1X48-18 MANUAL OF METHODS FOR PURE CULTURE STLTDY 

 resented by the hypothetical "half-reactions" 



Cl2+2e — ^ 2 Cl- 



l2+2e 2 I- 



-•; 



to show the participation of electrons. In the interaction, chlorine 

 is the electron-acceptor, and iodide the electron-donor. 



The chlorine, the iodine and a considerable number of other sys- 

 tems can be studied by means of electric cells in which such systems 

 can display their relative oxidation-reduction tendencies in terms of 

 electrode potentials. The latter permit evaluation of the change in 

 Gibbs free energy (see later) in the interaction of any two such 

 oxidation-reduction systems. 



Without going into details of derivation or refinements, we may state that the elec- 

 trode equation for a reversible oxidation-reduction system has the general form: 



RT [Reductant] / <• .• r tt j \ 

 p _y 7 \ I a lunction 01 pH and i ,„x 



t;, rrk -J ii I dissociation constants I ^ ' 



nt [OxidantJ \ / 



where Eh is the potential, in volts, referred to that of the normal hydrogen electrode; 

 Eo is a constant characteristic of the system at pH 0; R is the gas constant, 8.315 volt- 

 coulombs per degree per mole; T is the absolute temperature; n is the number of elec- 

 trons involved in the oxidation-reduction process; F is the faraday (96500 coulombs); 

 In is the logarithm to the base e; and brackets represent concentrations of the reduc- 

 tant and oxidant. At any fixed pH, the first and last terms on the right side of the 

 above equation may be combined as a constant, E'o, then, 



RT [Reductant] 



Eh = E'o In (7) 



nF [Oxidant] 



That is, Eh = E'o at anj- fixed pH when [Reductant] = [Oxidant]. 



It is apparent from equation 6, that the potential of such a system may be influenced 

 by the pH of the solution; and the potential of one system may vary relative to that of 

 another as the pH is varied. In fact, cases are known where system A can oxidize 

 system B at one pH level, and system B oxidize system A at another. Hence the 

 importance of comparing such potentials at the same pH, as well as the same tempera- 

 ture, and the desirability of specifying pH in connection with a statement of the Eh 

 of a system. 



Elaboration of the theory of reversible oxidation-reduction potentials can be found 

 in Clark (1928, 1948), Clark, Cohen, et al. (1928), and modern texts on electrochemistry, 

 such as Glasstone (1942). 



There are two methods of measuring oxidation-reduction potentials, the potentio- 

 metric method and the colorimetric. Each has its advantages and disadvantages; 

 but the potentiometric method is generally preferable for reasons that will appear 

 below. In either case, it is usually necessary to deaerate the container and the solution 

 to be measured by evacuation or by displacing gaseous and dissolved oxygen with an 

 inert gas such as purified nitrogen. Deoxygenation is often accomplished spontaneous- 

 ly in the depths of an actively growing culture of facultative bacteria. 



