Hydrogen reduction reversible potential

Sometimes the value of the redox potential attains even at low current densities such values that another simultaneous process is possible. For instance the cathodic reversible potential of the Ti++++/Ti+++ system with the same concentration of both kinds of ions is about tc° = 0.04 V if platinized platinum in a solution of sulphuric acid is used as a cathode the evolution of hydrogen commences at a potential also near zero and the ourrent efficiency with respect to the reduction of Ti++++ ions will be comparatively low. Much better results can be achieved by replacing the platinum by another material, suoh as lead or graphite whioh have an appreciable hydrogen overvoltage, whereby the deposition potential of hydrogen becomes more negative as oompared with the potential of the Ti++++/Ti+++ system. 

Simultaneous Discharge of Cations.—If a solution contains two cations, there is a possibility that simultaneous discharge may occur this problem is not only of interest in connection with the electrodeposition of alloys, but it is important in the deposition of single metals, since aqueous solutions always contain hydrogen ions. Were it not for a variety of complicating factors, such as the influence of one metal on the deposition potential of the other, the situation would, in principle, be relatively simple provided the discharge (reduction) potentials of the two ions were the same, simultaneous deposition would occur. For example, the reversible potential of a metal A in a solution of its ions of activity oa+, i.e., of the electrode A, A+, would be given by... 

Reversible Oxidation-Reduction Processes.— The fundamental principles concerned in the reduction of a reversible system at a cathode or in the oxidation at an unattackable anode have been already given in Chap. Xlll. If the potential of the cathode is made slightly more negative or that of the anode more positive than the reversible potential of the system, reduction or oxidation, respectively, will take place. As the current is increased there will be some polarization due to concentration changes, and eventually the limiting c.d. for the particular process will be attained any further increase will be accompanied by another reaction, e.g., evolution of hydrogen at a cathode or evolution of oxygen or chlorine at an anode. 

Equation (1) can serve as the basis for determining the highest reduction potential or the lowest oxidation potential of an electrocatalytic reaction. Figure 3 shows that spontaneous electroreductions result in a positive E° vs. NHE, often approached with the aid of a catalytic electrode (25, 26, 31, 33, 36). For the example of ethylene shown in Fig. 3, reduction can occur below its reversible potential away from equilibrium. If the ethylene electrode is combined with a hydrogen anode, an electrogenerative cell is formed (16,17,25,26,31,33,36), similar to a fuel cell, that can generate low-voltage. 

A very active metal such as sodium cannot be deposited from aqueous solutions except under special circumstances. The reversible potential for the reduction of iVa is —2.714 V. Even with a lead cathode an enormous current density would be required to bring the cathode below this potential the current efficiency for sodium deposition would be exceedingly small. Sodium can be deposited into mercury, which has a high hydrogen overpotential, if a highly alkaline solution is used. High current densities are required and the current efficiency is very low. Three factors influence the process. 

The total oxidation and reduction current densities will be equal at the point at which the anodic line for the metal dissolution reaction intersects the cathodic line for hydrogen evolution. The potential at which these lines intersect is the corrosion potential. The rate of the anodic reaction at the corrosion potential is the corrosion rate (corrosion current density). The corrosion potential always takes a value between the reversible potentials for the two partial reactions. 

FIGURE 15.4 Schematic Evans diagram illustrating the concepts of corrosion potential, cathodic protection, and anodic protection. (H+/H2) reversible potential for the hydrogen reduction reaction (HRR) %(h /h2)> exchange current density for the HRR on the metal surface con and Wr are the corrosion potential and corrosion current, respectively. See text for discussion of the other quantities. 

An interesting variation of the effect of galvanic coupling occurs with metals that exhibit active-passive transitions. When noble metals such as platinum, which are good catalysts for hydrogen reduction, are coupled to a metal with an active-passive transition below the reversible proton-hydrogen potential, spontaneous passivation may ensue (Fig. 7). Thus, a porous coating of noble metal on titanium, chromium, or stainless steels will result in anodic protection of the substrate. 

The paper fiirther revealed that the reduetion on Fe(II) of OH, formed from water oxidation, had a reversible potential of 0.64 V for Fe(II) (Equation 5.38), which was about the same as that calculated for 1-fold bond Pt (0.61 V at the B3LYP level). The paper suggested that the Fe-OH formation contributed to the observed overpotential on iron maerocycles in the same way that it did for the platinum eleetrode. In eomparison with Pt, the paper concluded, the important difference was the hydrogen bond interaction between (OHOH) bonded to Fe(II) and a nitrogen lone-pair orbital in the N4 chelate. This interaction prevented hydrogen peroxide from leaving as a two-electron reduction product and provided a path for reduction to water. 

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