NAD+ and FAD Couple Many Biological Oxidation and Reduction Reactions

In many chemical reactions, electrons are transferred from one atom or molecule to another; this transfer may or may not accompany the formation of new chemical bonds or the release of energy that can be coupled to other reactions. The loss of electrons from an atom or a molecule is called oxidation, and the gain of electrons by an atom or a molecule is called reduction. An example of oxidation is the removal of electrons from the sulfhydryl group–containing side chains of two cysteine amino acids to form a disulfide bond, described above in Section 2.2. Electrons are neither created nor destroyed in a chemical reaction, so if one atom or molecule is oxidized, another must be reduced. For example, oxygen draws electrons from Fe2+ (ferrous) ions to form Fe3+ (ferric) ions, a reaction that occurs as part of the process by which carbohydrates are degraded in mitochondria. Each oxygen atom receives two electrons, one from each of two Fe2+ ions:

2 Fe2+ + ½ O2 → 2 Fe2+ + O2–

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FIGURE 2-32 Conversion of succinate to fumarate. In this oxidation reaction, which occurs in mitochondria as part of the citric acid cycle, succinate loses two electrons and two protons. These protons and electrons are transferred to FAD, reducing it to FADH2.

Thus Fe2+ is oxidized and O2 is reduced. Such reactions in which one molecule is reduced and another is oxidized are often referred to as redox reactions. Oxygen is an electron acceptor in many redox reactions in cells under aerobic conditions.

Many biologically important oxidation and reduction reactions involve the removal or addition of hydrogen atoms (protons plus electrons) rather than the transfer of isolated electrons on their own. The oxidation of succinate to fumarate, which occurs in mitochondria, is an example (Figure 2-32). Protons are soluble in aqueous solutions (as H3O+), but electrons are not, so they must be transferred directly from one atom or molecule to another without a water-dissolved intermediate. In this type of oxidation reaction, electrons are often transferred to small electron-carrying molecules, sometimes referred to as coenzymes. The most common of these electron carriers are NAD+ (nicotinamide adenine dinucleotide), which is reduced to NADH, and FAD (flavin adenine dinucleotide), which is reduced to FADH2 (Figure 2-33). The reduced forms of these coenzymes can transfer protons and electrons to other molecules, thereby reducing them.

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FIGURE 2-33 The electron-carrying coenzymes NAD+ and FAD. (a) NAD+ (nicotinamide adenine dinucleotide) is reduced to NADH by the addition of two electrons and one proton simultaneously. In many biological redox reactions, a pair of hydrogen atoms (two protons and two electrons) is removed from a molecule. In some cases, one of the protons and both electrons are transferred to NAD+; the other proton is released into solution. (b) FAD (flavin adenine dinucleotide) is reduced to FADH2 by the addition of two electrons and two protons, as occurs when succinate is converted to fumarate (see Figure 2-32). In this two-step reaction, addition of one electron together with one proton first generates a short-lived semiquinone intermediate (not shown), which then accepts a second electron and proton.

To describe redox reactions, such as the reaction of ferrous ion (Fe2+) and oxygen (O2), it is easiest to divide them into two half-reactions:

Oxidation of Fe2+: 2 Fe2+ → 2 Fe3+ + 2 e–

Reduction of O2: 2 e– + ½ O2 → O2–

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In this case, the reduced oxygen (O2–) readily reacts with two protons to form one water molecule (H2O). The readiness with which an atom or a molecule gains an electron is its reduction potential (E). The tendency to lose electrons, the oxidation potential, has the same magnitude as the reduction potential for the reverse reaction, but has the opposite sign.

Reduction potentials are measured in volts (V) from an arbitrary zero point set at the reduction potential of the following half-reaction under standard conditions (25 °C, 1 atm, and reactants at 1 M):

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The value of E for a molecule or an atom under standard conditions is its standard reduction potential, E′0. A molecule or an ion with a positive E′0 has a higher affinity for electrons than the H+ ion does under standard conditions. Conversely, a molecule or ion with a negative E′0 has a lower affinity for electrons than the H+ ion does under standard conditions. Like the values of ΔG°′, standard reduction potentials may differ somewhat from those found under the conditions in a cell because the concentrations of reactants in a cell are not 1 M.

In a redox reaction, electrons move spontaneously toward atoms or molecules having more positive reduction potentials. In other words, a molecule having a more negative reduction potential can transfer electrons spontaneously to, or reduce, a molecule with a more positive reduction potential. In this type of reaction, the change in electric potential ΔE is the sum of the reduction and oxidation potentials for the two half-reactions. The ΔE for a redox reaction is related to the change in free energy ΔG by the following expression:

ΔG (cal/mol) = –n (23, 064) ΔE (volts) (2-11)

where n is the number of electrons transferred. Note that a redox reaction with a positive ΔE value will have a negative ΔG and thus will tend to proceed spontaneously from left to right.