CHAPTER OUTLINE
8.2 Formation, Stability, and Molecular Orbital View of Radicals
8.4 Reactions of Molecular Oxygen
8.5 Iron-Mediated Radical Reactions in Drug Metabolism
Box 8.1—Fenton chemistry in the action of antimalarial drugs
8.7 Case Study—Calicheamicin g1
8.1 Introduction
In this chapter we will examine the homolysis (homolytic cleavage) of σ bonds to form highly reactive radical species. When a bond breaks homolytically, the two electrons of the breaking bond end up on different atoms. The resulting radical species possess a single unpaired electron on an atom lacking a full octet of electrons. This makes radicals very electron deficient and unstable. They are often formed in low concentrations and are rarely stable enough to isolate, though they can serve as intermediates in chemical processes, as we will see. Biological systems take advantage of the high reactivity and transient nature of radicals to mediate a host of transformations required for life. Molecular oxygen exists as a diradical and we will see how it acts as a powerful oxidant, attacking organic molecules to initiate radical reactions. These oxygen-mediated radical reactions cause cellular damage and we will take a close look at how the antioxidant vitamins E and C (Figure 8.1) prevent this damage by acting as radical scavengers.
Figure 8.1 The antioxidant vitamins E and C.
8.2 Formation, Stability, and Molecular Orbital View of Radicals
The homolysis of a covalent bond to form two radicals is illustrated using two single-headed (“fishhook”) arrows, each of which indicates the movement of a single electron (Figure 8.2). Note that fishhook arrows are reserved for keeping track of electron count in radical reactions and are not interchangeable with the standard arrows used to indicate the movement of pairs of electrons in acid/base or nucleophile/electrophile chemistry. The homolysis of molecular hydrogen (H2) yields two free atoms of hydrogen. The process requires an amount of heat (104 kcal/mol) that is equal to the amount of heat produced when two free atoms of hydrogen combine to form a covalent bond. This is referred to as the bond dissociation energy. We will use bond dissociation energies to help understand the strength of bonds and the relative reactivity of radicals.
Figure 8.2 Homolytic cleavage of a bond is illustrated using “fishhook” arrows, which indicate the movement of a single electron. The homolysis of molecular hydrogen (H2) yields two hydrogen atom radicals.
Homolysis of the C–H bond in methane yields a methyl radical in a process that requires 105 kcal/mol of energy (Figure 8.3). The homolysis of a C–H bond at a primary (101 kcal/mol), secondary (98.5 kcal/mol), or tertiary (96.5 kcal/mol) substituted carbon atom requires sequentially less energy. This trend in dissociation energies corresponds to the relative stability of the resulting primary, secondary, and tertiary radicals, with the tertiary radical being the most stable.
Figure 8.3 The formation of methyl radical and the relative stability of carbon radicals.
Let us examine this stability trend in more detail by comparing methyl and tert-butyl radicals (Figure 8.4). In forming either radical, homolysis of the C–H bond converts the carbon center from tetrahedral to trigonal planar geometry. The carbon radical is sp2 hybridized with the unpaired electron located in the unhybridized p orbital. Replacing H atoms on the methyl radical with additional carbon substituents increases stability of the radical (Figure 8.3). The increasing stability is related to both steric and conjugative effects. For example, the methyl groups in 2-methylpropane will feel steric crowding to a much greater extent than the hydrogen atoms in methane. This steric crowding is relieved as bond angles increase in the transition from tetrahedral to planar geometry. The second effect relates to stability of the unpaired electron in the p orbital of the radical, which is electron deficient and thus looking to add some electron density. In the tert-butyl radical, the sp3 orbital of a neighboring C–H bond can donate electron density into the p orbital of the radical in a process called hyperconjugation. With three adjacent methyl groups, the tert-butyl radical can benefit from three such interactions. Although the additional stability provided is rather small (roughly 2 kcal/mol per methyl group), it adds up and results in the tertiary radical being the most stable radical in the series.
Figure 8.4 The electronic structure of carbon radicals.
Electron-deficient carbon radicals can also be stabilized by resonance delocalization. This fact is evident in the lower bond dissociation energy for a C–H bond located adjacent to a π bond (Figure 8.5). In terms of the orbitals involved, the electron-deficient p orbital of the radial is delocalized into the more electron rich p orbitals of the π bond. Resonance delocalization of allylic and benzylic radicals can be illustrated using resonance structures, as shown in Figure 8.5. Note that we must use fishhook arrows to keep track of electrons when drawing resonance structures of radical species. The stability of the allylic radical is roughly equal to that of the benzylic radical in spite of the benzylic radical having access to more resonance structures. The loss of some aromatic character in the benzylic radical explains this.
Figure 8.5 Formation of ethyl, allyl, and benzylic radicals with the corresponding bond dissociation energies shown.
8.3 Radical Reactions
In this section we will describe the three stages of a radical reaction, which include initiation, propagation, and termination. The chemical behavior of radicals is dominated by their high reactivity and electron-deficient character. Radicals will often react with the closest atom available, and can involve reaction with a σ or π bond (Figure 8.6). One common radical reaction is the abstraction of a hydrogen atom from a C–H σ bond of a nearby molecule. This produces a new radical species and is one of the ways a radical reaction can be propagated. A second important reaction of radicals is addition to a π bond to form a new C–C bond and a new radical species. On rare occasions, two radicals will be in close enough proximity to combine and form a new σ bond. The product of this process is no longer a radical and is representative of a termination step in a radical reaction sequence.
Figure 8.6 Radical reactions involving (a) breaking a C–H σ bond, (b) addition to a π bond, and (c) reaction with another radical.
As we will see in the following sections, radical reactions are important in a variety of biological processes, including the metabolism and clearance of many drugs. Before describing these more complex processes, let us examine in detail a relatively simple radical reaction—the chlorination of methane with Cl2 (Figure 8.7). As with most radical reactions, radical chlorination of an alkane proceeds through distinct initiation, propagation, and termination steps. We examine each of these steps separately below.
Figure 8.7 General reaction scheme for the halogenation of methane.
The initiation step in the chlorination of methane involves homolysis of the Cl–Cl bond in molecular chlorine (Cl2). This bond is a weak one (bond dissociation energy of just −58 kcal/mol) and will be more prone to homolytic cleavage than the stronger C–H bonds of methane. Homolysis can be promoted with either heat or light to produce a small concentration of chlorine atom radicals (Figure 8.8). This initiation step thus generates the reactive chlorine radical species that allows the rest of the steps in the overall process to proceed. Only a small fraction of the total Cl2 present need be converted to chlorine radicals in the initiation step, for reasons that will become clear as we examine the propagation stage of the reaction.
Figure 8.8 Homolysis of a Cl–Cl bond to form two chlorine atoms. This reaction represents the initiation step in the chlorination of methane.
During the propagation stage of a radical reaction, the initial radical species produced during initiation reacts to form a new bond and in the process generate a new radical species. There are two propagation steps to consider in the chlorination of methane, the first being when a chlorine atom abstracts a hydrogen atom from methane (step 1, Figure 8.9). The resulting methyl radical then attacks a Cl2 molecule yielding the chloromethane product and a chlorine atom radical (step 2). Note that the second step produces the reaction product while also generating a chlorine radical that can feed back into propagation step 1. This perpetual formation of chlorine radicals during propagation explains why very little homolysis of chlorine (by heat or light) is needed to initiate the reaction. It also explains why radical reactions are often referred to as chain reactions.
Figure 8.9 The two propagation steps in a radical chlorination reaction.
Termination occurs when two radicals combine to form a new covalent bond. These are called termination steps because they do not produce a new radical species to carry on the chain reaction (propagation). There are several possible termination reactions in the chlorination of methane (Figure 8.10). Note that the termination reaction of a chlorine and methyl radical produces the same chloromethane product that is also formed in step 2 of the propagation stage. Recombination of two methyl radicals on the other hand will produce a different reaction product (ethane). Remember, however, that the concentration of radical species remains very low over the course of a radical reaction. Once formed, a methyl radical is much more likely to react with Cl2 to form the desired product (and propagate the reaction) than it is to encounter and react with another methyl radical. The low concentration of radial species present allows the chain reaction to propagate and restricts the number of termination reactions.
Figure 8.10 Radical reaction termination steps in the chlorination of methane.
8.4 Reactions of Molecular Oxygen
The most stable form of molecular oxygen (O2) has two unpaired electrons occupying two degenerate molecular orbitals. These electrons have the same spin and are unable to form a bond, thus making oxygen a diradical. This diradical form of oxygen is referred to as triplet oxygen. Higher in energy by roughly 20 kcal/mol is singlet oxygen, which is much more reactive than triplet oxygen. Since the air you breathe is triplet oxygen and this diradical is what drives the biochemical processes of aerobic systems, we will focus our further discussion on triplet oxygen (Figure 8.11).
Figure 8.11 Triplet and singlet oxygen.
The diradical nature of oxygen allows it to drive many oxidation reactions ranging from the spoilage of food to oxidative damage in cells. A close look at how oxygen oxidizes unsaturated lipids will help us understand the role oxygen plays in these oxidative processes. Linoleic acid is a polyunsaturated fatty acid used in the biosynthesis of many bioactive compounds (arachidonic acid, prostaglandins) and is found in the lipids of cell membranes. As we have seen earlier (Figure 8.5), an allylic radical is stabilized by resonance delocalization with the adjacent π bond. We might therefore predict that the C–H bonds lying between the two double bonds of linoleic acid will be most prone to homolytic cleavage. Indeed, in the initiation step of the oxidation process, oxygen abstracts a hydrogen atom from the doubly allylic methylene group in linoleic acid to form an allylic radical that is stabilized by two neighboring double bonds (Figure 8.12).
Figure 8.12 Oxidation of linoleic acid by oxygen is a radical-mediated process.
In a propagation step, this resonance-stabilized radical combines with another molecule of oxygen to form a peroxy radical. The peroxy radical abstracts a hydrogen atom from another molecule of linoleic acid to propagate the chain reaction and produce linoleic acid hydroperoxide. The weak O–O bond in the hydroperoxide species can cleave homolytically to yield an alkoxy radical (Figure 8.12). The alkoxy radical further decomposes to form an unsaturated aldehyde and other decomposition products that can wreak havoc on the integrity of a cell.
The reason that these products of lipid oxidation are toxic to cells is that they are reactive electrophiles. We have seen in Chapters 6 and 7 that the side chains of certain amino acids such as cysteine (protein—SH) and lysine (protein—NH2) are potentially nucleophilic. Thus, the electrophilic products of lipid oxidation can covalently modify the nucleophilic side chains of proteins in nonspecific and detrimental ways. The biological function of these proteins becomes compromised and this contributes to the development of heart disease, cancer, emphysema, and many other chronic disease states.
You will note that the mechanism of oxygen-mediated lipid oxidation involves the three steps of initiation, propagation, and termination we outlined earlier for the chlorination of methane. Thus, a large amount of cellular damage can be initiated with a tiny amount of oxygen attacking a lipid molecule. Wouldn’t it be nice if your body had access to some “terminator” molecules that could help stop this damaging process before it had a chance to get going? In fact it does—our bodies use vitamins E and C in concert to help minimize the harmful effects of radical intermediates. Let us take a look at the structures of these vitamins to understand how they work.
Looking at the structure of vitamin E you might note that the long hydrocarbon chain is similar to that in linoleic acid. It is not surprising then that, like linoleic acid, large amounts of vitamin E are found in the lipid membrane. We also expect that the phenolic (OH) function of vitamin E should be quite acidic since the negative charge of the corresponding phenoxide anion will be delocalized into the aromatic ring. It is in fact the phenoxide form of vitamin E that is able to donate an electron to lipid radicals, reducing them to an anionic state (which is then rapidly protonated) and thereby interrupting the lipid degradation process that otherwise leads to reactive electrophilic species. Of course, in the process of donating one electron, vitamin E itself is converted to a radical. However, unlike the lipid radical, the vitamin E radical is much more stable and less reactive, due to the many resonance forms available to it (Figure 8.13). This radical form of vitamin E persists until it encounters a water-soluble reducing agent such as vitamin C at the surface of the cell membrane.
Figure 8.13 The reaction of vitamin E with linoleic acid alkoxy radical. The anionic form of vitamin E acts as a reducing agent, converting the alkoxy radical to a relatively inert alcohol.
As its other common name ascorbic acid suggests, vitamin C is an acid that exists significantly in an anionic form at physiological pH. This anion is able to donate an electron to the oxidized vitamin E radical, thus producing a vitamin C radical and regenerating vitamin E in its neutral/anionic form. Vitamin C is a strong antioxidant because several resonance forms stabilize the radical (Figure 8.14). Once formed, the vitamin C radical fractures into several smaller water-soluble compounds that are quickly excreted by your body. In this way, vitamins E and C work together as radical scavengers to rid the cell of toxic radical intermediates. Note that the essential structural feature of these vitamins is an acidic function (phenol or phenol-like O–H) that can donate an electron and then be stabilized in a radical form by resonance delocalization.
Figure 8.14 Formation of vitamin C radical by reduction of the vitamin E radical. Together vitamins E and C act as radical scavengers to protect the cell from potentially toxic radical species.
8.5 Iron-Mediated Radical Reactions in Drug Metabolism
Iron can be used to initiate a variety of radical reactions, both in the test tube and in biological systems. The reaction of hydrogen peroxide with ferrous sulfate is known as the Fenton reaction. In this reaction ferrous iron (2+ oxidation state) is oxidized by hydrogen peroxide to ferric iron (3+), with the consequent production of hydroxyl radical and hydroxide anion (Figure 8.15). Ferric iron (3+) is in turn reduced back to ferrous iron (2+) in reaction with hydrogen peroxide to form a hydroperoxyl radical and a proton. The overall process leads to the disproportionation of two equivalents of hydrogen peroxide into two highly reactive radical species (hydroxyl and hydroperoxyl radicals) and water. Fenton chemistry is also implicated in the action of antimalarial drugs such as artemisinin and arterolane (Box 8.1).
Figure 8.15 The Fenton reaction of ferrous iron with hydrogen peroxide.
The ferrous (2+) and ferric (3+) forms of iron involved in the Fenton reaction are also the two major forms of iron that exist under physiological conditions. The ability of iron to undergo one-electron reduction or oxidation is central to the useful chemistry performed by biological macromolecules that employ iron. However, the potential of iron to generate oxygen radical species also means that the transport and storage of iron is highly regulated in biology. One large family of enzymes that exploit iron chemistry is the cytochrome P450 enzymes (CYP enzymes, for short). The CYP superfamily includes mitochondrial enzymes involved in cellular respiration, important biosynthetic enzymes such as steroid hydroxylases, and not least, the microsomal CYP enzymes involved in drug metabolism.
Microsomal CYPs are abundant in the liver, where they function to oxidize organic xenobiotics (including many drugs), leading to their elimination from the body. Two of the most common reactions performed by microsomal CYPs are the oxidation of a C–H to a hydroxyl (C–OH) function and the epoxidation of an alkene or aromatic ring (Figure 8.16). Hydroxylation products of CYPs can be further converted to glucuronides (as introduced in Chapter 7), highly water-soluble conjugates that are rapidly eliminated from the body. Epoxidation products of CYPs can be converted to diols by epoxide hydrolase or can react as electrophiles in reaction with glutathione. In either case the result is a more hydrophilic metabolite that is more readily removed from the body.
Figure 8.16 Examples of hydroxylation and epoxidation reactions carried out by iron-dependent CYP enzymes. These “phase 1” metabolic processes are often followed by phase 2 metabolism involving conjugation to hydrophilic groups such as glucuronic acid or glutathione.
