Redox system reduction current density

Figure 1-29. Superposition of the current density potential curves of an Me/Me " and a redox electrode, which yields the polarization curve of anodic metal dissolution and cathodic reduction of the redox system Eq.m nd Fq, redox t Nernst potentials, r is the rest potential, i o,m

Fig. 13. (a) Schematic representation of the formation of mixed potential, M, at an inert electrode with two simultaneous redox processes (I) and (II) with formal equilibrium potentials E j and E2. Observed current density—potential curve is shown by the broken line, (b) Representation of the formation of corrosion potential, Econ, by simultaneous occurrence of metal dissolution (I), hydrogen evolution, and oxygen reduction. Dissolution of metal M takes place at far too noble potentials and hence does not contribute to EC0Ir and the oxygen evolution reaction. The broken line shows the observed current density—potential curve for the system. 

Sometimes the value of the redox potential attains even at low current densities such values that another simultaneous process is possible. For instance the cathodic reversible potential of the Ti++++/Ti+++ system with the same concentration of both kinds of ions is about tc° = 0.04 V if platinized platinum in a solution of sulphuric acid is used as a cathode the evolution of hydrogen commences at a potential also near zero and the ourrent efficiency with respect to the reduction of Ti++++ ions will be comparatively low. Much better results can be achieved by replacing the platinum by another material, suoh as lead or graphite whioh have an appreciable hydrogen overvoltage, whereby the deposition potential of hydrogen becomes more negative as oompared with the potential of the Ti++++/Ti+++ system. 

Under -> open-circuit conditions a possible passivation depends seriously on the environment, i.e., the pH of the solution and the potential of the redox system which is present within the electrolyte and its kinetics. For electrochemical studies redox systems are replaced by a -> potentiostat. Thus one may study the passivating properties of the metal independently of the thermodynamic or kinetic properties of the redox system. However, if a metal is passivated in a solution at open-circuit conditions the cathodic current density of the redox system has to exceed the maximum anodic dissolution current density of the metal to shift the electrode potential into the passive range (Fig. 1 of the next entry (- passivation potential)). In the case of iron, concentrated nitric acid will passivate the metal surface whereas diluted nitric acid does not passivate. However, diluted nitric acid may sustain passivity if the metal has been passivated before by other means. Thus redox systems may induce or only maintain passivity depending on their electrode potential and the kinetics of their reduction. In consequence, it depends on the characteristics of metal disso-... 

Equilibrium electrode potential — is the value of -> electrode potential determined exclusively by a single redox system ox/red in the absence of current and under complete equilibration. The rates of ox to red reduction and of red to ox oxidation processes are equal under these circumstances (see exchange current density). The value of equilibrium e.p. is determined by the - Nernst equation. Equilibrium e.p. presents a - redox potential in its fundamental sense. See also - reversibility. 

There is additionally the important problem involved in choosing the reduction or oxidahon potential of the electrolyte solutions from either cyclic voltammetry (CV) or linear sweep voltammetry (LSV). Since the oxidation or reduction reachon of cations or anions contained in the RTILs are electrochemically irreversible in general [8-10], the corresponding reduction or oxidation potential cannot be specifically obtained, unlike the case of the redox potential for an electrochemically reversible system. Figure 4.1 shows the typically observed voltammogram (LSV) for RTILs. Note that both the reduchon and oxidation current monotonically increase with the potential sweep in the cathodic and anodic directions, respectively. Since no peak is observed even at a high current density (10 mA cm ), a certain... 

Others are the reduction of Fe + and [Fe(CN)6] in solution. These systems are often used for chemical corrosion tests. Pitted metals expose a small area of a few intensively dissolving corrosion pits that are not protected by a passive layer and a large cathode of the passive metal surface. Because of the large size of the cathode, a much smaller cathodic current density is required for the compensating reduction of the redox system in comparison to the active metal dissolution within the pits. However, electronic conduction is still required across the passive layer. Figure 3 depicts the existing sections of a pitted metal surface with the related electrode reactions, the very small metal dissolution /pass, and the redox reaction ha,pass via the protecting oxide film and... 

Upon polarization of either electrode, the cell potential moves along the oxidation and reduction curves as shown in Fig. 1.1. When the current through the cell is f, the potential of the copper and zinc electrodes is Cj cu and e zn > and each of the electrodes have been polarized by (Ceq.cu i.Cu) and (Ceq.zn i,z )- Upon further polarization, the anodic and cathodic curves intersect at a point where the external current is maximized. The measured output potential in a corroding system, often termed the mixed potential or the corrosion potential (Tcorr)> h the potential at the intersection of the anodic and the cathodic polarization curves. The value of the current at the corrosion potential is termed the corrosion current (Icon) and can be used to calculate corrosion rate. The corrosion current and the corrosion potential can be estimated from the kinetics of the individual redox reactions such as standard electrode potentials and exchange current densities for a specific system. Electrochemical kinetics of corrosion and solved case studies are discussed in Chapter 3. 

In a photoanodic system, even at moderate current densities, the occurrence of sluggish counter electrode kinetics for the cathodic process will cause significant polarization losses and diminish the photovoltage. Minimization of these kinetic limitations necessitates a counter electrode with good catalytic properties. For example, as shown by [34], CoS on stainless steel or brass electrodes exhibits elec-trocatalytic properties toward polysulfide reduction and overpotentials as low as 1 mV cm mA has been realized. Composition of a particular redox electrolyte may have a bearing on the extent of counter electrode polarization [32, 39]. 

The question arises concerning which value of the anodic reverse y can then be expected for the same redox system at ZnO. Assuming the same maximum rate constant and the same A values as for the reduction of the [Fe(CN)g] / system, one obtains, according to Eq. (7.53), k = 1.5 X cm s , of course, for both A values. This is a very low value because the occupied states of the redox system are mainly distributed below the conduction band. The anodic current which is expected to be independent of the band bending is given by Eq. (7.52). Assuming 10 cm and = 6x 10 cm" (correspondingto 10 M), one obtains 10 A cm , that is, a current density which should be measurable. Morrison tried to measure it unfortunately, however, the currents were not reproducible. 

T] = E-Eq. a semi-logarithmic Tafel plot yields the lines of the current densities of anodic metal dissolution and cathodic reduction of the redox system, as presented for iron dissolution in 0.5 M H2SO4 in Fig. 1-30 (Kaesche, 1979). The intersection of both lines yields Er and the related corrosion current density 4 within the electrolyte. In the case of iron corrosion in sulfuric acid, the corrosion rates determined by the electrochemical evaluation of the Tafel plot and the chemical analysis of the dissolved species or the weight loss of the specimen for simple immersion tests agree sufficiently well (Kaesche, 1979). 

Figure 1-29 has shown the superposition of metal dissolution and a cathodic process which leads to zero total current density i and a corrosion rate I c at the rest potential Sr. If the Nernst potential of the metal/met-al-ion electrode and the involved redox system are sufficiently separated and the related I-E characteristic is sufficiently steep, the polarization curve only contains the anodic metal dissolution and the cathodic reduction of the redox system. The related opposite reactions are neglected. 

For large cathodic polarizations, the reduction of the redox system equals the total current density. Figure 1-30 presents the logarithm of the partial current densities for large anodic or cathodic n as an example. 

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