4.9 Catalysis

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166 CHAPTER 4 • INTRODUCTION TO ALKENES. STRUCTURE AND REACTIVITY

higher-energy transition state less stable intermediate ‡ lower-energy transition state more stable intermediate (CH3)2CHCH| 2 ‡ Br_ DG°‡ (CH )3C DG°‡ 3 | Br_

(CH3)2C A CH2 (CH3)2C A CH2 STANDARD FREE ENERGY STANDARD (CH3)2CHCH2 Br HBr HBr (CH3)3C Br + slower reaction L + faster reaction L reaction coordinate reaction coordinate

Figure 4.15 A reaction free-energy diagram for the two possible modes of HBr addition to 2-methylpropene. Hammond’s postulate states that the energy of each transition state is approximated by the energy of the corresponding carbocation. The formation of tert-butyl bromide (right panel) is faster than the formation of isobutyl bromide (left panel) because it involves the more stable carbocation intermediate and therefore the transition state of lower energy.

state of lower energy than the transition state for protonation to give a primary carbocation. The stabilities of the carbocations themselves do not determine which reaction is faster; the relative free energies of the transition states for carbocation formation determine the relative rates of the two processes. Only the validity of Hammond’s postulate allows us to make the connection between carbocation energy and transition-state energy. We need Hammond’s postulate because the structures of transition states are uncertain, whereas the structures of reactants, products, and reactive intermediates are known. There- fore, knowing that a transition state resembles a particular species (for example, a carbocation) helps us to make a good guess about the transition-state structure. In this text, we’ll fre- quently analyze or predict reaction rates by considering the structures and stabilities of reactive intermediates such as carbocations. When we do this, we are assuming that the transition states and the corresponding reactive intermediates have similar structures and energies; in other words, we are invoking Hammond’s postulate.

PROBLEM 4.31 Apply Hammond’s postulate to decide which reaction is faster: addition of HBr to 2-methylpropene or addition of HBr to trans-2-butene. Assume that the energy difference between the starting alkenes can be ignored. Why is this assumption necessary?

4.9 CATALYSIS

Some reactions take place much more rapidly in the presence of certain substances that are themselves left unchanged by the reaction. A substance that increases the rate of a reaction without being consumed is called a catalyst. A practical example of a catalyst is platinum in the catalytic converter on the modern automobile. The platinum catalyst in the converter 04_BRCLoudon_pgs4-6.qxd 11/26/08 8:39 AM Page 167

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‡ no catalyst

a catalyst increases the rate (decreases ∆G°‡)

∆G°‡ ‡

∆G°‡ + catalyst reactants STANDARD FREE ENERGY STANDARD a catalyst does not ∆G° change ∆G°

products

reaction coordinate

Figure 4.16 A reaction free-energy diagram comparing a hypothetical catalyzed reaction (red curve) to the un- catalyzed reaction (blue curve).

brings about the rapid oxidation (combustion) of hydrocarbon exhaust emissions. This reaction would not occur were it not for the catalyst; yet the catalyst is left unchanged by the combustion reaction. The catalyst increases the rate of the combustion reaction by many orders of magnitude. Here are some important points about catalysts. 1. A catalyst increases the reaction rate. This means that it lowers the standard free energy of activation for a reaction (Fig. 4.16). 2. A catalyst is not consumed. It may be consumed in one step of a catalyzed reaction, but if so, it is regenerated in a subsequent step. An implication of points 1 and 2 is that a catalyst that strongly accelerates a reaction can be used in very small amounts. Many expensive catalysts are practical for this reason. 3. A catalyst does not affect the energies of reactants and products. In other words, a catalyst does not affect the DG of a reaction and consequently also does not affect the equilibrium constant (Fig. 4.16).° 4. A catalyst accelerates both the forward and reverse of a reaction by the same factor. The last point follows from the fact that, at equilibrium, the rates of a reaction and its reverse are equal. If a catalyst does not affect the equilibrium constant (point 3) but increases the reaction rate in one direction, equality of rates at equilibrium requires that the rate of the reverse reaction must be increased by the same factor. When a catalyst and the reactants exist in separate phases, the catalyst is called a heterogeneous catalyst. The catalyst in the catalytic converter of an automobile is a heterogeneous catalyst because it is a solid and the reactants are gases. In other cases, a reaction in solution may be catalyzed by a soluble catalyst. A catalyst that is soluble in a reaction solution is called a homogeneous catalyst. 04_BRCLoudon_pgs4-6.qxd 11/26/08 8:39 AM Page 168

168 CHAPTER 4 • INTRODUCTION TO ALKENES. STRUCTURE AND REACTIVITY

A large number of organic reactions are catalyzed. In this section, we’ll introduce the idea of catalysis by considering three examples of catalyzed alkene reactions. The ﬁrst example, catalytic hydrogenation, is a very important example of heterogeneous catalysis. The second example, hydration, is an example of homogeneous catalysis. The last example involves catalysis of a biological reaction by an enzyme.

Catalyst Poisons Although in theory catalysts should function indeﬁnitely,in practice many catalysts,particularly heterogeneous catalysts, slowly become less effective. It is as if they “wear out.”One reason for this be- havior is that they slowly absorb impurities,called catalyst poisons,from the surroundings;these impurities impede the functioning of the catalyst. An example of this phenomenon also occurs with the catalytic converter.Lead is a potent poison of the catalyst in a catalytic converter.This fact,as well as the desire to eliminate atmospheric lead pollution, are the major reasons why leaded gasoline is no longer used in automotive engines in the United States.

A. Catalytic Hydrogenation of Alkenes When a solution of an alkene is stirred under an atmosphere of hydrogen, nothing happens. But if the same solution is stirred under hydrogen in the presence of a metal catalyst, the hydrogen is rapidly absorbed by the solution. The hydrogen is consumed because it undergoes an addition to the alkene double bond. H H Pt/C H M H ) (4.38) 2 %H y H + `M % $H cyclohexene cyclohexane

Pt/C CH3(CH2)5CH A CH2 H2 CH3(CH2)5CH2 CH3 (4.39) + L 1-octene octane These reactions are examples of catalytic hydrogenation, an addition of hydrogen to an alkene in the presence of a catalyst. Catalytic hydrogenation is one of the best ways to convert alkenes into alkanes. Catalytic hydrogenation is an important reaction in both industry and the laboratory. The inconvenience of using a special apparatus for the handling of a ﬂammable gas (dihydrogen) is more than offset by the great utility of the reaction. In the preceding reactions, the catalyst is written over the reaction arrows. Pt/C is read as “Platinum supported on carbon” or simply “Platinum on carbon.” This catalyst is a ﬁnely di- vided platinum metal that has been precipitated, or “supported,” on activated charcoal. A number of noble metals, such as platinum, palladium, and nickel, are useful as hydrogenation catalysts, and they are often used in conjunction with solid support materials such as alumina

(Al2O3), barium sulfate (BaSO4), or, as in the previous examples, activated carbon. Hydro- genation can be carried out at room temperature and pressure or, for especially difﬁcult cases, at higher temperature and pressure in a “bomb” (a closed vessel designed to withstand high pressures). Because hydrogenation catalysts are insoluble in the reaction solution, they are examples of heterogeneous catalysts. (Soluble hydrogenation catalysts are also known and, although important, are not so widely used; Sec. 18.6D.) Even though they involve relatively expensive noble metals, heterogeneous hydrogenation catalysts are very practical because they can 04_BRCLoudon_pgs4-6.qxd 11/26/08 8:39 AM Page 169

4.9 CATALYSIS 169

be ﬁltered off and reused. Furthermore, because they are exceedingly effective, they can be used in very small amounts. For example, typical catalytic hydrogenation reactions can be run with reactant-to-catalyst ratios of 100 or more. How do hydrogenation catalysts work? Research has shown that both the hydrogen and the alkene must be adsorbed on the surface of the catalyst for a reaction to occur. The catalyst is believed to form reactive metal–carbon and metal–hydrogen bonds that ultimately are broken to form the products and to regenerate the catalyst sites. Beyond this, the chemical details of catalytic hydrogenation are poorly understood. This is not a reaction for which a simple curved-arrow mechanism can be written. The mechanism of noble-metal catalysis is an active area of research in many branches of chemistry. The benzene ring is inert to conditions under which normal double bonds react readily:

Pt/C CH A CH2 H2 CH2 CH3 (4.40) cL + cL L styrene ethylbenzene

(Benzene rings can be hydrogenated, however, with certain catalysts under conditions of high temperature and pressure.) You will learn that many other alkene reactions do not affect the “double bonds” of a benzene ring. The relative inertness of benzene rings toward the conditions of alkene reactions was one of the great puzzles of organic chemistry that was ultimately explained by the theory of aromaticity, which is introduced in Chapter 15.

PROBLEMS 4.32 Give the product formed when each of the following alkenes reacts with a large excess of hydrogen in the presence of Pd/C. (a) 1-pentene (b) (E)-1,3-hexadiene

4.33 (a) Give the structures of ﬁve alkenes, each with the formula C6H12, that would give hexane as the product of catalytic hydrogenation.

(b) How many alkenes containing one double bond can react with H2 over a Pt/C catalyst to give methylcyclopentane? Give their structures. (Hint: See Study Problem 4.9, p. 153.)

B. Hydration of Alkenes The alkene double bond undergoes reversible addition of water in the presence of moderately

concentrated strong acids such as H2SO4, HClO4, and HNO3.

H3C 3 $ 1 M HNO3 $C A CH2 H OH H3CCC" H3 (4.41) + L LL H C (in excess; 3 solvent) "OH 2-methylpropene 2-methyl-2-propanol (tert-butyl alcohol)

The addition of the elements of water is in general called hydration. Hence, the addition of water to the alkene double bond is called alkene hydration. Hydration does not occur at a measurable rate in the absence of an acid, and the acid is not consumed in the reaction. Hence, alkene hydration is an acid-catalyzed reaction. Because the catalyzing acid is soluble in the reaction solution, it is a homogeneous catalyst. 04_BRCLoudon_pgs4-6.qxd 11/26/08 8:39 AM Page 170

170 CHAPTER 4 • INTRODUCTION TO ALKENES. STRUCTURE AND REACTIVITY

Notice that this reaction, like the addition of HBr, is regioselective. As in the addition of HBr, the hydrogen adds to the carbon of the double bond with the smaller number of alkyl sub- stituents. The more electronegative partner of the H OH bond, the OH group, like the Br in HBr addition, adds to the carbon of the double bondL with the greater number of alkyl sub- stituents. In this reaction, the manner in which the catalyst functions can be understood by considering the mechanism of the reaction, which is very similar to that of HBr addition. In the ﬁrst step of the reaction, which is the rate-limiting step, the double bond is protonated so as to give the more stable carbocation. Because water is present, the actual acid is the hydrated proton

(H3O|).

H3C H3C 1

$ $

$C A CH2 $C| CH3 H2O1 (4.42a) L + H3C H3C H OH2 L 1 | This is a Brønsted acid–base reaction. Because this is the rate-limiting step, the rate of the hy-

dration reaction increases when the rate of this step increases. The strong acid H3O| is more effective than the considerably weaker acid water in protonating a weak base (the alkene). If a strong acid is not present, the reaction does not occur because water alone is too weak an acid to protonate the alkene. In the next step of the hydration reaction, the Lewis base water combines with the carbocation in a Lewis acid–base association reaction:

| (CH3)3C| OH2 (CH3)3C OH2 (4.42b) 3 2 L 2 Finally, a proton is lost to solvent in another Brønsted acid–base reaction to give the alcohol

product and regenerate the catalyzing acid H3O|: ..

| .. .. 1 (CH3)3C OH (CH3)3C OH H3O| (4.42c) L L + "H H2O1 3 Notice three things about this mechanism. First, it consists entirely of Lewis acid–base and Brønsted acid–base reactions. Second, although the proton consumed in Eq. 4.42a is not the same as the one produced in Eq. 4.42c, there is no net consumption of protons. Finally, the base is Eq. 4.42c is water. Some students are tempted to use hydroxide ion in a situation like this because it is a stronger base. However, there is essentially no hydroxide in a 1 M nitric acid or sulfuric acid solution. Nor is hydroxide needed, because the acid on the left of Eq. + 4.42c is a strong acid. Whenever H3O acts as an acid, its conjugate base H2O acts as the base. (Read again about amphoteric compounds on p. 97 if this isn’t clear.) More generally, acids and their conjugate bases always act in tandem in acid–base catalysis. Because the hydration reaction involves carbocation intermediates, some alkenes give re- arranged hydration products. H OH

H3O H3C" C CHA CH2 H2O | H3C" C CH2 CH3 (4.43) LL + LL L "CH3 "CH3 04_BRCLoudon_pgs4-6.qxd 11/26/08 8:39 AM Page 171

4.9 CATALYSIS 171

PROBLEMS 4.34 Give the mechanism for the reaction in Eq. 4.43. Show each step of the mechanism separately with careful use of the curved-arrow notation. Explain why the rearrangement takes place. 4.35 The alkene 3,3-dimethyl-1-butene undergoes acid-catalyzed hydration with rearrangement. Use the mechanism of hydration and rearrangement to predict the structure of the hydration product of this alkene. 4.36 (a) Unlike the alcohol product Eq. 4.41, the product in Eq. 4.43 does not come to equilibrium with the starting alkene. However, it does come to equilibrium with two other alkenes. What are their structures? (b) Why isn’t the alkene starting material in Eq. 4.43 part of the equilibrium mixture?

The equilibrium constants for many alkene hydrations are close enough to unity that the hydration reaction can be run in reverse. The reverse of alkene hydration is called alcohol dehydration. The direction in which the reaction is run depends on the application of Le Châte- lier’s principle, which states that if an equilibrium is disturbed, it will react so as to offset the disturbance. For example, if the alkene is a gas (as in Eq. 4.41), the reaction vessel can be pres- surized with the alkene. The equilibrium reacts to the excess of alkene by forming more alcohol. Neutralization of the acid catalyst stops the reaction and permits isolation of the alcohol. This strategy is used particularly in industrial applications. One such application of alkene hydration is the commercial preparation of ethyl alcohol (ethanol) from ethylene:

H3PO4 (absorbed on solid support) H CA CH H O H C CH OH (4.44) 2 2 2 300 °C 3 2 + L L ethylene ethanol A high temperature is required because the hydration of ethylene is very slow at ordinary tem- peratures (see Problem 4.37). Recall (Sec. 4.8B) that increasing the temperature accelerates a reaction. This reaction was at one time a major source of industrial ethanol. Although it is still used, its importance has decreased as the fermentation of sugars from biomass (for example, corn) has become more prevalent. To run the hydration reaction in the reverse (dehydration) direction, the alkene is removed as it is formed, typically by distillation. (Alkenes have signiﬁcantly lower boiling points than alcohols, as we’ll further discuss in Sec. 8.3B.) The equilibrium responds by forming more alkene. Alcohol dehydration is more widely used than alkene hydration in the laboratory. We’ll consider this reaction in Sec. 10.1. Alkene hydration and alcohol dehydration illustrate two important points. First is one of the key points about catalysis: a catalyst accelerates the forward and reverse reactions of an equilibrium by the same factor. For example, because alkene hydration is acid-catalyzed, alcohol dehydration is acid-catalyzed as well. A second point is that alkene hydration and alcohol dehydration occur by the forward and reverse of the same mechanism. Generally, if a reaction occurs by a certain mechanism, the reverse reaction under the same conditions occurs by the exact reverse of that mechanism. This statement is called the principle of microscopic reversibility. Microscopic reversibility requires, for example, that if you know the mechanism of alkene hydration, then you know the mechanism of alcohol dehydration as well. A conse- quence of microscopic reversibility is that the rate-limiting transition states of a reaction and its reverse are the same. For example, if the rate-limiting step of alkene hydration is protonation of the double bond to form the carbocation intermediate (Eq. 4.42a), then the rate-limiting step of alcohol dehydration is the reverse of the same equation—deprotonation of the carbocation to give the alkene. 04_BRCLoudon_pgs4-6.qxd 11/26/08 8:39 AM Page 172

172 CHAPTER 4 • INTRODUCTION TO ALKENES. STRUCTURE AND REACTIVITY

PROBLEMS 4.37 Explain why the hydration of ethylene is a very slow reaction. (Hint: Think about the structure of the reactive intermediate and apply Hammond’s postulate.) 4.38 Isopropyl alcohol is produced commercially by the hydration of propene. Show the mecha- nistic steps of this process. If you do not know the structure of isopropyl alcohol, try to de- duce it by analogy from the structure of propene and the mechanism of alkene hydration.

C. Enzyme Catalysis Catalysis is not limited to the laboratory or chemical industry. The biological processes of nature involve thousands of chemical reactions, most of which have their own unique naturally occurring catalysts. These biological catalysts are called enzymes. (The structures of enzymes are discussed in Sec. 26.10.) Under physiological conditions, most important biological reactions would be too slow to be useful in the absence of their enzyme catalysts. Enzyme catalysts are important not only in nature; they are ﬁnding increasing use both in industry and in the laboratory. Many of the best characterized enzymes are soluble in aqueous solution and hence are homogeneous catalysts. However, other enzymes are immobilized within biological substruc- tures such as membranes and can be viewed as heterogeneous catalysts. An example of an important enzyme-catalyzed addition to an alkene is the hydration of fumarate ion to malate ion.

O $S O OH O

H C O_ fumarase $ L (an enzyme) S S A (4.45) $C C H2O _O C "CH CH2 C O_ + L L L LL C malate _O $H LS O fumarate This reaction is catalyzed by the enzyme fumarase. It is one reaction in the Krebs cycle, or cit- ric acid cycle, a series of reactions that plays a central role in the generation of energy in biological systems. Fumarase catalyzes only this reaction. The effectiveness of fumarase catalysis can be appreciated by the following comparison: At physiological pH and temperature (pH 7, 37 C), the enzyme-catalyzed reaction is about 1010 times faster than the same reaction =in the absence° of enzyme.

KEY IDEAS IN CHAPTER 4

I Alkenes are compounds containing carbon–carbon the s electrons and can be donated to Brønsted or double bonds. Alkene carbon atoms, as well as other Lewis acids. trigonal planar atoms, are sp2-hybridized. I In the IUPAC substitutive nomenclature of alkenes, I The carbon–carbon double bond consists of a s bond the principal chain, which is the carbon chain con- and a p bond.The p electrons are more reactive than taining the greatest number of double bonds, is