Anodes electrode potential

In order to explain the changing optical properties of AIROFs several models were proposed. The UPS investigations of the valence band of the emersed film support band theory models by Gottesfeld [94] and by Mozota and Conway [79, 88]. The assumption of nonstoichiometry and electron hopping in the model proposed by Burke et al. [87] is not necessary. Recent electroreflectance measurements on anodic iridium oxide films performed by Gutierrez et al. [95] showed a shift of optical absorption bands to lower photon energies with increasing anodic electrode potentials, which is probably due to a shift of the Fermi level with respect to the t2g band [67]. 

For some metallic electrodes, such as transition metals, metal ions dissolve directly from the metallic phase into acidic solutions tiiis direct dissolution of metal ions proceeds at relatively low (less anodic) electrode potentials. The direct dissolution of metal ions is inhibited by the formation of a thin oxide film on metallic electrodes at higher (more anodic) electrode potentials. At still higher electrode potentials this inhibitive film becomes electrochemically soluble (or apparently broken down) and the dissolution rate of the metal increases substantially. These three states of direct dissolution, inhibition by a film, and indirect dissolution via a film (or a broken film) are illustrated in Fig. 11-9. 

The energetic basis for the electron-transfer oxidation includes the thermodynamic potential of oxidation (E°ox) for the electron transfer from RH in Eq. (7). Such an electron detachment is commonly effected at an electrode, by an oxidant, or with light. The oxidation is driven electrochemically by the anodic electrode potential, which matches the E°m value. Likewise, the driving force in the chemical oxidation of RH is provided by the redox potential (fi°ed) of the electron acceptor or oxidant (A) according to Eq. (5). 

Figure 3.3.14 Experimental ORR activity of dealloyed Pt-Cu and Pt-Ni core-shell nanoparticle ORR catalysts compared to a pure-Pt nanoparticle catalyst. All three catalyst particles are supported on a high surface area carbon material indicated by the suffix 1C. The shift of the j-E curve of the core-shell catalysts indicates the onset of oxygen reduction catalysis at a more anodic electrode potential (equivalent to a lower overpotential) and hence represents improved ORR reactivity compared to pure Pt.

Potentiometric transducers measure the potential under conditions of constant current. This device can be used to determine the analytical quantity of interest, generally the concentration of a certain analyte. The potential that develops in the electrochemical cell is the result of the free-energy change that would occur if the chemical phenomena were to proceed until the equilibrium condition is satisfied. For electrochemical cells containing an anode and a cathode, the potential difference between the cathode electrode potential and the anode electrode potential is the potential of the electrochemical cell. If the reaction is conducted under standard-state conditions, then this equation allows the calculation of the standard cell potential. When the reaction conditions are not standard state, however, one must use the Nernst equation to determine the cell potential. Physical phenomena that do not involve explicit redox reactions, but whose initial conditions have a non-zero free energy, also will generate a potential. An example of this would be ion-concentration gradients across a semi-permeable membrane this can also be a potentiometric phenomenon and is the basis of measurements that use ion-selective electrodes (ISEs). 

Here, c02 and cm are oxygen and methanol concentrations in the respective catalyst layer, CQ2ref and cAfref are reference concentrations, r]c = cpe — (pc, rja = q>a — exchange current densities per unit volume (A cm-3), aa, ac are transfer coefficients, and ya, yc orders of reaction. 

Application of equation (4) to the organic cocktail oxidation profiles, given in Fig. 6(a), yields straight line plots. Two of these kinetic plots for current densities of 0.14 and 0.40 A cm are presented in Fig. 8. The linear relationships obtained demonstrate that pseudo first-order kinetics are obeyed. This relationship directly supports the heterogeneous bimolecular reaction mechanism proposed for the electrochemical oxidation of organics in the electrochemical reactor. The slopes of the linear plots yield the pseudo first-order rate constants, which are summarized in Table 2 for each value of the current density (i.e. the anodic electrode potential) used. It can be seen from the table that, with increasing current... 

Figure 10.15 Schematic potential-distance diagram for cathodic protection with sacrificial anodes. = electrical potential in the metalhc materials (structure, anodes and connections), Ea = electrode potential of the anode, = electrode potential of the structure (the cathode), AE = E - E, = electrical potential drop in the corrosive medium (the water, the electrolyte) from the anode to the cathode.

It can be seen from these relations that the reversible potential is dependent on temperature and pressure since the Gibbs free energy is a function of temperature and the activity coefficients are dependent on temperature, pressme for gases and ionic strengths for ionic electrolytes. The Nemst equation (1.7) is used to derive a formula for calculating the reversible cell potential as follows Anode electrode potential ... 

FIGURE 1.7. Voltage distribution in an electrochemical reactor = cell voltage E = anodic electrode potential 1 = voltage drop in anolyte = voltage drop in diaphragm = voltage drop in catholyte = cathodic overpotential. 

Anode electrode potential assumed equal to the electrode potential E at the cathode (calculations later showed that E could be adequately represented by E = O.OeiSi + 0.207 V with i in kA/m )... 

Standard anode electrode potential Standard cathode electrode potential Defined in Eq. (5.3) = nLI2d Effectiveness factor for the uniformity of current distribution in parallel plate cells... 

When the electromotive force — E ) is supplied, the corrosion cell is formed with the current flowing between the anode and the cathode. The cathodic electrode potential is shifted to the less noble direction and the anodic electrode potential is shifted to the more noble direction. The shifting of potentials are called cathodic and anodic polarization. The reaction rate... 

Once the fuel cell circuit is connected, there is a current flow, which causes the electrode to be polarized. This means that the anode electrode potential will move to a more positive value, and the cathode electrode potential will move to a more negative value, resulting in a decrease in the cell voltage, which is known as the voltage loss. The principle is shown in Figure 1.26. 

It is possible, nevertheless, to give a general statement about the product expected at the anode. Electrode potentials, as you have seen, depend on concentrations. It turns out that when the solution is concentrated enough in Cl, CI2 is the product but in dilute solution, O2 is the product. To see this, you would simply apply the Nemst equation to the CUICI2 half-reaction. 

The initial reaction takes place at the surface of the electrode and then the intermediates diffuse into the solution where they participate to secondary reactions. The oxidations take place at the anode with initial formation of radical cations as reactive intermediates and the reductions occur at the cathode, with formation of radical anions. At a sufficiently high positive (i.e. anodic) electrode potential, monomers undergo electrochemical oxidation, the polymerization process starts and cation radicals or other reactive species are formed. 

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. 

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