13.4 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 reactions, sometimes referred to as redox reactions. 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 responsible, directly or indirectly, for all work done by living organisms. In nonphotosynthetic organisms, the sources of electrons are reduced compounds (foods); in photosynthetic organisms, the initial electron donor is a chemical species excited by the absorption of light. The path of electron flow in metabolism is complex. Electrons move from various metabolic intermediates to specialized electron carriers in enzyme-catalyzed reactions. The carriers, in turn, donate electrons to acceptors with higher electron affinities, with the release of energy. Cells possess a variety of molecular energy transducers, which convert the energy of electron flow into useful work.

We begin by discussing how work can be accomplished by an electromotive force (emf), then consider the theoretical and experimental basis for measuring energy changes in oxidation reactions in terms of emf and the relationship between this force, expressed in volts, and the free-energy change, expressed in joules. We also describe the structures and oxidation-reduction chemistry of the most common of the specialized 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 battery containing two chemical species that differ in affinity 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 chemical 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 emf (typically a few volts) can accomplish work if an appropriate energy transducer — in this case a motor — is placed in the circuit. The motor can be coupled to a variety of mechanical devices to do 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 oxidized, the released electrons flow spontaneously through a series of electron-carrier intermediates to another chemical species, such as O2upper O Subscript 2. This electron flow is exergonic, because O2upper O Subscript 2 has a higher affinity for electrons than do the electron-carrier intermediates. The resulting emf provides energy to a variety 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 difference and a transmembrane electrical potential, accomplishing chemiosmotic and electrical work. The proton gradient thus formed has potential energy, sometimes 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 Piupper P Subscript i as protons flow spontaneously across the membrane. Similarly, membrane-localized enzymes in Escherichia coli convert emf to proton-motive force, which is then used to power flagellar motion. The principles of electrochemistry that govern energy changes in the macroscopic circuit with a motor and battery apply with equal validity to the molecular processes accompanying electron flow in living cells.

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 reaction separately. For example, the oxidation of ferrous ion by cupric ion,

Fe2++Cu2+  Fe3++Cu+Fe Superscript 2 plus Baseline plus Cu Superscript 2 plus Baseline right harpoon over left harpoon Fe Superscript 3 plus Baseline plus Cu Superscript plus

can be described in terms of two half-reactions:

  1. Fe2+  Fe3++eFe Superscript 2 plus Baseline right harpoon over left harpoon Fe Superscript 3 plus Baseline plus e Superscript minus
  2. Cu2++e  Cu+Cu Superscript 2 plus Baseline plus e Superscript minus Baseline right harpoon over left harpoon Cu Superscript plus

The electron-donating molecule in an oxidation-reduction reaction is called the reducing agent or reductant; the electron-accepting molecule is the oxidizing agent or oxidant. A given agent, such as an iron cation existing in the ferrous (Fe2+)left-parenthesis Fe Superscript 2 plus Baseline right-parenthesis or ferric (Fe3+)left-parenthesis Fe Superscript 3 plus Baseline right-parenthesis state, functions as a conjugate reductant-oxidant pair (redox pair), just as an acid and corresponding base function as a conjugate acid-base pair. Recall from Chapter 2 that in acid-base reactions we can write a general equation: proton donor  H++right harpoon over left harpoon upper H Superscript plus plus proton acceptor. In redox reactions we can write a similar general equation: electron donor (reductant)  e+right harpoon over left harpoon e Superscript minus plus electron acceptor (oxidant). In the reversible half-reaction (1) above, Fe2+Fe Superscript 2 plus is the electron donor and Fe3+Fe Superscript 3 plus is the electron acceptor; together, Fe2+Fe Superscript 2 plus and Fe3+Fe Superscript 3 plus constitute a conjugate redox pair. The mnemonic OIL RIG — oxidation is losing, reduction is gaining — may be helpful in remembering what happens to electrons in redox reactions.

The electron transfers in the oxidation-reduction reactions of organic compounds are not fundamentally different from those of inorganic species. Consider the oxidation of a reducing sugar (an aldehyde or a ketone) by cupric ion:

R bonded to C double bonded to O to the upper right and bonded to H to the lower right plus 4 O H minus plus 2 C u 2 plus undergo a reversible reaction to produce R bonded to C double bonded to O to the upper right and bonded to O H to the lower right plus C u 2 O plus 2 H 2 O.

This overall reaction can be expressed as two half-reactions:

R bonded to C double bonded to O to the upper right and bonded to H to the lower right plus 2 O H minus undergo a reversible reaction to produce R bonded to C double bonded to O to the upper right and bonded to O H to the lower right plus 2 e minus plus 2 H 2 O.
  1. 2Cu2++2e+2OH  Cu2O + H2O2 Cu Superscript 2 plus Baseline plus 2 e Superscript minus Baseline plus 2 OH Superscript minus Baseline right harpoon over left harpoon Cu Subscript 2 Baseline upper O plus upper H Subscript 2 Baseline upper O

Notice that because two electrons are removed from the aldehyde carbon, the second half-reaction (the one-electron reduction of cupric to cuprous ion) must be doubled to balance the overall equation.

Biological Oxidations Often Involve Dehydrogenation

The carbon in living cells exists in a range of oxidation states (Fig. 13-22). When a carbon atom shares an electron pair with another atom (typically H, C, S, N, or O), the sharing is unequal, in favor of the more electronegative atom. The order of increasing electronegativity is H<C<S<N<Oupper H less-than upper C less-than upper S less-than upper N less-than upper O. In oversimplified but useful terms, the more electronegative atom “owns” the bonding electrons it shares with another atom. For example, in methane (CH4)left-parenthesis CH Subscript 4 Baseline right-parenthesis, carbon is more electronegative than the four hydrogens bonded to it, and the C atom therefore owns all eight bonding electrons (Fig. 13-22). In ethane, the electrons in the CCupper C em-dash upper C bond are shared equally, so each C atom owns only seven of its eight bonding electrons. In ethanol, C-1 is less electronegative than the oxygen to which it is bonded, and the O atom therefore owns both electrons of the COupper C em-dash upper O bond, leaving C-1 with only five bonding electrons. With each formal loss of “owned” electrons, the carbon atom has undergone oxidation — even when no oxygen is involved, as in the conversion of an alkane (CH2CH2)left-parenthesis em-dash CH Subscript 2 Baseline em-dash CH Subscript 2 Baseline em-dash right-parenthesis to an alkene (CHCH)left-parenthesis em-dash CH box drawings double horizontal CH em-dash right-parenthesis. In this case, oxidation (loss of electrons) is coincident with the loss of hydrogen. In biological systems, as we noted earlier in the chapter, oxidation is often synonymous with dehydrogenation, and many enzymes that catalyze oxidation reactions are dehydrogenases. Notice that the more reduced compounds in Figure 13-22 (top) are richer in hydrogen than in oxygen, whereas the more oxidized compounds (bottom) have more oxygen and less hydrogen.

A figure shows the structures and levels of oxidation of 12 carbon compounds in the biosphere: methane, ethane, ethene, ethanol, acetylene, formaldehyde, acetaldehyde, acetone, formic acid, carbon monoxide, acetic acid, and carbon dioxide.

FIGURE 13-22 Different levels of oxidation of carbon compounds in the biosphere. To approximate the level of oxidation of these 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 expression of the degree of oxidation of that compound. As the red carbon undergoes oxidation (loses electrons), the number gets smaller.

Not all biological oxidation-reduction reactions involve carbon. For example, in the conversion of molecular nitrogen to ammonia, 6H+ +6e+N22NH36 upper H Superscript plus Baseline plus 6 e Superscript minus Baseline plus upper N Subscript 2 Baseline right-arrow 2 NH Subscript 3 Baseline, the nitrogen atoms are reduced.

Electrons are transferred from one molecule (electron donor) to another (electron acceptor) in one of four ways:

  1. Directly as electrons. For example, the Fe2+/Fe3+Fe Superscript 2 plus Baseline slash Fe Superscript 3 plus redox pair can transfer an electron to the Cu+/Cu2+Cu Superscript plus Baseline slash Cu Superscript 2 plus redox pair:
    Fe2++Cu2+  Fe3++Cu+Fe Superscript 2 plus Baseline plus Cu Superscript 2 plus Baseline right harpoon over left harpoon Fe Superscript 3 plus Baseline plus Cu Superscript plus
  2. As hydrogen atoms. Recall that a hydrogen atom consists of a proton (H+)left-parenthesis upper H Superscript plus Baseline right-parenthesis and a single electron (e)left-parenthesis e Superscript minus Baseline right-parenthesis. In this case we can write the general equation
    AH2  A+2e+2H+AH Subscript 2 Baseline right harpoon over left harpoon upper A plus 2 e Superscript minus Baseline plus 2 upper H Superscript plus

    where AH2AH Subscript 2 is the hydrogen/electron donor. (Do not mistake the above reaction for an acid dissociation, which involves a proton and no electron.) AH2AH Subscript 2 and A together constitute a conjugate redox pair (A/AH2)left-parenthesis upper A slash AH Subscript 2 Baseline right-parenthesis, which can reduce another compound B (or redox pair, B/BH2upper B slash BH Subscript 2) by transfer of hydrogen atoms:

    AH2+B  A+BH2AH Subscript 2 Baseline plus upper B right harpoon over left harpoon upper A plus BH Subscript 2 Baseline
  3. As a hydride ion (:H)left-parenthesis colon upper H Superscript minus Baseline right-parenthesis, which has two electrons. This occurs in the case of NAD-linked dehydrogenases, 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:
    RCH3+12O2RCH2OHupper R em-dash CH Subscript 3 Baseline plus one-half upper O Subscript 2 Baseline right-arrow upper R em-dash CH Subscript 2 Baseline em-dash OH

    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 or is part of a hydrogen atom or a hydride ion, or whether the electron transfer takes place in a reaction with oxygen to yield an oxygenated product.

Reduction Potentials Measure Affinity for Electrons

When two conjugate redox pairs are together in solution, 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 depends on the relative affinity of the electron acceptor of each redox pair for electrons. The standard reduction potential, E°upper E degree, a measure (in volts) of this affinity, can be determined in an experiment such as that described in Figure 13-23. Electrochemists have chosen as a standard of reference the half-reaction

H++e12H2upper H Superscript plus Baseline plus e Superscript minus Baseline right-arrow Overscript Endscripts one-half upper H Subscript 2

The electrode at which this half-reaction occurs (called a half-cell) is arbitrarily assigned an E°upper E degree of 0.00 V. When this hydrogen electrode is connected through an external circuit to another half-cell in which an oxidized species and its corresponding reduced species are present at standard concentrations (at 25 °C25 degree upper C, each solute at 1 m, each gas at 101.3 kPa), electrons tend to flow through the external circuit from the half-cell of lower E°upper E degree to the half-cell of higher E°upper E degree. By convention, a half-cell that takes electrons from the standard hydrogen cell is assigned a positive value of E°upper E degree, and one that donates electrons to the hydrogen cell, a negative value. When any two half-cells are connected, that with the larger (more positive) E°upper E degree will be reduced; it has the greater reduction potential.

A figure shows two beakers set up with a salt bridge and electrodes to measure the standard reduction potential (italicized E end italics prime degree symbol) of a redox pair.

FIGURE 13-23 Measurement of the standard reduction potential (E°)left-parenthesis upper E Superscript prime Baseline degree right-parenthesis of a redox pair. Electrons flow from the test electrode to the reference 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 (at 25 °C25 degree upper C), E°upper E prime degree for the hydrogen electrode is 0.414 Vnegative 0.414 upper 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 reference 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.

The reduction potential of a half-cell depends not only on the chemical species present but also on their activities, approximated by their concentrations. The Nernst equation relates standard reduction potential (E°)left-parenthesis upper E degree right-parenthesis to the actual reduction potential (E) at any concentration of oxidized and reduced species in a living cell:

E=E°+RTnFln [electron acceptor][electron donor]upper E equals upper E degree plus StartFraction upper R upper T Over n upper F EndFraction ln StartFraction left-bracket electron acceptor right-bracket Over left-bracket electron donor right-bracket EndFraction

(13-5)

where R and T have their usual meanings, n is the number of electrons transferred per molecule, and F is the Faraday constant, a proportionality constant that converts volts to joules (Table 13-1). At 298 K (25 °C)left-parenthesis 25 degree upper C right-parenthesis, this expression reduces to

E=E°+0.026 Vnln [electron acceptor][electron donor]upper E equals upper E degree plus StartFraction 0.026 upper V Over n EndFraction ln StartFraction left-bracket electron acceptor right-bracket Over left-bracket electron donor right-bracket EndFraction

(13-6)

Key convention

Many half-reactions of interest to biochemists involve protons. As in the definition of ΔG°upper Delta upper G prime degree, biochemists define the standard state for oxidation-reduction reactions as pH 7 and express a standard transformed reduction potential, E°upper E prime degree, the standard reduction potential at pH 7 and 25 °C25 degree upper C. By convention, ΔE°upper Delta upper E prime degree for any redox reaction is given as E°upper E prime degree of the electron acceptor minus E°upper E prime degree of the electron donor.

The standard reduction potentials given in Table 13-7 and used throughout this book are values for E°upper E prime degree 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, 25 °C25 degree upper C, and pH 7, is connected with the standard (pH 0) hydrogen electrode. Notice in Table 13-7 that when the conjugate pair 2H+/H22 upper H Superscript plus Baseline slash upper H Subscript 2 at pH 7 is connected with the standard hydrogen electrode (pH 0), electrons tend to flow from the pH 7 cell to the standard (pH 0) cell; the measured E°upper E prime degree for the 2H+/H22 upper H Superscript plus Baseline slash upper H Subscript 2 pair is –0.414 V.

TABLE 13-7 Standard Reduction Potentials of Some Biologically Important Half-Reactions

Half-reaction E°bold-italic upper E prime degree (V)

12O2+2H++2eH2Oone-half upper O Subscript 2 Baseline plus 2 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts upper H Subscript 2 Baseline upper O

0.816

Fe3++eFe2+Fe Superscript 3 plus Baseline plus e Superscript minus Baseline right-arrow Overscript Endscripts Fe Superscript 2 plus

0.771

NO3+2H++2eNO2+H2ONO Subscript 3 Superscript minus Baseline plus 2 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts NO Subscript 2 Superscript minus Baseline plus upper H Subscript 2 Baseline upper O

0.421

Cytochrome f (Fe3+)+ecytochrome f (Fe2+)Cytochrome f left-parenthesis Fe Superscript 3 plus Baseline right-parenthesis plus e Superscript minus Baseline right-arrow Overscript Endscripts cytochrome f left-parenthesis Fe Superscript 2 plus Baseline right-parenthesis

0.365

Fe(CN)63(ferricyanide)+eFe(CN)64Fe left-parenthesis CN right-parenthesis Subscript 6 Superscript 3 minus Baseline left-parenthesis ferricyanide right-parenthesis plus e Superscript minus Baseline right-arrow Overscript Endscripts Fe left-parenthesis CN right-parenthesis Subscript 6 Superscript 4 minus

0.36

Cytochrome a3 (Fe3+)+ecytochrome a3 (Fe2+)Cytochrome a 3 left-parenthesis Fe Superscript 3 plus Baseline right-parenthesis plus e Superscript minus Baseline right-arrow Overscript Endscripts cytochrome a 3 left-parenthesis Fe Superscript 2 plus Baseline right-parenthesis

0.35

O2+2H++2eH2O2upper O Subscript 2 Baseline plus 2 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts upper H Subscript 2 Baseline upper O Subscript 2

0.295

Cytochrome a (Fe3+)+e cytochrome a (Fe2+)Cytochrome a left-parenthesis Fe Superscript 3 plus Baseline right-parenthesis plus e Superscript minus Baseline right-arrow Overscript Endscripts cytochrome a left-parenthesis Fe Superscript 2 plus Baseline right-parenthesis

0.29

Cytochrome c (Fe3+)+e cytochrome c (Fe2+)Cytochrome c left-parenthesis Fe Superscript 3 plus Baseline right-parenthesis plus e Superscript minus Baseline right-arrow Overscript Endscripts cytochrome c left-parenthesis Fe Superscript 2 plus Baseline right-parenthesis

0.254

Cytochrome c1 (Fe3+)+e cytochrome c1 (Fe2+)Cytochrome c 1 left-parenthesis Fe Superscript 3 plus Baseline right-parenthesis plus e Superscript minus Baseline right-arrow Overscript Endscripts cytochrome c 1 left-parenthesis Fe Superscript 2 plus Baseline right-parenthesis

0.22

Cytochrome b (Fe3+)+e cytochrome b (Fe2+)Cytochrome b left-parenthesis Fe Superscript 3 plus Baseline right-parenthesis plus e Superscript minus Baseline right-arrow Overscript Endscripts cytochrome b left-parenthesis Fe Superscript 2 plus Baseline right-parenthesis

0.077

Ubiquinone+2H++2e ubiquinolUbiquinone plus 2 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts ubiquinol

0.045

Fumarate2+2H++2esuccinate2Fumarate Superscript 2 minus Baseline plus 2 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts succinate Superscript 2 minus

0.031

2H++2e  H22 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts upper H Subscript 2 Baseline (at standard conditions, pH 0)

0.000

Crotonyl-CoA+2H++2ebutyryl-CoACrotonyl hyphen CoA plus 2 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts butyryl hyphen CoA

0.015negative 0.015

Oxaloacetate2+2H++2emalate2Oxaloacetate Superscript 2 minus Baseline plus 2 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts malate Superscript 2 minus

0.166negative 0.166

Pyruvate+2H++2elactatePyruvate Superscript minus Baseline plus 2 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts lactate Superscript minus

0.185negative 0.185

Acetaldehyde+2H++2eethanolAcetaldehyde plus 2 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts ethanol

0.197negative 0.197

FAD+2H++2e FADH2FAD plus 2 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts FADH Subscript 2 Baseline

0.219negative 0.219a

Glutathione+2H++2e2 reduced glutathioneGlutathione plus 2 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts 2 reduced glutathione

0.23negative 0.23

S+2H++2eH2Supper S plus 2 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts upper H Subscript 2 Baseline upper S

0.243negative 0.243

Lipoic acid+2H++2edihydrolipoic acidLipoic acid plus 2 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts dihydrolipoic acid

0.29negative 0.29

NAD++H++2eNADHNAD Superscript plus Baseline plus upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts NADH

0.320negative 0.320

NADP++H++2eNADPHNADP Superscript plus Baseline plus upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts NADPH

0.324negative 0.324

Acetoacetate+2H++2e β-hydroxybutyrateAcetoacetate plus 2 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts beta hyphen hydroxybutyrate

0.346negative 0.346

α-Ketoglutarate+CO2+2H++2e isocitratealpha hyphen Ketoglutarate plus CO Subscript 2 Baseline plus 2 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts isocitrate

0.38negative 0.38

2H++2e H2 (at pH 7)2 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts upper H Subscript 2 Baseline left-parenthesis at pH 7 right-parenthesis

0.414negative 0.414

Ferredoxin (Fe3+)+e ferredoxin (Fe2+)Ferredoxin left-parenthesis Fe Superscript 3 plus Baseline right-parenthesis plus e Superscript minus Baseline right-arrow Overscript Endscripts ferredoxin left-parenthesis Fe Superscript 2 plus Baseline right-parenthesis

0.432negative 0.432

Data mostly from R. A. Loach, in Handbook of Biochemistry and Molecular Biology, 3rd edn (G. D. Fasman, ed.), Physical and Chemical Data, Vol. 1, p. 122, CRC Press, 1976.

aThis is the value for free FAD; FAD bound to a specific flavoprotein (e.g., succinate dehydrogenase) has a different E°upper E prime degree that depends on its protein environment.

Standard Reduction Potentials Can Be Used to Calculate Free-Energy Change

Why are reduction potentials so useful to the biochemist? When E values have been determined for any two half-cells, relative to the standard hydrogen electrode, we also know their reduction potentials relative to each other. 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 components of both half-cells are present in the same solution. Electrons tend to flow to the half-cell with the more positive E, and the strength of that tendency is proportional to ∆E, the difference in reduction potential. The energy made available by this spontaneous electron flow (the free-energy change, ∆G, for the oxidation-reduction reaction) is proportional to ∆E:

ΔG=nFΔE  or  ΔG°=nFΔE°upper Delta upper G equals minus n upper F upper Delta upper E or upper Delta upper G Superscript prime Baseline degree equals minus n upper F upper Delta upper E prime degree

(13-7)

where n is the number of electrons transferred in the reaction. With this equation we can calculate the actual free-energy change for any oxidation-reduction reaction from the values of ΔE°upper Delta upper E prime degree in a table of reduction potentials (Table 13-7) and the concentrations of reacting species.

WORKED EXAMPLE 13-3 Calculation of ΔG°upper Delta upper G prime degree and ΔG of a Redox Reaction

Calculate the standard free-energy change, ΔG°upper Delta upper G prime degree, for the reaction in which acetaldehyde is reduced by the biological electron carrier NADH:

Acetaldehyde+NADH+H+  ethanol+NAD+Acetaldehyde plus NADH plus upper H Superscript plus Baseline right-arrow Overscript Endscripts ethanol plus NAD Superscript plus

Then calculate the actual free-energy change, ∆G, when [acetaldehyde] and [NADH] are 1.00 m, and [ethanol] and [NAD+]left-bracket NAD Superscript plus Baseline right-bracket are 0.100 m. The relevant half-reactions and their E°upper E prime degree values are

  1. Acetaldehyde+2H++2e  ethanol E°=0.197 VAcetaldehyde plus 2 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts ethanol upper E Superscript prime Baseline degree equals negative 0.197 upper V
  2. NAD++2H++2e  NADH+H+     E°=0.320 VNAD Superscript plus Baseline plus 2 upper H Superscript plus Baseline plus 2 e Superscript minus Baseline right-arrow Overscript Endscripts NADH plus upper H Superscript plus Baseline upper E Superscript prime Baseline degree equals negative 0.320 upper V

Remember that, by convention, ΔE°upper Delta upper E prime degree is E°upper E prime degree of the electron acceptor minus E°upper E prime degree of the electron donor. It represents the difference between the electron affinities of the two half-reactions in the table of reduction potentials (Table 13-7). Note that the more widely separated the two half-reactions in the table, the more energetic the electron-transfer reaction when the two half-reactions occur together. By convention, in tables of reduction potentials, all half-reactions are represented as reductions, but when two half-reactions occur together, one of them must be an oxidation. Although that half-reaction will go in the opposite direction from that shown in Table 13-7, we do not change the sign of that half-reaction before calculating ΔE°upper Delta upper E prime degree, because ΔE°upper Delta upper E prime degree is defined as a difference of reduction potentials.

SOLUTION:

Because acetaldehyde is accepting electrons (n=2)left-parenthesis n equals 2 right-parenthesis from NADH, ΔE°=0.197 V(0.320 V)=0.123 Vupper Delta upper E Superscript prime Baseline degree equals negative 0.197 upper V minus left-parenthesis negative 0.320 upper V right-parenthesis equals 0.123 upper V. Therefore,

ΔG°=nFΔE°=2(96.5 kJ/Vmol) (0.123 V)=23.7 kJ/molupper Delta upper G Superscript prime Baseline degree equals minus n upper F upper Delta upper E Superscript prime Baseline degree equals minus 2 left-parenthesis 96.5 kJ slash upper V bullet mol right-parenthesis left-parenthesis 0.123 upper V right-parenthesis equals negative 23.7 kJ slash mol

This is the free-energy change for the oxidation-reduction reaction at 25 °C25 degree upper C and pH 7, when acetaldehyde, ethanol, NAD+NAD Superscript plus, and NADH are all present at 1.00 m concentrations.

To calculate ΔG when [acetaldehyde] and [NADH] are 1.00 m, and [ethanol] and [NAD+]left-bracket NAD Superscript plus Baseline right-bracket are 0.100 m, we can use Equation 13-4 and the standard free-energy change calculated above:

ΔG=ΔG°+RT ln[ethanol][NAD+][acetaldehyde][NADH]=23.7 kJ/mol+(8.315 J/molK)(298 K) ln(0.100 M)(0.100 M)(1.00 M)(1.00 M)=23.7 kJ/mol+(2.48 J/mol) ln 0.01=35.1 kJ/molStartLayout 1st Row 1st Column upper Delta upper G 2nd Column equals 3rd Column upper Delta upper G prime degree plus upper R upper T ln StartFraction left-bracket ethanol right-bracket left-bracket NAD Superscript plus Baseline right-bracket Over left-bracket acetaldehyde right-bracket left-bracket NADH right-bracket EndFraction 2nd Row 1st Column Blank 2nd Column equals 3rd Column negative 23.7 kJ slash mol plus left-parenthesis 8.315 upper J slash mol bullet upper K right-parenthesis left-parenthesis 298 upper K right-parenthesis ln StartFraction left-parenthesis 0.100 upper M right-parenthesis left-parenthesis 0.100 upper M right-parenthesis Over left-parenthesis 1.00 upper M right-parenthesis left-parenthesis 1.00 upper M right-parenthesis EndFraction 3rd Row 1st Column Blank 2nd Column equals 3rd Column negative 23.7 kJ slash mol plus left-parenthesis 2.48 upper J slash mol right-parenthesis ln 0.01 4th Row 1st Column Blank 2nd Column equals 3rd Column negative 35.1 kJ slash mol EndLayout

This is the actual free-energy change at the specified concentrations of the redox pairs.

A Few Types of Coenzymes and Proteins Serve as Universal Electron Carriers

The principles of oxidation-reduction energetics described above apply to the many metabolic reactions that involve electron transfers. For example, in many organisms, the oxidation of glucose supplies energy for the production of ATP. The complete oxidation of glucose

C6H12O6+6O2  6CO2+6H2Oupper C Subscript 6 Baseline upper H Subscript 12 Baseline upper O Subscript 6 Baseline plus 6 upper O Subscript 2 Baseline right-arrow Overscript Endscripts 6 CO Subscript 2 Baseline plus 6 upper H Subscript 2 Baseline upper O

has a ΔG°upper Delta upper G prime degree of 2,840 kJ/molnegative 2,840 kJ slash mol. This is a much larger release of free energy than is required for ATP synthesis in cells (50 to 60 kJ/mol; see Worked Example 13-2). Cells convert glucose to CO2CO Subscript 2 not in a single, high-energy-releasing reaction but rather in a series of controlled reactions, some of which are oxidations. The free energy released in these oxidation steps is of the same order of magnitude as that required for ATP synthesis from ADP, with some energy to spare. Electrons removed in these oxidation steps are transferred to coenzymes specialized for carrying electrons, such as NAD+NAD Superscript plus and FAD (described below).

The multitude of enzymes that catalyze cellular oxidations channel electrons from their hundreds of different substrates into just a few types of universal electron carriers. The reduction of these carriers in catabolic processes results in the conservation of free energy released by substrate oxidation. NAD, NADP, FMN, and FAD are water-soluble coenzymes that undergo reversible oxidation and reduction in many of the electron-transfer reactions of metabolism. The nucleotides NAD and NADP move readily from one enzyme to another; the flavin nucleotides FMN and FAD are usually very tightly bound to the enzymes, called flavoproteins, for which they serve as prosthetic groups. Lipid-soluble quinones such as ubiquinone and plastoquinone act as electron carriers and proton donors in the nonaqueous environment of membranes. Iron-sulfur proteins and cytochromes, which have tightly bound prosthetic groups that undergo reversible oxidation and reduction, also serve as electron carriers in many oxidation-reduction reactions. Some of these proteins are water-soluble, but others are peripheral or integral membrane proteins. The oxidation-reduction chemistry of quinones, iron-sulfur proteins, and cytochromes is discussed in Chapters 19 and 20.

Nicotinamide adenine dinucleotide (NAD; NAD+NAD Superscript plus in its oxidized form) and its close analog nicotinamide adenine dinucleotide phosphate (NADP; NADP+NADP Superscript plus when oxidized) are composed of two nucleotides joined through their phosphate groups by a phosphoanhydride bond (Fig. 13-24a). Because the nicotinamide ring resembles pyridine, these compounds are sometimes called pyridine nucleotides. The vitamin niacin is the source of the nicotinamide moiety in nicotinamide nucleotides.

A two-part figure shows the structures of N A D plus and N A D H in part a and a graph of the U V absorption spectra of N A D plus and N A D H in part b.

FIGURE 13-24 NAD and NADP. (a) Nicotinamide adenine dinucleotide, NAD+NAD Superscript plus, and its phosphorylated analog, NADP+NADP Superscript plus, undergo reduction to NADH and NADPH, accepting a hydride ion (two electrons and one proton) from an oxidizable substrate. The hydride ion is added to either the front or the back of the planar nicotinamide ring. (b) The UV absorption spectra of NAD+NAD Superscript plus and NADH. Reduction of the nicotinamide ring produces a new, broad absorption band with a maximum at 340 nm. The production of NADH during an enzyme-catalyzed reaction can be conveniently followed by observing the appearance of the absorbance at 340 nm (molar extinction coefficient ε340=6,200 M1 cm1epsilon 340 equals 6,200 upper M Superscript negative 1 Baseline cm Superscript negative 1).

Both coenzymes undergo reversible reduction of the nicotinamide ring (Fig. 13-24). As a substrate molecule undergoes oxidation (dehydrogenation), giving up two hydrogen atoms, the oxidized form of the nucleotide (NAD+ or NADP+)left-parenthesis NAD Superscript plus Baseline or NADP Superscript plus Baseline right-parenthesis accepts a hydride ion (:Hcolon upper H Superscript minus, the equivalent of a proton and two electrons) and is reduced (to NADH or NADPH). The second proton removed from the substrate is released to the aqueous solvent. The half-reactions for these nucleotide cofactors are

NAD++2e+2H+NADH+H+NADP++2e+2H+NADPH+H+StartLayout 1st Row NAD Superscript plus Baseline plus 2 e Superscript minus Baseline plus 2 upper H Superscript plus Baseline right-arrow Overscript Endscripts NADH plus upper H Superscript plus Baseline 2nd Row NADP Superscript plus Baseline plus 2 e Superscript minus Baseline plus 2 upper H Superscript plus Baseline right-arrow Overscript Endscripts NADPH plus upper H Superscript plus EndLayout

Reduction of NAD+NAD Superscript plus or NADP+NADP Superscript plus converts the benzenoid ring of the nicotinamide moiety (with a fixed positive charge on the ring nitrogen) to the quinonoid form (with no charge on the nitrogen). The reduced nucleotides absorb light at 340 nm; the oxidized forms do not (Fig. 13-24b). Biochemists use this difference in absorption to assay reactions involving these coenzymes. Note that the plus sign in the abbreviations NAD+NAD Superscript plus and NADP+NADP Superscript plus does not indicate the net charge on these molecules (in fact, both are negatively charged); rather, it indicates that the nicotinamide ring is in its oxidized form, with a positive charge on the nitrogen atom. In the abbreviations NADH and NADPH, the “H” denotes the added hydride ion. To refer to these nucleotides without specifying their oxidation state, we use NAD and NADP.

The total concentration of NAD++NADHNAD Superscript plus Baseline plus NADH in most tissues is about 105M10 Superscript negative 5 Baseline upper M; that of NADP++NADPHNADP Superscript plus Baseline plus NADPH is about 106M10 Superscript negative 6 Baseline upper M. In many cells and tissues, the ratio of NAD+NAD Superscript plus (oxidized) to NADH (reduced) is high, favoring hydride transfer from a substrate to NAD+NAD Superscript plus to form NADH. By contrast, NADPH is generally present at a higher concentration than NADP+NADP Superscript plus, favoring hydride transfer from NADPH to a substrate. This reflects the specialized metabolic roles of the two coenzymes: NAD+NAD Superscript plus generally functions in oxidations — usually as part of a catabolic reaction; NADPH is the usual coenzyme in reductions — nearly always as part of an anabolic reaction. A few enzymes can use either coenzyme, but most show a strong preference for one over the other. Also, the processes in which these two cofactors function are segregated in eukaryotic cells: for example, oxidations of fuels such as pyruvate, fatty acids, and α-keto acids derived from amino acids occur in the mitochondrial matrix, whereas reductive biosynthetic processes such as fatty acid synthesis take place in the cytosol. This functional and spatial specialization allows a cell to maintain two distinct pools of electron carriers, with two distinct functions.

More than 200 enzymes are known to catalyze reactions in which NAD+NAD Superscript plus (or NADP+NADP Superscript plus) accepts a hydride ion from a reduced substrate, or NADPH (or NADH) donates a hydride ion to an oxidized substrate. The general reactions are

AH2+NAD+A+NADH+H+A+NADPH+H+AH2+NADP+StartLayout 1st Row AH Subscript 2 Baseline plus NAD Superscript plus Baseline right-arrow Overscript Endscripts upper A plus NADH plus upper H Superscript plus Baseline 2nd Row upper A plus NADPH plus upper H Superscript plus Baseline right-arrow Overscript Endscripts AH Subscript 2 Baseline plus NADP Superscript plus EndLayout

where AH2AH Subscript 2 is the reduced substrate and A is the oxidized substrate. The general name for an enzyme of this type is oxidoreductase; they are also commonly called dehydrogenases. For example, alcohol dehydrogenase catalyzes the first step in the catabolism of ethanol, in which ethanol is oxidized to acetaldehyde:

CH3CH2OHEthanol +NAD+CH3CHO+NADH+H+AcetaldehydeCH Subscript 3 Baseline CH Subscript 2 Baseline OH Underscript Ethanol Endscripts plus NAD Superscript plus Baseline right-arrow Overscript Endscripts CH Subscript 3 Baseline CHO plus NADH plus upper H Superscript plus Baseline Underscript Acetaldehyde Endscripts

Notice that one of the carbon atoms in ethanol has lost a hydrogen; the compound has been oxidized from an alcohol to an aldehyde (refer again to Fig. 13-22 for the oxidation states of carbon).

The association between a dehydrogenase and NAD or NADP is relatively loose; the coenzyme readily diffuses from one enzyme to another, acting as a water-soluble carrier of electrons from one metabolite to another. For example, in the production of alcohol during fermentation of glucose by yeast cells, a hydride ion is removed from glyceraldehyde 3-phosphate by one enzyme (glyceraldehyde 3-phosphate dehydrogenase) and transferred to NAD+NAD Superscript plus. The NADH produced then leaves the enzyme surface and diffuses to another enzyme (alcohol dehydrogenase), which transfers a hydride ion to acetaldehyde, producing ethanol:

(1)Glyceraldehyde 3-phosphate+NAD+     3-phosphoglycerate+NADH+H+(2)Acetaldehyde+NADH+H+ ethanol+NAD+ _Sum:Glyceraldehyde 3-phosphate+acetaldehyde   3-phosphoglycerate+ethanolStartLayout 1st Row left-parenthesis 1 right-parenthesis Glyceraldehyde 3 hyphen phosphate plus NAD Superscript plus Baseline right-arrow Overscript Endscripts 2nd Row 3 hyphen phosphoglycerate plus NADH plus upper H Superscript plus Baseline 3rd Row ModifyingBelow left-parenthesis 2 right-parenthesis Acetaldehyde plus NADH plus upper H Superscript plus Baseline right-arrow Overscript Endscripts ethanol plus NAD Superscript plus Baseline With bar 4th Row upper S u m colon Glyceraldehyde 3 hyphen phosphate plus acetaldehyde right-arrow Overscript Endscripts 5th Row 3 hyphen phosphoglycerate plus ethanol EndLayout

Notice that in the overall reaction there is no net production or consumption of NAD+NAD Superscript plus or NADH; the coenzymes function catalytically and are recycled repeatedly without a net change in the total amount of NAD++NADHNAD Superscript plus Baseline plus NADH.

Both reduced and oxidized forms of NAD and NADP serve as allosteric effectors of proteins in catabolic pathways. As we describe in later chapters, the ratios NAD+/NADHNAD Superscript plus Baseline slash NADH and NADP+/NADPHNADP Superscript plus Baseline slash NADPH serve as sensitive gauges of a cell’s fuel supply, allowing rapid, appropriate changes in energy-yielding and energy-dependent metabolism.

NAD Has Important Functions in Addition to Electron Transfer

Some key cellular functions are regulated by enzymes that use NAD+NAD Superscript plus not as a redox cofactor but as a substrate in a coupled reaction in which the availability of NAD+NAD Superscript plus can be an indicator of the cell’s energy status. In DNA replication and repair, the enzyme DNA ligase is adenylylated and then transfers the AMP to a 55 prime phosphate in a nicked DNA (see Fig. 25-15); in bacteria, NAD+NAD Superscript plus serves as the source of the activating AMP group. A family of proteins called sirtuins regulate the activity of proteins in diverse cellular pathways by deacetylating the εepsilon-amino group of an acetylated Lys residue. The deacetylation is coupled to NAD+NAD Superscript plus hydrolysis, yielding O-acetyl-ADP-ribose and nicotinamide. Among the cellular processes regulated by sirtuins are inflammation, apoptosis, aging, and DNA transcription; deacetylation by a sirtuin alters the charge on histones, influencing which genes are expressed. The availability of NAD+NAD Superscript plus for these types of reactions may indicate that the cell is undergoing stress and that pathways designed to respond to stress should be activated.

NAD+NAD Superscript plus also plays an important role in cholera pathology (see Section 12.2). Cholera toxin has an enzymatic activity that transfers ADP-ribose from NAD+NAD Superscript plus to a G protein involved in regulating ion fluxes in the cells lining the gut. This ADP-ribosylation blocks water retention, causing the diarrhea and dehydration characteristic of cholera.

Dietary deficiency of niacin, the vitamin form of NAD and NADP, causes pellagra (Fig. 13-25). The pyridine-like rings of NAD and NADP are derived from the vitamin niacin (nicotinic acid; Fig. 13-26), which is synthesized from tryptophan. Humans generally cannot synthesize sufficient quantities of niacin, and this is especially so for individuals with diets low in tryptophan (maize, for example, has a low tryptophan content). Niacin deficiency, which affects all the NAD(P)-dependent dehydrogenases, causes the serious human disease pellagra (Italian for “rough skin”) and a related disease in dogs, called black tongue. Pellagra is characterized by the “three Ds”: dermatitis, diarrhea, and dementia, followed in many cases by death. A century ago, pellagra was a common human disease; in the southern United States, where maize was a dietary staple, about 100,000 people were afflicted and about 10,000 died as a result of this disease between 1912 and 1916. In 1920, Joseph Goldberger showed pellagra to be caused by a dietary insufficiency, and in 1937, Frank Strong, D. Wayne Woolley, and Conrad Elvehjem identified niacin as the curative agent for the dog version of pellagra, black tongue. Supplementation of the human diet with this inexpensive compound has nearly eradicated pellagra in the populations of the developed world, with one significant exception: people who drink excessive amounts of alcohol. In these individuals, intestinal absorption of niacin is much reduced, and caloric needs are often met with distilled spirits that are virtually devoid of vitamins, including niacin.

A photo shows a hand and lower arm with many dark spots. The skin suddenly becomes darker just above the wrist and is dark along the hand. It is lighter on the fingers and red on the knuckles. The veins are very dark on the lower wrist and hand.

FIGURE 13-25 Dermatitis associated with pellagra. Dermatitis involving the face, hands, and feet is an early sign of pellagra, a serious human disease that results from insufficient niacin in the diet. Untreated, pellagra leads to dementia and ultimately is fatal.

A figure shows the structures of niacin and three derivative nucleotides: nicotine, nicotinamide, and tryptophan.

FIGURE 13-26 Niacin (nicotinic acid) and its derivative nicotinamide. The biosynthetic precursor of these compounds is tryptophan. In the laboratory, nicotinic acid was first produced by oxidation of the natural product nicotine—thus the name. Both nicotinic acid and nicotinamide cure pellagra, but nicotine (from cigarettes or elsewhere) has no curative activity.

Flavin Nucleotides Are Tightly Bound in Flavoproteins

Flavoproteins are enzymes that catalyze oxidation-reduction reactions using either flavin mononucleotide (FMN) or flavin adenine dinucleotide (FAD) as coenzyme (Fig. 13-27). These coenzymes, the flavin nucleotides, are derived from the vitamin riboflavin. The fused ring structure of flavin nucleotides (the isoalloxazine ring) undergoes reversible reduction, accepting either one or two electrons in the form of one or two hydrogen atoms (each atom an electron plus a proton) from a reduced substrate. The fully reduced forms are abbreviated FADH2FADH Subscript 2 and FMNH2FMNH Subscript 2. When a fully oxidized flavin nucleotide accepts only one electron (one hydrogen atom), the semiquinone form of the isoalloxazine ring is produced, abbreviated FADHFADH Superscript bullet and FMNHFMNH Superscript bullet. Because flavin nucleotides have a chemical specialty that is slightly different from that of the nicotinamide coenzymes — the ability to participate in either one- or two-electron transfers — flavoproteins are involved in a greater diversity of reactions than the NAD(P)-linked dehydrogenases.

A figure shows the oxidized and reduced structures of F A D and F M N.

FIGURE 13-27 Oxidized and reduced FAD and FMN. FMN consists of the structure above the dashed red line across the FAD molecule (oxidized form). The flavin nucleotides accept two hydrogen atoms (two electrons and two protons), both of which appear in the flavin ring system (isoalloxazine ring). When FAD or FMN accepts only one hydrogen atom, the semiquinone, a stable free radical, forms.

Like the nicotinamide coenzymes (Fig. 13-24), the flavin nucleotides undergo a shift in a major absorption band on reduction (again, useful to biochemists who want to monitor reactions involving these coenzymes). Flavoproteins that are fully reduced (two electrons accepted) generally have an absorption maximum near 360 nm. When partially reduced (one electron), they acquire another absorption maximum at about 450 nm; when fully oxidized, the flavin has maxima at 370 and 440 nm.

The flavin nucleotide in most flavoproteins is bound rather tightly to the protein, and in some enzymes, such as succinate dehydrogenase, it is bound covalently. Such tightly bound coenzymes are properly called prosthetic groups. They do not transfer electrons by diffusing from one enzyme to another; rather, they provide a means by which the flavoprotein can temporarily hold electrons while it catalyzes electron transfer from a reduced substrate to an electron acceptor. One important feature of the flavoproteins is the variability in the standard reduction potential (E°)left-parenthesis upper E prime degree right-parenthesis of the bound flavin nucleotide. Tight association between the enzyme and prosthetic group confers on the flavin ring a reduction potential typical of that particular flavoprotein, sometimes quite different from the reduction potential of the free flavin nucleotide. FAD bound to succinate dehydrogenase, for example, has an E°upper E prime degree close to 0.0 V, compared with 0.219 Vnegative 0.219 upper V for free FAD; E°upper E prime degree for other flavoproteins ranges from 0.40 V to+0.06 V.negative 0.40 upper V to plus 0.06 upper V period Flavoproteins are often very complex; some have, in addition to a flavin nucleotide, tightly bound inorganic ions (iron or molybdenum, for example) capable of participating in electron transfers.

We examine the function of flavoproteins as electron carriers in Chapters 19 and 20, when we consider their roles in oxidative phosphorylation (in mitochondria) and photophosphorylation (in chloroplasts).

SUMMARY 13.4 Biological Oxidation-Reduction Reactions