OXIDATION-REDUCTION POTENTIALS 23 



and substituting the value of [e] from equation (4) above 

 we get 



^ ^^ Rl\ , RT , [Reductant] 

 ^ = ^-^ ^"S ^-^ ^"^ [Oxidant] 



„ , RT , [Reductant] , 



^ = ^^-^ l^g [Oxidant] . . • • i8) 



where A^i is another constant, smce K ^ log /c is a 



constant for any given temperature and reaction. The 

 electrode potential E can only be measured if it forms 

 one element of a cell of which the other is a standard 

 electrode, the hydrogen electrode being used as such in 

 these cases. The potential referred to the hydrogen 

 electrode as standard is denoted by Eh, and is given by 

 Eh =E—lc2 where k^ is the potential of the standard 

 hydrogen electrode. 



„ ^ , RT , [Reductant] , 



Hence Eh = A', =- log ^^ ., — — k^, 



^ nF ^ [Oxidant] '^ 



^ RT . [Reductant] ,^. 



^^ = ^--n-^^^g [Oxidant] ' ' ' ^^) 



where Eq= ki—k2, which is a constant for the system. 



It follows from a consideration of this equation that 

 the observed oxidation-reduction potential, Eh, depends 

 on Eo, which is a constant for the particular system under 

 consideration, and on the ratio of the concentrations of 

 the reduced and oxidised constituents of the system. 

 The more reduced substance there is present the lower 

 will be the Eh value, and the greater the proportion 

 of oxidised substance the higher will be the potential. 

 When the concentration of the reductant equals that of 

 the oxidant, that is, when the system is haK oxidised, 

 . . . . , „ „ . , . [Reductant] 



it IS obvious that Eh = Eo, smce the ratio tq -^^ y.fi "^ ■'■ 



and its logarithm =0. Thus if the potentials of different 



