Over potential hydrogen evolution

Adsorption of surface-active substances is attended by changes in EDL structure and in the value of the / -potential. Hence, the effects described in Section 14.2 will arise in addition. When surface-active cations [NR] are added to an acidic solution, the / -potential of the mercury electrode will move in the positive direction and cathodic hydrogen evolution at the mercury, according to Eq. (14.16), will slow down (Fig. 14.6, curve 2). When I ions are added, the reaction rate, to the contrary, will increase (curve 3), owing to the negative shift of / -potential. These effects disappear at potentiafs where the ions above become desorbed (at vafues of pofarization of less than 0.6 V in the case of [NR]4 and at values of polarization of over 0.9 V in the case of I ). 

The range of potentials over which potential sweep experiments are carried out in aqueous solutions is from around 0.00 to 1.5 V on the NHS. This is to avoid hydrogen evolution on the negative side and Oz evolution on the positive side (pH 0). If then, one had a steady current, the limitingly slow sweep rate would be one that covers 1500 mV in 27 s. On this basis, the lowest sweep rate should be about 50 mV s"1. 

Saturating the electrolyte with iron(lll) hydroxide (e.g., by addition of aqueous solutions of ferric nitrate) and simultaneously adding cobaltous salts leads to in situ formation of a mixed Fe(llI)/Co(ll)/Co(IIl) deposit, which exhibits catalytic activity comparable to that of Fe304 shown by the current voltage curve in Fig. 11. Such mixed oxidic catalyst coatings are composed of very small oxide crystals, which evidently are dissolved upon current interruption due to dissociative oxide dissolution. The transfer of dissolved metal ions to the cathode followed by cathodic deposition of the metal, however, can be completely prohibited, if the potential of the cathode due to optimal electrocatalysis of cathodic hydrogen evolution proceeds with an over-... 

Over the past several years, Gruen and coworkers have examined the SH response from iron electrodes in alkaline solutions [45, 53, 172]. In their work on polycrystalline iron, they concluded that the potential dependent SH response which was observed during surface oxidation could be attributed to two intermediate phases on the electrode surface between the passive film at oxidative potentials and the reduced metal at hydrogen evolution potentials [53]. They have recently extended this work to Fe(110). In this study [172], they examined the SH rotational anisotropy from this crystal under ambient conditions. They found that the experiments reveal the presence of both twofold and threefold symmetric species at the metal/oxide interface. When their data is fit to the theory of Tom et al. [68], they conclude that the measured three-fold symmetric oxide is found to be tilted by 5° from the Fe(110) plane. The two-fold symmetric structure is aligned with the Fe(110) surface. 

These studies of reduction of benzenoid aromatics reveal that the solvent, the electrolyte cation, the current density and the water content are all important variables. In general it is important to have a rather negative potential (large TAA+) and a proton source (water) present under conditions where hydrogen evolution or attack on the solvent does not occur. Under such conditions difunctional molecules can be selectively reduced by control over the number of Faradays/mole which are passed. This kind of predictable selectivity should give the electrochemical method real advantage over alkali metal reductions and the possibility to use materials other than liquid ammonia and alkali metal is quite attractive. 

Other disadvantages of the existing aqueous technology are economic in nature, such as the low current efficiency of the reduction of Cr(VI) in acid media. In addition, the difference in over-potential between chromium and hydrogen reduction results in the evolution of hydrogen gas, which can lead to hydrogen embrittlement in the substrate. 

The Tafel slope for this mechanism is 2.3RT/PF, and this is one of the few cases offering good evidence that P = a, namely, that the experimentally measured transfer coefficient is equal to the symmetry factor. A plot of log i versus E for the hydrogen evolution reaction (h.e.r.), obtained on a dropping mercury electrode in a dilute acid solution is shown in Fig. 4F. The accuracy shown here is not common in electrode kinetics measurements, even when a DME is employed. On solid electrodes, one must accept an even lower level of accuracy and reproducibility. The best values of the symmetry factor obtained in this kind of experiment are close to, but not exactly equal to, 0.500. It should be noted, however, that the Tafel line is very straight that is, P is strictly independent of potential over 0.6-0.7 V, corresponding to five to six orders of magnitude of current density. 

In order to suppress or inhibit zinc corrosion and hydrogen evolution at zinc anode, a number of additives are selected and used for the anode. Because of its high over potential to the hydrogen evolution reaction, mercury is an effective gassing suppresser and used to be widely employed in alkaline Zn/Mn02... 

Of special interest for the topic of the present chapter is the observation of Weaver that while the double-layer-corrected AS quantities are ligand sensitive, they are found to be independent of potential. This is not the case for the atom and electron transfer process involved in the hydrogen evolution reaction at Hg studied by Conway, et where an appreciable potential dependence of AS is observed, corresponding to conventionally anomalous variation of the Tafel slope with temperature. Unfortunately, in the work with the ionic redox reactions, as studied by Weaver, it is only possible to evaluate the variation of the transfer coefficient or symmetry factor with temperature with a limited variety of redox pairs since Tafel slopes, corresponding to any appreciable logarithmic range of current densities, are not always easily measurable. Alternatively, evaluation of a or /3 from reaction-order determination requires detailed double-layer studies over a range of temperatures. 

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