Integrated Process Synthesis – designing processes for reduced CO2 emissions

The Centre of Materials and Process Synthesis, University of the Witwatersrand, South Africa, is presenting a short course, “Integrated Process Synthesis – designing processes for reduced CO2 emissions”, under the auspices of the AIChE in Orlando on 12-14 November 2008. The approach presented in the course represents the work of Professors David Glasser and Diane Hildebrandt, who are amongst the leaders in the field of process synthesis.

For more information, please visit COMPS

A Primer on Chemical Reactions and Catalysis

From Preface of Basic Research Needs: Catalysis for Energy, 2008

Chemical transformations are essential to all living organisms—and also to the manufacture of many products including fuels, plastics, and pharmaceuticals. Without catalysts and catalytic technologies, the ease of transportation and the ready access to all of the materials needed for our daily lives would not be possible. The purpose of this primer is to show why catalysts are required for biological processes as well as those used in technology for the production of most fuels, chemicals, polymers, and pharmaceuticals. As we shall see, catalysts are the ultimate enablers of chemical transformation.

Catalysts Facilitate Molecular Transformations

The extent to which a chemical reaction could possibly transform one kind of molecule into another kind of molecule is governed by the principles of thermodynamics—some reactions are in principle possible, whereas others can, at most, occur to only an immeasurably small extent. But, the reactions that are thermodynamically possible may take place at such low rates as to be essentially stymied—we say that these reactions are limited by kinetics. When the reaction is thermodynamically possible but too slow to be useful, then a catalyst is needed. A catalyst increases the rate by intervening in the chemical change to open up a new, quicker pathway for change.

Thermodynamics and Chemical Reactions

Chemical reactions involve the transformation of reactant molecules to product molecules. A simple example is the combustion of hydrocarbons such as gasoline molecules to make carbon dioxide and water, a process that occurs at high temperature in the cylinder of an automobile engine. The Gibbs free energy change for this reaction, assuming hexane as a typical fuel, is downhill by ~ 68 MJ/L , and the energy released does the work to drive a typical automobile 12.5 miles. The fact that the free energy change is downhill tells us that the reaction is favorable and that once it occurs, the products have a lower potential for doing work than the reactants do. Processes that are characterized by such downhill changes in the Gibbs free energy can, in principle, occur spontaneously. The larger the magnitude of the change in the Gibbs free energy, the larger the ultimate fraction of the reactants that can be converted to products. Though the change in Gibbs free energy for hexane combustion is large, the reaction does not occur spontaneously. Thus, one can place liquid hexane in a glass and observe that it does not burst into flame when exposed to air at room temperature. The reason for the lack of combustion of hexane is that the molecules of hexane and oxygen are content to stay as they are for a very long time. To react, chemical bonds in both kinds of molecules must first break before new ones can form. To get these bonds to break, the temperature of the hexane-oxygen mixture is raised (as occurs in the automobile cylinder). The need for the high temperature is associated with a barrier along the pathway from reactant to product molecules, known as the activation barrier. When the reactant molecules are hot, they have the energy to cross the activation barrier, as the bonds between atoms in the reactant molecule are broken and the transformation of reactants to products ensues. The higher the barrier, the slower the reaction.

Kinetic Energy and Chemical Reactions

We emphasize that the reason reactions proceed more rapidly at higher temperatures is associated with the higher energy of the reactant molecules—we call this “kinetic energy.” A collection of molecules has a distribution of kinetic energies, some high, some low, but the average value of kinetic energy is determined by the temperature. If a reaction is to occur, some fraction of the molecules in the collection must have enough kinetic energy to overcome the barrier. If we think of reactant molecules as skateboarders at the bottom of a trough, then to surmount the walls of the trough and move over to a new trough, some fraction of the skateboarders will need to be moving fast enough (i.e., have sufficient kinetic energy) to surmount the barrier. Thus, the higher the temperature, the greater the fraction of the reactant molecules able to overcome the activation barrier and move over to the product side of the landscape. If the activation barrier is very high, the temperature required to achieve a useful rate of product formation will have to be so high that the vessel walls used to contain the reaction may fail—or the reaction could be so fast that it gets out of control (an explosion could occur!). Alternatively, the cost of the energy required to increase the temperature sufficiently for reaction to occur could become prohibitive. Furthermore, at high temperatures, some reactants may be fragile enough that they will decompose to useless products. Thus, raising the temperature needed to achieve a useful reaction rate can lead to various problems, and a better way is needed to get the reactants over the barrier to form products.

Why Catalysts Matter

Catalysts provide the better way. They alter the pathway for the reaction, so that the barrier becomes smaller. The catalyst works by interacting with the reactant molecules (forming chemical bonds with them) to alter the energy landscape for the reaction, leading to a lower activation barrier and, hence, a higher rate of reaction.

Because nature has to do most of its biological chemistry at near- ambient conditions, it has evolved an enormous set of catalysts, mostly enzymes, which are exquisitely tuned so that each one facilitates a single chemical reaction for a single reactant. When a series of reactions is to be carried out as, for example, in the metabolism of food, nature uses a different enzyme for each step in the series, and all the enzymes work in the same medium at the same temperature. Catalysts are also used to accelerate the chemical reactions used in the fuels and chemicals industry, but these catalysts are more primitive than nature’s catalysts. Thus, for example, if we wanted to reduce the temperature of hexane combustion, we could expose a hexane-oxygen mixture to a catalyst containing very small particles— nanoparticles—of the precious metal platinum. This same catalyst converts unburned gasoline in automobile exhaust converters, minimizing the pollution it would otherwise cause, and it simultaneously converts toxic carbon monoxide and nitrogen oxides in the exhaust to the non-toxic products carbon dioxide and nitrogen. Catalysts are also used to enhance the rate of a reaction to a preferred product relative to an undesired product. For example, silver catalyzes the oxidation of ethylene to ethylene oxide, the precursor to ethylene glycol, which is used as antifreeze in automobiles or as one of the monomers for making polyethylene terephthalate, the polymer used for making soft drink bottles. The beauty of a properly tuned silver catalyst is that it promotes the oxidation of ethylene to ethylene oxide rather than the combustion of ethylene to carbon dioxide and water. Thus, even though the thermodynamically preferred products are carbon dioxide and water, silver alters the reaction pathway so that more than 90 percent of the ethylene goes to ethylene oxide. The net effect is that ethylene is used efficiently to make the valuable product ethylene oxide and the undesired products, carbon dioxide and water, are minimized.

Driving Chemical Reactions that are Thermodynamically Uphill

Some reactions are characterized by a change in the Gibbs free energy of reaction that is uphill. For such reactions, thermodynamics teaches us that the reaction cannot occur to a significant extent, unless energy is supplied in the form of photons (e.g., sunlight) or electrons (e.g., from a hydroelectric generator). For example, plants are able to drive an uphill reaction converting carbon dioxide and water to the sugar glucose and oxygen by using sunlight via the process of photosynthesis. Alternatively, the same reactants can be converted electrochemically into carbon monoxide and hydrogen, a mixture that can be used with well-developed catalytic technology to manufacture diesel fuel. The kinetics of reactions that are uphill thermodynamically are often slow, even in the presence of light or electrons. But, the intervention of a catalyst opens a pathway for such reactions to occur at a higher rate with lower energy requirements for the photons or electrons. Catalysts of this type are referred to as photo- or electrocatalysts. Thus, for example, nature uses a series of enzymes to catalyze the photosynthesis of sugars from carbon dioxide and water, and platinum electrodes catalyze the conversion of the same reactants to carbon monoxide and hydrogen.

In Summary

Catalysts are required to facilitate chemical reactions so that they occur at useful rates and with preference to the desired product. If the rate of a reaction is too low, the size of the vessel in which the reaction takes place will be excessively large and expensive. If the product selectivity is low, the reactants are not used efficiently, and energy will be needed to separate the desired products from the undesired products. Thus, the availability of catalysts that make the reaction go fast (active catalysts); make the reaction go to the desired products (selective catalysts); and last a long time or regenerate themselves (stable or regenerable catalysts) allows us to carry out chemical reactions in the most efficient, economical, and environmentally responsible manner. Moreover, using catalysts to reduce the temperature at which reactions occur while achieving high conversions of reactants and high yields of desired products allows us to carry out the transformation with a maximum savings of the energy consumed. Virtually all of the products used by modern societies for fuels, chemicals, polymers, and pharmaceuticals, as well as for abatement of air and water pollution, depend on catalysts. It is notable that the catalysts discovered and developed by humankind are quite primitive relative to those that nature has evolved. However, advances made in the understanding of how catalysts work, together with advances in strategies for making them and the lessons learned from nature, are opening the way towards the design, preparation, and implementation of catalysts that will rival nature’s own and spare our precious energy and raw materials.

Catalysis in the Processing of Crude Oil

The refining of crude oil into petroleum products such as gasoline (trillions of barrels) and other chemicals (billions of pounds) is a major business that relies heavily on catalysis. The petroleum refining catalyst business is in excess of $650 million in the U.S. and over $ 1 billion worldwide.@ Over 90% of our chemical products are derived from petroleum. Once the crude oil is distilled, the products must be further treated with catalysts to produce valuable products. A major operation within refineries is catalytic cracking. Cracking is the production of smaller molecules, often for use in gasoline production, from the larger molecules distilled from crude oil. Over 380 million pounds of catalyst* are used for this operation every year. Other major steps in the refining of petroleum include hydrotreating (a major catalytic process used to remove unwanted sulfur and amine containing products within petroleum), reforming (the use of platinum based catalysts for rearranging petroleum products into gasoline), and alkylation (the use of large amounts of hydrofluoric or sulfuric acids to obtain branched chain molecules with higher octane numbers). Alkylation catalyst production exceeds 200 million pounds, worldwide.@ Beyond the volume and value of the catalysts themselves, there is a hugh value-added component obtained from the consumer products derived from petroleum based by-products.

Catalysts play an important economic role in extracting valuable products from petroleum efficiently. In the years to come, they will be a key part in developing new fuels to meet tougher emission control requirements and increasing our Nation’s energy efficiency.

16 September, 1992, J.N. Armor

Mobile Engine Emission Control Catalysts

4 June 1994

Since the 1960s. the U. S. Government (and now many other countries) required automobile manufacturers control the emission of nitrogen oxides (NOx), carbon monoxide (CO), and hydrocarbons produced by gasoline powered automobiles. Emissions regulations established for 1982 and later vehicles led to the development of the current three-way catalyst that simultaneously controls all three pollutants to the required levels. In the late 1980s, this was already a $500 million/year business in the U. S.*

Typical three way catalysts contain rhodium, platinum, and/or palladium metals with other additives that are all supported on an alumina support.# Generally, the supported catalyst is distributed onto a ceramic honeycomb that is then encased within a steel container mounted under the passenger compartment. Exhaust gases then diffuse into the pores and react with a catalyst and exit as non-pollutants. The catalyst reduces the pollutants within about 0.5 second and operates at about 1000°F. These durable systems operate efficiently for the life of the vehicle. Nevertheless, new changes to the regulations will demand further catalyst improvements. Improved catalysts are needed for controlling cold start emissions from lean fuel operated engines. Recently, B. J. Cooper summarized@ some of the technical challenges remaining in auto-emission control catalysis. Also, catalysts are needed for controlling emissions from diesel engines, especially with regard to soot control.

John N. Armor, PhD
Group Head
Catalysis Skill Center

* B. F. Greek, Chemical & Engineering News, (May 29, 1989) 29-56.
# K. C. Taylor, Chemtech, (September 1990) 551-555.
@ B. J. Cooper, Plat. Met. Rev. 38 (1994) 2-10.

Fluid Catalytic Cracking and Eger Murphree

Patent No. 2,451,804 Method of and Apparatus for Contacting Solids and Gases

Over half the world’s gasoline is currently produced by a process developed in 1942 by a group called the “Four Horsemen” of Exxon Research and Engineering Company. The world’s first commercial Fluid Catalytic Cracking facility began production for Exxon on May 25, 1942. The Fluid Cat Cracking process revolutionized the petroleum industry by more efficiently transforming higher boiling oils into lighter, usable products.

The four Exxon inventors responsible for this cracking process are Donald L. Campbell, Homer Z. Martin, Eger V. Murphree, and Charles Wesley Tyson.

When Exxon’s first commercial cat cracking facility went on-line in 1942, the U.S. had just entered World War II and was facing a shortage of high-octane aviation gasoline. This new process allowed the U.S. petroleum industry to increase output of aviation fuel by 6,000% over the next three years. Fluid Cat Cracking also aided the rapid buildup of butadiene production, which enhanced Exxon’s process for making synthetic butyl rubber–another new technology vital to the Allied war effort.

In the 1930s, Exxon began looking for a way to increase the yield of high-octane gasoline from crude oil. Researchers discovered that a finely powdered catalyst behaved like a fluid when mixed with oil in the form of vapor. During the cracking process, a catalyst will split hydrocarbon molecule chains into smaller pieces. These smaller, or cracked, molecules then go through a distillation process to retrieve the usable product. During the cracking process, the catalyst becomes covered with carbon; the carbon is then burned off and the catalyst can be re-used.

Campbell, Martin, Murphree, and Tyson began thinking of a design that would allow for a moving catalyst to ensure a steady and continuous cracking operation. The four ultimately invented a fluidized solids reactor bed and a pipe transfer system between the reactor and the regenerator unit in which the catalyst is processed for re-use. In this way, the solids and gases are continuously brought in contact with each other to bring on the chemical change.

This work culminated in a 100 barrel-per-day demonstration pilot plant located at Exxon’s Baton Rouge facility. The first commercial production plant processed 13,000 barrels of heavy oil daily, making 275,000 gallons of gasoline.

Considered essential to refinery operation, Fluid Cat Cracking produces gasoline as well as heating oil, fuel oil, propane, butane, and chemical feedstocks that are instrumental in producing other products such as plastics, synthetic rubbers and fabrics, and cosmetics. During today’s Fluid Cat Cracking process, a boxcar load of catalyst is mixed with a stream of oil vapor every minute. It is this mixture, behaving like a fluid, that moves continuously through the system as cracking reactions take place.

Fluid Cat Cracking currently takes place in over 370 Fluid Cat Cracking units in refineries around the world, producing almost 1/2 billion gallons of gasoline daily. It is considered one of the most important chemical engineering achievements of the 20th century. Fluid Cat Cracking technology continues to evolve as cleaner high-performance fuels are explored.

Donald L. Campbell was born August 5, 1904 in Clinton, Iowa. He has always been fascinated by inventing and solving problems. He first attended Iowa State University, then MIT and the Harvard Business School. During his 41 years at Exxon, 25 were spent in Exxon Research & Engineering. At his retirement in 1969, he held 30 patents and was the assistant to the vice president of New Areas of Research.

Homer Zettler Martin was born on November 20, 1910 in Chicago, Illinois. He received his B.S. in chemical engineering from the Illinois Institute of Technology and his M.S. and Ph.D. from Michigan. After joining Exxon in 1937, he became one of its most prolific inventors, with 82 patents upon his retirement in 1973. Martin died in Sun City, Arizona on September 1, 1993.

Eger Vaughan Murphree, born November 3, 1898 in Bayonne, New Jersey, moved as a youngster with his family to Kentucky. At Kentucky University, he graduated with degrees in chemistry and mathematics (1920), then went on for his master’s in chemistry (1921). After working as a high school teacher and football coach for a period of time, he attended MIT for two years. In 1924, he went to work at Solvay Process Company as a chemical engineer, and in 1930, joined what was then Standard Oil of New Jersey. From 1947 to 1962, he served as president of the Standard Oil Development Co., which was renamed Esso Research & Engineering in 1955. In 1956, he was given the job of directing military projects related to the guided-missile program; he served one year as special assistant to Defense Secretary Charles Wilson. Murphree, who was also a member of the committee that organized the Manhattan Project, was widely recognized as a leader in the fields of synthetic toluene, butadiene and hydrocarbon synthesis, fluid catalytic cracking, fluid hydroforming, and fluid coking. He died of a heart attack in 1962.

Charles Wesley Tyson, known as Wes to his friends, was born in 1900. In 1930, after receiving his bachelor’s and master’s degrees in chemical engineering from MIT, he joined Esso. In 1961, he was appointed special assistant to the vice president of Exxon Research & Engineering, and at his retirement in 1962, he held 50 patents. Tyson died in 1977.

Copyright 1999, National Inventors Hall of Fame, Akron, Ohio.

About Catalysis

Catalysts, in the definition developed by Berzelius and others in the last century, are materials which change the rate of attainment of chemical equilibrium without themselves being changed or consumed in the process.

Catalysis is an astonishing phenomenon. Some catalysts achieve astonishing activities, so that very small quantities of catalyst can convert thousands or millions of times their own weight of chemicals. Equally significant, however, is selectivity; usually thought of in terms of a catalyst accelerating one of a number of competing reactions, but also possible by virtue of a catalyst selecting one reagent out of a complex mixture.

Catalysis is the key to both life and lifestyle. It is an essential technology for chemical and materials manufacturing, for fuel cells and other energy conversion systems, for combustion devices, and for pollution control systems. Catalysts are widely used in food processing, and enhance the performance of other consumer products such as laundry detergents. The possibility of analysing and ultimately manipulating genes rests on the catalytic properties of RNA to replicate molecules containing biological information. New sensor systems use catalytic surfaces to detect specific molecules and announce their presence through the heat of a vigorous catalytic reaction. And while the tendency is to think of catalysis as a phenomenon for making things happen, the basis of many valuable drugs is the opposite phenomenon; Viagra and Quinapril combat impotence and hypertension by inhibiting enzymes, respectively PDE-V, a phosphodiesterase which breaks down the NO messenger cGMP, and ACE, the Angiotensin-Converting Enzyme.

The economic contribution from catalysis is as remarkable as the phenomenon itself. Estimates from just four years ago that 35% of global GDP depends on catalysis missed much of the emergent genetic business. Confining the analysis to the chemicals industry, with global sales of perhaps US$1.5 x 1012 the proportion of processes using catalysts is 80% and increasing. The catalyst market itself is US$1010, so that catalysis costs are much less than 1% of the sales revenue from the products which they help create. Small wonder that the catalyst market is increasing at 5% per annum.

The terms “catalyst” and “catalysis” have also translated from the world of science to everyday cliché. Our western society places a high value on the power to induce change, under the descriptor “progress”, and it is small wonder that “catalyst” is a tradename chosen for wine, perfume, magazines, management consultancies and advertising agencies. There is even a breed of comic-book superheroes.

The eastern tradition is different. Rather than depicting a catalyst as an agent of rapid breakdown and change, the Chinese characters for “catalyst” also apply to “marriage broker”.

This is a subtle and perceptive appreciation of how catalysts work. It also seems most appropriate given that the successful creation and application of catalytic processes is genuinely multidisciplinary. On a technical level it requires skills in chemistry, chemical engineering, materials technology, as well as the economics and practicalities of manufacturing processes. And it can best be induced by active and strategic collaboration between industry, universities and government.

Chris Adams
Institute of Applied Catalysis
See also: “Catalysing Business” by C J Adams, Chemistry and Industry, 1999, pp740-743