Cathodic reactions simultaneous

Paired synthesis. Paired electrochemical syntheses are processes in which both the anode and cathode reactions simultaneously contribute to the formation of the final products. The classic example is the simultaneous production of chlorine and sodium hydroxide in the chlor-alkali industry. Paired electrosynthesis can be generally classified in terms of the following ... 

Both the anodic and cathodic reactions lead to the same final product, (Fig. 2D), e.g. simultaneous anodic and cathodic production of glyoxylic acid [57]. 

Further, these anodic and cathodic reactions can occur spatially at adjacent locations on the stuface of a metal electrode rather than on two separated metal electrodes as shown in Fig. 11-1, where the anodic dissolution of iron and the cathodic reduction of hydrogen ions proceed simultaneously on an iron electrode in aqueous solution. The electrons produced in the anodic dissolution of iron are the same electrons involved in the cathodic reduction of hydrogen ions hence, the anodic reaction cannot proceed more rapidly than that the electrons can be accepted by the cathodic reaction and vice versa. Such an electrode at which a pair of anodic and cathodic reactions proceeds is called the mixed electrode . For the mixed electrodes, the anode (current entrance) and the cathode (current exit) coexist on the same electrode interface. The concept of the mixed electrode was first introduced in the field of corrosion science of metals [Evans, 1946 Wagner-Traud, 1938]. 

The concept of mixed potential can be extended to multi-electrode reactions on one electrode. If N electrode reactions happen on a single electrode simultaneously, the outside electrode current is up to zero. When N >2, all these electrode reactions constitute the multi-electrode reaction coupled system, in which a part of the electrode reactions belong to the anodic reaction and the others are cathodic reaction. Given that the current of the anodic reaction is positive value and that of the cathodic reaction is negative value, the following formula can be given. 

A bare surface of silicon can only exist in fluoride containing solutions. In reality, in these media, the electrode is considered to be passive due to the coverage by Si— terminal bonds. Nevertheless, the interface Si/HF electrolyte constitutes a basic example for the study of electrochemical processes at the Si electrode. In this system, the silicon must be considered both as a charge carrier reservoir in cathodic reactions, and as an electrochemical reactant under anodic polarization. Moreover, one must keep in mind that, according to the standard potential of the element, both anodic and cathodic charge transfers are involved simultaneously (corrosion process) in a wide range of potentials. 

The example just given assumed a constant pH of 7. Simultaneously with Fe2+(aq) release at the anode, however, hydroxide ion is formed by reduction of oxygen at the cathode (reaction 16.5), increasing the local pH, and diffuses out to mix with Fe2+ in the bulk aqueous phase. The solubility product of Fe(OH)2 is about 8 x 10-16, which means that concentrations of several micromoles Fe11 per liter could build up before Fe(OH)2 is precipitated. At the same time, increased pH will accelerate the oxidation... 

Consider now the processes caused by the formation of quasilevels. As was noted above, the shift of Fn relative to F is very small for majority carriers (electrons) and can usually be neglected precisely, this was done in constructing Fig. 16b. But for minority carriers (holes) the shift of Fp can be very large. The shifts of both Fnx F and Fp increase with the growing intensity of semiconductor illumination, so that for a certain illumination intensity Fp may reach the level of the electrochemical potential of anodic decomposition Fdec, p, and Fn—the level of a certain cathodic reaction (for example, reduction of water with hydrogen evolution FHljH20). These reactions start to proceed simultaneously, and their joint action constitutes the process of photocorrosion. 

The basic mechanism for the instability of ultrapure metals was suggested by Wagner and Traud in a classic paper in 1938.1 The essence of their view is that for corrosion to occur, there need not exist spatially separated electron-sink and -source areas on the corroding metal. Hence, impurities or other heterogeneities on the surface are not essential for the occurrence of corrosion. The necessary and sufficient condition for corrosion is that the metal dissolution reaction and some electronation reaction proceed simultaneously at the metal/environment interface. For these two processes to take place simultaneously, it is necessary and sufficient that the corrosion potential be more positive than the equilibrium potential of the M, + + ne M reaction and more negative than the equilibrium potential of the electronation (cathodic) reaction A + ne — D involving electron acceptors contained in the electrolyte (Fig. 12.8). 

It usually happens when the charge transfer step is very sluggish (k and j0 are very small), and large activation overpotential is needed to drive the reaction in any direction. In this case, the anodic and cathodic reactions are never simultaneously significant. 

Divided cells — Electrochemical cells divided by sintered glass, ceramics, or ion-exchange membrane (e.g., - Nafion) into two or three compartments. The semipermeable separators should avoid mixing of anolyte and - catholyte and/or to isolate the reference electrode from the studied solution, but simultaneously maintain the cell resistance as low as possible. The two- or three-compartment cells are typically used a) for preparative electrolytic experiments to prevent mixing of products and intermediates of anodic and cathodic reactions, respectively b) for experiments where different composition of the solution should be used for anodic and cathodic compartment c) when a component of the reference electrode (e.g., water, halide ions etc.) may interfere with the studied compounds or with the electrode. For very sensitive systems additional bridge compartments can be added. 

The kinetics is called irreversible in electrochemistry when the charge-transfer step is very sluggish, i.e., the standard rate constant (ks) and - exchange current density (j0) are very small. In this case the anodic and cathodic reactions are never simultaneously significant. In order to observe any current, the charge-transfer reaction has to be strongly activated either in cathodic or in anodic direction by application of -> overpotential. When the electrode process is neither very facile nor very sluggish we speak of quasireversible behavior. 

The separate anodic and cathodic processes will occur simultaneously but statistically independent of one another. The rate of each reaction will be governed by the electrical potential difference which exists across the metal-solution interface and the appropriate values of i0, a and 0 for each system. In the absence of an external disturbance, for instance an external current, a steady state will usually be reached where the sum of the rates of the cathodic reactions will equal the sum of the rates of the anodic reactions, viz., Sia =21. . The electrode potential will assume some value, Ej p, which is designated the mixed potential, and the electrode is considered as a poly-eieccrode (11, 16,17,18). 

Fig. ]2M Evans diagram for the corrosion of iron in the presence of two simultaneous cathodic reactions hydrogen evolution and oxygen reduction). The dotted line represents the sum of both catho-...

Kalu and Oloman [75] studied the simultaneous synthesis of alkaline hydrogen peroxide and sodium chlorate in a bench-scale flow-by single-cell electrochemical reactor. A schematic of the electrode conditions is shown in Fig. 18. Graphite felt was used as the cathode to synthesize peroxide from 0.5 -2.0 M NaOH chlorate was the product at a dimensionally stable anode (DSA). The anodic and cathodic reactions were as follows ... 

While an ovapotential may be applied electrically, we are interested in the overpotential that is reached via chemical equilibrium with a second reaction. As mentioned previously, the oxidation of a metal requires a corresponding reduction reaction. As shown in Figure 4.34, both copper oxidation, and the corresponding reduction reaction may be plotted on the same scale to determine the chemical equilibrium between the two reactions. The intersection of the two curves in Figure 4.34 gives the mixed potential and the corrosion current. The intersection point depends upon several factors including (the reversible potential of the cathodic reaction), cu2+/cu> Tafel slopes and of each reaction, and whether the reactions are controlled by Tafel kinetics or concentration polarization. In addition, other reduction and oxidation reactions may occur simultaneously which will influence the mixed potential. 

The electroless deposition of metals on a silicon surface in solutions is a corrosion process with a simultaneous metal deposition and oxidation/dissolution of silicon. The rate of deposition is determined by the reduction kinetics of the metals and by the anodic dissolution kinetics of silicon. The deposition process is complicated not only by the coupled anodic and cathodic reactions but also by the fact that as deposition proceeds, the effective surface areas for the anodic and cathodic reactions change. This is due to the gradual coverage of the metal deposits on the surface and may also be due to the formation of a silicon oxide film which passivates the surface. In addition, the metal deposits can act as either a catalyst or an inhibitor for hydrogen evolution. Furthermore, the dissolution of silicon may significantly change the surface morphology. 

Dissolution of PS. The dissolution of PS during PS formation may be due to two proeesses a proeess in the dark and a proeess under illumination. Both are essentially eorrosion proeesses by which the silicon in the PS is oxidized and dissolved with simultaneous reduction of the oxidizing species in the solution. The corrosion process is responsible for the formation of micro PS of certain thickness (stain film) as well as the dissolution of the existing PS. The material in the PS which is at a certain distance from the pore tips is little affected by the extanal bias due to the high resistivity of PS and is essentially at an open-circuit condition (OCP). This dissolution process, which is often referred to as chemical dissolution, is an electrochemical process because it involves charge transfer across the interface. The anodic and cathodic reactions in the microscopic corrosion cells depend on factors such as surface potential and carrier concentration on the surface which can be affected by illumination and the presence of oxidants in the solution. 

The ratio of the CH4-to-H2 direct decomposition rate to that of the CH4 steam-reforming reaction one was not found to be different from location to location in the porous electrode. Therefore, it was considered that the supply of H2O was sufficient in the present experiment. It was probable that a sufficient amount of CH4 and H2O diffused through the porous electrode and arrived at an electrodeelectrolyte interface. The direct decomposition reaction as well as the steam-reforming reaction simultaneously occurred at a narrow interface imder the condition of the sufficient supply of electric charge to the cathode electrode. As seen in Figure 9, some carbon deposited at the narrow interface between the electrolyte and electrode. This may be because the steam-reforming reaction was slower than the protonic conductivity of electrolyte and, consequently, the direct decomposition of CH4 provided electron and H+ ion to the interface. Since the carbon deposition was localized at the interface. Therefore, there was no effect on the mass and charge transfer in the present cell system. 

The kinetics of the hydrogen electrode reaction on dense porous graphite electrodes in molten KHSO4 from 245° C to 280°C [88-90] showed that the cathodic and anodic reactions are not strictly conjugated processes. The cathodic reaction was discussed in terms of conventional mechanisms, but the anodic reaction involves the simultaneous oxidation of hydrogen and graphite surface. The reaction exhibits a one-half power dependence on hydrogen pressure. 

For a certain illumination intensity, the hole quasilevel Fp at the semiconductor surface can reach the level of an anodic reaction (reaction of semiconductor decomposition in Fig. 9). In turn, the electron quasilevel F can reach, due to a shift of the Fermi level, the level of a cathodic reaction (reaction of hydrogen evolution from water in Fig. 9). Thus, both these reactions proceed simultaneously, which leads eventually to photocorrosion. Hence, nonequilibrium electrons and holes generated in a corroding semiconductor under its illumination are consumed in this case to accelerate the corresponding partial reactions. 

According to the electrochemical mechanism of electroless Ni-P plating, the electrons generated in the anodic half-reaction of hypophosphite oxidation (19.11) are involved in simultaneous cathodic reactions, that is, proton discharge from water (19.12), Ni(II) reduction... 

The cathode-to-anode area ratio is frequently a critical factor in corrosion. (This is true when well-defined cathodes and anodes exist. With mixed electrode behavior, where cathodic and anodic reactions occur simultaneously, separate areas are not readily distinguishable, and Aa is assumed equal to Ac.) Discussion of the influence of this ratio will be restricted to the case of a small total-corrosion-circuit resistance leading to the anodic and cathodic reactions occurring at essentially the same potential, Ecorr, as described previously. In Fig. 4.12, three different values of corrosion current, Icorr, and corrosion potential, Ecorr, are shown for three cathode areas relative to a fixed anode area of 1 cm2. For these cases, a reference electrode placed anywhere in the solution... 

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