Normal hydrogen electrode half-reactions

Fig. 11 Correlation between electrochemical potentials and OMTS bands for more than ten compounds including polyacenes, phthalo-cyanines, and porphyrins. OMTS data were acquired both from tunnel junctions and STM measurements. The standard potential relative to the normal hydrogen electrode associated with the half reaction M(solution) + e-(vac) —> M-(solution) is the y axis. The three outliers are assigned to the ring oxidation of porphyrins. (Reprinted with permission from [26])...

The hydrogen electrode. The hydrogen electrode is discussed first because it is the primary reference electrode used to define an internationally accepted scale of standard potentials in aqueous solution. By convention, the potential of an electrode half-reaction that is measured with respect to the normal hydrogen electrode (NHE also written as SHE, standard hydrogen electrode) is defined as the electrode potential of the half reaction. This convention amounts to an arbitrary assignment for the standard potential of the hydrogen electrode as zero at all temperatures. Thus, there is in effect a separate scale of electrode potentials at each temperature level. 

As indicated previously, it is desirable to consider the individual electrode reactions independently. One might suppose that this could be achieved by characterizing the individual electrodes as described in Section 3.1.3. However, for reasons of sound thermodynamics, another method has been established. It was decided to relate all electrode reactions to one common reference electrode. Electrochemists have chosen the H+/H2 reaction under standard conditions (ct 1+ = 1M p 12 = 1 bar) as such a general reference electrode. It is termed the normal hydrogen electrode or the standard hydrogen electrode (SHE). Thus, whenever E and E° values are presented for individual electrode reactions (half cells), it is understood that these values pertain to a complete cell in which the SHE constitutes the second electrode. 

Single-electrode potentials, corresponding to half-cell reactions, are often listed. These are actually potentials of the given electrodes relative to the normal hydrogen electrode as a standard. For example, in order to determine the electrode potential for the half-cell reaction... 

Unlike the table of the Electrochemical Series, which lists standard potentials, values for radicals are experimental values with experimental conditions given in the second column. Since the measurements leadingtopotentialsforionradicalsare very dependent on conditions, an attempt to report standard potentials for radicals would serve no useful purpose. For the same reason, the potentials are also reported as experimental values, usually a half-wave potential (i i/2 in polarography) or a peak potential (Ep in cyclic voltammetry). Unless otherwise stated, the values are reported vs. SCE (saturated calomel electrode). To obtain a value vs. normal hydrogen electrode, 0.241 Vhastohe added to the SCE values. All the ion radicals chosen for inclusion in the tables result from electrochemically reversible reactions. More detailed data on ion radicals can be found mHaeEncyclopedia of Electrochemistry of Elements, (A. J. Bard, Ed.), Vol. XI and XII in particular, Marcel Dekker, New York, 1978. 

One of the few periodic trends of the metals not to show a strong diagonal effect is the standard reduction potential. In fact, this trend follows more of a horizontal rule. The standard reduction potential, E°, is defined in Equation (5.21). The standard reduction potential for the normal hydrogen electrode (N.H.E.), or the half-reaction shown in Equation (5.22), is given a value of zero. Metal atoms with E s more n ative than the N.H.E. are easier to oxidize and harder to reduce. Metal atoms with s more positive than the N.H.E. are easier to reduce and harder to oxidize ... 

The electromotive series is a list of the elements in accordance with their electrode potentials. The measurement of what is commonly known as the "single electrode potential", the "half-reaction potential" or the "half-cell electromotive force" by means of a potentiometer requires a second electrode, a reference electrode, to complete the circuit. If the potential of the reference electrode is taken as zero, the measured E.M.P. will be equal to the potential of the unknown electrode on this scale. W. Ostwald prepared the first table of electrode potentials in 1887 with the dropping mercury electrode as a reference electrode. W. Nernst selected in 1889 the Normal Hydrogen Electrode as a reference electrode. G.N. Lewis and M. Randall published in 1923 their table of single electrode potentials with the Standard Hydrogen Electrode (SHE) as the reference electrode. The Commission of Electrochemistry of the I.U.P.A.C. meeting at Stockholm in 1953 defined the "electrode potential" of a half-cell with the SHE as the reference electrode. 

Eo = the specific standard half-cell potential for the reaction cited at unity activity for dissolved species and 1 atmosphere fugacity for gaseous substances at 25°C. The standard half cell potential for the reaction 2H + 2e" = H2 is defined as 0.0 V versus the normal hydrogen electrode (NHE) at 25°C. 

Each half-cell reaction has a specific standard potential reported as the potential of the reduction reaction vs. the normal hydrogen electrode (NHE). In an elecdochemical cell, there is a half-cell corresponding to the working electrode (WE), where the reactions under study take place, and a reference half-cell. Experimentally the cell potential is measured as the difference between the potentials of the WE half-cell and the reference electrode/ref-erence half-cell (see Chapter 4). The archetypal reference electrode is the NHE, also known as the standard hydrogen electrode (SHE) and is defined, by convention, as 0.000 V for any temperature. Although the NHE is not typically encountered due to difficulty of operation, all conventional electrodes are in turn referenced to this standard to define their absolute potential (i.e., the Ag/AgCl, 3 M KCl reference has a potential of 203 mV vs. the NHE). In practice, experimental results are either stated as being obtained vs. a specific reference electrode, or converted to potentials vs. NHE. 

The energetic requirements for water decomposition into H2 and O2 (eqs 1-5) or CO2 reduction (eqs 6-12) at pH 7 v.s. NHE (the normal hydrogen electrode) depend on the number of electrons in the redox half-reactions as shown below. 

Figure 1.3 Energy level diagrams illustrating reduction potentials of relevant half reactions. Ground state and excited state potentials of porphyrins 1 (red), 2 (green) and 3 (blue) were determined by cyclic voltammetry measurements together with absorption and emission spectra. Approximate potentials for the Ti02 and Sn02 conduction bands and the I3 /I" and BrJ/Br couples are also shown. The O2/H2O couple is 0.82 V at pH = 7. All potentials are reported in V vr. the normal hydrogen electrode (NHE).

The diagrams are normally constructed in relation to the hydrogen electrode with aH+ = 1 as reference, which would be a horizontal straight line. However, they can be modified for the reference to be another pH or another half-reaction. If we choose as reference the reduction of oxygen to water under standard conditions, for example, which has "e =+ 1.23 V, the modification corresponds to a clockwise rotation of the diagram. 

The half-cell electrode that is normally chosen to have a potential of zero is the standard hydrogen electrode (SHE). This half-cell consists of an inert platinum electrode immersed in 1 M HCl with hydrogen gas at 1 atm bubbling through the solution, as shown in Figure 18.6 t. When the SHE acts as the cathode, the following half-reaction occurs ... 

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