A “redox tower” showing the redox potential of common metabolic half reactions. Metabolic processes can be seen as moving electrons between molecules, often capturing some of the energy released as the electrons move from high energy to lower energy states as in glycolysis or respiration. Electrons donated by the “half-reactions” on top can be consumed in a half-reaction lower on the tower to complete a thermodynamically favorable reaction. For example, the net process of glycolysis involves the oxidation of glucose to pyruvate coupled to the reduction of NAD+ to NADH. Since the oxidation of glucose lies at the top of the tower and the reduction of NAD+ is below it, this electron flow is thermodynamically favorable. Comparing to the ATP hydrolysis scale bar we can also see that this electron flow is favorable enough to generate ATP. Aerobic respiration involves many intermediate electron transfers through the electron transport chain. Several of these transitions are shown, including the oxidation succinate to fumarate which is mechanistically coupled to the reduction of ubiquinone to ubiquinol in the inner mitochrondrial membranes. Each of these intermediate electron transfers must be thermodynamically favorable on its own in order for respiration to proceed. By comparing to the “ATP hydrolysis scale” we can see that the individual transformations in the electron transport chain are not energetic enough to generate ATP on their own. Yet they are favorable enough to pump a proton across the cell or mitochondrial membrane. This is the energetic basis for chemiosmosis: cells store quanta of energy too small for ATP synthesis in the proton gradient across a membrane. That energy is later used to generate ATP by converting the H+ gradient into phosphoanhydride bonds on ATP through the ATP synthase.

Redox potentials are used to characterize the free energy cost and direction of reactions involving electron transfer, one of the most ubiquitous and important of biochemical reactions. Such reduction-oxidation reactions are characterized by a free energy change that shares some conceptual features with that used to describe pKa in acid-base reactions where proton transfer is involved rather than electron transfer. In this vignette, one of the most abstract in the book, we discuss how the redox potential can be used as a measure of the driving force for a given oxidation-reduction reaction of interest.  By way of contrast, unlike the pH, there is no sense in which one can assign a single redox potential to an entire cell.

The redox potential, or more accurately the reduction potential, of a compound refers to its tendency to acquire electrons and thereby to be reduced. Some readers might remember the mnemonic “OILRIG” which reminds us that “oxidation is loss, reduction is gain”, where the loss and gain are of electrons. Consider a reaction that involves an electron transfer: Aox + ne ↔ Ared where n electrons are taken up by the oxidized form (Aox) to give the reduced form (Ared) of compound A. The redox potential difference ΔE between the electron donor and acceptor is related to the associated free energy change ΔG of the reaction via ΔG=nFΔE where n is the number of electrons transferred and F is Faraday’s constant (96,485 J/mol/V or ≈100 kJ/mol/V). By inspecting tabulated values of these potentials, it is possible to develop an intuition for the tendency for electron transfer and hence, of the direction of the reaction.

Though ATP is often claimed to be the energy currency of the cell, in fact,  for the energetic balance of the cell the carriers of reducing power are themselves no less important. The most important example of these carriers is the molecule NADH in its reduced or oxidized (NAD+) forms. We can use the redox potential to connect these two molecular protagonists, and estimate an upper bound on the number of ATP molecules that can be produced from the oxidation of NADH (produced, for example, in the TCA cycle). The NAD+/NADH pair has a redox potential of E = -0.32 V and it is oxidized by oxygen to give water (protons coming from the media) with a redox potential of E = +0.82 V. Both are shown in Figure 1 as part of a “redox tower” of key biological half reactions that can be linked to find the overall redox potential change and thus the free energy. For the reaction considered above of NADH oxidation by oxygen, the maximal associated free energy that can be extracted is thus

One confusing aspect of redox reactions is that the transfer can take several forms. In one case it is only electrons as in the reactions carried out by cytochromes in electron transfer chains. In another common case it is a combination of electrons and protons as in the cofactor NAD+/NADH where two electrons and one proton (H+) are transferred. Finally, there are the reactions where the same number of electrons and protons is transferred when one would naturally be tempted to discuss transfer of hydrogens. This is for example the case for the overall reaction of glucose oxidation where oxygen is reduced to water. Two hydrogens have thus been transferred, so should one discuss the transfer of electrons, hydrogens or protons? The definition of the redox potential (given above) focuses only on the electron “state”. What about the protons and what happens to these when one encounters a chain of electron transfer reactions where some intermediate compounds contain the hydrogen protons and some do not? The explanation resides in the surrounding water and their pH. The reaction occurs at a given pH, and the reacting compounds are in equilibrium with this pH and thus giving off or receiving a proton has no effect on the energetics. The aqueous medium serves as a pool where protons can be “parked” when the transfer reaction is solely of electrons (the analogy borrowed from the very accessible introductory biochemistry book “The chemistry of life” by Steven Rose). These parked protons can be borrowed back at subsequent stages as occurs in the final stage of oxidative respiration where cytochrome oxidase takes protons from the medium. Because one assumes that water is ubiquitous one does not need to account for protons except for knowing the prevailing pH which depicts the tendency to give or receive protons. This is the reason why we discuss electron donors and acceptors rather than hydrogen donors and acceptors.

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