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Advanced inorganic chemistry: Transition metals - homogeneous and heterogeneous catalysts

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Periodic Table - Transition Metal Chemistry - Doc Brown's Chemistry  Revising Advanced Level Inorganic Chemistry Periodic Table Revision Notes

 Appendix 6 Catalysis theory and practice – homogenous & heterogeneous

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The general theory of catalysed reactions is introduced followed by the theory of heterogeneous catalysis by transition metals or solid transition metal compounds is described and explained with suitable examples. Further examples are then given involving transition metal ions involved in examples of homogeneous catalysis. The ability of transition metals to physically and chemically adsorb gases is emphasised in heterogeneous catalysis AND the ease in change of oxidation state (oxidation number) where transition metal ions act as homogeneous catalysts.

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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:

    • (1) ==> (2) is represented by Ea1, 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 Ea2 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' for the production of sulfur trioxide in 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.

    • (i) SO2(g) + 1/2O2(g) ===> SO3(g) 

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

      • (ii) SO2 + V2O5 ==> SO3 + V2O4, then (iii) 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.

        • If you add equations (ii) + (iii) in an algebraic' manner, you get equation (i).

      • 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.

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

    • N2(g) + 3H2(g) === Fe/Fe2O3 catalyst ===> 2NH3(g)


  • 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 (Ea3 in diagram above) the overall reaction is:

      • (i) S2O82– (aq) + 2I(aq) ==> 2SO42– (aq) + I2 (aq)

        • [Eø = ?]

      • 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)

        • (ii) 2Fe3+(aq) + 2I(aq) ==> 2Fe2+(aq) + I2(aq) 

          • Eøreaction = 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 (Ea2 in diagram above)

        • (iii) 2Fe2+(aq) + S2O82–(aq) ==> 2SO42–(aq) + 2Fe3+(aq)

          • Eøreaction = 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: Eø 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, C2O42–, (from acid/salt, old name 'oxalic/oxalate').

      • 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).

      • [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 

        • and the carboxylate ion, OOCCH(OH)CH(OH)COO (a 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,

        • 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.

    • (iv) -

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

    • Examples to add

INORGANIC Part 10 3d block TRANSITION METALS sub–index:

10.1–10.2 Introduction to 3d–block Transition Metal chemistry

10.3 Chemistry of Scandium  *  10.4 Chemistry of Titanium

10.5 Chemistry of Vanadium  *  10.6 Chemistry of Chromium

10.7 Chemistry of Manganese  *  10.8 Chemistry of Iron

10.9 Chemistry of  Cobalt  *  10.10 Chemistry of Nickel

10.11 Chemistry of Copper  *  10.12 Chemistry of Zinc

10.13 Selected chemistry of other Transition Metals e.g. Ag and Pt

Appendix 1. Hydrated salts, acidity of hexa–aqua ions

Appendix 2. Complexes and ligands

Appendix 3. Complexes and isomerism

Appendix 4. Electron configuration and colour theory

Appendix 5. Redox equations, feasibility of reaction, Eø calculations

Appendix 6. Catalysis - types and effectiveness

Appendix 7. Redox equations - construction and balancing

Appendix 8. Stability constants of complexes and entropy changes

Appendix 9. Colorimetric analysis and determining a complex ion formula

Appendix 10 3d block – extended data table

Appendix 11 3d–block transition metal complexes, oxidation states & electrode potentials

Appendix 12 Hydroxide complex precipitate 'pictures', formulae and equations

Advanced Level Inorganic Chemistry Periodic Table Index: Part 1 Periodic Table history * Part 2 Electron configurations, spectroscopy, hydrogen spectrum, ionisation energies * Part 3 Period 1 survey H to He * Part 4 Period 2 survey Li to Ne * Part 5 Period 3 survey Na to Ar * Part 6 Period 4 survey K to Kr and important trends down a group * Part 7 s–block Groups 1/2 Alkali Metals/Alkaline Earth Metals * Part 8  p–block Groups 3/13 to 0/18 * Part 9 Group 7/17 The Halogens * Part 10 3d block elements & Transition Metal Series * Part 11 Group & Series data & periodicity plots * All 11 Parts have their own sub–indexes near the top of the pages

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