Water-oxygen half-cell

The explicit aims of boiler and feed-water treatment are to minimise corrosion, deposit formation, and carryover of boiler water solutes in steam. Corrosion control is sought primarily by adjustment of the pH and dissolved oxygen concentrations. Thus, the cathodic half-cell reactions of the two common corrosion processes are hindered. The pH is brought to a compromise value, usually just above 9 (at 25°C), so that the tendency for metal dissolution is at a practical minimum for both steel and copper alloys. Similarly, by the removal of dissolved oxygen, by a combination of mechanical and chemical means, the scope for the reduction of oxygen to hydroxyl is severely constrained. 

The permanganate ion is a powerftil oxidizing agent that oxidizes water to oxygen under standard conditions. Here are the half-cell reactions ... 

The oxygen/water half-cell reaction has been one of the most challenging electrode systems for decades. Despite enormous research, the detailed reaction mechanism of this complex multi-step process has remained elusive. Also elusive has been an electrode material and surface that significantly reduces the rate-determining kinetic activation barriers, and hence shows improvements in the catalytic activity compared to that of the single-noble-metal electrodes such as Pt or Au. 

One other important criterion for successful water cleavage that must be considered is the solution pH. Although the potential difference between the two half reactions for water decomposition is fixed at 1.23 V and is independent of pH, the half-cell reactions are dependent upon pH (Figure 4). Thus, by altering the pH of a solution it is sometimes possible to alter the half-cell potentials to be compatible with the redox properties of a photosensitizing catalyst. The oxidant must have a redox potential above the oxygen line, whilst the reductant must have a redox potential below the hydrogen line. The effect of pH is illustrated in subsequent sections of this chapter. 

To avoid these problems, engineers have focused on a cell in which the fuel is hydrogen gas, the oxidant is oxygen from the air, and the product is water vapor. One of the more promising hydrogen fuel cells is one in which the half-cell reactions are separated by a thin polymer sheet called a proton-exchange membrane (PEM). The PEM fuel cell operates at approximately 100°C, and the moist membrane itself is the electrolyte. 

Because this is more positive than the half-cell voltage (02,H30 (10 m) I H2O) = 0.815 V for pure watei CI2 has a greater tendency to be reduced than O2 therefore, CC has a lesser tendency to be oxidized than H2O, and the anode product is 02(g). If we try to increase the external potential above 1.229 V, all that will happen is that water will be electrolyzed at a greater rate to produce hydrogen and oxygen. Sodium and chlorine will not appear as long as sufficient water is present. 

The half cell reactions for hydrogen and oxygen form a starting point from which to consider redox systems in water. The Nernst equation for the reduction of oxygen may be written in terms of pH ... 

Fig. 1.7 Ranges of half-cell potentialsof some electrochemical reactionsof importance in corrosion. Vertical bars represent metal ion concentration of 1 molal (approximately 10%) down to 1 ppm. Dashed extensions may apply with precipitated and complexing species. The hydrogen and oxygen reactions depend on both pH and pressure of the gases. Values for the hydrogen are at one atmosphere pressure. Values for oxygen are for water in contact with air (aerated) giving 10 ppm dissolved oxygen and for water deaerated to 1 ppb dissolved oxygen.

If two or more electrochemical half-cell reactions can occur simultaneously at a metal surface, the metal acts as a mixed electrode and exhibits a potential relative to a reference electrode that is a function of the interaction of the several electrochemical reactions. If the metal can be considered inert, the interaction will be between species in the solution that can be oxidized by other species, which, in turn, will be reduced. For example, ferrous ions can be oxidized to ferric ions by dissolved oxygen and the oxygen reduced to water, the two processes occurring at different positions on the inert metal surface with electron transfer through the metal. If the metal is reactive, oxidation (corrosion) to convert metal to ions or reduction of ions in solution to the neutral metal introduces additional electrochemical reactions that contribute to the mixed electrode. 

The cycle of light-induced water decomposition can then be closed either by addition of Ru02 catalyst to the same solution or coupling to an oxygen producing half cell. 

We have seen earlier how hydrogen and oxygen may be combined and spontaneously form water molecules and that this reaction produces energy which may be used in fuel cells. The opposite process where hydrogen and oxygen are formed from water molecules is not spontaneous but requires an electrolytic process. The following half cell reaction takes place at the anode ... 

Approximate limits on the range of observed half-cell potentials in biochemical systems may be set from the following considerations. From Eqs. (3) and (6), the hydrogen cell potential at pH 7 is determined to be - 416 mV. This sets an approximate lower limit on E° for biochemical reductants, since reductants with lower potential in aqueous solution would reduce protons to hydrogen. (Note Reducing centers protected from exposure to water by protein can have lower potentials, however.) An upper limit on reduction potentials is set by the oxygen half-cell reaction ... 

Figure 1, The platinized Chi a cell. The Chi a-free electrode is used as a half cell in a liquid-junction photovoltaic cell. In photolytic reaction, only the platinized Chi a electrode is used in the production of molecular hydrogen and oxygen from water.

Interpretation. The presence of an above-critical amount of chloride ions at the rebars leads to depassivation and in the presence of oxygen and water to corrosion attack. From chloride profiles information on the transport of chlorides into the concrete (Chapter 6) can be obtained. In combination with results from potential mapping, the critical chloride content for the specific structure can be obtained. On chloride-contaminated structures an empirical correlation between chloride content and half-cell potential could be established, thus the chloride distribution can be roughly obtained from the potential map. 

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