Oxygen reduction reaction equilibrium electrode potential

The cathodic polarization curve is constructed using the oxygen electrode equilibrium potential and the cathodic slope, 6<- = —0.05V/decade. The equilibrium cathode potential, geq,c is calculated by applying the Nemst equation to the oxygen reduction reaction, Eq. (4.13), for an oxygen concentration of 1.0x10 mol/1 at pH= 11. 

The standard reduction potential of this reaction is around 0.21 V, much lower than the standard reduction potential of oxygen (R. 2.6) therefore, a mixed potential between 0.21 and 1.18 V is produced. Since the oxygen reduction reaction and Pt oxidation reaction dominate at the cathode due to kinetic reasons, the mixed potential is near the higher end. The other reason is hydrogen crossover through the PEM from the anode to the cathode. This is like an internal current flow in the cell and thus leads the electrodes (mainly the cathode) away from the 0 current thermodynamic equilibrium conditions. Due to the slow kinetics of Reaction 2.6, this internal current flow significantly lowers the cathode potential. Detailed estimation is given later in this chapter. For these two reasons, the OCV of a PEMFC is typically between 0.95 and 1.0 V. 

The mixed-potential model demonstrated the importance of electrode potential in flotation systems. The mixed potential or rest potential of an electrode provides information to determine the identity of the reactions that take place at the mineral surface and the rates of these processes. One approach is to compare the measured rest potential with equilibrium potential for various processes derived from thermodynamic data. Allison et al. (1971,1972) considered that a necessary condition for the electrochemical formation of dithiolate at the mineral surface is that the measmed mixed potential arising from the reduction of oxygen and the oxidation of this collector at the surface must be anodic to the equilibrium potential for the thio ion/dithiolate couple. They correlated the rest potential of a range of sulphide minerals in different thio-collector solutions with the products extracted from the surface as shown in Table 1.2 and 1.3. It can be seen from these Tables that only those minerals exhibiting rest potential in excess of the thio ion/disulphide couple formed dithiolate as a major reaction product. Those minerals which had a rest potential below this value formed the metal collector compoimds, except covellite on which dixanthogen was formed even though the measured rest potential was below the reversible potential. Allison et al. (1972) attributed the behavior to the decomposition of cupric xanthate. 

There existed oxidation-reduction reactions with the same reaction speed on the sulphide mineral surface in water. One is the self-corrosion of sulphide mineral. Another is the reduction of oxygen. If the equilibrium potential for the anodic reaction and the cathodic reaction are, respectively, E and, and the mineral electrode potential is E, the relationship among them is as follows ... 

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. 

Fe,v-Fe2+ system in the presence of a trace of dissolved oxygen. The measured zero-current potential is that value where the rate of 02 reduction at the electrode surface is equal to the rate of Fe2+ oxidation rather than the value of since at the latter point simultaneous 02 reduction produces excess cathodic current. In addition, because the net reaction of Em converts Fe2+ to Fe >>+, the measured potential exhibits a slow drift. Such mixed potentials are of little worth in determining equilibrium Eh values. 

Cathodic protection also can be accomplished by lowering the electrode potential to E M, the equilibrium potential for the metal to be protected, by an external power source. The circuit used to accomplish this is the same as shown in Fig. 2.12. With slight modification, it is again shown in Fig. 4.25 in which the metal to be protected is iron and the cathodic reaction supporting corrosion is either hydrogen-ion reduction, oxygen reduction, or both. 

As shown above, oxidation occurs when the electrode potential is higher, and reduction when it is lower, than the equilibrium potential. Hence, when we have, as in corrosion, two simultaneous reactions, of which one is oxidation (anodic reaction) and the other one is reduction (cathodic reaction), the real potential must lie between the equilibrium potentials of the two reactions. If we consider corrosion of iron in aerated water, with reduction of oxygen as the cathodic reaction, the potential has to be somewhere between the lines a and e in Figure 3.10. In acid (and usually in neutral) solutions the potential will lie in the corrosion region, in alkaline solutions in the passive region. With efficient oxygen supply, which for instance can be promoted by heavy convection of the solution, passivity may also be achieved in neutral water. 

The reaction rates themselves strongly depend on the conditions under which the reactions are conducted. In particular, cathodic oxygen reduction that, at temperatures below 150°C, is far from equilibrium comes closer to the equilibrium state as the temperature is raised. The reasons why the real value of the oxygen electrode s potential at low temperatures is far from the thermodynamic value and why cathodic oxygen reduction is so slow are not clear so far, despite a large number of studies that have been conducted to examine it. 

Rates of corrosion can also be measured using an electrochemical technique known as potentiodynamic polarization. The potential of the test metal electrode relative to a reference electrode (commonly the saturated calomel electrode SCE) is varied at a controlled rate using a potentiostat. The resultant current density which flows in the cell via an auxiliary electrode, typically platinum, is recorded as a function of potential. The schematic curve in fig. 2 is typical of data obtained from such a test. These data can provide various parameters in addition to corrosion rate, all of which are suitable for describing corrosion resistance. The corrosion potential F corr is nominally the open circuit or rest potential of the metal in solution. At this potential, the anodic and cathodic processes occurring on the surface are in equilibrium. When the sample is polarized to potentials more positive than Scon the anodic processes, such as metal dissolution, dominate (Anodic Polarization Curve). With polarization to potentials more negative than Scorr the cathodic processes involved in the corrosion reaction such as oxygen reduction and hydrogen evolution dominate (Cathodic Polarization Curve). These separate halves of the total polarization curve may provide information about the rates of anodic and cathodic processes. The current density at any particular potential is a measure of the... 

In practice, metal dissolution occurs in combination with the reduction of an oxidizing species. Hydrogen evolution and oxygen reduction are frequent counter reactions compensating partially or totally the metal dissolution of such a mixed electrode. Figure 1.40 presents schematically the reactions at an Me/Me + electrode and the H2/H+ electrode. Both electrodes show their Nemst equilibrium potentials and the anodic and cathodic branches of... 

E and E, . represent the equilibrium potential of mineral anodic dissolution and cathode reduction of oxygen, respectively. represents the mineral mixed potential in certain system. and Zg are current density of anodic and cathode reaction, respectively. When the discharge is the controlled step of electrode reaction, according to electrochemistry theory, the equation can be described as following ... 

It is therefore almost impossible to establish the equilibrium potential of the oxygen electrode in aqueous electrolytes (+1.23 V vs rev. hydrogen electrode) and only in a very careful experiment of Bockris and Hug (112) has this been achieved. Usually at zero current density a lower potential of approximately 0.95 V is obtained, which is established by mixed reactions, with the main reaction being the reduction of oxygen to hydroperoxide. 

As described in Chapter 3, and expressed by Nemst s equation (3.9), the equilibrium potential depends on the activity, or practically on the concentration of the species that take part in the reaction. The reduction of the oxygen eoneentration that oeeurs close to the metal surface when the electrode is polarized in the eathodie direetion, leading to the concentration profile shown in Figure 4.4, causes therefore a change in the equilibrium potential. This equilibrium potential shift is regarded as an overvoltage ... 

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