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Introduction

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Heterogeneous catalysis involving a solid phase catalyst particle and fluid phase reactant molecules

Heterogeneous catalysis refers to the type of catalysis where the phase of the catalyst differs from the phase of the reactants.[1] This contrasts with homogeneous catalysis where the reactants and catalyst exist in the same phase. Phase distinguishes between not only solid, liquid, and gas components, but also immiscible mixtures (e.g. oil and water), or anywhere an interface is present. Catalysts are useful because they increase the rate of a reaction[2] without themselves being consumed and are therefore reusable.

Heterogeneous catalysis typically involves solid phase catalysts and gas phase reactants.[3] In this case, there is a cycle of molecular adsorption, reaction, and desorption occurring at the catalyst surface. Thermodynamics, mass transfer, and heat transfer influence the rate (kinetics) of reaction.

Heterogeneous catalysis is very important because it enables faster, large-scale production and the selective product formation.[4] Approximately 35% of the world's GDP is influenced by catalysis.[5] The production of 90% of chemicals (by volume) is assisted by solid catalysts.[3] The chemical and energy industries rely heavily on heterogeneous catalysis. For example, the Haber-Basch process uses metal-based catalysts in the synthesis of ammonia, an important component in fertilizer; 144 million tons of ammonia were produced in 2016.[6]

Adsorption

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Adsorption is an essential step in heterogeneous catalysis. Adsorption is the process by which a gas (or solution) phase molecule (the adsorbate) binds to solid (or liquid) surface atoms (the adsorbent). The reverse of adsorption is is called desorption, the adsorbate splitting from adsorbent. In a reaction facilitated by heterogeneous catalysis, the catalyst is the adsorbent and the reactants are the adsorbate.

Types of Adsorption

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There are two types of adsorption: physisorption, weakly bound adsorption, and chemisorption, strongly bound adsorption. Many processes in heterogeneous catalysis lie between the two extremes. The Lennard-Jones model provides a basic framework for predicting molecular interactions as a function of atomic separation.[7]

Physisorption

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In physisorption, a molecule becomes attracted to the surface atoms via van der Waals forces. These include dipole-dipole interactions, induced dipole interactions, and London dispersion forces. Note that no chemical bonds are formed between adsorbate and adsorbent, and their electronic states remain relatively unperturbed. Typical energies for physisorption are from 3 to 10 kcal/mol.[3] In heterogeneous catalysis, when a reactant molecule physisorbs to a catalyst, it is commonly said to be in a precursor state, an intermediate energy state before chemisorption, a more strongly bound adsorption.[7] From the precursor state, a molecule can either undergo chemisorption, desorption, or migration across the surface.[8] This nature of the precursor state can influence reaction kinetics.[8]

Chemisorption

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When a molecule approaches close enough to surface atoms such that their electron clouds overlap, chemisorption can occur. In chemisorption, the adsorbate and adsorbent share electrons signifying the formation of chemical bonds. Typical energies for chemisorption range from 20 to 100 kcal/mol.[3] Two cases of chemisorption are:

  • Molecular adsorption: the adsorbate remains intact. An example is alkene binding by platinum.
  • Dissociation adsorption: one or more bonds break concomitantly with adsorption. In this case, the barrier to dissociation affects the rate of adsorption. An example of this the binding of H2 to a metal catalyst, where the H-H bond is broken upon adsorption.

Surface Reactions

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Catalytic Reaction Coordinate. (A) uncatalyzed (B) catalyzed (C) catalyzed with discreet intermediates (transition states)

The rate of a reaction is determined by the reactions slowed step, the rate-determining step. Catalysts reduce the energy activation barrier between the starting reactant and reactive intermediates (or transition states). Catalytic intermediates are the transition states between the overall reaction's reactant molecules and product molecules, although oftentimes they are too short-lived to observe. Each transition state represents a local minimum in the reaction coordinate (see right). The overall thermodynamics of the reaction (product and reactant energy) is unchanged, but the path between the endpoints is changed.

There are two main mechanisms for surface reactions in solid/gas heterogeneous catalytic systems.[3] For the reaction: A + B → C

  • Langmuir-Hinshelwood mechanism: The reactant molecules, A and B, both adsorb to the catalytic surface. While adsorbed to the surface, they migrate to each other and react to form a new molecule, C. Molecule C then desorbs from the surface.
  • Eley-Rideal mechanism: One reactant molecule, A, adsorbs to the catalytic surface. Without adsorbing, the other reactant molecule, B, reacts with A to form a new molecule, C, that is adsorbed to the surface. Molecule C then desorbs from the surface.

It is most common for heterogeneous catalysis to be described by the Langmuir-Hinshelwood model because both reactants are activated by the surface catalyst which follows faster reaction kinetics.[9]

Heterogeneous Catalyzed Reaction Cycle

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Catalytic hydrogenation of ethylene with Wilkinson's catalyst.

In heterogeneous catalysis, reactant molecules diffuse from the bulk fluid phase to adsorb to the catalyst surface. The adsorption site is not always a active catalyst site, so reactant molecules must migrate across the surface to an active site. At the active site, reactant molecules will react to form product molecule(s) by following a more energetically facile path through catalytic intermediates (see figure to the right). The product molecules then desorb from the surface and diffuse away. The catalyst itself remains intact and free to mediate further reactions. Transport phenomena such as heat and mass transfer, also play a role in the observed reaction rate.

Catalyst Design

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Catalysts are not active towards reactants across their entire surface; only specific locations possess catalytic activity, called active sites. The surface area of a solid catalyst has a strong influence on the number of available active sites. In industrial practice, solid catalysts are often porous to maximize surface area, commonly achieving 50-400 m2/g.[3] Some mesoporous silicates, such as the MCM-41, have surface areas greater than 1000 m2/g.[10] Porous materials are cost effective due to their high surface area-to-mass ratio and enhanced catalytic activity.

In many cases, a solid catalyst is dispersed on a supporting material to increase surface area (spread the number of active sites) and provide stability.[3] Usually catalyst supports are inert, high melting point materials, but they can also be catalytic themselves. Most catalyst supports are porous (frequently carbon, silica, zeolite, or alumina-based)[5] and chosen for their high surface area-to-mass ratio. For a given reaction, porous supports must be selected such that reactants and products can enter and exit the material.

Often, substances are intentionally added to the reaction feed or on the catalyst to influence catalytic activity, selectivity, and/or stability. These compounds are called promoters. For example, alumina (Al2O3) is added during ammonia synthesis to providing greater stability by slowing sintering processes on the Fe-catalyst.[3]

Zeolite structure ZSM-5. A common catalyst support material in hydrocracking. Acts as a catalyst in hydrocarbon alkylation and isomerization.

Catalyst Deactivation

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Catalyst deactivation is defined as a loss in catalytic activity and/or selectivity over time.

Substances that decrease reaction rate are called poisons. Poisons chemisorb to catalyst surface and reduce the number of available active sites for reactant molecules to bind to.[11] Common poisons include Group V, VI, and VII elements (e.g. S, O, P, Cl), some toxic metals (e.g. As, Pb), and adsorbing species with multiple bonds (e.g. CO, unsaturated hydrocarbons)[7][11]. For example, sulfur disrupts the production of methanol by poisoning the Cu/ZnO catalyst.[12] Substances that increase reaction rate are called promoters. For example, the presence of alkali metals in ammonia synthesis increases the rate of N2 dissociation[12].

The presence of poisons and promoters can alter the activation energy of the rate-limiting step and affect a catalyst's selectivity for the formation of certain products. Depending on the amount, a substance can be favorable or unfavorable for a chemical process. For example, in the production of ethylene, a small amount of chemisorbed chlorine will act as a promoter by improving Ag-catalyst selectivity towards ethylene over CO2, while too much chlorine will act as a poison.[7]

Other mechanisms for catalyst deactivation include:

  • Sintering: when heated, dispersed catalytic metal particles can migrate across the support surface and form crystals. This results in a reduction of catalyst surface area.
  • Fouling: the deposition of materials from the fluid phase onto the solid phase catalyst and/or support surfaces. This results in active site and/or pore blockage.
  • Coking: the deposition of heavy, carbon-rich solids onto surfaces due to the decomposition of hydrocarbons[11]
  • Vapor-solid reactions: formation of an inactive surface layer and/or formation of a volatile compound that exits the reactor.[11] This results in a loss of surface area and/or catalyst material.
  • Solid-state transformation: solid-state diffusion of catalyst support atoms to the surface followed by a reaction that forms an inactive phase. This results in a loss of catalyst surface area.
  • Erosion: continual attrition of catalyst material common in fluidized-bed reactors.[13] This results in a loss of catalyst material.

In industry, catalyst deactivation costs billions every year due to process shutdown and catalyst replacement.[11]

Industrial Examples

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Process flow diagram illustrating the use of catalysis in the synthesis of ammonia (NH3)

In industry, many design variables must be considered including reactor design. The conventional heterogeneous catalysis reactors include batch, continuous, and fluidized-bed reactors, while more recent setups include fixed-bed, microchannel, and multi-functional reactors.[7] Other variables to consider are reactor dimensions, surface area, catalyst type, catalyst support, as well as reactor operating conditions such as temperature, pressure, and reactant concentrations.

Some large-scale industrial processes incorporating heterogeneous catalysts are listed below.[5]

Table

Nanoscience Examples

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Other non- solid catalyst - fluid reactant examples

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References

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  1. ^ Schlögl, Robert (2015-03-09). "Heterogeneous Catalysis". Angewandte Chemie International Edition. 54 (11): 3465–3520. doi:10.1002/anie.201410738. hdl:11858/00-001M-0000-0025-0A33-6. PMID 25693734.
  2. ^ Chemistry, International Union of Pure and Applied. "IUPAC Gold Book - catalyst". goldbook.iupac.org. Retrieved 2019-02-12.
  3. ^ a b c d e f g h Gadi., Rothenberg (2008). Catalysis : concepts and green applications. Weinheim [Germany]: Wiley-VCH. ISBN 9783527318247. OCLC 213106542.
  4. ^ Information., Lawrence Berkeley National Laboratory. United States. Department of Energy. Office of Scientific and Technical (2003). The impact of nanoscience on heterogeneous catalysis. Lawrence Berkeley National Laboratory. OCLC 727328504.
  5. ^ a b c Ma, Zhen; Zaera, Francisco (2006-03-15), "Heterogeneous Catalysis by Metals", in King, R. Bruce; Crabtree, Robert H.; Lukehart, Charles M.; Atwood, David A. (eds.), Encyclopedia of Inorganic Chemistry, John Wiley & Sons, Ltd, doi:10.1002/0470862106.ia084, ISBN 9780470860786, retrieved 2019-03-17
  6. ^ "United States Geological Survey, Mineral Commodity Summaries" (PDF). USGS. January 2018.
  7. ^ a b c d e Meurig), Thomas, J. M. (John (19 November 2014). Principles and practice of heterogeneous catalysis. Thomas, W. J. (Second, revised ed.). Weinheim, Germany. ISBN 9783527683789. OCLC 898421752.{{cite book}}: CS1 maint: location missing publisher (link) CS1 maint: multiple names: authors list (link)
  8. ^ a b Bowker, Michael (2016-03-28). "The Role of Precursor States in Adsorption, Surface Reactions and Catalysis". Topics in Catalysis. 59 (8–9): 663–670. doi:10.1007/s11244-016-0538-6. ISSN 1022-5528. S2CID 101660153.
  9. ^ Petukhov, A.V. (1997). "Effect of molecular mobility on kinetics of an electrochemical Langmuir-Hinshelwood reaction". Chemical Physics Letters. 277 (5–6): 539–544. doi:10.1016/s0009-2614(97)00916-0. ISSN 0009-2614.
  10. ^ Kresge, C. T.; Leonowicz, M. E.; Roth, W. J.; Vartuli, J. C.; Beck, J. S. (1992). "Ordered mesoporous molecular sieves synthesized by a liquid-crystal template mechanism". Nature. 359 (6397): 710–712. doi:10.1038/359710a0. ISSN 0028-0836. S2CID 4249872.
  11. ^ a b c d e Bartholomew, Calvin H (2001). "Mechanisms of catalyst deactivation". Applied Catalysis A: General. 212 (1–2): 17–60. doi:10.1016/S0926-860X(00)00843-7.
  12. ^ a b K., Nørskov, Jens (25 August 2014). Fundamental concepts in heterogeneous catalysis. Studt, Felix., Abild-Pedersen, Frank., Bligaard, Thomas. Hoboken, New Jersey. ISBN 9781118892022. OCLC 884500509.{{cite book}}: CS1 maint: location missing publisher (link) CS1 maint: multiple names: authors list (link)
  13. ^ Forzatti, P (1999-09-14). "Catalyst deactivation". Catalysis Today. 52 (2–3): 165–181. doi:10.1016/s0920-5861(99)00074-7. ISSN 0920-5861.
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