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Advanced level chemistry kinetics notes: Homogeneous-heterolytic catalyst mechanisms

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



  • 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.

  • 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'.

  • 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'.

  • 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.

  • 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.

    • 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!

  • 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.

    • e.g. in the laboratory preparation of oxygen from the MnO2(s) catalysed decomposition of hydrogen peroxide solution, the residual water seems stained brown due to very fine particles of MnO2(s).

  • Some examples of comparing the activation energies of uncatalysed and catalysed reactions.

    • The decomposition of hydrogen peroxide

      • 2H2O2(aq)  ====> 2H2O(l)  +  O2(g)

      • Three activation energies (Ea) are quoted

      • (i) uncatalysed: 75 kJmol-1

      • (ii) catalysed with colloidal platinum: 49 kJmol-1

      • (iii) catalysed by the enzyme catalase in you body: 23 kJmol-1

      • Both catalysts are effective in considerably reducing the activation energy and hence greatly increasing the speed of the decomposition reaction.

    • The decomposition of hydrogen iodide.

      • 2HI(g)  ====>  H2(g)  +  I2(g)

      • Activation energies: (i) uncatalysed 183 kJmol-1, (ii) Au catalysed 105 kJmol-1, (iii) Pt catalysed 58 kJmol-1

      • The surface of both transition metals act as an efficient catalyst.

    • The synthesis of ammonia

      • N2(g)  +  3H2(g)  ====>  2NH3(g)

      • The uncatalysed reaction activation energy is very high at 350 kJmol-1.

      • This value reflects the high bond enthalpies of the species involved.

        • Bond enthalpies: NN 944 kJmol-1, H-H 436 kJmol-1.

      • A transition metal tungsten (W) catalyst reduces this to 162 kJmol-1.

      • I couldn't find a value for the actually used iron based catalyst?

      • Anybody know the value? I get the impression from diagrams on the internet that the best catalysts have reduced it to ~103 kJmol-1?

  • 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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    • A HETEROGENEOUS CATALYST IS IN A DIFFERENT PHASE (often solid state') THAN THE REACTANTS (often gaseous or liquid/solution)

    • The catalyst and reactants are in different phases (usually solid catalyst and liquid/gaseous reactant)

    • 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'.

    • 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.

    • The diagram above illustrates a typical heterogeneous catalysis e.g. hydrogenation of alkenes with hydrogen and a nickel catalyst.

  • The strength of adsorption is crucial to having a 'fruitful' catalyst surface, 

    • 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 ....

    • If the catalyst–reactant bonding is too strong, most reactant/product molecules will be too strongly 'chemisorped' inhibiting reaction progress e.g. tungsten (W),

    • 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,

      • but even silver, can act as a catalyst for some reactions.

    • 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.

      • *Palladium can catalyse the spontaneous combustion/combination of hydrogen and oxygen at room temperature!

      • In catalytic converters the very expensive metals Pt, Rh and Pd are used.

        • 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.

        • Unfortunately it would appear that the cheaper alternative metals chemisorb the 'catalytic poisons' more strongly the more expensive Pt and Rh.

        • Once poisoned, the catalyst in a converter cannot be regenerated and its a new costly converter!

  • 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. e.g. the platinum–rhodium metal is distributed on a temperature resistant ceramic support in catalytic converters of motor vehicle exhausts.

  • 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.

    • Sulphur poisons the iron catalyst in the Haber Process for making ammonia (probably chemisorbed to form iron sulphide).

    • 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.

      • 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!

  • topBelow is a diagram of a two stage reaction profile for a catalytic cycle (Ea = activation energy)

    • This can apply to heterogeneous catalysis or homogeneous catalysis.

    • 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.

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  • In heterogeneous catalysis the catalyst and reactants are in different phases (usually solid catalyst and liquid/gaseous reactant) e.g.

  • 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.

    • 2H2O2(aq)  ==> 2H2O(l) + O2(g) 

  • 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.

    • N2(g) + 3H2(g) ==> 2NH3(g) 

    • * 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

  • 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.

    • 2NO(g) + 2CO(g) ==> N2(g) + 2CO2(g) 

      • 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)

      • The CO is from the inefficient combustion of the hydrocarbon fuel,

        • CxHy + (x/2 + y/4)O2 ==> xCO + y/2H2

      • and the nitrogen(II) oxide is 'naturally' formed at high temperature in the engine (as it is in lightning strikes!).

        • N2(g) + O2(g) ==> 2NO(g)

    • These transition metal catalysts can also oxidise unburned hydrocarbons from inefficient combustion.

      • CxHy + (x + y/4)O2 ==> xCO2 + y/2H2

    • 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!

  • 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.

    • –CH=CH– + H2 == Ni catalyst ==> –CH2–CH2

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    • The diagram above illustrates this typical heterogeneous catalysis.

  • Case study 5.2b.5: Solid heterogeneous catalysts are really important in the petrochemical industry e.g.

    • 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.

      • e.g. hexane ==> methylpentanes or dimethylbutanes

    • 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.

      • e.g. hexane ==> cyclohexane (+H2) ==> benzene (+3H2)

    • 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.

      • e.g. higher alkanes ==> lower alkanes and alkenes

      • 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.

  • 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.

    • The catalysing of the conversion of sulphur dioxide into sulphur trioxide is explained via change in oxidation state changes.

    • 2SO2(g) + O2(g) ==> 2SO3(g) 

    • The mechanism, somewhat simplified, it goes via the catalytic cycle ...

      • (i) SO2 + V2O5 ==> SO3 + V2O4, then (ii) V2O4 + 1/2O2 ==> V2O5 

      • The vanadium changes oxidation state from +5 to +4 and back to +5 in the cycle.

5.2c. Homogeneous catalysis and theory


    • (usually a solution of a catalyst and a solution of the reactants – contrast with a heterogeneous catalyst)

  • 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.

  • 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.

  • Case study 5.2c.1: Oxidation of iodide ions by peroxodisulfate described in section 7.2

  • 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').

  • 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.

  • 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'

    • 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)

    • 2MnO4(aq) + 16H+(aq) + 5C2O42–(aq) ==> 2Mn2+(aq) + 8H2O(l) + 10CO2(g) 

    • or

    • 2MnO4(aq) + 6H+(aq) + 5H2C2O4(aq) ==> 2Mn2+(aq) + 8H2O(l) + 10CO2(g) 

    • 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.

      • 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.

    • 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!

    • 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.

    • 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!

    • 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.

    • 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)

    • H2C2O4(aq) + MnO4(aq) ==> Mn2+(aq) + CO2(g) + H2O(l)

      • 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.

      1. C2O42–(aq) + Mn2+(aq) ==> [MnIIC2O4](aq)

        • 1. The formation an Mn(II)–complex with ethanedioate ion ligand which is the basis of the faster catalysed route of lower activation energy. This reacts with the manganate(VII) ion, followed by the reduction of Mn(III) and Mn(IV) species by the ethanedioate ion or a complex of it.

      2. [MnIIC2O4](aq) + MnVIIO4(aq) ==> 2MnIVO2(colloidal) + CO2(g)

        • ==> MnIII/IV complexes like [MnIII(C2O4)3]3–/[MnIVO2.xH2O]

          • ==> Mn2+(aq) + CO2(g) + H2O(l) + [MnIII(C2O4)2](aq)

        • 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).

        • 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.

      3. [MnIII(C2O4)2] ==> [MnIIC2O4] + 2CO2 + e

        • 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.

        • 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.

  • 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.

    • Starting with the pink hexa–aqua Co2+ ion, which is a Co(II) complex 

      • and the carboxylate ion, OOCCH(OH)CH(OH)COO (bidentate 2– anionic ligand)

    • [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,

    • [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–aqua 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.

  • Transition metal ions are often at the 'heart' of many biological catalysts.

  • Case study 5.2c.4: Protein built enzymes form complexes with the metal ions, the protein acting as a multi–dentate ligand,.

    • e.g. transition metal ions like Fe2+/Fe3+, Cu2+, Co2+

    • and non–transition metal ions like Zn2+, Mg2+, Ca2+

    • e.g. the enzymes peroxidase/catalase rapidly decompose hydrogen peroxide.

    • The reaction involves a protein–Fen+ complex undergoing ligand and oxidation state changes to break the hydrogen peroxide down into water and oxygen,

      • H2O2(aq) ==> 2H2O(l) + O2(g) 

      • 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.

      • Note: If the enzyme is immobilised on some support, and the solution of substrate molecules moves over the surface, it is strictly speaking a case of heterogeneous catalysis?

  • Other examples of homogeneous catalysis NOT involving transition metal ions.

  • 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.

    • RCOOH + R'OH doc b RCOOR' + H2O

  • Case study 5.2c.5: Friedel–Craft reactions: The alkylation and acylation of aromatic hydrocarbons is catalysed by aluminium chloride.

  • Case study5.2c.6: The 'catalytic cycle' of ozone destruction by chlorine atoms from CFC's is described on another page.

    • It is an example of a gaseous phase homogenous catalysis.

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