Catalysts

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CATALYSTS

A catalyst is a substance whose presence increases the rate of a chemical reaction. The exercise of using catalysts is called catalysis. Today the vast majority of all commercially important chemical reactions involve catalysts, especially in the fields of energy production, petrochemicals manufacture, pharmaceuticals synthesis, and environmental protection.

HISTORY

Catalysis was practiced long before it was recognized as a scientific discipline. The earliest example of catalytic reactions was in the generation of alcoholic beverages through biocatalysis dating from the Neolithic Age. About 2,500 years ago a base-catalyzed (potash lye) process was used to manufacture soap in the Mediterranean area. Although the details are not known, in the 1500s alchemists claimed to have prepared sulfuric acid (sulfuric ether and oil of vitriol they called it) by a mysterious process that would probably be classified as a heterogeneous catalytic reaction today.

During the first third of the nineteenth century, several systematic observations led researchers to conclude that the mere presence of metals induced chemical transformations in fluids that would otherwise not have occurred. Early on, Thenard had observed that ammonia could be decomposed when passed through a red-hot porcelain tube, but only if the tube contained iron, copper, silver, gold, or platinum. Humphrey Davy observed that a warm Pt wire glowed red when placed into a gas-air mixture, and the gas was combusted. His cousin Edmond Davy was able to combust alcohol when exposed to finely-divided Pt particles even at room temperature. Döbereiner combined these discoveries with a hydrogen generator to produce a commercial lighter. Michael Faraday commented on Döbereiner's work in 1823 and, during three months of experiments in 1835, demonstrated catalyst poisoning. Faraday considered catalysts to be just one manifestation of ordinary chemical reactions. Eilhardt Mitscherlich summarized these and other strange results in 1834 and attributed the phenomena to being in "contact" with the substances. Five years before the word was catalysis was coined, in 1831 Peregrine Phillips obtained the first patent in field for an improved method of making sulfuric acid.

In 1836 Jons Jakob Berzelius considered eight seemingly unrelated experimental results and concluded that there was a common thread among them. The commonality he defined as catalysis. In doing this, Berzelius proposed that a "catalytic force" was responsible for catalytic action. The concept of catalysis is today considered by most researchers to be due to Berzelius, probably because of the popularity of his annual Handbook of Chemistry where he published his definition of catalytic action. For the next one hundred years many referred to the phenomenon as "contact catalysis" or "contact action," as proposed by Mitscherlich.

Justus von Liebig was another leader in the training of chemists, and many of his students were placed in influential positions throughout the scientific world during the mid-1800s. Liebig was a forceful personality who defended his "turf" with vigor. His concept of catalysis was strongly influenced by purification, a subject poorly understood at that time. Making an analogy with spoilage, Liebig proposed that catalytic action is based on an induced vibration. Just as one rotten apple will eventually cause all apples in a barrel to rot, so Liebig considered that a substance that vibrates at just the right frequency will induce vibrations in certain other molecules through contact and thereby enhance the rate of their reaction. Liebig used this concept as his basis of catalysis and to explain many other phenomena; he even considered catalysis as being analogous to a perpetual motion machine. Liebig's view was seriously considered for nearly half a century.

Wilhelm Ostwald was also defining physical chemistry during the 1880s. As an editor of a new journal devoted to physical chemistry, he wrote brief critical comments about many papers. Reviewing a paper that used Liebig's vibrational theory to explain results, Ostwald provided a new definition of a catalyst that was widely accepted and led to his being awarded the Nobel Prize in 1909.

Ostwald first came to catalysis through his work on the acceleration of homogeneous reactions by acids. This work was popular at the time although ultimately it would be shown to be incorrect because he believed that the acid, acting as a catalyst, did not enter into the chemical change which it influenced but rather acted by its mere presence (contact catalysis).

Discarding Liebig's theory as worthless because it could not be subjected to experimental verification, Ostwald contended that a catalyst merely sped up a reaction that was already occurring at a very slow rate. He also indicated that a catalyst cannot change the equilibrium composition. By analogy, he considered a catalyst to be like oil to a machine, or a whip to a horse; both caused the rate to increase.

Catalysis was soon divided into two classes: positive ones that accelerate reactions and negative ones that suppress reactions. It is now recognized that what was viewed for many years as negative catalysis was actually a case of catalyst poisoning; that is, some material is so strongly adsorbed that it effectively reduces the number of catalytic sites available to the reactant, thereby decreasing the reaction rate. This led to an understanding that while the catalytic steps themselves regenerate the original catalyst, most catalytic reactions are accompanied by side reactions that irreversibly decrease the catalytic activity. Some of these include sintering of highly dispersed metal particles; chemical poisoning by reactants, products, or impurities in the feed stream; physical blockage of active sites; or mechanical damage to catalytic particles. In spite of these effects, Haensel once calculated that each Pt atom in a petroleum reforming catalyst could convert a staggering 10 million hydrocarbon molecules into higher octane fuels during its lifetime before the catalyst had to be regenerated.

A recent definition of catalysis that is based on thermodynamics was advanced by the Subcommittee on Chemical Kinetics, Physical Chemistry Division, IUPA:

"A catalyst is a substance that increases the rate of a reaction without modifying the overall standard Gibbs energy change in the reaction; the process is called catalysis, and a reaction in which a catalyst is involved is known as a catalyzed reaction."

While this definition does not address the question of "how" catalysts effect rate increases, it does ensure that a catalyst cannot cause the equilibrium composition to deviate from that of the uncatalyzed reaction.

In 1947 Sir Hugh S. Taylor summarized the state of catalysis in a "Science in Progress" article as follows:

"Catalysis has been employed in science to designate a substance which by its mere presence facilitates or enhances the rate of chemical reactions. As such it was a cloak for ignorance. When the states of an over-all catalytic process can be described in terms of a well-defined succession of chemical and physical processes the details of which are well understood or are quite plausible, then the necessity for employing such a word as catalysis to mask our ignorance no longer exists. . ."

HOMOGENEOUS CATALYTIC REACTION MECHANISMS

Compared with uncatalyzed reactions, catalysts introduce alternative pathways that, in nearly all cases, involve two or more consecutive reaction steps. Each of these steps has a lower activation energy than does the uncatalyzed reaction. We can use as an example the gas phase reaction of ozone and oxygen atoms. In the homogeneous uncatalyzed case, the reaction is represented to occur in a single irreversible step that has a high activation energy:

When chlorine acts as a catalyst, the reaction can be considered as two steps with the Cl being depleted in Reaction 2 and regenerated in Reaction 3: The activation energies of Reactions 2 and 3 are each much lower than the activation energy of the uncatalyzed case [1]. Thus, the kinetic definition could be stated along the following lines: A catalyst effects the rate increase by altering the homogeneous reaction pathway to a polystep reaction pathway, wherein each catalyzed step has a lower activation energy than the single homogeneous reaction, thereby increasing the rate of reactant conversion above that of the uncatalyzed reaction.

HETEROGENEOUS CATALYTIC REACTION MECHANISMS

Heterogeneous catalytic reactions always involve more than one phase with an interface separating them. The chemical reactions occur at that interface, as shown in Figure 1. A fluid molecule (e.g., gaseous) to be converted must react with a surface (usually solid) to form a surface adsorbed species. That species then reacts either with another adsorbed molecule (or a molecule from the fluid phase, or it may act unimolecularly as in Figure 1) to be transformed into an adsorbed product molecule, which then desorbs into the fluid phase. Each step (dashed lines) must have an activation energy that is lower than the homogeneous barrier height (solid curve). The depth of the potential energy curve indicates the strength with which each of the species is adsorbed. If the energy decrease is very large, the molecules are strongly adsorbed. When the strongly adsorbed molecule is the reaction product, it may be difficult to remove and can cover the active sites to the point where the reaction rate is actually inhibited by its presence.

Heterogeneous catalytic systems offer the advantage that separation of the products from the catalyst is usually not a problem. The reacting fluid passes through a catalyst-filled reactor in the steady state, and the reaction products can be separated by standard methods. A recent innovation called catalytic distillation combines both the catalytic reaction and the separation process in the same vessel. This combination decreases the number of unit operations involved in a chemical process and has been used to make gasoline additives such as MTBE (methyl tertiary butyl ether).

KINETICS

All catalytic reactions involve chemical combination of reacting species with the catalyst to form some type of intermediate complex, the nature of which is the subject of abundant research in catalysis. The overall reaction rate is often determined by the rate at which these complexes are formed and decomposed. The most widely-used nonlinear kinetic equation that describes homogeneous reactions involving enzyme catalysts was developed by Leonor Michaelis and Maude Menten: where A is the reacting species, E the enzyme catalyst, C_i is the fluid phase concentration of each species i, k is the temperature dependent reaction rate constant, and C_M is the Michaelis constant that has the same dimensions as do the concentration terms. _A is the fraction of enzyme molecules tied up in the intermediate complex at any time. Note that when C_A is much smaller than C_M (weak binding) the rate depends linearly on both the enzyme concentration C_E and the reactant concentration C_A. However, if C_A is much larger than C_M (strong binding), then _A approaches unity and the reaction rate depends only on the concentration of the enzyme and is independent of the reactant concentration.

A similar nonlinear equation for heterogeneous catalytic systems was developed empirically by Olaf Hougen and Kenneth Watson and derived on a more scientific basis by Irving Langmuir and Cyril Hinshelwood. When applied to fluid reactants and solid catalysts, the nonlinear equation in its simplest form becomes where again k is the reaction rate constant,_A is the fraction of active sites covered with adsorbed A, and K_A is the adsorption equilibrium constant (a large value means A is strongly adsorbed).

If the three-parameter Michaelis-Menten equation is divided by C_M, it becomes the same as the three-parameter Langmuir-Hinshelwood equation where 1/C_M = K_A. Both these rate equations can become quite complex when more than one species is competing with the reactant(s) for the enzyme or active sites on the solid catalyst.

TRANSPORT EFFECTS

It is not unusual for the full chemical potential of a reaction to be diminished by slower transport processes (i.e., to be transport limited). In fast liquid phase enzyme reactions, mechanical stirring rates can have a strong influence on the observed kinetics that may be limited by the rate of contacting of the reactants and enzymes. Most heterogeneous catalytic reactions take place on catalysts with surface areas of 100 to 1,000 m2/g. These high surface areas are usually generated by preparing catalysts in the form porous pellets (a few mm in diameter) containing a network of interconnecting pores that may be in the range of a few nanometers in diameter. Diffusion into these small pores by reacting molecules whose size is the same order of magnitude can be extremely slow. Assuming all pores are uniform cylinders, a typical silica-alumina cracking catalyst with surface area of 300 m²/g and pore volume of 0.25 cm³/g would contain pores the order of 33 Å diameter and more than 17,000,000 miles/g total length if they were all connected end to end.

For reactions taking place in a flow reactor packed with catalyst particles, each reacting molecule must negotiate a series of seven consecutive steps to accomplish its conversion. It must diffuse across the external boundary layer surrounding the pellet; diffuse inside the pores to an active site; adsorb on, react, and desorb from the active sites; then the liberated product molecules must diffuse back out of the pellet and across the boundary layer before being counted in the product stream. Any one of these sequential steps could be the bottleneck that limits the overall performance of the catalyst. Moreover, heat generated (or absorbed) during the reactions must be accounted for in order to avoid damage to the catalyst and/or hazards to personnel and the environment. This is why reaction engineering plays such an important role in optimizing catalytic processes.

MULTIFUNCTIONAL CATALYSTS

Catalytic processes frequently require more than a single chemical function, and these bifunctional or polyfunctional materials must be prepared in a way to assure effective communication among the various constituents. For example, naphtha reforming requires both an acidic function for isomerization and alkylation and a hydrogenation function for aromatization and saturation. The acidic function is often a promoted porous metal oxide (e.g., alumina) with a noble metal (e.g., platinum) deposited on its surface to provide the hydrogenation sites. To avoid separation problems, it is not unusual to attach homogeneous catalysts and even enzymes to solid surfaces for use in flow reactors. Although this technique works well in some environmental catalytic systems, such attachment sometimes modifies the catalytic specificity of the homogeneous catalyst due to the geometric constraints imposed on the molecules by the solid. With so many factors contributing to the interdisciplinary field of catalysis, it is not surprising that almost all branches of physical science, math, and engineering must be included in the successful development of a catalytic process.

IMPORTANT COMMERCIAL PROCESSES

Industrial catalytic applications comprise four major categories: chemicals manufacturing (25% of money spent on catalytic processes), environmental protection (23%), petroleum processing (26%), and polymers production (26%). In 2003 the total sales of catalysts worldwide is predicted to be $8.9 billion (not including the value of the precious metals and substrates used and includes only manufacturing fees). It has been estimated that about 20 percent of all the world's manufactured products have been touched somewhere along the line by one or more catalytic processes. The field of catalysis has obviously blossomed during the twentieth century, and without any doubt it will be a major factor in the world economy during the foreseeable future.

Burtron H. Davis Joe W. Hightower

See also: Faraday, Michael; Thermodynamics.