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Potential redox equilibrium

We consider a transfer of redox electrons at semiconductor electrodes polarized at an overvoltage t relative to the equilibrium redox potential (the Fermi level cfcredox)). The transfer current of redox electrons is given in Eqn. 8-54 by the arithmetic sum of the electron current via the conduction band, in(ti) - (0(11) > and the hole current via the valence band, ij(ii) - i (Ti) ... [Pg.258]

Further, the total overvoltage, ii, is the difference between the polarization potential E(=- aod the equilibrium redox potential (= - BvmvTm /e)... [Pg.348]

We can conclude from these thermodynamic considerations that it is possible to estimate the redox potentials of excited molecules, if we know the equilibrium redox potentials for the molecules in the ground state, as well for reduction as for oxidation, and add or subtract from these redox potentials the excitation energy AE of the lowest singlet or triplet state. For most dye molecules the reduction redox potential is experimentally more easily accessible than the oxidation redox potential. In such cases we have found that an estimation can be made by assuming that the ionisation energy of the dye molecule in crystalline state is similar to the ionisation energy in a polar solvent and gives an approximate value for the absolute redox potential. Such estimations are especially useful for a comparison of molecules with similar structure. [Pg.37]

The equilibrium redox potential, the free energy change per mole electron for a given reduction, represents the oxidizing intensity of the couple at equilibrium. It is conveniently expressed for many applications in terms of the parameter, pE, as proposed by Jorgensen (8) and popularized by Sillen (14). This parameter is defined by the relation,... [Pg.278]

Thermodynamic analysis was performed to determine the equilibrium redox potential for the Cu(I) Cu(II) conversion in the HCl(aq) solution. These data are important for estimating the voltage efficiency of the electrolyser and understanding the phase equilibria in the anolyte over the experimental ranges of temperature and applied potential. [Pg.254]

Fig. 25. Dependence of faradaic resistance measured at the equilibrium redox potential on the polycrystalline film resistivity for (1) Fe(CN)63, 4 and (2) quinone/hydroquinone systems. Reprinted from [110], Copyright (1997), with permission from Elsevier Science. Fig. 25. Dependence of faradaic resistance measured at the equilibrium redox potential on the polycrystalline film resistivity for (1) Fe(CN)63, 4 and (2) quinone/hydroquinone systems. Reprinted from [110], Copyright (1997), with permission from Elsevier Science.
The equilibrium redox potential can be calculated from the following Nemst equation ... [Pg.37]

The E value reflects the stabilization energy of the negative charge by surrounding solvent molecules. Thus its variation by solvation can be used as one of the solvent parameters. Since reaction 1 is nothing but a one-electron oxidation reaction, E can also be called optical oxidation potential The standard oxidation potential value is often diffrcult to be determined because many redox reactions are not reversible. Therefore, the E value should be a good alternative as a measure of redox reactivity for which the equilibrium redox potential is not known. [Pg.409]

While H" exists as a hydrated species in water, c does not. As we shall see, pe is related to the equilibrium redox potential (volts, hydrogen scale). The electron, as discussed here and used as a component in our equilibrium calculations, is different from the solvated electron, which is a transient reactant in photolyzed solutions. [Pg.429]

Figure 8(a) shows a cyclic voltammetric curve obtained at BDD electrode in 0.5 M H2SO4. The fact that the separation between the cathodic and the anodic peaks (AEp) is very high (about 0.9 V) indicates that the Q/H2Q system is irreversible at the boron-doped diamond electrode. Furthermore, the apparent equilibrium redox potential of the couple Q/H2Q(Eo = 0.65 V) is much closer to the anodic peak potential than to the cathodic one. [Pg.897]

If the redox center in the gap is exposed to half of the bias voltage drop, y= f, and the maximum is at the equilibrium redox potential, = 0- This holds a diagnostic clue regarding the mechanism of single-molecule electronic tunneling. We shall return later to data analysis based on this view. The precise location of the maximum depends, however, rather sensitively on the potential distribution in the turmeling gap, as reflected in the correlation between the parameters and r(Eq. (2.10)). [Pg.97]

Figure 7-18. A sequence of in situ STM images of azurin adsorbed on Au(lll) in 50 mM NTLtAc (pH 4.6). Taken for increasing positive substrate potentials -0.05 (a), 0.0 (b), +0.05 (c) and +0.10 V (d). Potentials quoted are relative to the equilibrium redox potential of the centre. Note that the image contrast has decreased in images (c) and especially (d), which are above the equihbrium potential at the copper redox site (see the text). Reprinted from ref 57 with permission. Figure 7-18. A sequence of in situ STM images of azurin adsorbed on Au(lll) in 50 mM NTLtAc (pH 4.6). Taken for increasing positive substrate potentials -0.05 (a), 0.0 (b), +0.05 (c) and +0.10 V (d). Potentials quoted are relative to the equilibrium redox potential of the centre. Note that the image contrast has decreased in images (c) and especially (d), which are above the equihbrium potential at the copper redox site (see the text). Reprinted from ref 57 with permission.
Figure 21.3. An illustrative diagram to show the way in which the environment has changed as it moved inevitably from reducing to oxidizing conditions because of the rise in oxygen in the atmosphere (as shown in Figure 21.4). The standard oxidation/reduction potentials at pH = 7.0 are used on a scale of H2/IF at -0.4 V and H2O/O2 at -1-O.8 V. The environment is assumed to change close to the equilibrium redox potential set by the partial pressure of oxygen. Further data are provided in references [5, 6]. Figure 21.3. An illustrative diagram to show the way in which the environment has changed as it moved inevitably from reducing to oxidizing conditions because of the rise in oxygen in the atmosphere (as shown in Figure 21.4). The standard oxidation/reduction potentials at pH = 7.0 are used on a scale of H2/IF at -0.4 V and H2O/O2 at -1-O.8 V. The environment is assumed to change close to the equilibrium redox potential set by the partial pressure of oxygen. Further data are provided in references [5, 6].
The two distributions overlap at the Fermi energy. As will be shown in Chapter 3, this energy is equal to the equilibrium redox potential and marks the electrochemical potential of the electrons in the electrolyte. [Pg.55]

Figure 2.33 DOS of elections for reduced and oxidized ions as function of energy in an electrol)de. Equal concentrations of oxidized and reduced ions are assumed the reference point for the energy is the Fermi energy, defined as energy of equal DOS of occupied and reduced states. The Fermi energy is equal to the equilibrium redox potential (Chapter 3). Figure 2.33 DOS of elections for reduced and oxidized ions as function of energy in an electrol)de. Equal concentrations of oxidized and reduced ions are assumed the reference point for the energy is the Fermi energy, defined as energy of equal DOS of occupied and reduced states. The Fermi energy is equal to the equilibrium redox potential (Chapter 3).
Figure 4.8 Electron energy distribution function D(E ) as function of the energy — Fq>. (a) In the metal, potential cp, (b) for an adsorbed ion in the adsorption layer (inner Helmholtz layer), potential (p, and (c) for the metal ion in the electrolyte, potential cp. Ep is the Fermi energy. In equilibrium is Fpnj = Spad = Fpei. The Fermi energy of the electrolyte is equal to the equilibrium redox potential (Nemst potential). ... Figure 4.8 Electron energy distribution function D(E ) as function of the energy — Fq>. (a) In the metal, potential cp, (b) for an adsorbed ion in the adsorption layer (inner Helmholtz layer), potential (p, and (c) for the metal ion in the electrolyte, potential cp. Ep is the Fermi energy. In equilibrium is Fpnj = Spad = Fpei. The Fermi energy of the electrolyte is equal to the equilibrium redox potential (Nemst potential). ...
In other words, when f is sufficiently large, the ratio Tox/Tred is virtually constant and the change in open-circuit potential, A Foe, is produced by the thermally induced change in the junction potentials, Ar(dFTj/dr), and by the thermally induced change in the equilibrium redox potential, Ar(dF°Vdr+ (Fie- ° )/F). When the condition (1 + co) /yW 1 is not met, the ILIT-induced charge-transfer will effect a change in the ratio Fox/Fred- Wc will examine this in greater detail in Sec. IV.F. [Pg.127]

At first sight, the reaction would evolve in the reverse sense of the given one. However, at the equivalence point, the reaction is displaced toward the right since the bromine concentration is quasi-null and the equilibrium redox potential of the solution is weaker than the standard potential of the couple Br2/Br. ... [Pg.412]

We have chosen to work with only the first of these charge displacements antimycin A was added to eliminate reaction III and the experiments were carried out at an equilibrium redox potential of about Eh +400 mV to eliminate reaction II. There were three reasons for this choice of experimental conditions. First because reaction I is exceedingly fast, the extent of the accompanying change in AV can be measured without interference from dissipative ionic fluxes. Second, the driving force of reaction I (1.40 0.37 V, see Clayton, 1978) greatly exceeds the values of the lC "-diffusion potential that can be attained (0.14 v, see below) so... [Pg.342]


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See also in sourсe #XX -- [ Pg.55 , Pg.61 , Pg.62 ]




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