Basics of catalysis

Catalysis is the foundation of the chemistry industry and is widely used in large scale synthesis of bulk chemicals and fine chemicals. It is the heart of the fertiliser, petroleum, polymer, inorganic and pharmaceutical industries amongst others and is of growing importance in environmental control including pollutant and waste mitigation, pollution ablation as well as in the generation of new alternative energy and fuel sources. A general review of catalysis is beyond the scope of this book but the reader is referred to a number of excellent books. Chorkendorff and Niemantsverdriet have recently reviewed the general area of catalysis96 and Ertl et al. have edited a comprehensive summary of the state-of-the art.97 Morris and others provide a preface to summarise recent work.98 The industrial perspective has been well reviewed by Rase99 whilst Cybulski and Moulijn have presented details of modern reactors and process design.100 The import subjects of catalyst preparation and synthesis, which are pivotal in determining cost-effectiveness of the process and the efficacy of the catalyst itself, have been thoroughly summarised.101 Finally, the theory of catalysis has been detailed in depth.102 Catalytic mechanisms are normally defined by the adsorption and specifically the chemisorption of molecules at the catalyst surface.103

Catalysis has been divided into two separate subject areas. The first of these is homogeneous catalysis where the reactants and the catalyst are in the same phases, e. g. ion catalysed reactions in solution. The second area is heterogeneous catalysis, principally used in the manufacture of very large quantities of chemicals, organic and inorganic materials.104 Here, the catalyst and reactants are in different phases most usually a solid catalyst and gas phase reactants. This is most useful for high throughput and is used for many of the very large scale processes carried out in industry including sulphuric acid, ammonia, methanol, nitric acid, ethylene oxide, cyanide synthesis as well as petroleum reforming and cracking, gasification, steam-reforming and water-gas shift processes.

As every high school student would know, a catalyst increases the rate of a chemical reaction without itself being consumed or altered in the reaction. This is a result of the catalyst interacting with the reactants to lower the activation energy barrier to the reaction. It does not alter the energetics (i. e. the free energy) of the reaction and so does not change the equilibrium between reactants and products directly. In the most obvious use of a catalyst, the rate of reaction is increased allowing increased rates of production at a particular temperature. Alternatively, the catalyst may be used by offsetting the increased rate of reactions that can be achieved (which can be orders of magnitude greater) against the process temperature used. This may be used to simply move a reaction into a feasible temperature range where reactor engineering becomes practical. It can also be used to reduce the energy input required by endothermic processes allowing more economic operation. A simple example would be methane combustion which can be catalysed by a number of precious metals as well as lanthanide oxide materials.22-24 The catalyst is used in gas turbines to lower combustion temperatures, thereby reducing the oxidation of nitrogen to nitrogen oxides. Similar reactions where that catalyst would be used to simply increase the rate of reaction in the formation of the thermodynamically stable product are the oxidation of CO and hydrocarbons to CO2 as shown below.105

Practically, catalysis is often used in more complex ways so as to produce significant amounts of a product that would not be obtainable in an un-catalysed process. An example is the partial oxidation of ethylene to ethylene oxide:106

H H yos

Y = c’ + 0.502 ^ H — C — C — H


This is the reaction which is mildly exothermic and the ethylene oxide formed is a partial oxidation product. The catalyst is not used to increase the rate of formation but rather to provide an operable process window at lower temperatures because at higher temperatures combustion of ethylene is highly favoured and the ethylene oxide would be a very short-lived and unrecoverable intermediate. Industrially, the reaction is catalysed by silver supported on alumina although various promoters for the epoxidation reaction as well as total oxidation inhibitors are used as part of the catalyst formulation to ensure high selectivity. The reaction operates at around 250°C, a low enough temperature for the desired product to be separated.

As mentioned earlier, the Haber-Bosch process has been used for the synthesis of ammonia from nitrogen and hydrogen for many years:11

N2 + 3H2 ^ 3NH3

This is an exothermic, thermodynamically favoured process but is kinetically limited. The reaction is catalysed by a potassium promoted iron catalyst. In simplistic terms, the catalyst is used to allow low temperature dissociation of nitrogen, the main activation energy barrier in the reaction. However, this is an equilibrium process and as the temperature is increased, the equilibrium favours the reactants (le Chateliers principle). In this way, although a catalyst does not alter the equilibrium in a chemical process, in this case the catalyst allows a low temperature (400°C) to be used and, thus, a higher equilibrium concentration of the product ammonia to be achieved. An un-catalysed mixture of hydrogen and nitrogen would reach the same equilibrium concentrations but would take considerably longer and not be consistent with the continuous process needed for large volume manufacture. The use of catalysts to effectively control equilibrium and rate is further illustrated by the water gas shift reaction:107

CO + H2O ^ CO2 + H2

This process is also mildly exothermic and equilibrium limited and like ammonia synthesis, the equilibrium favours reactants at higher temperatures but is kinetically limited requiring a catalyst to be used to achieve reasonable rates at temperatures below 1000°C. In order to achieve a high rate of reaction it is carried out in two different stages. The first is high temperature shift (HTS) using an alumina supported nickel catalyst at around 400°C which yields high rates of reaction. The second, low temperature shift (LTS) process uses an alumina supported copper based catalyst at around 200°C — the lower temperature allows almost complete recovery of the valuable hydrogen product at the equilibrium limit. The shift reaction(s) also demonstrates how important the synthesis of the catalyst is in these high volume processes since that catalyst must have sufficient activity and lifetime to be cost-effective. The subject of catalyst design and manufacture is briefly outlined in this review because in catalytic pyrolysis the catalysts used face very challenging environments.

In very general terms, the catalyst functions by reactant molecules interacting with the surface and becoming ‘activated’ in some way. Understanding how and where these reactions occur on the catalyst surface has been the subject of intense research and most researchers use an active site concept where specific arrangements of surface atoms or defects in the surface provide locations for adsorption/activating of reacting species.108 Very often, catalytic activity can be a unique property of a certain metal, oxide or combination. In order to maximise the number of collisions with the surface, it is usual to provide the active component with as high a surface area as is possible. This is normally achieved using an inorganic ‘ support’ over which the active component is ‘dispersed’.109 This allows very high surface areas of what are sometimes quite expensive materials (e. g. precious metals) to be achieved at relatively low amounts (below 5% by weight of the support) and at sustainable costs. The inorganic support is designed to be thermally robust and plays little part in the catalysis reactions although metal-support interactions between an active metal component and a support oxide have been well documented.110 In many cases the support is designed to be porous in order to provide as great a surface area as possible.111,112 Despite the fact that theoretically catalysts are unchanged by the reaction it promotes, their performance is often assessed by their lifetime in practical use. The lifetime is defined by the period in which its rate lies within an acceptable performance level (i. e. rate of production). Very often, the rate decreases with time and although this can be compensated by increasing reactor temperature, there is a point where practically an upper operating temperature is reached. A too higher temperature might be manifest as an unacceptable amount of side products — i. e. the reaction selectivity is compromised. Alternatively, the temperature may rise beyond process variables to protect plant and safety. This decrease in activity largely results from two different processes occurring during use. These are sintering and poisoning.113-115

Sintering is a complex process whereby loss of surface area is observed through thermal treatment which promotes mass transport and particle agglomeration or loss of pore structure. Pores play an important role in catalysis as they allow access to internal surfaces and can promote size controlled reactions where the molecules restrain the molecules that can enter the pore system. The process of sintering is thermodynamically favoured because it results in lower surface area and consequent decreases in the free energy of the system. Highly dispersed and supported active materials (as crystallites or particles) grow by a diffusion mechanism into larger particles. Other high temperature processes may also be responsible for sintering including solid-state reactions, new phase precipitation (e. g. in the thermal phase transformation of high surface area y-alumina to low surface area a-alumina) and the crystallisation of amorphous silica supports.

Poisoning generally describes the adsorption of strongly held species at the active catalyst sites. These passivate the surface to the desired catalyst mechanism. This results in deactivation as either reduction in production rate or as loss of selectivity to the required product. Common poisons include S, P and Cl. These are normally adsorbed from contaminants in the gas phase and sacrificial adsorbents can be used to reduce the concentration of these.116 More importantly for pyrolysis catalysts, carbon is also as an important poison and arises from hydrocarbon at the catalyst surface. Extensive carbon formation can leave to ‘coking’ where thick carbon deposits are formed. This coking can also be useful in oil chemistries because it can lead to useful liquid (resids) formation.117