Biochemistry: The Chemical Reactions of Living Cells, Volume 1 - D. Metzler 1980

Energetics of Biochemical Reactions
Electrode Potentials and Free Energy Changes for Oxidation-Reduction Reactions

We live in an environment containing a powerful oxidizing agent, O₂. In cellular Respiration processes, oxygen is reduced to H₂O, a very weak reducing agent. At the other end of the oxidizing scale is H⁺, a very weak oxidizing agent that certain Bacteria are capable of converting into a strong reducing agent, H₂. The O₂—H₂O and H⁺—H₂ couples are two biologically crucial redox systems. Between them lie numerous other pairs of compounds that participate in intracellular redox reactions and play a vital role in METABOLISM.

Two standard Methods are widely used to quantitatively evaluate the oxidizing and reducing properties of conjugate redox pairs. On the one hand, one can list the ∆G values for The oxidation of the reduced form by oxygen O₂ and its conversion into the oxidized form. A compound with a large negative ∆G value will be a strong reducing agent. An example is H2, for which the Free energy of combustion at pH 7 (Table 3-3) is —237 kJ∙mol^-1. Weak reducing agents, such as Fe³⁺, are characterized by small ∆G oxidation values (—8.5 kJ∙mol^-1 for the reaction 2Fe²⁺—>2Fe^3-). The free energies of oxidation for biological hydrogen carriers (which will be discussed in Chapter 8) typically fall between the aforementioned values for H₂ and Fe²⁺.

The second approach for expressing the same information involves The Use of electrode potentials (Table 3-7). Electrode potentials are also important because their direct measurement sometimes provides a way to study intracellular redox reactions.