Mixed solvent standard electrode potentials

AG ° (i) = standard free energy change for the transfer of i from water into the mixed solvent N = Avogadro s number e = electronic charge Dw = dielectric constant of water Ds = dielectric constant of the mixed solvent Th2o = radius of the water molecule Ew° = standard electrode potential in water EB° = standard electrode potential in the mixed solvent M = cation X = anion... 

The values of E, log y , Mxy> A, and B from Equations 5-10 substituted into Equation 4, make it possible to calculate Em0 at known molalities of hydrobromic acid, solvent compositions, and temperatures. By plotting values of Em° at a given solvent composition and temperature vs. molality, one can find the standard electrode potential E° of the Ag-AgBr electrode at that solvent composition and temperature from the value of Em° extrapolated to infinite dilution. This method has been used successfully in water and in organic solvent-water mixtures of higher dielectric constants, but if the mixed solvents have low dielectric constants, ca. 50 or below, the curvatures of the Em0f vs. m plots are sufficient to prevent accurate determinations of Em0 and hence of E°. 

A cognate issue is the establishment of a universal standard electrode potential scale in nonaqueous and mixed solvents, based on the aqueous standard hydrogen electrode (SHE), for which (H, aq/Hj) = 0 at all temperatures. The standard potentials E° are obtained on extrapolation of the EMF of a suitable cell to zero of the concentration of the electroactive electrolyte, the one that responds to the electrodes irrespective of the eventual presence of a constant inert background electrolyte, in the solutions of the two half-cells. A proper procedure for such an extrapolation that assures accuracy has been described by Mussini et al. [15]. The standard electrode potentials for a cation/metal pair M /M in a solvent S vs. the SHE is related to its standard potential in water and the standard molar Gibbs energy of transfer of the cation from water to the solvent (or solvent mixture) ... 

Studies of pzc in mixed solvents were also carried out by Blaszczyk etal n using the dipping method. They worked in mixtures offormamide and NMF and estimated the shift of the standard potential of the hydrogen electrode, of the surface dipole potential atHg, and of the liquid junction potential. 

The standard potential of the silver-silver bromide electrode has been determined from emf measurements of cells with hydrogen electrodes and silver-silver bromide electrodes in solutions of hydrogen bromide in mixtures of water and N-methylacetamide (NMA). The mole fractions of NMA in the mixed solvents were 0.06, 0.15, 0.25, and 0.50, and the dielectric constants varied from 87 to 110 at 25°C. The molality of HBr covered the range 0.01-0.1 mol kg 1. Data for the mixed solvents were obtained at nine temperatures from 5° to 45°C. The results were used to derive the standard emf of the cell as well as the mean ionic activity coefficients and standard thermodynamic constants for HBr. The information obtained sheds some light on the nature of ion-ion and ion-solvent interactions in this system of high dielectric constant. 

The Activity Coefficients of Hydrogen Chloride and the Standard Potential of the Silver-Silver Chloride Electrode in Mixed Solvents. 

Reference Redox Systems in Nonaqueous Systems and the Relation of Electrode Potentials in Nonaqueous and Mixed Solvents to Standard Potentials in Water... 

This chapter is concerned primarily with the computation of potentials of a cell using the hydrogen electrode as a probe for studying ionic equilibrium processes in mixed-organic-aqueous solvent systems. Computation of a number of other thermodynamic functions of the ionic process under investigation or of the solvent used is rather straightforward once the standard potential of the measuring cell has been calculated. 

Measurement of pH in a nonaqueous solvent when the electrode is standardized with an aqueous solution has little significance in terms of possible hydrogen ion activity because of the unknown liquid-junction potential, which can be rather large, depending on the solvent. Measurements made in this way are usually referred to as apparent pH. pH scales and standards for nonaqueous solvents have been suggested using an approach similar to the one for aqueous solutions. These scales have no rigorous relation to the aqueous pH scale, however. You are referred to the book by Bates (Ref. 3) for a discussion of this topic. See also M. S. Frant, How to Measure pH in Mixed Nonaqueous Solutions, Today s Chemist at Work, American Chemical Society, June, 1995, p. 39. 

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