not quite in line
with other layers of atoms (dislocations). One theory of
'active sites' suggests the substrate reactant molecules are
more strongly bound, aiding bond scission, if surrounded
by more of the catalyst atoms in a 'hole' or on a
'step/terrace', but the same effect applies to catalyst
poisons too!
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In heterogeneous
catalysis the catalyst
and reactants are in different phases (usually solid catalyst and
liquid/gaseous reactant) e.g.
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Case study
5.2b.1: The
black insoluble powder, manganese(IV) oxide, MnO2(s),
catalyses the decomposition of hydrogen peroxide solution into
water and oxygen.
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Case study
5.2b.2: Iron,
Fe(s), catalyses the combination of nitrogen and
hydrogen gases in the important industrial Haber synthesis of
ammonia, important in the manufacture of nitric acid*
and artificial fertiliser salts.
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N2(g)
+ 3H2(g) ==> 2NH3(g)
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*
In the 1st stage of making nitric acid the ammonia is oxidised to nitrogen(II)
oxide by mixing it with oxygen over a hot platinum
catalyst (another heterogeneous catalysis) which is then
reacted with oxygen and water to form nitric acid solution.
See
GCSE ammonia page
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Case study
5.2b.3: Platinum/rhodium/palladium
metals, Pt(s)/Rh(s)/Pd(s),on a ceramic
support, catalyse the following reactions in car exhausts
inside the catalytic converter.
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2NO(g)
+ 2CO(g) ==> N2(g) + 2CO2(g)
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The
NO and CO are adsorbed onto the catalyst surface,
bonds broken and reformed prior to the products
nitrogen and carbon dioxide leaving the catalyst
surface in a similar way to the hydrogenation
illustrated above. Remember that the bonding to the
catalyst surface (chemisorption/adsorption)
must be strong enough to weaken reactant bonds but
weak enough to allow the products to 'escape' (desorption)
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The
CO is from the inefficient combustion of the
hydrocarbon fuel,
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and
the nitrogen(II) oxide is 'naturally' formed at high
temperature in the engine (as it is in lightning
strikes!).
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These
transition metal catalysts can also oxidise unburned
hydrocarbons from inefficient combustion.
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Note:
Pt, Rh and Pd are very expensive metals and copper and
nickel are cheaper alternatives but they are vulnerable to
catalytic poisoning by traces of sulphur dioxide in the
exhaust gases. Once poisoned, the catalyst in a converter
cannot be regenerated, so, its a new costly converter!
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Case study
5.2b.4: Nickel, Ni(s), catalyses the addition of hydrogen
to an alkene double bond, e.g. in the hydrogenation of
unsaturated vegetable oils to make more saturated margarine
with a slightly higher softening point making it more
spreadable.
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Case study
5.2b.5: Solid heterogeneous catalysts are really important in the
petrochemical
industry e.g.
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Isomerisation:
These reactions convert linear alkane vapours into
branched alkanes of the same carbon number over a
platinum–aluminium oxide (Pt/Al2O3)
catalyst at 150oC. Branched alkanes have a
higher octane rating than linear alkanes, so a better
petrol fuel components.
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Reforming:
Converting straight chain alkane vapour into cyclic
alkanes and aromatic hydrocarbons can be achieved by using a
Pt/Al2O3
catalyst at 500oC. Aromatic hydrocarbons are
important chemical feedstock to make many useful aromatic
compounds.
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Cracking:
Catalytic cracking of vapourised hydrocarbons at e.g. 500oC
using zeolites to make more lower alkanes for
petrol or alkenes. Alkenes are important intermediates
in making many useful compounds from anti–freeze to
plastics.
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e.g.
higher alkanes ==> lower alkanes and
alkenes
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Note:
Zeolites tend to become 'poisoned' with carbon–soot
deposits in the high temperature cracking reactions
and this blocks the adsorption of the hydrocarbons.
However, in this case, the catalyst can be regenerated
in a separate container through which very hot air is
passed to burn off the carbon–soot deposits.
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Case study
5.2b.6:
Vanadium(V) oxide, V2O5,
is used as a catalyst in the 'Contact Process'
in the production of sulphur
trioxide for the manufacture
of sulphuric acid.
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The catalysing of the
conversion of sulphur dioxide into sulphur trioxide is explained via change in oxidation state
changes.
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2SO2(g)
+ O2(g) ==> 2SO3(g)
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The mechanism,
somewhat simplified, it goes
via the catalytic cycle ...
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(i) SO2 + V2O5 ==> SO3
+ V2O4, then (ii) V2O4
+ 1/2O2 ==> V2O5
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The vanadium
changes oxidation state from +5 to +4 and back to +5 in the cycle.
5.2c.
Homogeneous
catalysis and theory
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A HOMOGENEOUS CATALYST IS IN THE
SAME PHASE
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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.
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Catalysis
can be due to temporary changes in
the oxidation state and ligand(s) of a
transition metal ion and results in a 'catalytic cycle'.
In other words, the reaction occurs via some intermediate species
and the original catalyst is reformed e.g.
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Case study
5.2c.1:
Oxidation of iodide ions by
peroxodisulfate described in section 7.2
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Case study
5.2c.2: The
autocatalysis by Mn2+ ions when the oxidising agent
potassium manganate(VII), KMnO4, is used to titrate
the ethanedioate ion, C2O42–,
or the acid H2C2O4 /
(COOH)2,
(old names 'oxalate/oxalic acid').
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The titration,
particularly at the start, is too slow, so the initial
mixture is heated to 60oC before starting the
KMnO4 addition. One reason why the initial
reaction is slow, is because the two ions involved (C2O42–
and MnO4–) are both negative, so on collision
the repulsion force factor is high, so the activation energy is high
giving low kinetic activity.
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However, as the titration
proceeds, the decolourisation of the added manganate(VII) speeds up
as the conical flask is swirled. The reason for this, is that one
of the reaction products, the Mn2+ ion, can itself act as
a catalyst for this redox reaction, hence the 'auto–catalysis'
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The overall
balanced redox reaction is ... based on ... 2 x Mn(+7) ==> 2 x Mn(+2)
and 10 x C(+3)
==>10 x C(+4)
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2MnO4–(aq)
+ 16H+(aq) + 5C2O42–(aq)
==> 2Mn2+(aq) + 8H2O(l) +
10CO2(g)
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or
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2MnO4–(aq)
+ 6H+(aq) + 5H2C2O4(aq)
==> 2Mn2+(aq) + 8H2O(l) +
10CO2(g)
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The
intensely purple coloured manganate(VII) ion is
reduced by the ethanedioate ion to the very pale pink
(virtually colourless) manganese(II) ion in the presence
of dilute sulphuric acid.
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The
intense colour of the manganate(VII) ion is not due to the
splitting of the 3d electron levels by the 'field splitting'
ligands, and the subsequent absorption of visible light
photons, causing electronic promotion from the lower to
higher upper 3d sub–level, which is the usual case for
transition metal complexes. The MnO4–
ion is known as a charge transfer complex ion,
because the colour is due to the electronic transitions of
oxygen's 2p electrons. These are temporarily raised from a
full oxygen 2p orbital to the vacant 3d or 4s levels of the
Mn(VII) ion.
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Initially,
on face value, the reactant collision is between two anions
(–) which will
have a high activation energy due to extra
repulsion as well as the outer electron–outer electron repulsion
that exists between any two atoms/molecules/ions, hence the slow
start to the reaction. However, in the acid medium, the organic
species is likely to be mainly in the form of the unionised
acid H2C2O4 but
the activation energy is obviously high until the first traces
of the catalyst Mn2+ are formed, when the lower
energy pathway kicks in!
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In the titration,
the first few drops of manganate(VII) seem to take a
relatively long time to 'decolourise' (from purple to
colourless), but as the titration proceeds, the
decolourisation becomes faster because the Mn2+(aq)
complex
ions act as a catalyst in the oxidation of the
ethanedioate ion and recent research has shown that
colloidal manganese(IV) oxide, MnO2,is
involved too (MnO2 in bulk is insoluble). If you
carry out the reaction without the addition of acid you
actually see brown–black MnO2 formed as a
precipitate, which is a neutral or alkaline solution
reduction product of MnO4–.
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At GCE
AS–A2 level this is often quoted as an example of
auto–catalysis BUT it is an
extremely complicated reaction. The reaction
is still a subject of research and the latest theories
involve at least nine mechanism steps and the
formation of colloidal manganese(IV) oxide, MnO2,
and complexes of MnO2 combined with
ethanedioic acid. This reaction
is seriously complicated!
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Incidentally,
if you mix potassium manganate(VII), acidified
ethanedioic acid and manganese(II) sulphate to
deliberately speed up the reaction apparently you see a red MnIII
complex formed, maybe [Mn(C2O4)3]3–
which will be a part of one pathway in the multi–catalytic cycle.
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Well,
make of it what you can! but to sum up for GCE AS–A2
level, and not undergraduate level purposes the sequence in
terms of the initial/intermediate reactants and products is
basically ... (the equations are
NOT meant to be balanced or
the full complex structures always shown, oxidation states are also
shown in the compounds/complexes as Roman numeral superscripts)
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H2C2O4(aq) + MnO4–(aq)
==> Mn2+(aq) + CO2(g) + H2O(l)
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Initially
there is a
very slow uncatalysed reaction to give the initial trace
of the Mn2+ based catalyst, presumably by a different,
but equally complex sequence, such as that described for the auto–catalytic cycle
below, involving, amongst others, Mn(III) complexes.
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C2O42–(aq)
+ Mn2+(aq) ==> [MnIIC2O4](aq)
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[MnIIC2O4](aq)
+ MnVIIO4–(aq) ==> 2MnIVO2(colloidal) + CO2(g)
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==> MnIII/IV complexes
like [MnIII(C2O4)3]3–/[MnIVO2.xH2O]
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2. The
formation of colloidal manganese(IV) oxide (as a
hydrated complex) as well as 'soluble' Mn(III) complexes
is also involved in producing the final reaction
products. However, the bulk formation of MnO2
is suppressed by acidification of the solution, if you
don't swirl fast enough in the titration, you get a
black–brown stain of MnO2 (colloidal or
precipitate).
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All these reactions involve changes in ligand and
oxidation state of Mn so characteristic of transition metal
chemistry. A change in ligand, or other chemical
state, can change the relative half–cell potential (or
stability) for the interchange of two oxidation states
of the same metal ion, in this case involving MnII/III/IV/VII
species.
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[MnIII(C2O4)2]–
==> [MnIIC2O4]
+ 2CO2 + e–
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3.
Shows, in principle, one possible reduction of Mn(III) to Mn(II) in
which, by electron transfer within the complex, an
ethanedioate ion is oxidised to carbon dioxide (e–
loss) and the MnIII–ethanedioate
complex is reduced back to the MnII complex (e–
gain) first
formed in step (1), so completing the autocatalytic
cycle.
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The concentration of Mn2+ or other MnIII
or MnIV species
increases as the reaction proceeds, so this increase in
'catalytic species' accounts for the
observed increase in rate of the manganate(VII) 'decolourisation' as the
titration proceeds, at least until near the end, when the
reactant/catalytic intermediate concentrations are then becoming quite low.
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Case study
5.2c.3: 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, the
equations are NOT
meant to be balanced.
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Starting
with the
pink hexa–aqua Co2+ ion, which is a Co(II)
complex
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[Co(H2O)6]2+(aq) ==> [Co(OOCCH(OH)CH(OH)COO)3]4–(aq)
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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?
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[Co(OOCCH(OH)CH(OH)COO)3]4–(aq) == via
H2O2 ==> [Co(OOCCH(OH)CH(OH)COO)3]3–(aq)
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[Co(OOCCH(OH)CH(OH)COO)3]3–(aq) ==> [Co(H2O)6]2+(aq),H2O(l),HCOO–(aq),CO2 (aq/g)
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the
green Co(III) complex then breaks down to
give the products,
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and
you see the bubbles of carbon dioxide and the 'return'
of the
pink hexa–aqua Co2+ complex ion.
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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.
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Transition metal
ions are often at the 'heart' of many biological catalysts.
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Case study
5.2c.4: Protein built
enzymes form complexes with the metal ions, the
protein acting as a
multi–dentate ligand,.
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e.g.
transition metal ions like Fe2+/Fe3+, Cu2+,
Co2+,
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and
non–transition metal ions like Zn2+, Mg2+,
Ca2+,
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e.g. the enzymes
peroxidase/catalase rapidly decompose hydrogen peroxide.
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The reaction
involves a protein–Fen+ complex undergoing ligand and oxidation
state changes to break the hydrogen peroxide down into water and
oxygen,
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H2O2(aq)
==> 2H2O(l) + O2(g)
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and the
enzymes are at least a thousand times more effective than, the
'magic' black powder catalyst, manganese(IV) oxide, MnO2(s),
used in the laboratory preparation of oxygen. We have a long
way to go to beat enzymes using completely 'synthetic' catalysts.
Biotechnology
is developing lots of organic synthesis reactions which all go via
enzyme catalysts.
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Other examples
of homogeneous catalysis NOT involving transition metal ions.
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Case study
5.2c.4: Esterification:
Acids, proton donors (H+), catalyse the
formation of an ester from a carboxylic acid and alcohol, they
also catalyse the reverse reaction of hydrolysis.
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Case study
5.2c.5: Friedel–Craft
reactions: The alkylation and acylation of aromatic
hydrocarbons is catalysed by aluminium chloride.
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Case study5.2c.6: The
'catalytic
cycle' of ozone destruction by chlorine atoms from CFC's
is described on another page.
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