Cathodic process evolution

It follows from equation 1.45 that the corrosion rate of a metal can be evaluated from the rate of the cathodic process, since the two are faradai-cally equivalent thus either the rate of hydrogen evolution or of oxygen reduction may be used to determine the corrosion rate, providing no other cathodic process occurs. If the anodic and cathodic sites are physically separable the rate of transfer of charge (the current) from one to the other can also be used, as, for example, in evaluating the effects produced by coupling two dissimilar metals. There are a number of examples quoted in the literature where this has been achieved, and reference should be made to the early work of Evans who determined the current and the rate of anodic dissolution in a number of systems in which the anodes and cathodes were physically separable. 

The hydrogen evolution reaction (h.e.r.) and the oxygen reduction reaction (equations 1.11 and 1.12) are the two most important cathodic processes in the corrosion of metals, and this is due to the fact that hydrogen ions and water molecules are invariably present in aqueous solution, and since most aqueous solutions are in contact with the atmosphere, dissolved oxygen molecules will normally be present. 

An attempt to measure the pzc of Pb + Na alloys has been reported by Kiseleva etal.m The system was studied in the context of the cathodic processes during hydrogen evolution that are believed to result in the incorporation of alkali metal atoms. 

The relation between E and t is S-shaped (curve 2 in Fig. 12.10). In the initial part we see the nonfaradaic charging current. The faradaic process starts when certain values of potential are attained, and a typical potential arrest arises in the curve. When zero reactant concentration is approached, the potential again moves strongly in the negative direction (toward potentials where a new electrode reaction will start, e.g., cathodic hydrogen evolution). It thus becomes possible to determine the transition time fiinj precisely. Knowing this time, we can use Eq. (11.9) to find the reactant s bulk concentration or, when the concentration is known, its diffusion coefficient. 

Cathodic hydrogen evolution is one of the most common electrochemical reactions. It is the principal reaction in electrolytic hydrogen production, the auxiliary reaction in the production of many substances forming at the anode, such as chlorine, and a side reaction in many cathodic processes, particularly in electrohydrometallurgy. It is of considerable importance in the corrosion of metals. Its special characteristic is the fact that it can proceed in any aqueous solution particular reactants need not be added. The reverse reaction, which is the anodic ionization of molecular hydrogen, is utilized in batteries and fuel cells. 

Our investigations showed that in mixed melts of eutectic composition carbamide-NH4(K)Cl, the oxidation and reduction of melt constituents take place mainly independently of each other. The anodic process at platinum electrodes in the range of potentials below 0.9V is associated with the direct oxidation of carbamide to secondary and tertiary amide compounds, accumulation of ammonium ions in the melt, and evolution of the same gaseous products as in carbamide electrolysis [8], The cathodic process is accompanied by the formation of ammonia, CO, and C02, i.e. of the same products as in pure- carbamide electrolysis. In contrast to carbamide melt, a large amount of hydrogen appears in the cathode gases of the mixed melt, and in the anode gases of the carbamide-KCl melt, the presence of chlorine has been established at potentials above 0.9V. In the... 

Provided the reaction is, in some sense, reversible, so that equilibrium can be attained, and provided the reactants and products arc all gas-phase, solution or solid-state species with well-defined free energies, it is possible to define the free energies for all such reactions under any defined reaction conditions with respect to a standard process this is conventionally chosen to be the hydrogen evolution/oxidation process shown in (1.11). The relationship between the relative free energy of a process and the emf of a hypothetical cell with the reaction (1.11) as the cathode process is given by the expression AC = — nFE, or, for the free energy and potential under standard conditions, AG° = — nFEl where n is the number of electrons involved in the process, F is Faraday s constant and E is the emf. 

The flow-through cathode is the result of a tailored-to-the-process evolution of the GDE structure, which is available also in two additional configurations double-sided (originally developed for fuel cell servicing) and single-sided (see Fig. 9.7). The double-sided type is particularly suited for the electrochemical process where the product should not be released on to the back surface of the cathode, as in the case of oxygen-depolarised chlor-alkali electrolysis, discussed in Section 9.3. 

The predominant cathodic process for Cu corrosion will therefore be oxygen reduction in aerated solutions and hydrogen evolution in deaerated solutions. 

The only cathodic process in HF solutions is the hydrogen evolution reaction (HER), which is important in that it is involved in almost all reactions at both anodic and cathodic potentials. The silicon electrode can be passivated by hydrogen termination... 

Passage of 1.0 mol of electrons (one faraday, 96,485 A s) will produce 1.0 mol of oxidation or reduction—in this case, 1.0 mol of Cl- converted to 0.5 mol of Cl2, and 1.0 mol of water reduced to 1.0 mol of OH- plus 0.5 mol of H2. Thermodynamically, the electrical potential required to do this is given by the difference in standard electrode potentials (Chapter 15 and Appendix D) for the anode and cathode processes, but there is also an additional voltage or overpotential that originates in kinetic barriers within these multistep gas-evolving electrode processes. The overpotential can be minimized by catalyzing the electrode reactions in the case of chlorine evolution, this can be done by coating the anode with ruthenium dioxide. 

Alloying to modify the overpotential of the metal surface for H2 evolution or O2 absorption can help control corrosion, although it is not always obvious whether these cathodic processes should be suppressed (i0 lowered) or stimulated to produce the desired corrosion resistance. In the case of titanium (see Section 16.6), for example, palladium was alloyed in to catalyze H2 evolution and to force the metal into a passive condition. 

The chemistry of electrochemical reaction mechanisms is the most hampered and therefore most in need of catalytic acceleration. Therefore, we understand that electrochemical catalysis does not, in principle, differ much fundamentally and mechanistically from chemical catalysis. In addition, apart from the fact that charge-transfer rates and electrosorption equilibria do depend exponentially on electrode potential—a fact that has no comparable counterpart in chemical heterogeneous catalysis—in many cases electrocatalysis and catalysis of electrochemical and chemical oxidation or reduction processes follow very similar if not the same pathways. For instance as electrochemical hydrogen oxidation and generation is coupled to the chemical splitting of the H2 molecule or its formation from adsorbed hydrogen atoms, respectively, electrocatalysts for cathodic hydrogen evolution—... 

Obviously the contribution of the pore walls—according to the current density distribution—to cathodic hydrogen evolution becomes negligible beyond 10 fim pore depth so that for a perfect, undivided Raney-nickel coating of 100 fim thickness, only 7 to 8% utilization is anticipated. This is the reason why the fissures and cracks, the so-called tertiary structure of the catalyst, formed in Raney-nickel coatings by the leaching process are so important for improving its utilization. 

It has been established that most cathode metals are to some extent soluble in chromic acid solutions, and ions will enter the solution in the highest available oxidation state [e.g. copper(II), gold(III)]. Polarization of the cathode will then cause reduction to lower oxidation states [kinetic factors will prevent the prior reduction of chromate(VI)], then new low-valent species may then initiate a chemical reduction of the chromium(Vl). Chromium deposition occurs within the potential range for the evolution of dihydrogen and, indeed, the latter is the dominant cathode process with the result that typically cathode current efficiencies of only 10-20% are achieved (see equation 9). 

The advantages of the liquid surface and large overpotential for hydrogen evolution make mercury the material of choice for cathodic processes, unless the use of mercury is specifically contraindicated by some incompatibility with the system. Incompatibility can arise from strong specific absorption, as with some sulfur-containing compounds, or in high-temperature systems such as fused salts because of the low boiling point of mercury (356.6°C). 

We have just mentioned that one reason for a limited range of potentials in a particular SSE is the reactivity of the components of the SSE toward oxidation and reduction. It is also obvious that the limiting cathodic process in protic solvents, nos 1-9 in Table 4, must be reduction of protons or the equivalent, the proton donor. The unfavourable cathodic limit for reduction of protons can, however, be vastly improved by the use of mercury as the cathode material and a tetraalkylammonium salt as SSE (nos. 1 and 3). The reason for mercury being such a favourable material is its large overpotential (see Section 10) for the reduction of protons (hydrogen evolution reaction). We have already commented (p. 24) on the fact that the reduction of protons occurs many orders of magnitude faster on certain metals than on others, and this manifests itself by the overpotential, i.e., in order to make the reaction go at a measurable rate one has to increase the electrode potential from the equilibrium potential. Table 6 shows overpotentials for hydrogen evolution and... 

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. 

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