A Primer on Chemical Reactions and Catalysis

From Preface of Basic Research Needs: Catalysis for Energy, 2008
http://www.sc.doe.gov/bes/reports/list.html

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.