Equilibrium potential oxygen reaction

Dependence of oxygen-reaction equilibrium potential on oxygen-gas partial pressure... 

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. 

At almost all electrodes, reaction (15.20) occurs with appreciable polarization in both the cathodic and anodic directions the exchange CD of this reaction is very low 10 ° to 10 mA/cm. For this reason the equilibrium potential of this reaction is not established at the electrodes. The OCP measured in oxygen is between 0.85 and 1.1 V (RHE) that is, it is 0.15 to 0.4 V more negative than the equilibrium potential. 

At mercury and graphite electrodes the kinetics of reactions (15.21) and (15.22) can be studied separately (in different regions of potential). It follows from the experimental data (Fig. 15.6) that in acidic solutions the slope b 0.12 V. The reaction rate is proportional to the oxygen partial pressure (its solution concentration). At a given current density the electrode potential is independent of solution pH because of the shift of equilibrium potential, the electrode s polarization decreases by 0.06 V when the pH is raised by a unit. These data indicate that the rate-determining step is addition of the first electron to the oxygen molecule ... 

For metal electrodes, the anodic 03Q n reaction proceeds at electrode potentials more anodic than the equilibrium potential Bo of the reaction as shown in Fig. 10-14. For n-type semiconductor electrodes, the anodic photoexdted oxygen reaction proceeds at electrode potentials where the potential E of the valence band edge (predsely, the potential pEp of the quasi-Fermi level of interfadal holes, pCp = — CpEp) is more anodic than the equilibrium oxygen potential Eq, even i/the observed electrode potential E is less anodic than the equilibrium oxygen potential E03. The anodic hole transfer of the o Qgen reaction, hence, occurs at photoexdted n-type semiconductor electrodes even in the range of potential less anodic than the equilibriiun potential Eq of the reaction as shown in Fig. 10-14. 

The potential Eo ) at which the anodic photoexdted oxygen reaction occurs is lower by an amount AE than the equilibrium potential 03, which is the potential for the onset of the anodic o tygen reaction without photoexdtation. This potential sifts, is nearly equivalent to the band gap e, as indicated in Eqn. 10-17 ... 

Fig. 10-14. Energy levels and polarization curves (current vs. potential) for anodic transfer ofphotoexdted holes in oxygen reaction (2 HgO. -t- 4h O24 4 H. ) on a metal electrode and on an n-type semiconductor electrode j = anodic reaction current ep(02 20)- Fermi level of oxygen electrode reaction dCpi, = gain of photoenergy q = potential for the onset of anodic photoexdted ox en reacti . 4 pi, (=-Ae.. le) = shift of potential for the onset of anodic oxygen reaction from equilibrium oxygen potential in the negative direction due to gain of photoenergy in an n-type electrode Eib = flat band potential of an n-type electrode.

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 ... 

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 ... 

The equilibrium potential observed is depicted in the Figure 1 where we illustrate the potential-current behavior for the cathodic reaction (the oxygen reduction) coupled with a possible unspecified anodic oxidation of organic impurities. At electrochemical equilibrium i. = i and no net current flows in the circuit,... 

Figure 6.17. Free energy diagram for oxygen reduction over Pt(lll) at low surface coverage at zero cell potential (U=0), at the equilibrium potential (zero overpotential) (U=1.23) and at the highest potential where all reaction steps are exothermic (adapted from [117]).

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. 

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. 

Metal M is more noble than M and is not corroded because the oxygen reduction reaction is far too slow to bring M to a potential where it can be corroded. Hydrogen evolution is a faster reaction but cannot corrode M because its equilibrium potential is more negative than E and therefore the reaction cannot proceed on M under those conditions. 

Theoretically, it can be any reaction with an equilibrium potential that is more positive than the equilibrium potential of the metal-dissolution reaction. In practice, it is a reaction of the type of A+we D, where A is an electron-acceptor species present in the electrolyte that is in contact with the corroding metal. In aqueous electrolytes, the electron acceptors invariably present are HaO+ ions and dissolved oxygen, the corresponding electronation reactions being... 

The standard equilibrium potential of the oxygen electrode, corresponding to the reaction... 

The standard equilibrium potential at the anode related to reaction (XXIV-7 is 7c° = 0.356 V. As oxygen is evolved owing to overvoltage from neutral solutions as late as the potential is about 1.2 V and from alkaline solution at about 0.8 V, the oxidation of ferrocyanide to ferricyanide can proceed with a 100 per... 

Figure 3.2.2 The free energy diagram of oxygen evolution on 110 facets of rutile RuC>2 at three different potentials V - 0 V versus the RHE the equilibrium potential, Ee, U - 1.23 V and the smallest potential where all reaction steps are downhill, AGOER, U- 1.60 V. The 3.2 eV difference between OH and HOO is shown by the arrow on the right the smaller arrow to the left indicates the universal descriptor for OER.

The cell reaction is the transfer of oxygen from one side to the other. (See also Lambda probe). In the case of an - electrochemical equilibrium (subentry of -> equilibrium) the measured -> open circuit potential (subentry of - potential) or -> equilibrium potential (subentry of -> potential) Ueq or E (emf) can be calculated by the -> Nernst equation ... 

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