Standard hydrogen electrode formal potential

Formal potential — Symbol Efr (SI Unit V), has been introduced in order to replace the standard potential of -> cell reaction when the values of - activity coefficients of the species involved in the cell reaction are unknown, and therefore concentrations used in the equation expressing the composition dependence of ceii instead of activities. It also involves the activity effect regarding the -+ standard hydrogen electrode, consequently in this way the formal electrode potential is also defined. Formal potentials are similar to conditional (apparent) equilibrium constants (-> equilibrium constant), in that, beside the effect of the activity coefficients, side reaction equilibria are also considered if those are not known or too complex to be taken into account. It follows that when the logarithmic term which contains the ratio of concentrations in the -> Nernst... 

Formal potentials are empirically derived potentials that compensate for the types of activity and competing equilibria effects that we have just described. The formal potential of a system is the potential of the half-cell with respect to the standard hydrogen electrode measured under conditions such that the ratio of analytical concentrations of reactants and products as they appear in the Nernst equation is exactly unity and the concentrations of other species in the system are all carefully specified. For example, the formal potential for the half-reaction... 

In general, formal potentials are reported with reference to the standard hydrogen electrode, cf. Section 11.1.6.5, as exemplified in Tables V.2 and V.3 of the uranium NEA review [92GRE/FUG]. In that case, the appearing in the reduction reaction is already at standard conditions. For example, experimental data are available on the formal potentials for reactions ... 

In general, formal potentials are reported with reference to the standard hydrogen electrode, cf. Section 11.1.6.5, as exemplified in Tables V.2 and V.3 of the uranium... 

By definition, both redox potentials in Eq. (3) have to be referred to a common reference potential, in this case that of the standard hydrogen electrode (SHE). Obviously, some thermodynamic assumptions have to be introduced in order to evaluate [ Eo2/r2]she. which denotes the formal redox potential of a redox couple in the organic phase versus the SHE in water. 

The for a half-reaction is the potential of that reaction versus the standard hydrogen electrode, with all species at unit activity. Most reduction potentials are not determined under such conditions, so it is expedient to define a formal reduction potential. This is a reduction potential measured under conditions where the reaction quotient in the Nernst equation is one and other nonstandard conditions are described solvent, electrolyte, pN, and so on. Formal reduction potentials are represented by °. Reduction potentials determined by cyclic voltammetry are usually formal potentials. The difference between standard and formal potentials is not expected to be great. Other definitions of the formal potential are offered. ... 

Fig. 14.6 Standard (or formal) reduction potentials of actinium and the actinide ions in acidic (pH 0) and basic (pH 14) aqueous solutions (values are in volts vs standard hydrogen electrode).

The ease of reduction of Q may be quantified by the reduction potential of the couple Q/Q , usually denoted by E and expressed in volts relative to the Normal Hydrogen Electrode (NHE). Only if both oxidant Q and reductant Q are in their thermodynamic standard states will the potential E equate to the standard reduction potential of the couple and symbolized as or E . In the present context, interest is focused on effective potentials under physiological conditions, i.e. in water at pH values close to 7. Reduction potentials such as polarographic half-wave potentials obtained in non-aqueous media, especially aprotic solvents such as dimethylformamide or acetonitrile and relative to a standard calomel electrode (for example), will be numerically quite different. should not be confused with a formal potential Eq, which is usually defined as the potential when the ratio of the total concentrations of oxidized and reduced... 

The pH scale has been defined operationally, and standard reference solutions based on a conventional scale of hydrogen ion activity have been selected (i, 2). Measurements of the pH of seawater made with different electrodes and instruments are satisfactorily reproducible when standardized in the same way (3). The results obtained, however, do not always have a clear interpretation. Formally, this diflSculty can be attributed to the residual liquid junction potential involved in the measurement. The primary standards are necessarily dilute buffer solutions (ionic strength, I 0.1) whereas seawater normally has an ionic strength exceeding 0.6. This difference in the concentrations and mobilities of the ions coming in contact with the concentrated solution of potassium chloride of which the salt bridge-liquid junction is composed gives rise to a potential difference that is indeterminate. Consequently, the meas-m ed pH is in error by an unknown amount and does not fall exactly on the scale fixed by the primary standards. 

Shortcomings of the choice of the equilibrium state as the electrical reference point in the evaluation of the temperature effect on the rate of electrode reactions, and consequently of the overpotential as an experimental substitute for A(A0) in the WE-RE cell at various temperatures, have been discussed in the previous section. Hence, another reference point should be sought. From a theoretical point of view, the choice is unambiguous—it is the zero point on the relative electrode potential scale, defined by the SHE convention. Basically, this is also an equilibrium state, but of a single reaction selected by convention, namely, the reduction of two hydrogen ions to molecular hydrogen. The value of A0 at the interface when this reaction is held at equilibrium, assuming all species involved are in standard thermodynamic states, is fixed by the SHE convention as zero. The same convention associates additional properties with this reference state temperature, solvent, and solute Independence. Formally, the properties of the SHE satisfy the principal theoretical requirements for the electrical reference point in the evaluation of the effect of temperature on the rate of electrode reactions. 

---