Oxidative Phosphorylation Depends on Electron Transfer

II. Transducing and Storing Energy
18. Oxidative Phosphorylation
18.1. Oxidative Phosphorylation in Eukaryotes Takes Place in Mitochondria
Figure 18.5. Overlapping Gene Complements of Mitochondria. The genes present within each oval are those present
within the organism represented by the oval. Only rRNA- and protein-coding genes are shown. The genome of
Reclinomonas contains all the protein-coding genes found in all the sequenced mitochondrial genomes. [After M. W.
Gray, G. Burger, and B. F. Lang. Science 283(1999): 1476 1481.]
II. Transducing and Storing Energy
18. Oxidative Phosphorylation
18.2. Oxidative Phosphorylation Depends on Electron Transfer
In Chapter 17, the primary function of the citric acid cycle was identified as the generation of NADH and FADH2 by the
oxidation of acetyl CoA. In oxidative phosphorylation, NADH and FADH2 are used to reduce molecular oxygen to
water. The highly exergonic reduction of molecular oxygen by NADH and FADH2 occurs in a number of electrontransfer reactions, taking place in a set of membrane proteins known as the electron-transport chain.
18.2.1. High-Energy Electrons: Redox Potentials and Free-Energy Changes
High-energy electrons and redox potentials are of fundamental importance in oxidative phosphorylation. In oxidative
phosphorylation, the electron transfer potential of NADH or FADH2 is converted into the phosphoryl transfer potential
of ATP. We need quantitative expressions for these forms of free energy. The measure of phosphoryl transfer potential is
already familiar to us: it is given by ∆ G°´ for the hydrolysis of the activated phosphate compound. The corresponding
expression for the electron transfer potential is E´ 0, the reduction potential (also called the redox potential or oxidationreduction potential).
The reduction potential is an electrochemical concept. Consider a substance that can exist in an oxidized form X and a
reduced form X-. Such a pair is called a redox couple. The reduction potential of this couple can be determined by
measuring the electromotive force generated by a sample half-cell connected to a standard reference half-cell (Figure
18.6). The sample half-cell consists of an electrode immersed in a solution of 1 M oxidant (X) and 1 M reductant (X-).
The standard reference half-cell consists of an electrode immersed in a 1 M H+ solution that is in equilibrium with H2
gas at 1 atmosphere pressure. The electrodes are connected to a voltmeter, and an agar bridge establishes electrical
continuity between the half-cells. Electrons then flow from one half-cell to the other. If the reaction proceeds in the
direction
the reactions in the half-cells (referred to as half-reactions or couples) must be
Thus, electrons flow from the sample half-cell to the standard reference half-cell, and the sample-cell electrode is taken
to be negative with respect to the standard-cell electrode. The reduction potential of the X:X - couple is the observed
+
voltage at the start of the experiment (when X, X-, and H+ are 1 M). The reduction potential of the H :H couple is
2
defined to be 0 volts.
The meaning of the reduction potential is now evident. A negative reduction potential means that the reduced form of a
substance has lower affinity for electrons than does H2, as in the preceding example. A positive reduction potential
means that the reduced form of a substance has higher affinity for electrons than does H2. These comparisons refer to
standard conditions namely, 1 M oxidant, 1 M reductant, 1 M H+, and 1 atmosphere H2. Thus, a strong reducing agent
(such as NADH) is poised to donate electrons and has a negative reduction potential, whereas a strong oxidizing agent
(such as O ) is ready to accept electrons and has a positive reduction potential.
2
The reduction potentials of many biologically important redox couples are known (Table 18.1). Table 18.1 is like those
presented in chemistry texts except that a hydrogen ion concentration of 10-7 M (pH 7) instead of 1 M (pH 0) is the
standard state adopted by biochemists. This difference is denoted by the prime in E´ 0. Recall that the prime in ∆ G°´
denotes a standard free-energy change at pH 7.
The standard free-energy change ∆ G°´ is related to the change in reduction potential ∆ E´ 0 by
in which n is the number of electrons transferred, F is a proportionality constant called the faraday [23.06 kcal mol-1 V-1
(96.48 kJ mol-1 V-1)], ∆ E´ 0 is in volts, and ∆ G°´ is in kilocalories or kilojoules per mole.
The free-energy change of an oxidation-reduction reaction can be readily calculated from the reduction potentials of the
reactants. For example, consider the reduction of pyruvate by NADH, catalyzed by lactate dehydrogenase.
The reduction potential of the NAD+:NADH couple, or half-reaction, is -0.32 V, whereas that of the pyruvate: lactate
couple is -0.19 V. By convention, reduction potentials (as in Table 18.1) refer to partial reactions written as reductions:
oxidant + ereductant. Hence,
To obtain reaction a from reactions b and c, we need to reverse the direction of reaction c so that NADH appears on the
left side of the arrow. In doing so, the sign of E´0 must be changed.
For reaction b, the free energy can be calculated with n = 2.
Likewise, for reaction d,
Thus, the free energy for reaction a is given by
18.2.2. A 1.14-Volt Potential Difference Between NADH and O2 Drives Electron
Transport Through the Chain and Favors the Formation of a Proton Gradient
The driving force of oxidative phosphorylation is the electron transfer potential of NADH or FADH2 relative to that of
O2. How much energy is released by the reduction of O2 with NADH? Let us calculate ∆ G°´ for this reaction. The
pertinent half-reactions are
Subtracting reaction b from reaction a yields
The standard free energy for this reaction is then given by
This is a substantial release of free energy. Recall that ∆ G°´ = - 7.5 kcal mol-1 ( - 31.4 kJ mol-1) for the hydrolysis of
ATP. The released energy is used initially to generate a proton gradient that is then used for the synthesis of ATP and the
transport of metabolites across the mitochondrial membrane.
How can the energy associated with a proton gradient be quantified? Recall that the free-energy change for a species
moving from one side of a membrane where it is at concentration c 1 to the other side where it is at a concentration c 2 is
given by
in which Z is the electrical charge of the transported species and ∆ V is the potential in volts across the membrane
(Section 13.1.2). Under typical conditions for the inner mitochondrial membrane, the pH outside is 1.4 units lower than
inside [corresponding to log10 (c 2/c 1) of 1.4] and the membrane potential is 0.14 V, the outside being positive. Because
Z = +1 for protons, the free-energy change is (2.303 × 1.98 × 10-3 kcal mol-1 K-1 × 310 K × 1.4) + ( + 1 × 23.06 kcal mol1 V-1 × 0.14 V) = 5.2 kcal mol-1 (21.8 kJ mol-1). Thus, each proton that is transported out of the matrix to the cytosolic
side corresponds to 5.2 kcal mol-1 of free energy.
18.2.3. Electrons Can Be Transferred Between Groups That Are Not in Contact
As will be discussed shortly, the electron-carrying groups in the protein constituents of the electron-transport chain are
flavins, iron-sulfur clusters, quinones, hemes, and copper ions. How are electrons transferred between electron-carrying
groups that are frequently buried in the interior of a protein in fixed positions and are therefore not directly in contact?
Electrons can move through space, even through a vacuum. However, the rate of electron transfer through space falls off
rapidly as the electron donor and electron acceptor move apart from each other, decreasing by a factor of 10 for each
increase in separation of 0.8 Å. The protein environment provides more-efficient pathways for electron conduction:
typically, the rate of electron transfer decreases by a factor of 10 every 1.7 Å (Figure 18.7). For groups in contact,
electron-transfer reactions can be quite fast with rates of approximately 1013 s-1. Within proteins in the electron-transport
chain, electron-carrying groups are typically separated by 15 Å beyond their van der Waals contact distance. For such
separations, we expect electron-transfer rates of approximately 104 s-1 (i.e., electron transfer in less than 1 ms), assuming
that all other factors are optimal. Without the mediation of the protein, an electron transfer over this distance would take
approximately 1 day.
Another important factor in determining the rate of electron transfer is the driving force, the free-energy change
associated with the reaction (Figure 18.8). Like the rates of most reactions, those of electron-transfer reactions tend to
increase as the free-energy change for the reaction becomes more favorable. Interestingly, however, each electrontransfer reaction has an optimal driving force; making the reaction more favorable beyond this point decreases the rate of
the electron-transfer process. This so-called inverted region is of tremendous importance for the light reactions of
photosynthesis, to be discussed in Chapter 19. For the purposes of the electron-transport chain, the effects of distance
and driving force combine to determine which pathway, among the set of those possible, will be used at each stage in the
course of a reaction.
II. Transducing and Storing Energy
18. Oxidative Phosphorylation
18.2. Oxidative Phosphorylation Depends on Electron Transfer
Figure 18.6. Measurement of Redox Potential. Apparatus for the measurement of the standard oxidation-reduction
potential of a redox couple. Electrons, but not X or X-, can flow through the agar bridge.
II. Transducing and Storing Energy
18. Oxidative Phosphorylation
18.2. Oxidative Phosphorylation Depends on Electron Transfer
Table 18.1. Standard reduction potentials of some reactions
Oxidant
Reductant
n
E´ 0 (V)
Succinate + CO2
α-Ketoglutarate
2
- 0.67
Acetate
Ferredoxin (oxidized)
2 H+
Acetaldehyde
Ferredoxin (reduced)
H2
2
1
2
- 0.60
- 0.43
- 0.42
NAD+
NADH + H+
2
- 0.32
NADP+
Lipoate (oxidized)
Glutathione (oxidized)
FAD
NADPH + H+
Lipoate (reduced)
Glutathione (reduced)
FADH2
2
- 0.32
2
2
2
- 0.29
- 0.23
- 0.22
Acetaldehyde
Pyruvate
Fumarate
Cytochrome b (+3)
Dehydroascorbate
Ubiquinone (oxidized)
Cytochrome c (+3)
Fe (+3)
Ethanol
Lactate
Succinate
Cytochrome b (+2)
Ascorbate
Ubiquinone (reduced)
Cytochrome c (+2)
Fe (+2)
H2O
2
2
2
1
2
2
1
1
2
- 0.20
- 0.19
0.03
0.07
0.08
0.10
0.22
0.77
0.82
1/
2
O2+ 2 H+
Note: E´ is the standard oxidation-reduction potential (pH 7, 25°C) and n is the number of electrons transferred. E´ refers to the
0
partial reaction written as
0