《生物化学原理》（英文版）BIOENERGETICS AND METABOLISM

506 Chapter 13 Principles of Bioenergetics Inorganic Polyphosphate Is a Potential fers a survival advantage. Deletion of the genes for Phosphoryl group donor polyphosphate kinases diminishes the ability of certain pathogenic bacteria to invade animal tissues. The en- composed of many tens or hundreds of P, residues ymes may therefore prove to be vulnerable targets in Inorganic polyphosphate (polyP)is a linear polymer the development of new antimicrobial drug linked through phosphoanhydride bonds. This polymer, No gene in yeast encodes a PPK-like protein, but present in all organisms, may accumulate to high levels four genes-unrelated to bacterial PPK in some cells In yeast, for example, the amount of polyP essary for the synthesis of polyphosphate. The mecha- tributed uniformly throughout the cell, a concentration to be quite different from that in prokaryotes. of 200 mM( Compare this with the concentrations of other phosphoryl donors listed in Table 13-5.) Biochemical and Chemical Equations Are Not ldentical P-0-P-0-P-0 Biochemists write metabolic equations in a simplified way, and this is particularly evident for reactions in- volving ATP. Phosphorylated compounds can exist in Inorganic polyphosphate(polyP) several ionization states and, as we have noted, the dif- One potential role for polyP is to serve as a phos- 2 mM Mg t ATP exists in the forms ATP. HAand ferent species can bind Mg2+. For example, at pH phagen, a reservoir of phosphoryl groups that can be H2ATP2-, MgHATP, and Mg2ATP In thinking about the used to generate ATP, as creatine phosphate is used in biological role of ATP, however, we are not always in- muscle. PolyP has about the same phosphoryl group terested in all this detail, and so we consider ATP as an transfer potential as PP; The shortest polyphosphate. PP(n=2), can serve as the energy source for active entity made up of a sum of species, and we write its hy- transport of H in plant vacuoles. For at least one form drolysis as the biochemical equation of the enzyme phosphofructokinase in plants, PPi is the ATP+H20→ADP+P1 phosphoryl group donor, a role played by AtP in ani- where ATP, ADP, and Pi are sums of species mals and microbes(p. XXX). The finding of high con- corresponding apparent equilibrium constant. centrations of polyP in volcanic condensates and steam JADPJIP VATPL depends on the pH and the concentra vents suggests that it could have served as an energy tion of free Mg2+. Note that H* and Mg t do not ap In prokaryotes, the enzyme polyphosphate ki. pear in the biochemical equation because they are held source in prebiotic and early cellular evolution. nase.l(PPK-1) catalyzes the reversible reaction constant. Thus a biochemical equation does not balance H, Mg, or charge, although it does balance all other el- ements involved in the reaction(C, N, 0, and P in the ATP polyN ADP polyP+ equation above We can write a chemical equation that does balance histidine intermediate (recall the mechanism of nucle. is hydrolyzed at a pH above 8. 5 in the absence oflo oside diphosphate kinase, described above). A second the chemical reaction is represented by enzyme, polyphosphate kinase. 2(PPK-2), catalyzes ATP-+H,- ADP-+HPOz-+H+ the reversible synthesis of GTP (or ATP) from poly phosphate and GDP (or ADP The corresp ADPI[HPOA I(H*MATPI depends only on tem- perature, pressure, and ionic strength PPK-2 is believed to act primarily in the direction of in biochemistry. Chemical equations are needed when GTP and ATP synthesis, and PPK-l in the direction of we want to account for all atoms and charges in a re- polyphosphate synthesis. PPK-I and PPK-2 are present action, as when we are considering the mechanism of a in a wide variety of prokaryotes, including many patho- chemical reaction. Biochemical equations are used to genic bacteria. determine in which direction a reaction will proceed In prokaryotes, elevated levels of polyP have been spontaneously, given a specified pH and [Mg*or to shown to promote expression of a number of genes in- calculate the equilibrium constant of such a reaction. volved in adaptation of the organism to conditions of Throughout this book we use biochemical equa- starvation or other threats to survival. In Escherichia tions, unless the focus is on chemical mechanism, and oli, for example, polyP accumulates when cells are we use values of AG and Kea as determined at pH 7 starved for amino acids or P, and this accumulation con- and 1 mM mg2

Inorganic Polyphosphate Is a Potential Phosphoryl Group Donor Inorganic polyphosphate (polyP) is a linear polymer composed of many tens or hundreds of Pi residues linked through phosphoanhydride bonds. This polymer, present in all organisms, may accumulate to high levels in some cells. In yeast, for example, the amount of polyP that accumulates in the vacuoles would represent, if dis￾tributed uniformly throughout the cell, a concentration of 200 mM! (Compare this with the concentrations of other phosphoryl donors listed in Table 13–5.) One potential role for polyP is to serve as a phos￾phagen, a reservoir of phosphoryl groups that can be used to generate ATP, as creatine phosphate is used in muscle. PolyP has about the same phosphoryl group transfer potential as PPi . The shortest polyphosphate, PPi (n  2), can serve as the energy source for active transport of H in plant vacuoles. For at least one form of the enzyme phosphofructokinase in plants, PPi is the phosphoryl group donor, a role played by ATP in ani￾mals and microbes (p. XXX). The finding of high con￾centrations of polyP in volcanic condensates and steam vents suggests that it could have served as an energy source in prebiotic and early cellular evolution. In prokaryotes, the enzyme polyphosphate ki￾nase-1 (PPK-1) catalyzes the reversible reaction ATP polyPn ADP polyPn1 DG  20 kJ/mol by a mechanism involving an enzyme-bound phospho￾histidine intermediate (recall the mechanism of nucle￾oside diphosphate kinase, described above). A second enzyme, polyphosphate kinase-2 (PPK-2), catalyzes the reversible synthesis of GTP (or ATP) from poly￾phosphate and GDP (or ADP): GDP polyPn1 GTP polyPn PPK-2 is believed to act primarily in the direction of GTP and ATP synthesis, and PPK-1 in the direction of polyphosphate synthesis. PPK-1 and PPK-2 are present in a wide variety of prokaryotes, including many patho￾genic bacteria. In prokaryotes, elevated levels of polyP have been shown to promote expression of a number of genes in￾volved in adaptation of the organism to conditions of starvation or other threats to survival. In Escherichia coli, for example, polyP accumulates when cells are starved for amino acids or Pi , and this accumulation con￾Mn2 3:::4 Mg2 3:::4 O P O O O P O O O P O O O P O O O P O O O Inorganic polyphosphate (polyP) fers a survival advantage. Deletion of the genes for polyphosphate kinases diminishes the ability of certain pathogenic bacteria to invade animal tissues. The en￾zymes may therefore prove to be vulnerable targets in the development of new antimicrobial drugs. No gene in yeast encodes a PPK-like protein, but four genes—unrelated to bacterial PPK genes—are nec￾essary for the synthesis of polyphosphate. The mecha￾nism for polyphosphate synthesis in eukaryotes seems to be quite different from that in prokaryotes. Biochemical and Chemical Equations Are Not Identical Biochemists write metabolic equations in a simplified way, and this is particularly evident for reactions in￾volving ATP. Phosphorylated compounds can exist in several ionization states and, as we have noted, the dif￾ferent species can bind Mg2. For example, at pH 7 and 2 mM Mg2, ATP exists in the forms ATP4, HATP3, H2ATP2, MgHATP, and Mg2ATP. In thinking about the biological role of ATP, however, we are not always in￾terested in all this detail, and so we consider ATP as an entity made up of a sum of species, and we write its hy￾drolysis as the biochemical equation ATP H2O 8n ADP Pi where ATP, ADP, and Pi are sums of species. The corresponding apparent equilibrium constant, K eq  [ADP][Pi ]/[ATP], depends on the pH and the concentra￾tion of free Mg2. Note that H and Mg2 do not ap￾pear in the biochemical equation because they are held constant. Thus a biochemical equation does not balance H, Mg, or charge, although it does balance all other el￾ements involved in the reaction (C, N, O, and P in the equation above). We can write a chemical equation that does balance for all elements and for charge. For example, when ATP is hydrolyzed at a pH above 8.5 in the absence of Mg2, the chemical reaction is represented by ATP4 H2O 8n ADP3 HPO4 2 H The corresponding equilibrium constant, K eq  [ADP3][HPO4 2][H]/[ATP4], depends only on tem￾perature, pressure, and ionic strength. Both ways of writing a metabolic reaction have value in biochemistry. Chemical equations are needed when we want to account for all atoms and charges in a re￾action, as when we are considering the mechanism of a chemical reaction. Biochemical equations are used to determine in which direction a reaction will proceed spontaneously, given a specified pH and [Mg2], or to calculate the equilibrium constant of such a reaction. Throughout this book we use biochemical equa￾tions, unless the focus is on chemical mechanism, and we use values of G and K eq as determined at pH 7 and 1 mM Mg2. 506 Chapter 13 Principles of Bioenergetics

13.3 Biological Oxidation-Reduction Reactions 507 SUMMARY 13.2 Phosphoryl Group Transfers The carriers in turn donate electrons to acceptors with and ATP higher electron affinities, with the release of Cells contain a variety of molecular energy transducers and anabolism. It is the energy currency of the wo convert the energy of electron flow into useful atP is the chemical link between catabolism which exergonic conversion of ATP to cussion w ith a description of the ADP and P, or to AMP and PPi, is coupled to general types of metabolic reactions in which electrons many endergonic reactions and processes are transferred. After considering the theoretical and a Direct hydrolysis of ATP is the source of experimental basis for measuring the energy changes in oxidation reactions in terms of electromotive force. w energy in the conformational changes that produce muscle contraction but, in general, it discuss the relationship between this force, expressed is not ATP hydrolysis but the transfer of a in volts, and the free-energy change, expressed in joules We conclude by describing the structures and oxidation phosphoryl, pyrophosphoryl, or adenylyl group reduction chemistry of the most common of the spe- from atP to a substrate or enzyme molecule cialized electron carriers, which you will encounter that couples the energy of ATP breakdown to repeatedly in later chapters endergonic transformations of substrates a Through these group transfer reactions, ATP The Flow of Electrons Can Do Biological Work provides the energy for anabolic reactions, including the synthesis of informational Every time we use a motor, an electric light or heater molecules, and for the transport of molecules or a spark to ignite gasoline in a car engine, we use the and ions across membranes against flow of electrons to accomplish work. In the circuit that concentration gradients and electrical potential powers a motor, the source of electrons can be a bat gradients tery containing two chemical species that differ in affin- ity for electrons. Electrical wires provide a pathway for I Cells contain other metabolites with large, negative, free energies of hydrolysis, including electron flow from the chemical species at one pole of the battery, through the motor, to the chemical species phosphoenolpyruvate, 1,3-bisphosphoglycerate, at the other pole of the battery. Because the two chem- and phosphocreatine. These high-energy ical species differ in their affinity for electrons, electrons ompounds, like ATP, have a high phosphoryl flow spontaneously through the circuit, driven by a force group transfer potential; they are good donors proportional to the difference in electron affinity, the of the phosphoryl group. Thioesters also have electromotive force (emf). The electromotive force high free energies of hydrolysis (typically a few volts) can accomplish work if an ap- a Inorganic polyphosphate, present in all cells propriate energy transducer-in this case a motor-is may serve as a reservoir of phosphoryl groups placed in the circuit. The motor can be coupled to a va with high group transfer potential riety of mechanical devices to accomplish useful work Living cells have an analogous biological"circuit, with a relatively reduced compound such as glucose as the source of electrons. As glucose is enzymatically ox 13.3 Biological Oxidation-Reduction idized, the released electrons flow spontaneously Reactions through a series of electron-carrier intermediates to an other chemical species, such as O2. This electron flow The transfer of phosphoryl groups is feature is exergonic, because O2 has a higher affinity for elec- of metabolism. Equally important is kind of trons than do the electron-carrier intermediates. the transfer, electron transfer in oxidation-reduction reac- resulting electromotive force provides energy to a vari- tions. These reactions involve the loss of electrons by ety of molecular energy transducers(enzymes and other one chemical species, which is thereby oxidized, and the proteins) that do biological work. In the mitochondrion, of electrons in oxidation-reduction reactions is respon- flow to the production of a transmembrane ph d s gain of electrons by another, which is reduced. The flow for example, membrane-bound enzymes couple electro sible, directly or indirectly, for all work done by living ence, accomplishing osmotic and electrical work. The organisms. In nonphotosynthetic organisms, the sources proton gradient thus formed has potential energy, some. of electrons are reduced compounds(foods); in photo- times called the proton-motive force by analogy with synthetic organisms, the initial electron donor is a chem. electromotive force. Another enzyme, ATP synthase in ical species excited by the absorption of light. The path the inner mitochondrial membrane, uses the proton of electron flow in metabolism is complex. Electrons motive force to do chemical work: synthesis of ATP from move from various metabolic intermediates to special- ADP and P; as protons flow spontaneously across the ized electron carriers in enzyme-catalyzed reactions. membrane. Similarly, membrane-localized enzymes in

SUMMARY 13.2 Phosphoryl Group Transfers and ATP ■ ATP is the chemical link between catabolism and anabolism. It is the energy currency of the living cell. The exergonic conversion of ATP to ADP and Pi , or to AMP and PPi , is coupled to many endergonic reactions and processes. ■ Direct hydrolysis of ATP is the source of energy in the conformational changes that produce muscle contraction but, in general, it is not ATP hydrolysis but the transfer of a phosphoryl, pyrophosphoryl, or adenylyl group from ATP to a substrate or enzyme molecule that couples the energy of ATP breakdown to endergonic transformations of substrates. ■ Through these group transfer reactions, ATP provides the energy for anabolic reactions, including the synthesis of informational molecules, and for the transport of molecules and ions across membranes against concentration gradients and electrical potential gradients. ■ Cells contain other metabolites with large, negative, free energies of hydrolysis, including phosphoenolpyruvate, 1,3-bisphosphoglycerate, and phosphocreatine. These high-energy compounds, like ATP, have a high phosphoryl group transfer potential; they are good donors of the phosphoryl group. Thioesters also have high free energies of hydrolysis. ■ Inorganic polyphosphate, present in all cells, may serve as a reservoir of phosphoryl groups with high group transfer potential. 13.3 Biological Oxidation-Reduction Reactions The transfer of phosphoryl groups is a central feature of metabolism. Equally important is another kind of transfer, electron transfer in oxidation-reduction reac￾tions. These reactions involve the loss of electrons by one chemical species, which is thereby oxidized, and the gain of electrons by another, which is reduced. The flow of electrons in oxidation-reduction reactions is respon￾sible, directly or indirectly, for all work done by living organisms. In nonphotosynthetic organisms, the sources of electrons are reduced compounds (foods); in photo￾synthetic organisms, the initial electron donor is a chem￾ical species excited by the absorption of light. The path of electron flow in metabolism is complex. Electrons move from various metabolic intermediates to special￾ized electron carriers in enzyme-catalyzed reactions. The carriers in turn donate electrons to acceptors with higher electron affinities, with the release of energy. Cells contain a variety of molecular energy transducers, which convert the energy of electron flow into useful work. We begin our discussion with a description of the general types of metabolic reactions in which electrons are transferred. After considering the theoretical and experimental basis for measuring the energy changes in oxidation reactions in terms of electromotive force, we discuss the relationship between this force, expressed in volts, and the free-energy change, expressed in joules. We conclude by describing the structures and oxidation￾reduction chemistry of the most common of the spe￾cialized electron carriers, which you will encounter repeatedly in later chapters. The Flow of Electrons Can Do Biological Work Every time we use a motor, an electric light or heater, or a spark to ignite gasoline in a car engine, we use the flow of electrons to accomplish work. In the circuit that powers a motor, the source of electrons can be a bat￾tery containing two chemical species that differ in affin￾ity for electrons. Electrical wires provide a pathway for electron flow from the chemical species at one pole of the battery, through the motor, to the chemical species at the other pole of the battery. Because the two chem￾ical species differ in their affinity for electrons, electrons flow spontaneously through the circuit, driven by a force proportional to the difference in electron affinity, the electromotive force (emf). The electromotive force (typically a few volts) can accomplish work if an ap￾propriate energy transducer—in this case a motor—is placed in the circuit. The motor can be coupled to a va￾riety of mechanical devices to accomplish useful work. Living cells have an analogous biological “circuit,” with a relatively reduced compound such as glucose as the source of electrons. As glucose is enzymatically ox￾idized, the released electrons flow spontaneously through a series of electron-carrier intermediates to an￾other chemical species, such as O2. This electron flow is exergonic, because O2 has a higher affinity for elec￾trons than do the electron-carrier intermediates. The resulting electromotive force provides energy to a vari￾ety of molecular energy transducers (enzymes and other proteins) that do biological work. In the mitochondrion, for example, membrane-bound enzymes couple electron flow to the production of a transmembrane pH differ￾ence, accomplishing osmotic and electrical work. The proton gradient thus formed has potential energy, some￾times called the proton-motive force by analogy with electromotive force. Another enzyme, ATP synthase in the inner mitochondrial membrane, uses the proton￾motive force to do chemical work: synthesis of ATP from ADP and Pi as protons flow spontaneously across the membrane. Similarly, membrane-localized enzymes in 13.3 Biological Oxidation-Reduction Reactions 507

E. coli convert electromotive force to proton-motive Biological Oxidations Often Involve Dehydrogenati fo The principles of electrochemistry that govem en- The carbon in living cells exists in a range of oxidation ce,which is then used to power flagellar motion ergy changes in the macroscopic circuit with a motor states(Fig. 13-13). When a carbon atom shares an elec- and battery apply with equal validity to the molecular tron pair with another atom(typically H, C, S, N, or O) processes accompanying electron flow in living cells. We the sharing is unequal in favor of the more electroneg turn now to a discussion of those principles H<C<S<N<O In oversimplified but useful terms, Oxidation-Reductions Can Be Described the more electronegative atom"owns the bonding elec- as half-Reactions it shares with another atom. For example, in methane(CHA), carbon is more electronegative than th Although oxidation and reduction must occur together, four hydrogens bonded to it, and the C atom therefore it is convenient when describing electron transfers to "owns"all eight bonding electrons(Fig. 13-13). In consider the two halves of an oxidation-reduction reac- ethane the electrons in the C-c bond are shared tion separately. For example, the oxidation of ferrous equally, so each C atom owns only seven of its eight ion by cupric ion. bonding electrons. In ethanol, C-1 is less electronega Fe2++ Cu2+=Fe++Cut tive than the oxygen to which it is bonded, and the O atom therefore "owns" both electrons of the c-o bond, can be described in terms of two half-reactions leaving C-l with only five bonding electrons. With each formal loss of electrons, the carbon atom has undergone oxidation-even when no oxygen is involved, as in the conversion of an alkane(CH2-CH2)to an alkene The electron-donating molecule in an oxidation- (CH=CH-) In this case, oxidation (oss of elec reduction reaction is called the reducing agent or reduc- trons)is coincident with the loss of hydrogen. In bio- tant;the electron-accepting molecule is the oxidizing logical systems, oxidation is often synonymous with de agent or oxidant. a given agent, such as an iron cation hydrogenation, and many enzymes that catalyze existing in the ferrous(Fe2+)or ferric(Fe+)state, func- oxidation reactions are dehydrogenases. Notice that ions as a conjugate reductant-oxidant pair (redox pair), the more reduced compounds in Figure 13-13(top)are just as an acid and corresponding base function as a con- richer in hydrogen than in oxygen, whereas the more jugate acid-base pair Recall from Chapter 2 that in acid- oxidized compounds(bottom) have more oxygen and base reactions we can write a general equation: proton less hydrog donor= H+ proton acceptor. In redox reactions we Not all biological oxidation-reduction reactions in- can write a similar general equation: electron donor= volve carbon. For example, in the conversion of molec- e+electron acceptor In the reversible half-reaction(1) ular nitrogen to ammonia, 6HT 6e N2+2NH3, above, Fe?+ is the electron donor and Fe+ is the elec- the nitrogen atoms are reduced. tron acceptor; together, Fe and Fe constitute a con Electrons are transferred from one molecule (elec- jugate redox pair. tron donor) to another (electron acceptor) in one of The electron transfers in the oxi four different ways reactions of organic compounds are not fundamentally different from those of inorganic species. In Chapter 7 1. Directly as electrons. For example, the Fe2+/Fe3+ we considered the oxidation of a reducing sugar (an redox pair can transfer an electron to the aldehyde or ketone)by cupric ion(see Fig. 7-10a) Cut/Cu redox pair Recall that a hydroge consists of a proton(H)and a single electron(e) This overall reaction can be expressed as two half- In this case we can write the general equation (1)R-C、+20H=R-C、+2e-+H20 (Do not mistake the above reaction for an acid dissociation the h arises from the removal of a (2)2Cu2++2e+20H=Cu20+H2O hydrogen atom, H+ e AH2 and a together Because two electrons are removed from the aldehyde constitute a conjugate redox pair(A/AH2), which carbon, the second half-reaction (the one-electron re- can reduce another compound b (or redox pair, duction of cupric to cuprous ion) must be doubled to B/BH2) by transfer of hydrogen atoms balance the overall equation. AH2+B=A+BH

E. coli convert electromotive force to proton-motive force, which is then used to power flagellar motion. The principles of electrochemistry that govern en￾ergy changes in the macroscopic circuit with a motor and battery apply with equal validity to the molecular processes accompanying electron flow in living cells. We turn now to a discussion of those principles. Oxidation-Reductions Can Be Described as Half-Reactions Although oxidation and reduction must occur together, it is convenient when describing electron transfers to consider the two halves of an oxidation-reduction reac￾tion separately. For example, the oxidation of ferrous ion by cupric ion, Fe2 Cu2 34 Fe3 Cu can be described in terms of two half-reactions: (1) Fe2 34 Fe3 e (2) Cu2 e 34 Cu The electron-donating molecule in an oxidation￾reduction reaction is called the reducing agent or reduc￾tant; the electron-accepting molecule is the oxidizing agent or oxidant. A given agent, such as an iron cation existing in the ferrous (Fe2) or ferric (Fe3) state, func￾tions as a conjugate reductant-oxidant pair (redox pair), just as an acid and corresponding base function as a con￾jugate acid-base pair. Recall from Chapter 2 that in acid￾base reactions we can write a general equation: proton donor 34 H proton acceptor. In redox reactions we can write a similar general equation: electron donor 34 e electron acceptor. In the reversible half-reaction (1) above, Fe2 is the electron donor and Fe3 is the elec￾tron acceptor; together, Fe2 and Fe3 constitute a con￾jugate redox pair. The electron transfers in the oxidation-reduction reactions of organic compounds are not fundamentally different from those of inorganic species. In Chapter 7 we considered the oxidation of a reducing sugar (an aldehyde or ketone) by cupric ion (see Fig. 7–10a): This overall reaction can be expressed as two half￾reactions: (1) (2) 2Cu2 2e 2OH 34 Cu2O H2O Because two electrons are removed from the aldehyde carbon, the second half-reaction (the one-electron re￾duction of cupric to cuprous ion) must be doubled to balance the overall equation. R C H O 2OH 2e  R C H2O OH O R C H O 4OH 2Cu2 R C Cu2O 2H2O OH O Biological Oxidations Often Involve Dehydrogenation The carbon in living cells exists in a range of oxidation states (Fig. 13–13). When a carbon atom shares an elec￾tron pair with another atom (typically H, C, S, N, or O), the sharing is unequal in favor of the more electroneg￾ative atom. The order of increasing electronegativity is H
C
S
N
O. In oversimplified but useful terms, the more electronegative atom “owns” the bonding elec￾trons it shares with another atom. For example, in methane (CH4), carbon is more electronegative than the four hydrogens bonded to it, and the C atom therefore “owns” all eight bonding electrons (Fig. 13–13). In ethane, the electrons in the COC bond are shared equally, so each C atom owns only seven of its eight bonding electrons. In ethanol, C-1 is less electronega￾tive than the oxygen to which it is bonded, and the O atom therefore “owns” both electrons of the COO bond, leaving C-1 with only five bonding electrons. With each formal loss of electrons, the carbon atom has undergone oxidation—even when no oxygen is involved, as in the conversion of an alkane (OCH2OCH2O) to an alkene (OCHUCHO). In this case, oxidation (loss of elec￾trons) is coincident with the loss of hydrogen. In bio￾logical systems, oxidation is often synonymous with de￾hydrogenation, and many enzymes that catalyze oxidation reactions are dehydrogenases. Notice that the more reduced compounds in Figure 13–13 (top) are richer in hydrogen than in oxygen, whereas the more oxidized compounds (bottom) have more oxygen and less hydrogen. Not all biological oxidation-reduction reactions in￾volve carbon. For example, in the conversion of molec￾ular nitrogen to ammonia, 6H 6e N2 n 2NH3, the nitrogen atoms are reduced. Electrons are transferred from one molecule (elec￾tron donor) to another (electron acceptor) in one of four different ways: 1. Directly as electrons. For example, the Fe2/Fe3 redox pair can transfer an electron to the Cu/Cu2 redox pair: Fe2 Cu2 34 Fe3 Cu 2. As hydrogen atoms. Recall that a hydrogen atom consists of a proton (H) and a single electron (e). In this case we can write the general equation AH2 34 A 2e 2H where AH2 is the hydrogen/electron donor. (Do not mistake the above reaction for an acid dissociation; the H arises from the removal of a hydrogen atom, H e.) AH2 and A together constitute a conjugate redox pair (A/AH2), which can reduce another compound B (or redox pair, B/BH2) by transfer of hydrogen atoms: AH2 B 34 A BH2 508 Chapter 13 Principles of Bioenergetics

13.3 Biological Oxidation-Reduction Reactions 509 3. As a hydride ion (H"), which has two electrons. potential of 0.00 V When this hydrogen electrode is con- This occurs in the case of NAD-linked dehydroge- nected through an external circuit to another half-cell in which an oxidized species and its corresponding re- 4. Through direct combination with oxygen. In this duced species are present at standard concentrations case, oxygen combines with an organic reductant (each solute at 1 M, each gas at 101.3 kPa), electrons tend nd is covalently incorporated in the product, to flow through the external circuit from the half-cell of in the oxidation of a hydrocarbon to an alcohol R-CH3+O2→R-CH2OH The hydrocarbon is the electron donor and the oxygen atom is the electron acceptor Methane H: C: H All four types of electron transfer occur in cells. The neutral term reducing equivalent is commonly used to designate a single electron equivalent participating in an oxidation-reduction reaction, no matter whether this H: C: C: H equivalent is an electron per se, a hydrogen atom, or a hy (alkane) dride ion, or whether the electron transfer takes place in a reaction with oxygen to yield an oxygenated product. Because biological fuel molecules are usually enzymati- Ethene cally dehydrogenated to lose two reducing equivalents at a time, and because each oxygen atom can accept two re- ducing equivalents, biochemists by convention regard the unit of biological oxidations as two reducing equivalents Ethanol H: C:C:O:H passing from substrate to oxygen. Reduction Potentials Measure Affinity for electrons H: C::: C:H When two conjugate redox pairs are together in solu tion, electron transfer from the electron donor of one pair to the electron acceptor of the other may proceed spontaneously. The tendency for such a reaction de Formaldehyd pends on the relative affinity of the electron acceptor of each redox pair for electrons. The standard reduc tion potential, E, a measure(in volts) of this affin ity, can be determined in an experiment such as that H: C: C described in Figure 13-14. Electrochemists have cho- H sen as a standard of reference the half-reaction The electrode at which this half-reaction occurs(called H: C: C: C:H 2 a half-cell) is arbitrarily assigned a standard reduction FIGURE 13-13 Oxidation states of carbon in the biosphere. The (carboxylic id) oxidation states are illustrated with some representative compounds. Focus on the red carbon atom and its bonding electrons. When this carbon is bonded to the less electronegative H atom, both bonding Carbon electrons(red) are assigned to the carbon. When carbon is bonded to another carbon, bonding electrons are shared equally, so one of the two electrons is assigned to the red carbon. When the red carbon is bonded to the more electronegative O atom, the bonding electrons Acetic acid H: C: C are assigned to the oxygen. The number to the right of each compound (carboxylic s the number of electrons owned" by the red carbon, a rough ex- pression of the oxidation state of that carbon. When the red carbon undergoes oxidation (loses electrons), the number gets smaller. Thus Carbon 0::C::0 the oxidation state increases from top to bottom of the list

potential of 0.00 V. When this hydrogen electrode is con￾nected through an external circuit to another half-cell in which an oxidized species and its corresponding re￾duced species are present at standard concentrations (each solute at 1 M, each gas at 101.3 kPa), electrons tend to flow through the external circuit from the half-cell of 13.3 Biological Oxidation-Reduction Reactions 509 Methane 8 H H H H C Ethane (alkane) 7 H H H C H H C H Ethanol (alcohol) 5 H H H C H H C O H Acetylene (alkyne) H H C C 5 Ethene (alkene) C C 6 H H H H Acetaldehyde (aldehyde) 3 H H H C O C H Formaldehyde 4 H H C O Carbon monoxide C O 2 Carbon dioxide O C O 0 Formic acid (carboxylic acid) 2 H H C O O Acetic acid (carboxylic acid) 1 H H H C C H O O Acetone (ketone) 2 H H H C H H C O C H FIGURE 13–13 Oxidation states of carbon in the biosphere. The oxidation states are illustrated with some representative compounds. Focus on the red carbon atom and its bonding electrons. When this carbon is bonded to the less electronegative H atom, both bonding electrons (red) are assigned to the carbon. When carbon is bonded to another carbon, bonding electrons are shared equally, so one of the two electrons is assigned to the red carbon. When the red carbon is bonded to the more electronegative O atom, the bonding electrons are assigned to the oxygen. The number to the right of each compound is the number of electrons “owned” by the red carbon, a rough ex￾pression of the oxidation state of that carbon. When the red carbon undergoes oxidation (loses electrons), the number gets smaller. Thus the oxidation state increases from top to bottom of the list. 3. As a hydride ion (:H), which has two electrons. This occurs in the case of NAD-linked dehydroge￾nases, described below. 4. Through direct combination with oxygen. In this case, oxygen combines with an organic reductant and is covalently incorporated in the product, as in the oxidation of a hydrocarbon to an alcohol: RXCH3  1 2 O2 88n RXCH2XOH The hydrocarbon is the electron donor and the oxygen atom is the electron acceptor. All four types of electron transfer occur in cells. The neutral term reducing equivalent is commonly used to designate a single electron equivalent participating in an oxidation-reduction reaction, no matter whether this equivalent is an electron per se, a hydrogen atom, or a hy￾dride ion, or whether the electron transfer takes place in a reaction with oxygen to yield an oxygenated product. Because biological fuel molecules are usually enzymati￾cally dehydrogenated to lose two reducing equivalents at a time, and because each oxygen atom can accept two re￾ducing equivalents, biochemists by convention regard the unit of biological oxidations as two reducing equivalents passing from substrate to oxygen. Reduction Potentials Measure Affinity for Electrons When two conjugate redox pairs are together in solu￾tion, electron transfer from the electron donor of one pair to the electron acceptor of the other may proceed spontaneously. The tendency for such a reaction de￾pends on the relative affinity of the electron acceptor of each redox pair for electrons. The standard reduc￾tion potential, E, a measure (in volts) of this affin￾ity, can be determined in an experiment such as that described in Figure 13–14. Electrochemists have cho￾sen as a standard of reference the half-reaction H e 88n  1 2  H2 The electrode at which this half-reaction occurs (called a half-cell) is arbitrarily assigned a standard reduction

10 Device for where r and thave their usual n is the num measuring emf ber of electrons transferred per molecule, and J is the Faraday constant(Table 13-1). At 298 K(25C), this expression reduces to E=po+0.026V, [electron acceptor H2gas→ Many half-reactions of interest to biochemists in- pressure Salt bridge olve protons. As in the definition of AG, biochemists define the standard state for oxidation-reduction reac. tions as pH 7 and express reduction potential as Eo, the standard reduction potential at pH 7. The standard re- duction potentials given in Table 13-7 and used through out this book are values for eo and are therefore valid only for systems at neutral pH. Each value represents the potential difference when the conjugate redox pair, at 1 M concentrations and pH 7, is connected with the standard (pH 0) hydrogen electrode. Notice in Table 13-7 that when the conjugate pair 2H/H2 at pH 7 is connected with the standard hydrogen electrode (pl Reference cell of Test cell containing 0), electrons tend to flow from the pH 7 cell to the stan- dard (pH O) cell; the measured E for the 2H/H2 pair in which H2 gas is-0.414V. at 101.3 kPa is the redox pair to equilibrated at he electrode Standard reduction potentials can be used ith IMH to Calculate the Free-Energy Change The usefulness of reduction potentials stems from the FIGURE 13-14 Measurement of the standard reduction potential fact that when E values have been determined for any two half-cells, relative to the standard hydrogen elec- erence electrode, or vice versa. he ultimate reterence hall-cell is the trode, their reduction potentials relative to each other hydrogen electrode, as shown here, at pH O. The electromotive force are also known. We can then predict the direction in (emf)of this electrode is designated 0.00 V. At pH 7 in the test cell, which electrons will tend to flow when the two half-cells E for the hydrogen electrode is-0. 414 V. The direction of electron ow depends on the relative electron"pressure"or potential of the are connected through an external circuit or when com two cells. A salt bridge containing a saturated KCl solution provides onents of both half-cells are present in the same solu a path for counter-ion movement between the test cell and the refer. tion. Electrons tend to flow to the half-cell with the more IL. From the observed emf and the known emf of the reference positive E, and the strength of that tendency is pro- cell, the experimenter can find the emf of the test cell containing the portional to the difference in reduction potentials, AE dox pair. The cell that gains electrons has, by convention, the me The energy made available by this spontaneous positive reduction potential electron flow(the free-energy change for the oxidation reduction reaction) is proportional to AE △G=-n△Eor△G°=-n△E°(13-6) lower standard reduction potential to the half-cell of Here n represents the number of electrons transferred higher standard reduction potential. By convention, the in the reaction. With this equation we can calculate the half-cell with the stronger tendency to acquire electrons free-energy change for any oxidation-reduction reaction is assigned a positive value of E from the values of E in a table of reduction potentials The reduction potential of a half-cell depends not ( Table 13-7)and the concentrations of the species par only on the chemical species present but also on their ticipating in the reaction activities, approximated by their concentrations. About Consider the reaction in whi a century ago. Walther Nernst derived an equation that reduced by the biological electron carrier NADH elates standard reduction potential (E)to the reduc tion potential(E) at any concentration of oxidized and Acetaldehyde+NADH+H→→ ethanol+NAD reduced species in the cell The relevant half-reactions and their e values are E=E+RTIn electron acceptor (13-4) ()Acetaldehyde+ 2H+ 2e - ethanol

lower standard reduction potential to the half-cell of higher standard reduction potential. By convention, the half-cell with the stronger tendency to acquire electrons is assigned a positive value of E . The reduction potential of a half-cell depends not only on the chemical species present but also on their activities, approximated by their concentrations. About a century ago, Walther Nernst derived an equation that relates standard reduction potential (E ) to the reduc￾tion potential (E) at any concentration of oxidized and reduced species in the cell: E  E  R nℑ T  ln (13–4)  [electron acceptor] [electron donor] where R and T have their usual meanings, n is the num￾ber of electrons transferred per molecule, and is the Faraday constant (Table 13–1). At 298 K (25 C), this expression reduces to E  E  0.02 n  6 V ln (13–5) Many half-reactions of interest to biochemists in￾volve protons. As in the definition of G , biochemists define the standard state for oxidation-reduction reac￾tions as pH 7 and express reduction potential as E , the standard reduction potential at pH 7. The standard re￾duction potentials given in Table 13–7 and used through￾out this book are values for E and are therefore valid only for systems at neutral pH. Each value represents the potential difference when the conjugate redox pair, at 1 M concentrations and pH 7, is connected with the standard (pH 0) hydrogen electrode. Notice in Table 13–7 that when the conjugate pair 2H/H2 at pH 7 is connected with the standard hydrogen electrode (pH 0), electrons tend to flow from the pH 7 cell to the stan￾dard (pH 0) cell; the measured E for the 2H/H2 pair is 0.414 V. Standard Reduction Potentials Can Be Used to Calculate the Free-Energy Change The usefulness of reduction potentials stems from the fact that when E values have been determined for any two half-cells, relative to the standard hydrogen elec￾trode, their reduction potentials relative to each other are also known. We can then predict the direction in which electrons will tend to flow when the two half-cells are connected through an external circuit or when com￾ponents of both half-cells are present in the same solu￾tion. Electrons tend to flow to the half-cell with the more positive E, and the strength of that tendency is pro￾portional to the difference in reduction potentials, E. The energy made available by this spontaneous electron flow (the free-energy change for the oxidation￾reduction reaction) is proportional to E: G  n E or G  n E (13–6) Here n represents the number of electrons transferred in the reaction. With this equation we can calculate the free-energy change for any oxidation-reduction reaction from the values of E in a table of reduction potentials (Table 13–7) and the concentrations of the species par￾ticipating in the reaction. Consider the reaction in which acetaldehyde is reduced by the biological electron carrier NADH: Acetaldehyde NADH H 88n ethanol NAD The relevant half-reactions and their E values are: (1) Acetaldehyde 2H 2e 88n ethanol E  0.197 V  [electron acceptor] [electron donor] 510 Chapter 13 Principles of Bioenergetics Salt bridge (KCl solution) Reference cell of known emf: the hydrogen electrode in which H2 gas at 101.3 kPa is equilibrated at the electrode with 1 M H Test cell containing 1 M concentrations of the oxidized and reduced species of the redox pair to be examined H2 gas (standard pressure) Device for measuring emf FIGURE 13–14 Measurement of the standard reduction potential (E ) of a redox pair. Electrons flow from the test electrode to the ref￾erence electrode, or vice versa. The ultimate reference half-cell is the hydrogen electrode, as shown here, at pH 0. The electromotive force (emf) of this electrode is designated 0.00 V. At pH 7 in the test cell, E for the hydrogen electrode is 0.414 V. The direction of electron flow depends on the relative electron “pressure” or potential of the two cells. A salt bridge containing a saturated KCl solution provides a path for counter-ion movement between the test cell and the refer￾ence cell. From the observed emf and the known emf of the reference cell, the experimenter can find the emf of the test cell containing the redox pair. The cell that gains electrons has, by convention, the more positive reduction potential.