THE MEASUREMENT OF PH AND TITRATABLE ACIDITY 91 



The chlorine, the iodine, and a considerable number of other systems 

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

 their relative oxidation-reduction tendencies in terms of electrode poten- 

 tials. 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: 



F — F — ^^] [r^<l^<^tant] / a function of pH and \ 

 nF [oxidant] \dissociation constants/ 



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; -K 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 



„ j^, i2T [reductant] , 



^'=^^--^^'' [oxidant] ^^^ 



That is. Eh — E'o at any fixed pH when [reductant] = [oxidant]. 



It is apparent from Eq. (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 knovm 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 electrochemis- 

 try, 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 solu- 

 tion to be measured by evacuation or by displacing gaseous and dissolved oxygen 

 with an inert gas such as purified nitrogen. Deoxygenation is often accompHshed 

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



The Potentiometric Method 



Electrode vessel. This may be a test tube with a constriction and 

 bulb at its lower end or a more elaborate container depending on the 

 requirements of the experiment. Such vessels are described by Clark, 

 Cohen, et al. (1928), Borsook and Schott (1931), Allyn and Baldwin 

 (1932), and Hewitt (1936). 



Electrodes. A ''noble/' or "unattackable," metal is the electrode of 

 choice. A coil of bright platinum wire has been frequently employed, 



