Oxygen/water half-cell reaction

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. 

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 ... 

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. 

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 ... 

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. 

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 ... 

At the anode of a PEMFC, hydrogen is oxidized, creating protons and electrons. The polymer membrane provides proton-conducting pathways, whereas the electrons are forced through an external circuit by a potential difference between the anode and cathode. Within the CCL, oxygen is reduced to water in the presence of protons and electrons. The respective half-cell reactions, typically catalyzed by cost-intensive platinum, are... 

In all cases, metal atoms in the solid (s) metallic state (in the electronic conductor) are converted to an oxidized state. These species remain either on the metallic surface or are ejected into the aqueous (aq) solution (the ionic conducting phase). Typical cathodic half-cell reactions in aqueous media involve the reduction of oxygen, protons, ind water at the interface between the metallic and ionic conducting phases. This process involves consumption of electrons liberated during oxidation ... 

In the anodic reaction, the metal that is more easily oxidized loses electrons to produce metal ions. In (a), this is iron in (b), it is zinc. In the cathodic reaction, oxygen gas, which is dissolved in a thin film of water on the metal, is reduced to OH. Rusting of iron occurs in (a), but it does not in (b). When iron corrodes, Fe and OH ions from the half-cell reactions initiate these further reactions. 

The electrons allow the reduction of oxygen in the presence of water to form four hydroxide ions, which replenish the hydroxide ions used up at the anode. By combining the two half-reactions, the overall cell reaction is determined. As you can see, it is the same as the equation for the burning of hydrogen in oxygen. 

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.

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.

The electrolysis of water, shown in Figure 3.4, leads to the cell reaction in which water is broken down into its elements, H2 and O2. Recall that hydrogen gas and oxygen gas combine spontaneously to form water and are used to power fuel cells, which produce electricity. Therefore, the reverse process (electrolysis of water) is nonspontaneous and requires electrical energy. The two half-reactions occur at the anode and cathode. 

In the previous section, we discussed fuel cell thermodynamics. However, in reality, fuel cell operation with an external load is much more practical than in a thermodynamic state. When a H2/air PEM fuel cell outputs power, the half-electrochemical reactions will proceed simultaneously on both the anode and the cathode. The anode electrochemical reaction expressed by Reaction (l.I) will proceed from H2 to protons and electrons, while the oxygen from the air will be reduced at the cathode to water, as expressed by electrochemical Reaction (l.II). For these two reactions, although the hydrogen oxidation reaction (HOR) is much faster than the oxygen reduction reaction (ORR), both have limited reaction rates. Therefore, the kinetics of both the HOR and the ORR must be discussed to achieve a better understanding of the processes occurring in a PEM fuel cell. 

One of the most widely used applications of electrolytic cells is in electrolysis, the decomposition of a compound. Water may be decomposed into hydrogen and oxygen. Aluminum oxide may be electrolyzed to produce aluminum metal. In these situations, several questions may be asked How longw xW it take how much can be produced what current must be used Given any two of these quantities, the third may be calculated. To answer these questions, the balanced half-reaction must be known. Then the following relationships can be applied ... 

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