Corrosion redox system reduction

Spent fuels vary in microstructure, and phase and elemental distribution depending on the in-core reactor operating conditions and reactor history. The chemical stability of spent U oxide fuel is described by local pH and Eh conditions, redox being the most important parameter. However, the redox system will also evolve with time as various radionuclides decay and the proportion of oxidants and reductants generated at the fuel/water interface changes with the altering a-, (J-, y-radiation field and with the generation of other corrosion products that can act as... 

Corrosion inhibitor - corrosion inhibitors are chemicals which are added to the electrolyte or a gas phase (gas phase inhibitors) which slow down the - kinetics of the corrosion process. Both partial reactions of the corrosion process may be inhibited, the anodic metal dissolution and/or the cathodic reduction of a redox-system [i]. In many cases organic chemicals or compounds after their reaction in solution are adsorbed at the metal surface and block the reactive centers. They may also form layers with metal cations, thus growing a protective film at the surface like anodic oxide films in case of passivity. Benzo-triazole is an example for the inhibition of copper cor-... 

This model has been proved experimentally by studying the competition of the anodic decomposition reaction and the oxidation of Cu at p-GaAs in the dark and at n-GaAs under illumination [93]. This is a suitable redox system, because reduction and oxidation occur via the valence band, and because the anodic oxidation of Cu proceeds independently from the corrosion. Accordingly, the total current is given by... 

All refined metals have a tendency to revert to a thermodynamically more stable form such as those in which they occur naturally on earth. Thus one of the corrosion products of iron is iron oxide (Fea03(s)) which is one form of iron ore. Almost all types of corrosion can be explained in terms of electrochemistry (oxidation-reduction reactions) for this reason we will consider corrosion as an example of the application of redox chemistry or electrochemistry to a practical situation. We will not present a detailed quantitative analysis of corrosion and the design of corrosion-control systems. Other texts should be consulted for this type of information. 

Corrosion is an electrochemical process that consists of at least two reactions that compensate each other electronically in open circuit conditions. The anodic metal dissolution is counterbalanced by the cathodic reduction of a redox system within the electrolyte. Various processes may serve as cathodic counterreactions. The most common are oxygen reduction and hydrogen evolution in acidic electrolytes. 

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... 

A very important part is the metal/elec-trolyte interface. This is the site where anodic metal dissolution and the compensating cathodic reduction of redox systems take place. At this site, adsorption and layer formation occur which might drastically reduce the corrosion rate by inhibition of the corresponding electrode reactions. The adjacent electrolyte is the medium where preceding or consecutive electrode reactions occur, and diffusion or migration of corro-... 

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. 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

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. 

Under many conditions, metal dissolution is reduced by the formation of surface layers. These layers usually complicate the kinetics of free metal dissolution as well as the reduction of a redox system. Besides precipitated layers of oxides, hydroxides, and salts of the metal of interest, passive layers and films of organic materials are a common and very important means of corrosion protection. 

The detailed mechanism is often not fully understood. A blocking of active sites is often postulated for anodic dissolution or cathodic reduction of redox systems. However, the conclusions were mostly drawn indirectly from electrokinetic data, and a detailed investigation with surface analytical methods is still missing. The chemical reaction steps and their inhibition at surface sites should be further investigated systematically by methods like XPS, AES, STM, SFM, etc. Most of the knowledge is still empiric. However, the results of electrochemical investigations are very valuable for a better corrosion protection of materials in industrial applications. Some few principles of inhibition are discussed in the following sections. 

Electronic properties can be studied by photoelectrochemistry using redox reactions at passivated metal surfaces in the dark or with illumination. Their knowledge is very valuable to understand layer formation and corrosion phenomena for open circuit conditions, i.e., under the influence of redox systems within the electrolyte. Furthermore layer formation includes often oxidation of intermediates, i.e., oxidation of lower valent to higher valent cations, which requires electronic conduction through the already existing film. Similarly layer reduction needs electron transfer across the film. 

Under certain conditions, it will be impossible for the metal and the melt to come to equilibrium and continuous corrosion will occur (case 2) this is often the case when metals are in contact with molten salts in practice. There are two main possibilities first, the redox potential of the melt may be prevented from falling, either because it is in contact with an external oxidising environment (such as an air atmosphere) or because the conditions cause the products of its reduction to be continually removed (e.g. distillation of metallic sodium and condensation on to a colder part of the system) second, the electrode potential of the metal may be prevented from rising (for instance, if the corrosion product of the metal is volatile). In addition, equilibrium may not be possible when there is a temperature gradient in the system or when alloys are involved, but these cases will be considered in detail later. Rates of corrosion under conditions where equilibrium cannot be reached are controlled by diffusion and interphase mass transfer of oxidising species and/or corrosion products geometry of the system will be a determining factor. 

Ti, or PEEK (polyether ether ketone) to allow measurements under very corrosive conditions. The separated phases pass AMX gadgets for on-line detection (radiometric, spectrophotometric, etc.) or phase sampling for external measurements (atomic absorption, spectrometric, etc.), depending on the system studied. The aqueous phase is also provided with cells for pH measurement, redox control (e.g., by reduction cells using platinum black and hydrogen, metal ion determination, etc.) and temperature control (thermocouples). 

Moreover, almost in all the early steps, the redox potential of the clusters, which decreases with the nuclearity, is quite negative. Therefore the growth process undergoes another competition with a spontaneous corrosion by the solvent and the radiolytic protons, corrosion which may even prevent the formation of clusters, as mostly in the case of nonnoble metals. Monomeric atoms and oligomers of these elements are so fragile to reverse oxidation by the medium that H2 is evolved and the zerovalent metal is not formed [11]. For that reason, it is preferable in these systems to scavenge the protons by adding a base to the solution and to favor the coalescence by a reduction faster than the oxidation [53]. 

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. 

The redox chemistry of these systems is relatively well defined contaminant reduction results in oxidative dissolution of Fe° by a reaction that is equivalent to corrosion of Fe° by organic oxidants. Metals such as Zn and Sn can reduce contaminants by similar reactions. 

The understanding gained by considering the Evans diagrams allows us to measure the corrosion current in a straightforward manner. First we must realize that the corrosion potential is in fact the open-circuit potential of a system undergoing corrosion. It represents steady state, but not equilibrium. It resembles the reversible potential in that it can be very stable. Following a small perturbation, the system will return to the open-circuit corrosion potential just as it returns to the reversible potential. It differs from the equilibrium potential in that it does not follow the Nemst equation for any redox couple and there is both a net oxidation of one species and a net reduction of another. 

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