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.