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 E_a, compared to the uncatalysed reaction (see 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

HETEROGENEOUS CATALYSIS: (e.g. diagram above for 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.

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• A two stage reaction profile for a catalytic cycle (E_a = 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.

  • E_a1 is the activation energy leading to the formation of an intermediate complex.

  • E_a2 is the activation energy for the change of the intermediate complex into products.

  • E_a3 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:

    • (1) ==> (2) is represented by E_a1, the absorption of the reactant molecules onto the catalyst surface.

    • (3) represents the minimum potential energy trough where the molecules are absorbed onto the catalyst surface.

    • (2) ==> (3-5) is represented by E_a2 the formation of the products form the intermediate absorbed states of the molecules.

• Vanadium(V) oxide, V₂O₅, is used as a heterogeneous catalyst in the 'Contact Process' in the production of sulfur trioxide for the manufacture of sulfuric acid.

  • 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.

  • 2SO_2(g) + O_2(g) ==> 2SO_3(g) 

  • The mechanism, somewhat simplified, 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 catalytic cycle, a classic combination of two characteristics of transition metals – variable oxidation state and catalytic properties.

• An iron/iron(III) oxide mixture is used as a heterogeneous catalyst in the Haber synthesis of ammonia from hydrogen and nitrogen.

  • N_2(g) + 3H_2(g) === Fe/Fe₂O_{3 catalyst} ===> 2NH_3(g)

• -

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.

  • 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.

  • e.g. (i) Either iron(II) Fe²⁺ ions or iron(III) Fe³⁺ ions catalyse the oxidation of iodide ions by peroxodisulfate

    • uncatalysed (E_a3 in diagram above) the overall reaction is:

    • (i) S₂O₈^2– _(aq) + 2I^– _(aq) ==> 2SO₄^2– _(aq) + I_{2 (aq)} 

      • [E^ø = ?]

    • 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 (E_a1 in diagram above)

      • (ii) 2Fe³⁺_(aq) + 2I^–_(aq) ==> 2Fe²⁺_(aq) + I_2(aq) 

        • E^ø_reaction = V, just a single step?

        • Note: If no iron(III) ions are present, but an iron(II) salt is added, iron(III) ions are generated via equation (iii) and so the catalytic cycle of reactions (ii) and (iii) can begin.

    • followed by (E_a2 in diagram above)

      • (iii) 2Fe²⁺_(aq) + S₂O₈^2–_(aq) ==> 2SO₄^2–_(aq) + 2Fe³⁺_(aq) 

        • E^ø_reaction = V, several steps I would think?

    • 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.

    • 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: E^ø arguments can be used to check on the feasibility of the reaction or mechanism steps.

  • (ii) The autocatalysis by Mn²⁺ ions when the oxidising agent potassium manganate(VII), KMnO₄, is used to titrate the ethanedioate ion, C₂O₄^2–, (from acid/salt, old name 'oxalic/oxalate').

    • 2MnO₄^–_(aq) + 16H⁺_(aq)  + 5C₂O₄^2–_(aq) ==> 2Mn²⁺_(aq) + 8H₂O_(l) + 10CO_2(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).

    • [Mn^II(H₂O)₆]²⁺_(aq) ==> [Mn^II(C₂O₄)₃]^4–_(aq)  ==>  [Mn^III(C₂O₄)₃]^3–_(aq) ==> [Mn(H₂O)₆]²⁺_(aq) + CO_2(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 E_a1 and E_a2 represents the formation of the intermediate complex ions of manganese(II) and manganese(III), which considerably reduce the activation energy E_a3.

  • (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, the equations are NOT meant to be balanced.

  • Starting with the pink hexa–aqa 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),CO_{2 (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 Co^II–acid complex is stable as regards 'breakdown', but is readily oxidised to the Co^III–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.

  • (iv) -

• Transition metal ions at the 'heart' of many enzymes – biological catalysts or 'molecule carriers'

  • Examples to add