Thermodynamic oxygen evolution

If the potential of a metal surface is moved below line a, the hydrogen reaction line, cathodic hydrogen evolution is favored on the surface. Similarly a potential below line b, the oxygen reaction line, favors the cathodic oxygen reduction reaction. A potential above the oxygen reaction line favors oxygen evolution by the anodic oxidation of water. In between these two lines is the region where water is thermodynamically stable. 

Several significant electrode potentials of interest in aqueous batteries are listed in Table 2 these include the oxidation of carbon, and oxygen evolution/reduction reactions in acid and alkaline electrolytes. For example, for the oxidation of carbon in alkaline electrolyte, E° at 25 °C is -0.780 V vs. SHE or -0.682 V (vs. Hg/HgO reference electrode) in 0.1 molL IC0 2 at pH [14]. Based on the standard potentials for carbon in aqueous electrolytes, it is thermodynamically stable in water and other aqueous solutions at a pH less than about 13, provided no oxidizing agents are present. 

These reactions can be easily combined, if necessary, with an anodic reaction such as oxygen evolution to estimate the thermodynamic standard free energy, AG ... 

The stability of the charge that is stored is limited by the thermodynamically favored reactions connected with both hydrogen and oxygen evolution [348]... 

As the system is thermodynamically unstable with respect to hydrogen and oxygen evolution, lead-acid cells are subject to self-discharge ... 

Oxygen evolution occurs at a much lower potential and is therefore favored thermodynamically ... 

Aqueous solutions all evolve H2 when the cathode potential is made sufficiently negative. However, it may be possible to have an aqueous solution that contains inexpensively dissolved substances (e.g., S02) that become oxidized at potentials much less than that of water itself. Then, looked at thermodynamically, the reversible potential of the reaction in the cell would be less than that of water. In addition, the i0 value for oxygen evolution (i0 1010 A cm-2 at 25 °C) is particularly low and the anode overpotential particularly high. Substitution of, e.g., S02 oxidation could be achieved at a lesser overpotential than with 02 evolution. 

Molten salts are favored electrolytes in cases where aqueous solutions cannot be used because the decomposition voltage of water is lower than that of the salt in question. Although high overvoltages for hydrogen and oxygen evolution allow us to extend the use of aqueous electrolytes somewhat beyond the limits set by thermodynamics, there are many substances which cannot be electrowon or plated from aqueous solutions. 

The region of stability of water is 1.229 V independent of pH, since the reversible potentials for hydrogen and oxygen evolution change with pH in the same manner. This, incidentally, is the potential region in which the f l cell can operate. Thermodynamic considerations... 

At active electrodes there is a strong electrode (M)-hydroxyl radical ( OH) interaction. In this case, the adsorbed hydroxyl radicals may interact with the anode with possible transition of the oxygen from the hydroxyl radical to the anode surface, forming the so-called higher oxide (Eq. 78). This may be the case when higher oxidation states on the surface electrode are available above the thermodynamic potential for oxygen evolution (1.23 V/RHE). 

Most reviewers5 8 now argue that photosynthetic oxygen evolution results from a sequential four-step electron-transfer process in which oxidizing equivalents from chi fl+- are accumulated in a "charge-storing" complex to accomplish the concerted four-electron oxidation of two H2O molecules to one O2 molecule. The photooxidant (chi + ) and reductant (pheo O are one-electron transfer agents, and the matrix is the lipoprotein thylakoid membrane. Hence, evaluation and consideration of the one-electron redox potentials for PS 11 components within a lipoprotein matrix are necessary in order to assess the thermodynamic feasibility of proposed mechanistic sequences. 

The consequence of all these effects is that oxygen, or hydrogen, evolution reactions follow quite different kinetics than would be expected from thermodynamical considerations. For example, oxygen evolution on a platinum electrode starts at a potential significantly higher than that predicted by the Nernst equation. The reason for this is the formation of surface oxide and its associated dynamics as elucidated, besides others, by B. Conway [4]. 

In the most favorable case (complex 1), it was possible to observe the production of a photoanodic current in the presence of plain water/LiC104 0.1 M at pH 5 (HCIO4). However, despite a reasonable negative free energy difference for water oxidation (ca. 0.46 V, considering that the thermodynamic potential for oxygen evolution at pH 5 is 0.95 V vs NHE), the photocurrent was extremely small, ca. 

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