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The values below are standard apparent red … The values below are standard apparent reduction potentials (E°') for electro-biochemical half-reactions measured at 25 °C, 1 atmosphere and a pH of 7 in aqueous solution. The actual physiological potential depends on the ratio of the reduced (Red) and oxidized (Ox) forms according to the Nernst equation and the thermal voltage. When an oxidizer (Ox) accepts a number z of electrons ( e−) to be converted in its reduced form (Red), the half-reaction is expressed as: Ox + z e− → Red The reaction quotient (Qr) is the ratio of the chemical activity (ai) of the reduced form (the reductant, aRed) to the activity of the oxidized form (the oxidant, aox). It is equal to the ratio of their concentrations (Ci) only if the system is sufficiently diluted and the activity coefficients (γi) are close to unity (ai = γi Ci): The Nernst equation is a function of Qr and can be written as follows: At chemical equilibrium, the reaction quotient Qr of the product activity (aRed) by the reagent activity (aOx) is equal to the equilibrium constant (K) of the half-reaction and in the absence of driving force (ΔG = 0) the potential (Ered) also becomes nul. The numerically simplified form of the Nernst equation is expressed as: Where is the standard reduction potential of the half-reaction expressed versus the standard reduction potential of hydrogen. For standard conditions in electrochemistry (T = 25 °C, P = 1 atm and all concentrations being fixed at 1 mol/L, or 1 M) the standard reduction potential of hydrogen is fixed at zero by convention as it serves of reference. The standard hydrogen electrode (SHE), with [ H+] = 1 M works thus at a pH = 0. At pH = 7, when [ H+] = 10−7 M, the reduction potential of H+ differs from zero because it depends on pH. Solving the Nernst equation for the half-reaction of reduction of two protons into hydrogen gas gives: 2 H+ + 2 e− ⇌ H2 In biochemistry and in biological fluids, at pH = 7, it is thus important to note that the reduction potential of the protons ( H+) into hydrogen gas H2 is no longer zero as with the standard hydrogen electrode (SHE) at 1 M H+ (pH = 0) in classical electrochemistry, but that versus the standard hydrogen electrode (SHE). The same also applies for the reduction potential of oxygen: O2 + 4 H+ + 4 e− ⇌ 2 H2O For O2, = 1.229 V, so, applying the Nernst equation for pH = 7 gives: For obtaining the values of the reduction potential at pH = 7 for the redox reactions relevant for biological systems, the same kind of conversion exercise is done using the corresponding Nernst equation expressed as a function of pH. The conversion is simple, but care must be taken not to inadvertently mix reduction potential converted at pH = 7 with other data directly taken from tables referring to SHE (pH = 0).ken from tables referring to SHE (pH = 0).
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The values below are standard apparent red … The values below are standard apparent reduction potentials (E°') for electro-biochemical half-reactions measured at 25 °C, 1 atmosphere and a pH of 7 in aqueous solution. The actual physiological potential depends on the ratio of the reduced (Red) and oxidized (Ox) forms according to the Nernst equation and the thermal voltage. When an oxidizer (Ox) accepts a number z of electrons ( e−) to be converted in its reduced form (Red), the half-reaction is expressed as: Ox + z e− → Red The Nernst equation is a function of Qr and can be written as follows: 2 H+ + 2 e− ⇌ H2 O2 + 4 H+ + 4 e− ⇌ 2 H2O 2 H+ + 2 e− ⇌ H2 O2 + 4 H+ + 4 e− ⇌ 2 H2O
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Table of standard reduction potentials for half-reactions important in biochemistry
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