Fuel equilibrium electrode potential

In the given form, the Butler-Volmer equation is applicable rather broadly, for flat model electrodes, as well as for heterogeneous fuel cell electrodes. In the latter case, concentrations in Eq. (2.13) are local concentrations, established by mass transport and reaction in the random composite structure. At equilibrium,/f = 0, concentrations are uniform. These externally controlled equilibrium concentrations serve as the reference (superscript ref) for defining the equilibrium electrode potential via the Nernst equation. 

The electrode reactions taking place at the electrodes of direct methanol fuel cells, the overall current-producing reactions, and the corresponding thermodynamic values of equilibrium electrode potentials EP and electromotive force (EMF) of the direct methanol fuel cell are given as follows ... 

Voltage drop caused by mixed potential and fuel crossover. The mixed potential at electrodes is due to unavoidable parasitic reactions that tend to lower the equilibrium electrode potential. For a mixed cathode potential, a local-cell mechanism has been put forward to explain the Pt-02 reaction mechanism at the electrode in an 02-saturated acidic solution [20, 21], The mixed potential is composed of both the cathodic O2/H2O reaction potential (O2+ 4H -l- 4e 2H2O,... 

If the redox state of the fuel salt is characterized by an uranium ratio [U(IV)]/[U(III)] < 1, the alloy specimens get a more negative stationary electrode potential than equilibrium electrode potentials of some uranium intermetaUic compounds and alloys with nickel and molybdenum. This leads to a spontaneous behavior of alloy formation processes on the specimen siur-face and further diffusion of uranium deep into the metaUic phase. As a consequence, films of intermetaUic compounds and alloys of nickel, molybdenum, and tungsten with uranium are formed on the aUoy specimen surfaces, and IGC does not take place. 

This is considerably higher than that of an H2-O2 fuel cell (i.e., 83%). However, under normal operating conditions, at a current density j, the electrode potentials deviate from their equilibrium values as a result of large overpotentials, r, at both electrodes (Fig. 5) ... 

Initially, both electrodes are at equilibrium. Since the anode has accumulated electrons and the cathode has depleted electrons, electrons begin to flow from electrode from the anode to the cathode. The thermodynamic driving force for the electron flow is the electrode potential difference, which for the fuel cell reaction is 1.23 Y at standard conditions. In addition to electron flow, H + ions produced at the anode diffuse through the bulk solution and react at the cathode. The reaction is able to continue as long as H2 is fed at the anode and 02 at the cathode. Hence, the cell is not at equilibrium. The shift in electrode potential from equilibrium is called the overpotential (>/). 

When a fuel cell is under operation, the electric current passes through both anode and cathode sides, and the electrode potentials of both deviate from their equilibrium values. These deviations are called anode and cathode overpotentials, respectively. At both sides, Equation 1.37 holds. Thus, at the anode side,... 

A common source of error in fuel cell modeling is poor use of input data which must be relevant for the studied condition. For example, the exchange current density for the charge transfer reaction have to be valid for the reference concentrations and reference electrode potential used for the calculation of the concentration overpotential, since this will determine the convergence to the correct equilibrium currents. If possible, values for transport properties must also follow the same reference state to avoid unnecessary sources of inconsistency. 

FIGURE 1.8 Potential distribution in a fuel cell with planar Pt electrodes, (a) Equilibrium (open-circuit) conditions (b) under load, assuming infinite membrane conductivity, and (c) under load, assuming finite membrane conductivity. In cases (b) and (c), the anode polarization is assumed to remain constant, implying a negligible shift of the anode electrode potential. 

The main function of a fuel cell electrode is to convert a chemical flux of reactants into fluxes of charged particles, or vice versa, at the electrochemical interface. Electrochemical kinetics relates the local interfacial current density j to the local interfacial potential drop between metal and electrolyte phases, illustrated in Figure 1.8. A deviation of the potential drop from equilibrium corresponds to a local overpotential q at the interface, which is the driving force for the interfacial reaction. The reaction rate depends on overpotential, concentrations of active species, and temperature. For the remainder of this section, it is assumed that the metal electrode material is an ideal catalyst, that is, it does not undergo chemical transformation and serves as a sink or source of electrons. The basic question of electrochemical kinetics is how does the rate of interfacial electron transfer depend on the metal phase potential ... 

What reaUy distinguishes electrodeposition from other deposition techniques is the applied potential. The applied potential controls the departure from equilibrium and, therefore, the rate of the reaction. The electrodeposition of metals only requires that the electrode potential be driven negative of the equilibrium potential. The difference between the applied potential and the equilibrium potential is called the overpotential. Because the electrode must be poised at a potential for deposition to occur, the substrate must be a conductor or semiconductor. In contrast, vapor phase deposition does not require a conducting substrate. This substrate constraint is not a limitation in the electrodeposition of catalysts for fuel cells, batteries, and photoelectrochanical solar cells. 

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