Hydrogen Evolution and Ionization

FIGURE 15.1 Polarization curves for the reduction of ions in 5X10 MNa2S20g + 0.009 M NaF solution at different metals. 

Cathodic hydrogen evolution is one of the most common electrochemical reactions. It is the principal reaction in electrolytic hydrogen production, the auxiliary reaction in the production of many substances forming at the anode, such as chlorine, and a side reaction in many cathodic processes, particularly in electrohydrometallurgy. It is of considerable importance in the corrosion of metals. Its special characteristic is the fact that it can proceed in any aqueous solution particular reactants need not be added. The reverse reaction, which is the anodic ionization of molecular hydrogen, is utilized in batteries and fuel cells. 

Most metals (other than the alkali and alkaline-earth metals) are corrosion resistant when cathodically polarized to the potentials of hydrogen evolution, so that this reaction can be realized at many of them. It has thus been the subject of innumerable studies, and became the fundamental model in the development of current kinetic concepts for electrochemical reactions. Many of the principles 

In the past, elevated voltages in electrolysis cells (a cell overvoltage) had been attributed mainly to polarization of the hydrogen evolution reaction. Hence the term hydrogen overvoltage became common for this kind of polarization. 

The range of current densities encountered in cathodic hydrogen evolution is extremely wide. Corrosion of metals may become important at rates of hydrogen evolution on the order of 10 to 10 mA/cm, whereas in industrial electrolyzers, where this reaction occurs at the cathode, the current densities attain values of 10 mA/cm (10 A/m ) and more. 

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FIGURE 28.2 Relation between the exchange current densities of hydrogen evolution and ionization at different metals and the electron work functions. 

In the present chapter we want to look at certain electrochemical redox reactions occurring at inert electrodes not involved in the reactions stoichiometrically. The reactions to be considered are the change of charge of ions in an electrolyte solution, the evolution and ionization of hydrogen, oxygen, and chlorine, the oxidation and reduction of organic compounds, and the like. The rates of these reactions, often also their direction, depend on the catalytic properties of the electrode employed (discussed in greater detail in Chapter 28). It is for this reason that these reactions are sometimes called electrocatalytic. For each of the examples, we point out its practical value at present and in the future and provide certain kinetic and mechanistic details. Some catalytic features are also discussed. 

The above-mentioned circumstances have the result that the potentiostatic method, which is used with success for investigating simple electrode reactions (evolution and ionization of hydrogen or oxygen), makes it possible to obtain information on the stationary oxidation process only with appropriate selection of potential scanning rate (e. g., for investigation of the oxidation of ethylene the rate should be less than 10" V/sec) [172]. 

It was demonstrated by R. Parsons and H. Gerischer that the adsorption energy of the hydrogen atom determines not only the rate of the Volmer reaction (5.7.1) but also the relative rates of all three reactions (5.7.1) to (5.7.3). The relative rates of these three reactions decide over the mechanism of the overall process of evolution or ionization of hydrogen and decide between possible rate-determining steps at electrodes from different materials. 

It was also observed, in 1973, that the fast reduction of Cu ions by solvated electrons in liquid ammonia did not yield the metal and that, instead, molecular hydrogen was evolved [11]. These results were explained by assigning to the quasi-atomic state of the nascent metal, specific thermodynamical properties distinct from those of the bulk metal, which is stable under the same conditions. This concept implied that, as soon as formed, atoms and small clusters of a metal, even a noble metal, may exhibit much stronger reducing properties than the bulk metal, and may be spontaneously corroded by the solvent with simultaneous hydrogen evolution. It also implied that for a given metal the thermodynamics depended on the particle nuclearity (number of atoms reduced per particle), and it therefore provided a rationalized interpretation of other previous data [7,9,10]. Furthermore, experiments on the photoionization of silver atoms in solution demonstrated that their ionization potential was much lower than that of the bulk metal [12]. Moreover, it was shown that the redox potential of isolated silver atoms in water must... 

The primary photochemical step, resulting from the absorption of the photon, must produce transition to a spectroscopically recognizable upper state—excitation or ionization, in which the Franck-Condon principle is observed. Which of these is involved in the process of hydrogen evolution could be clearly ascertained for one of the simple systems in which Reaction 1 occurs 02-free aqueous solutions of iodide ions. The absorption spectrum of such solutions is well known, showing two bands with maxima at 226 and 194 m/z. Photolytic light at 254 m/z is absorbed near the onset of the 226 m/z max. band. 

The more massive post-AGB stars (M > 0.6M ) with higher luminosity contract more quickly (L > 1O4L0, tK 10000y). When their surface temperature exceeds 30 000 K, surface hydrogen is fully ionized, the Lyman continuum opacity drops and ultraviolet photons are able to ionize the much-expanded shell, which now becomes visible as a planetary nebula. Schematically we might represent the evolution as follows ... 

Sites where dissolution is occurring are assumed to be generally anodic to the equilibrium potential and sites where oxygen reduction or hydrogen evolution is occurring are assumed to be cathodic. At the anodic sites, because of local interactions between the metal atoms and the electrolyte, the electron cloud is assumed to be oriented toward the electrode. The capacitance of the double layer is thereby decreased and makes the metal atom cores more susceptible to ionization and dissolution. 

The solid lines in Figure 18.5 represent electrochemical equilibria. The dashed vertical line corresponds to the chemical self-ionization shown in Eq. (18.18), at equal concentrations of the two ions. The lines bounding the shaded area are the reversible potentials for oxygen and hydrogen evolution as functions of pH. 

Stabilization processes may differ in nature. For hydrogen evolution reactions, the conversion of adsorbed hydrogen into H2 as a result of activationless electrochemical desorption[180,181] can be one of the ways of stabilization. Since the activationless ionization probability is decreased because of a low proton tunneling probability, adsorbed hydrogen is stable for some time which is sufficient for H3O ions to approach it. These ions and the adsorbed hydrogen enter into the reaction of electrochemical desorption competing with the reverse process, viz. ionization. 

Figure 1T2 shows anodic d cathodic polarization curves for the partial CD of dissolution 4 and deposition 4 of the metal and for the partial CD of ionization 4 and evolution 4 of hydrogen, as well as curves for the overall reaction current densities involving the metal (4) and the hydrogen (4). The spontaneous dissolution current density 4 evidently is determined by the point of intersection. A, of these combined curves. 

Similar size effects have been observed in some other electrochemical systems, but by far not in all of them. At platinized platinum, the rate of hydrogen ionization and evolution is approximately an order of magnitude lower than at smooth platinum. Yet in the literature, examples can be found where such a size effect is absent or where it is in the opposite direction. In cathodic oxygen reduction at platinum and at silver, there is little difference in the reaction rates between smooth and disperse electrodes. In methanol oxidation at nickel electrodes in alkaline solution, the reaction rate increases markedly with increasing degree of dispersion of the nickel powders. Such size effects have been reported in many papers and were the subject of reviews (Kinoshita, 1982 Mukerjee, 1990). 

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