Catalysis

Type of physical science: Chemistry

Field of study: Chemical reactions

Catalysts are substances that increase the rates of chemical reactions without being consumed by them and without affecting their equilibrium point. Catalysts are of great practical use in the chemical industry, as well as providing important insights into the mechanisms and rates of chemical reactions.

Overview

A catalyst is a substance that increases the rate of a chemical reaction while being unchanged by it. Thus, the catalyst is neither a reactant nor a product of the reaction that it catalyzes. For that reason, a small amount of catalyst can lead to the conversion of very large amounts of reactants into products. Moreover, a catalyst does not affect any attributes of the overall reaction having to do with energy. If the reaction is accompanied by the release of heat energy, the same amount of heat is released as in a corresponding uncatalyzed reaction.

Likewise, reactions end by approaching an equilibrium ratio of reactant to product concentrations, the ratio being a characteristic of a particular reaction and closely related to its associated energy change; this final equilibrium ratio is the same regardless of whether the reaction is accelerated by a catalyst.

Because the catalyst is neither reactant nor product, it might appear not to be a participant in the reaction. In fact, the catalyst does take part in a reaction in a transitory manner; it reacts with reactants at the beginning and is released as an unchanged catalyst, along with products, at the end. In other words, the catalyst plays a cyclic role and is regenerated in its original state after the reaction is complete.

To see how a catalyst's participation in a reaction can lead to an increase in rate, it is necessary to see what limits the rates of chemical reactions in general. An important limit is connected with the energy of the reaction; indeed, it is an energy change that determines whether a reaction can occur at all. First, all reactions go downhill with respect to energy: The equilibrium mixture of products and reactants is at a lower energy than reactants alone. By analogy, a ball sitting on a smooth hillside can roll downward spontaneously, going from a higher to a lower energy. A reaction spontaneously goes in the direction of its equilibrium mixture. This mixture, then, corresponds to the bottom of the hill; neither will the ball roll spontaneously up the hillside nor will the reaction leave equilibrium once it is there. If a reaction had a smooth path (in energy) from reactants to products, however, all reactions would go equally rapidly, and catalysts would not accelerate rates. In fact, there is a bump in the energy "hill," and this bump must be surmounted for the reaction to occur (or for the ball to roll). The energy of the bump is called the activation energy. An increased temperature raises reaction rates, and does so by providing heat energy to drive reactants over the barrier and on toward products. In contrast, a catalyst makes a reaction go faster without a temperature increase by making the energy barrier lower. Thus, more molecules can get over the now-diminished barrier and be converted into products. The catalyst decreases the energy barrier by providing an alternative pathway for the reaction. In general, the alternative pathway has more steps than the uncatalyzed reaction, but all the steps exhibit activation energies substantially lower than that of the uncatalyzed reaction. To put it another way, a catalyst allows a reaction that is different from the uncatalyzed one, but that has the same reactants and products and is more rapid, because of the lower activation energy barriers.

Thus, a reaction traverses a path with three identifiable states: reactants, intermediate state on the top of the energy barrier, and products. The reactants are high in energy, the intermediate state (also called "transition state") is higher still, and the energy of the products is lower than either one. The catalyst lowers the energy barrier by forming a complex with reactants such that its activation energy is lower than that of the transition state in the uncatalyzed reaction.

It should be added that the transition state differs from the original reactant molecule in its overall geometry and in the strength of its bonds. The transition state is extremely unstable, with a lifetime of the order of 10^-12 seconds.

There are two major categories of catalyst, heterogeneous and homogeneous. A heterogeneous catalyst is in a different physical state from the one in which the reaction occurs.

For example, the reaction might take place in the liquid phase and the catalyst might be a solid, such as a metallic surface or a powder. In a heterogeneous catalyst, the reactant and the catalyst come together, often with the formation of chemical bonds, and the bound reactant has a lower activation energy than the unbound molecule. In homogeneous catalysis, the catalyst and the reaction occur in the same physical state. For example, the oxidation of sulfur dioxide by oxygen to form sulfuric acid is catalyzed by nitrogen oxides; all these compounds are in the gas phase.

Homogeneous and heterogeneous catalysts work in exactly the same manner, with reactants and catalyst forming a transition state complex, often through the creation of a chemical bond.

Often catalysts are either acids or bases, compounds able either to donate or to receive hydrogen ions. In such cases, the transition state is achieved through either addition or removal of a hydrogen ion, an alteration that makes the structure more reactive. Commonly, any acid or base (including a hydrogen ion or hydroxyl ion) can serve as catalyst, in which case the process is called either general acid or general base catalysis. The acid or base is regenerated at the end of the reaction and is free to continue the catalytic process.

Often, acid catalysis employs the ability of hydrogen ions to draw electrons toward themselves, thereby making transition state bonds more reactive. Positively charged metal ions can also be electron attractors and, hence, catalysts. Because many metal ions have more than one positive charge, they can be formidable electron attractors and, therefore, powerful catalysts.

Metal ions also form complexes with a variety of compounds and ions, and in many cases, the transition state consists of such a complex between metal and portions of reactant molecules.

Enzymes are defined simply as biological catalysts. They catalyze the vast array of chemical reactions that occur in living organisms, and have long been the object of intensive study, to the extent that chemists may know more about their mechanisms than those of any other catalysts. Interest in enzymes comes partly from their fundamental importance in understanding living processes, partly because of obvious medical implications, and partly because they have important industrial applications. Enzymes are usually proteins, large polymers made up of a mixture of about twenty amino acids, but can also be ribonucleic acids (RNA). Enzymes made of RNA only catalyze reactions having to do with RNA transformations.

Enzymes are legendary in their efficiency as catalysts. An enzyme called catalase stimulates the breakdown of hydrogen peroxide to yield water plus oxygen. One molecule of catalase can catalyze the breakdown of more than one million molecules of hydrogen peroxide per second, and other enzymes are equally active. Enzymes work by forming a complex between the enzyme and reactant molecules. Transition state intermediates generally occur within this complex, and when the complex breaks apart, products are released. The mechanisms by which enzymes increase the rate of reactions are usually the same as in nonbiological catalysis, including, for example, general acid and base, and metal-ion catalysis. Because they are enormous molecules with molecular weights often in the hundreds of thousands, enzymes have an additional way of making the much smaller reactant molecules more reactive. When a reactant binds to the surface of an enzyme, it does so at a region called the active site, often inducing a change in the shape of the whole enzyme. This change stretches or bends bonds in the reactant molecule, leading to greater reactivity of the enzyme-reactant complex and, hence, an increased reaction rate.

Applications

Because of their ability to increase the rates of chemical reactions without requiring increased temperature, catalysts are widely used in a variety of industrial chemical reactions.

They are particularly useful in situations in which excess heat might lead to destruction of the reaction products, or in which the products are flammable or explosive at elevated temperatures, as in the case of some hydrocarbon reactions. The lower reaction temperatures made possible by catalysts also yield economic benefits in the form of energy savings and longer lives of reaction facilities.

An example of a process employing catalysis in the refining of petroleum is the so-called "cracking" of hydrocarbons, which means their cleavage into smaller molecules. In practice, cracking leads to conversion of heavy crude oils into such lighter fractions as gasoline and alkenes, which are themselves important industrial intermediates. Such reactions can be induced to occur by elevating temperature to, say, 600 degrees Celsius, with the hydrocarbon in the gas phase, at a pressure of about 15 atmospheres. They can also be carried out at much lower temperatures and pressures through the use of catalysts. In one of the first successful commercial catalytic cracking processes, the catalyst was a solid and consisted of a mixture of aluminum silicate, iron oxide, calcium oxide, and magnesium oxide. Catalysis can be enhanced by further addition of manganous, copper, and vanadium oxides. From the complex nature of the catalyst mixture, it can be deduced that the reactions occurring are, themselves, extremely complex. The smaller hydrocarbons produced by cleavage are recombined in various ways, and a variety of other chemical reactions occur.

There are numerous other examples of the industrial application of catalysis, about as many as there are industrial chemical reactions. Both homogeneous and heterogeneous catalysis are used, cracking obviously being an instance of the latter. Frequently, homogeneous catalysis is carried out using hydrobromic, sulfuric, or phosphoric acid, compounds that increase the rates of industrial-scale molecular rearrangements, dehydrations (removal of water), hydrations (addition of water), and esterifications (linking acidic and alcoholic groups).

Just as catalysts are used in the production of gasoline, so are they used again in its combustion in motor vehicles. Thus, many automobiles are equipped with a catalytic converter in order to diminish the amount of exhaust pollutants, in particular, carbon monoxide, unburned hydrocarbons, and nitrogen oxides. The first two are eliminated by oxidation, employing as a catalyst a mixture of platinum and palladium bonded to a solid matrix. The catalyst is contained in a metal cylinder located in the exhaust system as close to the engine as possible, in order that the exhaust gases remain as hot as possible. The desirability of fuel low in lead additives stems from the ability of lead to coat the platinum-palladium surfaces, thus "poisoning" the catalyst.

Finally, industry has increasingly used enzymes to carry out commercially important reactions. Enzymes make attractive catalysts as they are very specific with regard to the reaction that they catalyze. Thus, they frequently lead to a single product that does not require further purification--an important consideration, as purification can be costly. One disadvantage of enzymes in the chemical industry is their sensitivity to heat; many are unstable for any extended period even at room temperature. It has been possible to circumvent this difficulty through the isolation of enzymes from bacteria that can tolerate high temperatures (such as bacteria isolated from hot springs). Enzymes from these organisms are often remarkably resistant to high temperatures and, therefore, excellent candidates for industrial applications.

Context

The term "catalysis" is derived from the Greek and means "to loosen or break down."

The word was first coined by the Swedish chemist Jons Jakob Berzelius in 1835. The concept was based on observations such as those of Sir Humphry Davy, who earlier showed that platinum speeds the combustion of various gases. Similarly, Michael Faraday showed that metallic platinum accelerated the reaction of the gases hydrogen and oxygen to form water, and noted that the metal's surface needed to be perfectly clean for the rate to be enhanced. The insights of Berzelius were not particularly influential until the mid-nineteenth century, as the whole idea of reaction rate had not yet been fully developed. One of the first studies of the rate of a chemical reaction employed the conversion of cane sugar (which is a double molecule) into its two constituent molecules, the conversion being measured by a change in the optical properties of the sugar solution. This was done in the 1850's, when it was also observed that the presence of acid in the solution increased the rate of the reaction--a case of acid catalysis. This was the first really quantitative study of a catalytic process. By the end of the century, it was realized that a catalyst failed to affect the equilibrium of a reaction and that it increased the rate to the same extent in either reaction direction. By the same time, numerous examples of catalysis had been investigated, and several were being employed in the infant chemical industry. For example, the commercial production of chlorine gas from hydrochloric acid employed copper salts as a catalyst as early as 1870, and the "hardening" of liquid fats by hydrogen was carried out with nickel catalysts in the 1890's.

Biological catalysis by enzymes was implicit in fermentations that had been carried out for millennia but became an object of scientific study only in the last half of the nineteenth century. For example, in 1897, the German chemist Eduard Buchner showed that a cell-free filtrate from broken yeast cells carried out the transformation of sugar into carbon dioxide, a process that we now know requires the action of more than a dozen different enzymes. By 1913, a detailed theory of enzyme-catalyzed reaction rates had been developed by biochemists Lenore Michaelis and Maude Menten, and by 1926, James Sumner had crystallized an enzyme that catalyzed the break-down of urea. Such crystallization was important in making it clear that enzymes were, after all, ordinary (albeit large) molecules, devoid of any "living" qualities that had been earlier attributed to them.

The study of catalysts has significantly enlarged our understanding of reaction rates and of the fundamental mechanisms through which reactions occur. In particular, catalysis has shed important light on the temperature dependence of reactions and on the importance of transition states in reaction pathways. Understanding of biological catalysts, the enzymes, has led to our present extensive knowledge about the chemistry of life and to significant advances in nearly all areas of medicine.

Principal terms

ACID-BASE CATALYST: a catalyst in which the active substance is either an acid (hydrogen ion donor) or a base (hydrogen ion acceptor)

ACTIVATION ENERGY: the energy of an intermediate state in a chemical reaction, also known as the transition state; a catalyst works by providing a reaction route with a lower activation energy

CATALYSIS: increase in rate of a chemical reaction upon addition of a catalyst

CATALYST: a substance that enhances the rate of a reaction without being either a reactant or a product, and without affecting the final equilibrium point of the reaction

ENZYME: a biological catalyst, usually a protein but sometimes a ribonucleic acid molecule

HETEROGENEOUS CATALYST: a catalyst that is in a different physical state from that in which a reaction occurs; for example, the reactants might be liquids and the catalyst, a solid powder

HOMOGENEOUS CATALYST: a catalyst that is in the same physical state as that in which the reaction occurs, as when the reactants and the catalyst all are liquids

REACTANT: chemical that is the starting point of a reaction, in other words, that is transformed by the reaction into products

SURFACE CATALYSIS: catalysis at the interface between a solid catalyst and a liquid, or gas, containing the reactants

TRANSITION STATE: an intermediate state in a reaction pathway, having a higher energy than either reactants or products; this energy is called the activation energy and represents a barrier to completion of the reaction

Bibliography

Bender, M. L., and L. J. Brubacher. CATALYSIS AND ENZYME ACTION. New York: McGraw-Hill, 1973. A lucid introduction to all forms of chemical catalysis, with special attention paid to enzyme mechanisms. The authors have contributed greatly to research in the field, and the book, although somewhat out of date, is perhaps the best introduction available for a reader with a little chemical knowledge. The book is also an excellent source of references up to the time of its publication.

Burwell, R. L. "The Mechanism of Heterogeneous Catalysis." CHEMICAL AND ENGINEERING NEWS 43 (August 22, 1966): 56. A good, nontechnical discussion of a major category of catalysis, with considerable attention paid to practical applications.

Fersht, Allan. ENZYME STRUCTURE AND MECHANISM. 2d ed. New York: W. H. Freeman, 1985. This may be the best book written about enzyme catalysis. It is also an excellent introduction to chemical catalysis in general, with an excellent chapter on that topic and a particularly clear treatment of acid-base catalysis.

Haensel, Vladimir, and Robert L. Burwell, Jr. "Catalysis." SCIENTIFIC AMERICAN 225 (December, 1971): 46-58. A clear discussion of catalysis, with particular emphasis on heterogeneous catalysis. Much of the article discusses economic applications, especially in the refining of petroleum products and their use as starting material in other industrial processes.

Leisten, J. A. "Homogeneous Catalysis: A Reexamination of Definitions." CHEMICAL EDUCATION 41 (1964): 23-27. A clear article that should be understandable to a reader with some knowledge of chemistry. It asks (and answers) a number of fundamental questions about catalysis and gives several alternative definitions that shed considerable light on the process.

Mills, G. A. "Catalysis." In Vol. 4 of ENCYCLOPEDIA OF CHEMICAL TECHNOLOGY, edited by Martin Grayson. 26 vols. 3d ed. New York: Wiley, 1978-1984. A detailed but not excessively technical discussion of chemical catalysis, with examples from petroleum technology.

Acids and Bases

Dynamics of Chemical Reactions

Chemical Reactions and Collisions