Electrode potentials, reversible standard

The standard electrode potential of M /Mcan be evaluated accurately and conveniently providing it forms a reversible electrode and can be... 

It must be emphasised that standard electrode potential values relate to an equilibrium condition between the metal electrode and the solution. Potentials determined under, or calculated for, such conditions are often referred to as reversible electrode potentials , and it must be remembered that the Nernst equation is only strictly applicable under such conditions. 

Electrical units 503, 519 Electrification due to wiping 77 Electro-analysis see Electrolysis and Electrogravimetry Electrochemical series 63 Electro-deposition completeness of, 507 Electrode potentials 60 change of during titration, 360 Nernst equation of, 60 reversible, 63 standard 60, (T) 62 Electrode reactions 505 Electrodeless discharge lamps 790 Electrodes antimony, 555 auxiliary, 538, 545 bimetallic, 575... 

The electrical double layer at pc-Zn/fyO interfaces has been studied in many works,154 190 613-629 but the situation is somewhat ambiguous and complex. The polycrystalline Zn electrode was found to be ideally polarizable for sufficiently wide negative polarizations.622"627 With pc-Zn/H20, the value of Eg was found at -1.15 V (SCE)615 628 (Table 14). The values of nun are in reasonable agreement with the data of Caswell et al.623,624 Practically the same value of Eff was obtained by the scrape method in NaC104 + HjO solution (pH = 7.0).190 Later it was shown154,259,625,628 that the determination of Eo=0 by direct observation of Emin on C,E curves in dilute surface-inactive electrolyte solutions is not possible in the case of Zn because Zn belongs to the group of metals for which E -o is close to the reversible standard potential in aqueous solution. 

The negative standard potential means that the Zn2+/Zn electrode is the anode in a cell with H+/H2 as the other electrode and, therefore, that the reverse of the cell reaction, specifically... 

Oxygen electrode. In principle, a classical oxygen electrode in a liquid electrolyte would be possible if an electrode material were known on the surface of which the redox system 02/0H is electrochemically reversible however, Luther26 measured its standard potential from the following cell without a liquid junction ... 

In the chronopotentiogram of mixtures (see Fig. 3.58), the reactive components will yield different inflection points if their standard potentials show a sufficient difference (at least 0.1 V). To take a simple example, let us consider a reversible electrode process for both ox and ox2 in a solution of the supporting electrolyte. Then eqn. 3.72 is simply valid for the first reacting oxj with up to the first inflection point however, beyond this point the last traces of exj... 

An alternative electrochemical method has recently been used to obtain the standard potentials of a series of 31 PhO /PhO- redox couples (13). This method uses conventional cyclic voltammetry, and it is based on the CV s obtained on alkaline solutions of the phenols. The observed CV s are completely irreversible and simply show a wave corresponding to the one-electron oxidation of PhO-. The irreversibility is due to the rapid homogeneous decay of the PhO radicals produced, such that no reverse wave can be detected. It is well known that PhO radicals decay with second-order kinetics and rate constants close to the diffusion-controlled limit. If the mechanism of the electrochemical oxidation of PhO- consists of diffusion-limited transfer of the electron from PhO- to the electrode and the second-order decay of the PhO radicals, the following equation describes the scan-rate dependence of the peak potential ... 

Before depicting these examples, we examine two questions. One deals with the cyclic voltammetric responses of systems reversibly exchanging two electrons with the electrode as a function of the standard potential separation, having in mind the use of these signals to determine the difference between the two standard potentials. The other concerns the response of a molecule containing two or more identical and independent reducible or oxidizable groups. 

Let us now suppose that the waveform of figure 16.3 is applied to study the reversible oxidation of a species R to R in a given solvent. The reaction occurs at the working electrode (anode), and /i°(R/R ) is the standard potential of the R/R- couple. Because the standard potential of the reference electrode in our cell is known accurately relative to the standard potential of the SHE (E° = 0 by definition), we can write the cell reaction and the Nernst equation as... 

Example 5.1. Calculate the reversible electrode potential for Cu immersed in a CUSO4 solution having concentrations of 0.01, 0.001, and 0.0001 mol/L at 25°C, neglecting ion-ion interaction (using concentrations instead of activities). The standard electrode potential for Cu/Cu electrode is 0.337 V. 

Samhoun and David [169] have studied the reduction of Cf(III) by radiopolarography in 0.1 M LiCl at pH 2 and found a reversible electrode process attributed to the Cf(III)/Cf(0) couple at 1/2 = —1.508 V versus SHE with an estimated standard potential of —2.030 V. These results were called into question in a subsequent paper by Musikas etal. [173] in which they determined... 

The value of the constant V, and hence the values of standard potentials, depend on the choice of the reference electrode and on the character of electrode reaction, which takes place on it With the reference electrode potential conventionally taken as zero, we can choose, for example, the normal hydrogen electrode (NHE), i.e., an electrode, for which the equilibrium at the interface is attained due to the reversible redox reaction H+ + e = H2, provided the activity of H+ ions in the solution is 1 mol/liter and the pressure of gaseous hydrogen above the solution is 1 atm. Many of the measured potentials are given below relative to the saturated calomel electrode (SCE) its potential relative to the NHE is 0.242 V. 

In addition, electrode reactions are frequently characterized by an irreversible, i.e., slow, electron transfer. Therefore, overpotentials have to be applied in preparative-scale electrolyses to a smaller or larger extent. This means not only a higher energy consumption but also a loss in selectivity as other functions within the molecule can already be attacked. In the case of indirect electrolyses, no overpotentials are encountered as long as reversible redox systems are used as mediators. It is very exciting that not only overpotentials can be eliminated but frequently redox catalysts can be applied with potentials which are 600 mV or in some cases even up to 1 Volt lower than the electrode potentials of the substrates. These so-called redox reactions opposite to the standard potential gradient can take place in two different ways. In the first place, a thermodynamically unfavorable electron-transfer equilibrium (Eq. (3)) may be followed by a fast and irreversible step (Eq. (4)) which will shift the electron-transfer equilibrium to the product side. In this case the reaction rate (Eq. (5)) is not only controlled by the equilibrium constant K, i.e., by the standard potential difference be-... 

At a platinum electrode that is slightly activated by platinization, the dissolved hydrogen in various solvents is oxidized nearly reversibly by H2= 2H+ + 2c. We can determine by cyclic voltammetry the standard potential of this process. If the standard potentials in various solvents are compared using a common potential scale, the Gibbs energies of transfer of H+ can be obtained [44], With electrodes other than platinum, it is difficult to observe reversible oxidation of dissolved hydrogen [44 b],... 

While many of the standard electroanalytical techniques utilized with metal electrodes can be employed to characterize the semiconductor-electrolyte interface, one must be careful not to interpret the semiconductor response in terms of the standard diagnostics employed with metal electrodes. Fundamental to our understanding of the metal-electrolyte interface is the assumption that all potential applied to the back side of a metal electrode will appear at the metal electrode surface. That is, in the case of a metal electrode, a potential drop only appears on the solution side of the interface (i.e., via the electrode double layer and the bulk electrolyte resistance). This is not the case when a semiconductor is employed. If the semiconductor responds in an ideal manner, the potential applied to the back side of the electrode will be dropped across the internal electrode-electrolyte interface. This has two implications (1) the potential applied to a semiconducting electrode does not control the electrochemistry, and (2) in most cases there exists a built-in barrier to charge transfer at the semiconductor-electrolyte interface, so that, electrochemical reversible behavior can never exist. In order to understand the radically different response of a semiconductor to an applied external potential, one must explore the solid-state band structure of the semiconductor. This topic is treated at an introductory level in References 1 and 2. A more complete discussion can be found in References 3, 4, 5, and 6, along with a detailed review of the photoelectrochemical response of a wide variety of inorganic semiconducting materials. 

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