Cathodic-reactant reduction

Knowledge of the parameters of the individual electrode reactions permits writing expressions for the individual oxidation or reduction curves (see the section Complete Polarization Curves for a Single Half-Cell Reaction in Chapter 3). Thus, the expression for the cathodic-reactant reduction reaction ... 

Fig. 6.2 Schematic experimental polarization curves (solid curves) assuming Tafel behavior for the individual oxidation and cathodic-reactant reduction polarization curves (dashed curves)...

The thermodynamics of electrochemical reactions can be understood by considering the standard electrode potential, the potential of a reaction under standard conditions of temperature and pressure where all reactants and products are at unit activity. Table 1 Hsts a variety of standard electrode potentials. The standard potential is expressed relative to the standard hydrogen reference electrode potential in units of volts. A given reaction tends to proceed in the anodic direction, ie, toward the oxidation reaction, if the potential of the reaction is positive with respect to the standard potential. Conversely, a movement of the potential in the negative direction away from the standard potential encourages a cathodic or reduction reaction. 

From these two examples, which as will be seen subsequently, present a very oversimplified picture of the actual situation, it is evident that macroheterogeneities can lead to localised attack by forming a large cathode/small anode corrosion cell. For localised attack to proceed, an ample and continuous supply of the electron acceptor (dissolved oxygen in the example, but other species such as the ion and Cu can act in a similar manner) must be present at the cathode surface, and the anodic reaction must not be stifled by the formation of protective films of corrosion products. In general, localised attack is more prevalent in near-neutral solutions in which dissolved oxygen is the cathode reactant thus in a strongly acid solution the millscale would be removed by reductive dissolution see Section 11.2) and attack would become uniform. 

Twin polarized platinum microelectrodes are conveniently used for endpoint detection for oxidation-reduction titrations. Consider a titration curve for oxidation-reduction titration where both reactants behave reversibly at the electrodes. An example of this kind of titration is titration of iron(II) with cerium(IV) (Fig. 14A). At the starting point of the titration, no current is observed because no suitable cathode reactant is available. With addition of cerium(IV), a mixture of iron(II) and iron(III) is produced, which permits the passage of current. Beyond the midpoint in the titration, iron(III) becomes in excess, and the current is then regulated by decreasing iron(II) concentration. At the equivalence point, the current approaches zero because iron(III) are present, and the applied potential is not great enough to cause these to react at the electrode. Beyond the equivalence point, the current rises again because both cerium(III) and cerium(IV) are present to react at the electrodes. 

The electrons transferred during the redox reaction move through an external circuit, exiting from the anode after oxidation, and entering into the cathode for reduction. The two semi-reactions can occur because the two separate spaces are inter-connected by a conductive hquid or solid phase (electrolyte) able to transfer ionic species, thus permitting the closing of the electric circuit. Then the electrolyte has to be ionically conductive, whereas the electrodes have to be electrically conductive and, in the case of gaseous reactants, sufficiently porous to allow the transfer of reactants and products to and from the reaction sites (see Sect. 3.2). 

The rate of a corrosion reaction is affected by pH (via H reduction and hydroxide formation), the partial pressnre of O, (the solubility/concentration of oxygen in solution), fluid agitation, and electrolyte condnctivity. Corrosion processes are analyzed using the thermodynamics of electrode reactions, mass transfer of the cathode reactants O2 and/or H, and the kinetics of metal dissolution reactions [157, 158]. 

Fig. 4.14 Representation of a mixed electrode with anodic reactant, M, and cathodic reactant, X. (a) Freely corroding condition, (b) Net external oxidation current, (c) Net external reduction current...

The above relationship is equally applicable if either the metal oxidation-rate curve or the reduction-rate curve for the cathodic reactant does not obey Tafel behavior. To illustrate this point, three additional schematic pairs of individual anodic and cathodic polarization curves are examined. In Fig. 6.3, the metal undergoes active-passive oxidation behavior and Ecorr is in the passive region. In Fig. 6.4, where the total re-... 

By using boiled water, the dissolved oxygen is expelled and hence, there should be no corrosion as the cathode reactant has been eliminated from the electrolyte. Unless the boiled water is kept in sealed containers, air (oxygen) will slowly dissolve into the water and corrosion of the metal or alloy will re-commence. As an alternative, using hot demineralised or distilled water will reduce the concentration of dissolved oxygen and hence corrosion, but this must be counter-balanced by the rise in reaction rates with temperature. In open conservation tanks, a temperature of 70°C is required to notice a significant reduction in rates of corrosion of metals. Small copper alloy artefacts from the Mary Rose were treated in this way using water at 80°C for 30 days. At the end of this period, the chloride levels in the water dropped to below 1 ppm. 

Two types of ohmic losses occur in fuel cells. These are potential losses due to electron transport through electrodes, bipolar plates, and collector plates and potential loss due to proton transport through the membrane. The magnitudes of these potential losses depend on the materials used in the construction of the fuel cells and its operating conditions [27]. Membrane conductivity increases with membrane water content. Reduction in the thickness of the membrane between anode and cathode may be thought of as an expedient way to eliminate ohmic overpotential. However, thin membrane may cause the problem of crossover or intermixing of anodic and cathodic reactants [27]. 

According to mixed-potential theory, any overall electrochemical reaction can be algebraically divided into half-cell oxidation and reduction reactions, and there can be no net electrical charge accumulation [J7], For open-circuit corrosion in the absence of an applied potential, the oxidation of the metal and the reduction of some species in solution occur simultaneously at the metal/electrolyte interface, as described by Eq 14, Under these circumstances, the net measurable current density, t pp, is zero. However, a finite rate of corrosion defined by t con. occurs at anodic sites on the metal surface, as indicated in Fig. 1. When the corrosion potential, Eco ., is located at a potential that is distincdy different from the reversible electrode potentials (E dox) of either the corroding metal or the species in solution that is cathodically reduced, the oxidation of cathodic reactants or the reduction of any metallic ions in solution becomes negligible. Because the magnitude of at E is the quantity of interest in the corroding system, this parameter must be determined independendy of the oxidation reaction rates of other adsorbed or dissolved reactants. 

Figure 13.6. Durability of Pt/C and Pt Coi. /C-based MEAs tested iu a short stack (active area = 465 cm ) under H2-air at a ceU temperature of 80 °C and total reactant pressures of 150 kPUabs, with both anode and cathode humidities at 100% and anode and cathode reactant stoichiometries of s = 2/2. Data are shown with the stack under a constant load of 0.20 A/cm over 1000 h. Data were averaged over four cells of each type of MEA [1], (Reprinted from Applied Catalysis B Environmental, 56, Gasteiger HA, Kocha SS, Sompalh B, Wagner FT. Activity benchmarks and requirements for Pt, Pt-alloy, and non-Pt oxygen reduction catalysts for PEMFCs, 9-35, 2005, with permission from Elsevier.)...

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