In most of the processes studied within the physical sciences, the lesson again and again is that nature provides no "free lunch"; in other words, it is not possible to get something for nothing. A chemical reaction, for instance, involves the creation of substances different from those that reacted in the first place, but the number of atoms involved does not change. In view of nature's inherently conservative tendencies, then, the idea of a catalyst—a substance that speeds up a reaction without being consumed—seems almost like a magic trick. But catalysts are very real, and their presence in the human body helps to sustain life. Similarly, catalysts enable the synthesis of foods, and catalytic converters in automobiles protect the environment from dangerous exhaust fumes. Yet the presence of one particular catalyst in the upper atmosphere poses such a threat to Earth's ozone layer that production of certain chemicals containing that substance has been banned.
HOW IT WORKS
Reactions and Collisions
In a chemical reaction, substances known as reactants interact with one another to create new substances, called products. In the present context, our concern is not with the reactants and products themselves, but with an additional entity, an agent that enables the reaction to move forward at faster rates and lower temperatures.
According to the collision model generally accepted by chemists, chemical reactions are the result of collisions between molecules. Collisions that are sufficiently energetic break the chemical bonds that hold molecules together; as a result, the atoms in those molecules are free to recombine with other atoms to form new molecules. Hastening of a chemical reaction can be produced in one of three ways. If the concentrations of the reactants are increased, this means that more molecules are colliding, and potentially more bonds are being broken. Likewise if the temperature is increased, the speeds of the molecules themselves increase, and their collisions possess more energy.
Energy is an important component in the chemical reaction because a certain threshold, called the activation energy (E a), must be crossed before a reaction can occur. A temperature increase raises the energy of the collisions, increasing the likelihood that the activation-energy threshold will be crossed, resulting in the breaking of molecular bonds.
Catalysts and Catalysis
It is not always feasible or desirable, however, to increase the concentration of reactants, or the temperature of the system in which the reaction is to occur. Many of the processes that take place in the human body, for instance, "should" require high temperatures—temperatures too high to sustain human life. But fortunately, our bodies contain proteins called enzymes, discussed later in this essay, that facilitate the necessary reactions without raising temperatures or increasing the concentrations of substances.
An enzyme is an example of a catalyst, a substance that speeds up a reaction without participating in it either as a reactant or product. Catalysts are thus not consumed in the reaction. The catalyst does its work—catalysis—by creating a different path for the reaction, and though the means whereby it does this are too complex to discuss in detail here, the process of catalyst can at least be presented in general terms.
Imagine a graph whose x-axis is labeled "reaction progress," while the y-axis bears the legend "energy." There is some value of y equal to the normal activation energy, and in the course of experiencing the molecular collisions that lead to a reaction, the reactants reach this level. In a catalyzed reaction, however, the level of activation energy necessary for the reaction is represented by a lower y-value on the graph. The catalyzed substances do not need to have as much energy as they do without a catalyst, and therefore the reaction can proceed more quickly—without changing the temperature or concentrations of reactants.
A Brief History of Catalysis
Long before chemists recognized the existence of catalysts, ordinary people had been using the process of catalysis for a number of purposes: making soap, for instance, or fermenting wine to create vinegar, or leavening bread. Early in the nineteenth century, chemists began to take note of this phenomenon.
In 1812, Russian chemist Gottlieb Kirchhof was studying the conversion of starches to sugar in the presence of strong acids when he noticed something interesting. When a suspension of starch in water was boiled, Kirchhof observed, no change occurred in the starch. However, when he added a few drops of concentrated acid before boiling the suspension (that is, particles of starch suspended in water), he obtained a very different result. This time, the starch broke down to form glucose, a simple sugar, while the acid—which clearly had facilitated the reaction—underwent no change.
Around the same time, English chemist Sir Humphry Davy (1778-1829) noticed that in certain organic reactions, platinum acted to speed along the reaction without undergoing any change. Later on, Davy's star pupil, the great British physicist and chemist Michael Faraday (1791-1867), demonstrated the ability of platinum to recombine hydrogen and oxygen that had been separated by the electrolysis of water. The catalytic properties of platinum later found application in catalytic converters, as we shall see.
AN IMPROVED DEFINITION.
In 1835, Swedish chemist Jons Berzelius (1779-1848) provided a name to the process Kirchhof and Davy had observed from very different perspectives: catalysis, derived from the Greek words kata ("down") and lyein ("loosen.") As Berzelius defined it, catalysis involved an activity quite different from that of an ordinary chemical reaction. Catalysis induced decomposition in substances, resulting in the formation of new compounds—but without the catalyst itself actually entering the compound.
Berzelius's definition assumed that a catalyst manages to do what it does without changing at all. This was perfectly adequate for describing heterogeneous catalysis, in which the catalyst and the reactants are in different phases of matter. In the platinum-catalyzed reactions that Davy and Faraday observed, for instance, the platinum is a solid, while the reaction itself takes place in a gaseous or liquid state. However, homogeneous catalysis, in which catalyst and reactants are in the same state, required a different explanation, which English chemist Alexander William Williamson (1824-1904) provided in an 1852 study.
In discussing the reaction observed by Kirchhof, of liquid sulfuric acid with starch in an aqueous solution, Williamson was able to show that the catalyst does break down in the course of the reaction. As the reaction takes place, it forms an intermediate compound, but this too is broken down before the reaction ends. The catalyst thus emerges in the same form it had at the beginning of the reaction.
Enzymes: Helpful Catalysts in the Body
In 1833, French physiologist Anselme Payen (1795-1871) isolated a material from malt that accelerated the conversion of starch to sugar, as for instance in the brewing of beer. Payen gave the name "diastase" to this substance, and in 1857, the renowned French chemist Louis Pasteur (1822-1895) suggested that lactic acid fermentation is caused by a living organism.
In fact, the catalysts studied by Pasteur are not themselves separate organisms, as German biochemist Eduard Buchner (1860-1917) showed in 1897. Buchner isolated the catalysts that bring about the fermentation of alcohol from living yeast cells—what Payen had called "diastase," and Pasteur "ferments." Buchner demonstrated that these are actually chemical substances, not organisms. By that time, German physiologist Willy Kahne had suggested the name "enzyme" for these catalysts in living systems.
Enzymes are made up of amino acids, which in turn are constructed from organic compounds called proteins. About 20 amino acids make up the building blocks of the many thousands of known enzymes. The beauty of an enzyme is that it speeds up complex, life-sustaining reactions in the human body—reactions that would be too slow at ordinary body temperatures. Rather than force the body to undergo harmful increases in temperature, the enzyme facilitates the reaction by opening up a different reaction pathway that allows a lower activation energy.
One example of an enzyme is cytochrome, which aids the respiratory system by catalyzing the combination of oxygen with hydrogen within the cells. Other enzymes facilitate the conversion of food to energy, and make possible a variety of other necessary biological functions.
Because numerous interactions are required in their work of catalysis, enzymes are very large, and may have atomic mass figures as high as 1 million amu. However, it should be noted that reactions are catalyzed at very specific locations—called active sites—on an enzyme. The reactant molecule fits neatly into the active site on the enzyme, much like a key fitting in a lock; hence the name of this theory, the "lock-and-model."
Catalysis and the Environment
The exhaust from an automobile contains many substances that are harmful to the environment. As a result of increased concerns regarding the potential damage to the atmosphere, the federal government in the 1970s mandated the adoption of catalytic converters, devices that employ a catalyst to transform pollutants in the exhaust to less harmful substances.
Platinum and palladium are favored materials for catalytic converters, though some nonmetallic materials, such as ceramics, have been used as well. In any case, the function of a catalytic converter is to convert exhausts through oxidation-reduction reactions. Nitric oxide is reduced to molecular oxygen and nitrogen; at the same time, the hydrocarbons in petroleum, along with carbon monoxide, are oxidized to form carbon dioxide and water. Sometimes a reducing agent, such as ammonia, is used to make the reduction process more effective.
A DANGEROUS CATALYST IN THE ATMOSPHERE.
Around the same time that automakers began rolling out models equipped with catalytic converters, scientists and the general public alike became increasingly concerned about another threat to the environment. In the upper atmosphere of Earth are traces of ozone, a triatomic (three-atom) molecular form of oxygen which protects the planet from the Sun's ultraviolet rays. During the latter part of the twentieth century, it became apparent that a hole had developed in the ozone layer over Antarctica, and many chemists suspected a culprit in chlorofluorocarbons, or CFCs.
CFCs had long been used in refrigerants and air conditioners, and as propellants in aerosol sprays. Because they were nontoxic and noncorrosive, they worked quite well for such purposes, but the fact that they are chemically unreactive had an extremely negative side-effect. Instead of reacting with other substances to form new compounds, they linger in Earth's atmosphere, eventually drifting to high altitudes, where ultraviolet light decomposes them. The real trouble begins when atoms of chlorine, isolated from the CFC, encounter ozone.
Chlorine acts as a catalyst to transform the ozone to elemental oxygen, which is not nearly as effective as ozone for shielding Earth from ultra-violet light. It does so by interacting also with monatomic, or single-atom oxygen, with which it produces ClO, or the hypochlorite ion. The end result of reactions between chlorine, monatomic oxygen, hypochlorite, and ozone is the production of chlorine, hypochlorite, and diatomic oxygen—in other words, no more ozone. It is estimated that a single chlorine atom can destroy up to 1 million ozone molecules per second.
Due to concerns about the danger to the ozone layer, an international agreement called the Montreal Protocol, signed in 1996, banned the production of CFCs and the coolant Freon that contains them. But people still need coolants for their homes and cars, and this has led to the creation of substitutes—most notably hydrochlorofluorocarbons (HCFCs), organic compounds that do not catalyze ozone.
Other Examples of Catalysts
Catalysts appear in a number of reactions, both natural and artificial. For instance, catalysts are used in the industrial production of ammonia,nitric acid (produced from ammonia), sulfuric acid, and other substances. The ammonia process, developed in 1908 by German chemist Fritz Haber (1868-1934), is particularly noteworthy. Using iron as a catalyst, Haber was able to combine nitrogen and hydrogen under pressure to form ammonia—one of the world's most widely used chemicals.
Eighteen ninety-seven was a good year for catalysts. In that year, it was accidentally discovered that mercury catalyzes the reaction by which indigo dye is produced; also in 1897, French chemist Paul Sabatier (1854-1941) found that nickel catalyzes the production of edible fats. Thanks to Sabatier's discovery, nickel is used to transform inedible plant oils to margarine and shortening.
Another good year for catalysts—particularly those involved in the production of polymers—was 1953. That was the year when German chemist Karl Ziegler (1898-1973) discovered a resin catalyst for the production of polyethylene, which produced a newer, tougher product with a much higher melting point than polyethylene as it was produced up to that time. Also in 1953, Italian chemist Giulio Natta (1903-1979) adapted Ziegler's idea, and developed a new type of plastic he called "isotactic" polymers. These could be produced easily, and in abundance, through the use of catalysts.
One of the lessons of chemistry, or indeed of any science, is that there are few things chemists can do that nature cannot achieve on a far more wondrous scale. No artificial catalyst can compete with enzymes, and no use of a catalyst in a laboratory can compare with the grandeur of that which takes place on the Sun. As German-American physicist Hans Bethe (1906-) showed in 1938, the reactions of hydrogen that form helium on the surface of the Sun are catalyzed by carbon—the same element, incidentally, found in all living things on Earth.
WHERE TO LEARN MORE
"Bugs in the News: What the Heck Is an Enzyme?" University of Kansas (Web site). <http://falcon.cc.ukans.edu/~jbrown/enzyme.html> (June 9, 2001).
"Catalysis." University of Idaho Department of Chemistry (Web site). <http://www.chem.uidaho.edu/~honors/rate4.html> (June 9, 2001).
"Catalysts" (Web site). <http://edie.cprost.sfu.ca/~rhlogan/catalyst.html> (June 9, 2001).
Ebbing, Darrell D.; R. A. D. Wentworth; and James P. Birk. Introductory Chemistry. Boston: Houghton Mifflin, 1995.
"Enzymes." Strategis (Web site). <http://strategis.ic.gc.ca/SSG/tc00048e.html> (June 9, 2001).
"Enzymes: Classification, Structure, Mechanism." The Hebrew University (Web site). <http://www.md.huji.ac.il/MedChem/Mechanism-Chymotrypsin/> (June 9, 2001).
Oxlade, Chris. Chemistry. Illustrated by Chris Fairclough. Austin, TX: Raintree Steck-Vaughn, 1999.
"Ozone Depletion" (Web site). <http://www.energy.rochester.edu/iea/1992/p1/2-3.htm> (June 9, 2001).
"University Chemistry: Chemical Kinetics: Catalysis." University of Alberta Department of Chemistry (Web site). <http://www.chem.ualberta.ca/courses/plambeck/p102/p0217x.htm> (June 9, 2001).
Zumdahl, Steven S. Introductory Chemistry: A Foundation, 4th ed. Boston: Houghton Mifflin, 2000.
The minimal energy required to convert reactants intoproducts, symbolized Ea
A mixture of water and a substance that is dissolved in it.
A substance that speeds upa chemical reaction without participating in it, either as a reactant or product. Catalysts are thus not consumed in the reaction.
A process whereby the chemical properties of a substance are changed by a rearrangement of the atoms in the substance.
The theory that chemical reactions are the result of collisions between molecules strong enough to break bonds in the reactants, resulting in are formation of atoms.
A reaction in which the catalyst and the reactants are in different phases of matter.
A reaction in which catalyst and reactants are in the same phase of matter.
The substance or substances that result from a chemical reaction.
A substance that interacts with another substance in a chemical reaction, resulting in a product.
In chemistry and other sciences, the term "system" usually refers to any set of interactions isolated from the rest of the universe. Anything outside of the system, including all factors and forces irrelevant to a discussion of that system, is known as the environment.
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.
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 . 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).
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, Ci is the fluid phase concentration of each species i, k is the temperature dependent reaction rate constant, and CM 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 CA is much smaller than CM (weak binding) the rate depends linearly on both the enzyme concentration CE and the reactant concentration CA. However, if CA is much larger than CM (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 KA is the adsorption equilibrium constant (a large value means A is strongly adsorbed).
If the three-parameter Michaelis-Menten equation is divided by CM, it becomes the same as the three-parameter Langmuir-Hinshelwood equation where 1/CM = KA. 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.
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.
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.
Boudart, M., and Diega-Mariadassou, G. (1984). Kinetics of Heterogenous Catalytic Reactions. Princeton, NJ: Princeton University Press.
Hill, C. G. (1977). An Introduction to Chemical Engineering Kinetics and Reactor Design. New York: John Wiley & Sons.
Prettre, M. (1963). Catalysis and Catalysts, tr. D. Antin. New York: Dover Publications.
Roberts, W. W. (2000). "Birth of the Catalytic Concept." Catalysis Letters 67:1.
Satterfield, C. G. (1970). Mass Transfer in Heterogeneous Catalysis. Cambridge, MA: MIT Press.
Hydrochlorofluorocarbons (HCFCs) are compounds consisting of hydrogen, chlorine, fluorine, and carbon atoms. HCFCs and hydrofluorocarbons (HFCs) were created in the 1980s as substitutes for chlorofluorocarbons (CFCs) for use in refrigeration and a wide variety of manufacturing processes. Because all three of these classes of compounds either destroy the stratospheric ozone layer essential to life on Earth or contribute to global warming, international agreements have been signed to eliminate their production and use by either the year 2000 (CFCs) or 2040 (HCFCs and HFCs).
Thomas Midgley, an organic chemist working at the Frigidaire division of General Motors, created chlorofluorocarbons in 1928 as a safe and inexpensive coolant for use in refrigerators and air conditioners. CFCs are nonflammable, nontoxic, noncorroding gases. In addition to their widespread use as coolants, they were used in the manufacturing such products as contact lenses, telephones, artificial hip joints, foam for car seats and furniture, and computer circuit boards. CFCs have also been used as a propellant in aerosol products.
By 1974, however, researchers discovered that CFCs emitted to the atmosphere slowly accumulated in the stratosphere, higher than about 15 mi (25 km) above Earth’s surface. CFCs are degraded in the stratosphere by solar ultraviolet radiation, and this releases chlorine radicals that attack ozone molecules. Although ozone in the lower atmosphere is a harmful pollutant, in the stratosphere it acts to shield organisms at the surface of Earth from the harmful effects of solar ultraviolet radiation.
When ultraviolet radiation in the stratosphere degrades CFCs or HCFCs, the chlorine released acts to consume ozone molecules, which contain three oxygen atoms, into separate chlorine-oxygen and two-oxygen molecules (the latter is known as oxygen gas). Because the chlorine atoms can persist in the stratosphere for more than a century, they are recycled through the ozone-degrading reactions. One chlorine atom can destroy up to 100,000 molecules of stratospheric ozone.
The use of CFCs as aerosol propellants was banned in the United States, Canada, Switzerland, and the Scandinavian countries in 1978, as the dangers posed by their use were increasingly understood. By the early 1980s, companies such as DuPont, the world’s largest manufacturer of CFCs, were creating alternate, less-damaging compounds, including HCFCs and HFCs.
HCFC compounds react differently from CFCs because HCFCs contain a hydrogen atom, which causes these chemicals to decompose photochemically before they reach the stratosphere. HFCs do not contain chlorine and thus do not attack the ozone layer. HCFCs and HFCs survive in the atmosphere for 2 to 40 years, compared with about 150 years for CFCs.
As a result of their shorter persistence and different molecular composition, HCFC and HFC compounds
Chlorofluorocarbons (CFCs) —Chemical compounds containing chlorine, fluorine and carbon. CFCs were a key component in the development of refrigeration, air conditioning, and foam products.
Greenhouse gases —Gases that contribute to the warming of Earth’s atmosphere. Examples include carbon dioxide, HCFCs, CFCs, and HFCs.
Hydrofluorocarbons (HFCs) —Chemical compounds that contain hydrogen, fluorine, and carbon atoms.
Montreal Protocol on Substances that Deplete the Ozone Layer —An agreement signed by 43 countries in 1987, and amended and signed by 90 nations in 1990, to eliminate the production and use of compounds that destroy the ozone layer.
Ozone —A gas made up of three atoms of oxygen. Pale blue in color, it is a pollutant in the lower atmosphere, but essential for the survival of life on Earth’s surface when found in the upper atmosphere because it blocks dangerous ultraviolet solar radiation.
Ozone layer —A layer of ozone in the stratosphere that shields the surface of Earth from dangerous ultraviolet solar radiation.
Stratosphere —A layer of the upper atmosphere above an altitude of 5–10.6 mi (8–17 km) and extending to about 31 mi (50 km), depending on season and latitude. Within the stratosphere, air temperature changes little with altitude, and there are few convective air currents.
Troposphere —The layer of air up to 15 mi (24 km) above the surface of Earth, also known as the lower atmosphere.
Ultraviolet radiation —Radiation similar to visible light but of shorter wavelength, and thus higher energy.
have replaced CFCs in most major uses, including the production of foams for insulation, furniture, and vehicle seats, and as a coolant in refrigerators and air conditioners.
HCFCs and HFCs are more expensive to manufacture than CFCs and still negatively affect Earth’s atmosphere to some degree. Although HCFCs destroy 98% less ozone in the stratosphere than do CFCs, HCFCs and HFCs are still greenhouse gases that may contribute to global warming. In comparison to a more common greenhouse gas, CFCs are about 4,100 times more efficient in their global warming potential, while HFCs are 350 times more effective.
CFCs and HCFCs have contributed to the quality of modern life, particularly as valuable components in refrigeration and computer technology. However, their impact on the atmosphere has prompted several countries to agree to stop producing them. The Montreal Protocol on Substances that Deplete the Ozone Layer was signed by 43 countries in 1987 to limit and eventually eliminate the production and use of CFCs. When additional evidence emerged that the ozone layer was being damaged more quickly than originally thought, more than 90 countries signed an amendment to the Montreal Protocol in 1990. In the year 2000, CFCs were banned from use and guidelines included new phase-outs for HCFCs and HFCs by the year 2020 if possible, and no later than 2040.
Research results suggest that there is a need to develop acceptable alternatives to HCFCs. In laboratory tests, male rats exposed to 5,000 parts per million (ppm) of HCFCs over a two-year period (equivalent to what humans working occupationally with the compound might experience over 30-40 years) developed tumors in the pancreas and testes. The tumors were benign and did not result in death for the tested rats. Nevertheless, this research resulted in the recommended eight-hour occupational exposure levels to HCFCs for humans being reduced from 100 ppm to 10 ppm.
Two possible alternatives to HCFCs are already being used successfully. Refrigerators that use propane gas, ammonia, or water as coolants are being tested in research laboratories, and use up to 10% less energy than typical models using CFCs as a coolant. Telephone companies are experimenting with crushed orange peels and other materials to clean computer circuit boards, as substitutes for another important use of CFCs and HCFCs. Certain microorganisms are also being tested that degrade HCFCs and HFCs, which could help in controlling emissions of these compounds during manufacturing processes involving their use.
Anslyn, E.V. and D.A. Dougherty. Modern Physical Organic Chemistry. Herndon, VA: University Science Books, 2005.
Hydrochlorofluorocarbons (HCFCs) are compounds made up of hydrogen , chlorine , fluorine, and carbon atoms . HCFCs and their cousins, hydrofluorocarbons (HFCs), were created in the 1980s as substitutes for chlorofluorocarbons (CFCs) for use in refrigeration and a wide variety of manufacturing processes. Because all three of these classes of compounds either destroy the stratospheric ozone layer essential to life on Earth , and/or contribute to an unnatural warming of the planet's climate, international agreements have been signed to eliminate their production and use by either the year 2000 (CFCs) or 2040 (HCFCs and HFCs).
Thomas Midgley, an organic chemist working at the Frigidaire division of General Motors, created chlorofluorocarbons in 1928 as a safe and inexpensive coolant for use in refrigerators and air conditioners. CFCs are non-flammable, non-toxic, non-corroding gases. In addition to their widespread use as coolants, they were used in the manufacturing of hundreds of products, such as contact lenses, telephones, artificial hip joints, foam for car seats and furniture, and computer circuit boards. CFCs have also been used as a propellant of aerosol products.
By 1974, however, researchers discovered that CFCs emitted to the atmosphere slowly traveled to the upper-altitude layer known as the stratosphere, higher than about 15 mi (25 km) above Earth's surface. The CFCs are degraded in the stratosphere by solar ultraviolet radiation , and this releases chlorine radicals that attack ozone molecules. Although ozone in the lower atmosphere is a harmful pollutant, in the stratosphere it acts to shield organisms at the surface of Earth from the harmful effects of solar ultraviolet radiation.
When ultraviolet radiation in the stratosphere degrades CFCs or HCFCs, the chlorine released acts to consume ozone molecules, which contain three oxygen atoms, into separate chlorine-oxygen and two-oxygen molecules (the latter is known as oxygen gas). Because the chlorine atoms can persist in the stratosphere for more than a century, they are recycled through the ozone-degrading reactions; one chlorine atom can destroy up to 100,000 molecules of stratospheric ozone.
The use of CFCs as aerosol propellants was banned in the United States, Canada, Switzerland, and the Scandinavian countries in 1978, as the dangers posed by their use were increasingly understood. By the early 1980s, companies such as DuPont, the world's largest manufacturer of CFCs, were creating alternate, less-damaging compounds, including HCFCs and HFCs.
The good news and the bad news
HCFC compounds react differently from CFCs. This is because the HCFCs contain a hydrogen atom, which causes these chemicals to decompose photochemically before they reach the stratosphere. HFCs do not contain chlorine and thus do not attack the ozone layer. HCFCs and HFCs survive in the atmosphere for two to 40 years, compared with about 150 years for CFCs.
As a result of their shorter persistence and different molecular composition, HCFC and HFC compounds are expected to replace CFCs in most major uses, including the production of foams for insulation, furniture, and vehicle seats, and as a coolant in refrigerators and air conditioners.
HCFCs and HFCs are much more expensive to manufacture than CFCs, and they still negatively affect Earth's atmosphere to some degree. Although HCFCs destroy 98% less ozone in the stratosphere than do CFCs, HCFCs and HFCs are still greenhouse gases that may contribute to global warming . In comparison to carbon dioxide , a more common greenhouse gas, CFCs are about 4,100 times more efficient in their global warming potential, while HFCs are 350 times more effective.
The future of HCFCs
CFCs and HCFCs have contributed to our quality of life, particularly as valuable components in refrigeration and computer technology. However, their impact on the atmosphere has prompted several countries to agree to stop producing them. The Montreal Protocol on Substances that Deplete the Ozone Layer was signed by 43 countries in 1987 to limit and eventually eliminate the production and use of CFCs. When additional evidence emerged that the ozone layer was being damaged more quickly than originally thought, more than 90 countries signed an amendment to the Montreal Protocol in 1990. In the year 2000, CFCs were banned from use, and guidelines included new phaseouts for HCFCs and HFCs by the year 2020 if possible, and no later than 2040.
Other research is increasing the need to develop acceptable alternatives to HCFCs. In laboratory tests, male rats exposed to 5,000 parts per million (ppm) of HCFCs over a two-year period (equivalent to what humans working occupationally with the compound might experience over 30-40 years) developed tumors in the pancreas and testes. The tumors were benign and did not result in death for the tested rats. Nevertheless, this research resulted in the recommended eight-hour occupational exposure levels to HCFCs for humans being reduced from 100 ppm to 10 ppm.
Two possible alternatives to HCFCs are already being used successfully. Refrigerators that use propane gas, ammonia , or water as coolants are being tested in research laboratories, and are using up to 10% less energy than typical models using CFCs as a coolant. Telephone companies are experimenting with crushed orange peels and other materials to clean computer circuit boards, as substitutes for another important use of CFCs and HCFCs. Certain microorganisms are also being tested that degrade HCFCs and HFCs, which could help in controlling emissions of these compounds during manufacturing processes involving their use.
Duden, Jane. The Ozone Layer. New York: Crestwood House, 1990.
Fisher, David E. Fire and Ice. New York: HarperCollins, 1990.
Fisher, Marshall. The Ozone Layer. New York: Chelsea House Publishers, 1992.
Gay, Kathlyn. Air Pollution. New York: Franklin Watts, 1991.
Jahn, F., M. Cook, and M. Graham. Hydrocarbon Exploration and Production. Developments in Petroleum Science. Vol. 46. The Netherlands: Elsevier Science, 2000.
MacKenzie, Debora. "Cheaper Alternatives for CFCs." NewScientist (June 30, 1990): 39-40.
Wallington, Timothy J., et al. "The Environmental Impact of CFC Replacement-HFCs and HCFCs." Environmental Science & Technology 28 (1994): 320A-326A.
. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
- Chlorofluorocarbons (CFCs)
—Chemical compounds containing chlorine, fluorine and carbon. CFCs were a key component in the development of refrigeration, air conditioning, and foam products.
- Greenhouse gases
—Gases that contribute to the warming of the earth's atmosphere. Examples include carbon dioxide, HCFCs, CFCs, and HFCs.
- Hydrofluorocarbons (HFCs)
—Chemical compounds that contain hydrogen, fluorine, and carbon atoms.
- Montreal Protocol on Substances that Deplete the Ozone Layer
—An agreement signed by 43 countries in 1987, and amended and signed by 90 nations in 1990, to eliminate the production and use of compounds that destroy the ozone layer.
—A gas made up of three atoms of oxygen. Pale blue in color, it is a pollutant in the lower atmosphere, but essential for the survival of life on Earth's surface when found in the upper atmosphere because it blocks dangerous ultraviolet solar radiation.
- Ozone layer
—A layer of ozone in the stratosphere that shields the surface of Earth from dangerous ultraviolet solar radiation.
—A layer of the upper atmosphere above an altitude of 5–10.6 mi (8–17 km) and extending to about 31 mi (50 km), depending on season and latitude. Within the stratosphere, air temperature changes little with altitude, and there are few convective air currents.
—The layer of air up to 15 mi (24 km) above the surface of the earth, also known as the lower atmosphere.
- Ultraviolet radiation
—Radiation similar to visible light but of shorter wavelength, and thus higher energy.
The term hydrochlorofluorocarbon (HCFC) refers to halogenated hydrocarbons that contain chlorine and/or fluorine in place of some hydrogen atoms in the molecule. They are chemical cousins of the chlorofluorocarbons (CFCs), but differ from them in that they have less chlorine. A special subgroup of the HCFCs is the hydrofluorocarbons (HFCs), which contain no chlorine at all.
A total of 53 HCFCs and HFCs are possible.
The HCFCs and HFCs have become commercially and environmentally important since the 1980s. Their growing significance has resulted from increasing concerns about the damage being done to stratospheric ozone by CFCs.
Significant production of the CFCs began in the late 1930s. At first, they were used almost exclusively as refrigerants. Gradually other applications—especially as propellants and blowing agents—were developed. By 1970, the production of CFCs was growing by more than 10% per year, with a worldwide production of well over 662 million lb (300 million kg) of one family member alone, CFC-11.
Environmental studies began to show, however, that CFCs decompose in the upper atmosphere . Chlorine atoms produced in this reaction attack ozone molecules (O3), converting them to normal oxygen (O2). Since stratospheric ozone provides protection for humans against solar ultraviolet radiation , this finding was a source of great concern. By 1987, 31 nations had signed the Montreal Protocol, agreeing to cut back significantly on their production of CFCs.
The question became how nations were to find substitutes for the CFCs. The problem was especially severe in developing nations where CFCs are widely used in refrigeration and air-conditioning systems. Countries like China and India refused to take part in the CFC-reduction plan unless
developed nations helped them switch over to an equally satisfactory substitute.
Scientists soon learned that HCFCs were a more benign alternative to the CFCs. They discovered that compounds with less chlorine than the amount present in traditional CFCs were less stable and often decomposed before they reached the stratosphere . By mid 1992, the United States Environmental Protection Agency (EPA) had selected 11 chemicals that they considered to be possible replacements for CFCs. Nine of those compounds are HFCs and two are HCFCs.
The HCFC-HFC solution is not totally satisfactory, however. Computer models have shown that nearly all of the proposed substitutes will have at least some slight effect on the ozone layer and the greenhouse effect . In fact, the British government considered banning one possible substitute for CFCs, HCFC-22, almost as soon as the compound was developed. In addition, one of the most promising candidates, HCFC-123, was found to be carcinogenic in rats.
Finally, the cost of replacing CFCs with HCFCs and HFCs is expected to be high. One consulting firm, Metroeconomica, has estimated that CFC substitutes may be six to 15 times as expensive as CFCs themselves.
[David E. Newton ]
Johnson, J. "CFC Substitutes Will Still Add to Global Warming." New Scientist 126 (April 14, 1990): 20.
MacKenzie, D. "Cheaper Alternatives for CFCs." New Scientist 126 (June 30, 1990): 39–40.
Pool, R. "Red Flag on CFC Substitute." Nature 352 (July 11, 1991): 352.
Stone, R. "Ozone Depletion: Warm Reception for Substitute Coolant." Science 256 (April 3, 1992): 22.