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The tendency of acids and bases to donate or accept protons was described. In acid-base systems, one compound acts as the proton (H+) donor and the other as the proton acceptor, the donor being the acid and the acceptor the base.
The two form a conjugate acid-base pair. In a similar manner, oxidizing and reducing agents function in pairs.
In this case, they are called redox pairs or redox couples. The member of the pair that donates the electron is called the reducing agent or reductant and the electron acceptor is called the oxidizing agent or oxidant.
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The terms donating and accepting fail to convey the nature of the amounts or energy change involved in the reaction. In effect, the reducing agent has a certain capability to retain electrons, as does the oxidizing agent. In a redox couple, one member attracts electrons more strongly than the other, and in effect the oxidizing agent can pull the electrons away from the reducing agent. This capability to gain (or lose) electrons can be measured and is called the oxidation- reduction potential or redox potential and is expressed in volts.
For convenience, the potential of most redox couples is measured against a reference standard, which usually is a hydrogen electrode. Most measurements are made under standard conditions using 1.0 M concentrations of oxidant and reductant at 25 °C and at pH 7.0. The hydrogen electrode is equilibrated with H2 gas at 1 atmosphere and the [H+] is 1.0 M at 25°C.
When the pH is adjusted to 7.0 (i.e., [H+] = 10-7 M), the redox potential of the reference hydrogen electrode is – 0.42 V. The redox potentials of the electron transport system intermediates are given in Table 16-2. The greater the E’Q value, the more strongly the oxidant accepts electrons. However, an oxidant having a lower Eq than another compound becomes the reductant of or electron donor to that compound.
The Henderson-Hasselbach equation describes the relationship between pH and the dissociation constant. In a similar way, the Nernst equation shows for a redox pair the relationship between the standard redox potential (Eq), the observed potential at any concentration, and the concentration ratio of the oxidant and reductant.
Where E0‘ is the standard redox potential, Eh is the observed electrode potential, R is the gas constant (8.31 J/deg/mole, T is the absolute temperature in degrees Kelvin, n is the number of electrons transferred, and F is the faraday (96,406 kJ/V). When electrons are transferred and the constants are combined, the Nernst equation becomes
Where n = the number of electrons involved.
