1. Biegler ond R. Woods CSlRO Division of Mineral Chemistry P. 0. Box 124 Port Melbourne, Victoria 3207, Australia
The Standard Hydrogen Electrode A misrepresented concept
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The concept of standard electrode potentials, and that of a standard scale on which to refer such potentials, are basic to elementary treatments of galvanic cells. The conventional scale used is the standard hydrogen scale, and on this scale the potential is said to he with respect to a standard hydrogen electrode (S.H.E.). Not only is this scale used for standard potentials, hut also experimental measurements of single electrode potentials in a wide variety of electrochemical fields-e.g., electrode kinetics, corrosion, electrochemical equilibrium (Pourbaix) diagrams, and electrode potentials in biological systemsare frequently reported with respect to the S.H.E. The S.H.E. is often the only reference electrode mentioned in such reports. An experimenter wishing to carry out measurements in one of these fields could be, and in our experience often is, led to believe that a practical S.H.E. exists. This belief would be reinforced by reading standard electrochemical texts which, in various different ways, produce the impression that the S.H.E. has been and can he made. Several texts, in discussing the question of single electrode potentials, will tell him that the potential of an electrode A is the emf obtained by comhining A with a standard hydrogen electrode. In a recent treatise on electrochemistry ( I ) he will see a diagram with the caption "the nonpolarizable standard hydrogen electrode," showing a platinized platinum electrode in an unspecified hydrogen-saturated solution which is labelled " a ~ += 1." Another text on electrode kinetics (2) tells him that "the standard hydrogen electrode consists of a platinized platinum sheet immersed in an acid with hydrogen ion activity of 1 mole/l (pH = O), through which is bubbled hydrogen at 1 atm pressure" and then refers to a diagram. Both these descriptions look very much like recipes for constructing an S.H.E.; not only do they specify the inert metal to he used hut they add a requirement for platinization. Our experimenter, convinced by what he has read that an S.H.E. is the proper reference electrode to use, will wish either to purchase one or to construct one himself. Frustration awaits him. The catalogues from electrode manufacturers will offer him a hydrogen electrode hut not a "standard" one. The "recipes" he has seen, such as those quoted ahove, fail to provide adequate information regarding the electrolyte to he used. And he will he unable to locate the elusive electrode in the index of the standard text on reference electrodes (3). The answer to his dilemma is, of course, that the exper-
'In the recent literature, especially that relating to electrode kinetics in fuel cells, a hydrogen electrode in the supporting eketrolyte has become a popular reference electrode for experiments in these solutions and is often referred to as the R.H.E.(reversible hydrogen electrode). So long as the solution compasitian is properly specified,potentials measured in this way can be repeated by other experimmters, but they cannot easily be referred to the standard hydrogen scale. 604 /Journal of Chemical Education
imenter has been led up the garden path. Although the standard hydrogen scale is real enough, the standard hydrogen electrode is hypothetical. It is because of the misleading way in which the S.H.E. is often presented that we wish to give here a clear exposition of the S.H.E. concept and of the way in which the practical scale is achieved. Let us first recall how single electrode standard potentials are obtained from standard emfs of cells. Such standard emfs can be determined rieorouslv. that a " . orovided . reversibie cell without liquid junction can he set up, e.g., the classical Harned cell (4) consisting of a hydrogen electrode and a silver/silver-chloride electrode in hydrochloric acid electrolvte. In this case the emf is measured a t different hydrochforic acid concentrations and the standard emf obtained by suitable extrapolation. This value is the emf of a hypothetical cell in which all reactants (Ha, H+, C1-, AgCI, Ag) are at unit activity. All thermodynamic standard emf values are obtained by procedures involving extrapolation of experimental results. By considering that a standard emf is the difference between the single potentials of the two electrodes in each cell, a series of relative single electrode potentials can he set up from a set of measurements of standard emfs. In order to give numerical values to this series one half-cell has to be assigned a particular value. By convention the value zero is given to the standard potential of the hydrogen half-cell, thus leading to the standard hydrogen scale. It is important to remember that this manipulation involves only standard emfs and therefore only hypothetical cells. Its validity rests on the fact that in such cells all reactants, including single ions, are a t unit activity. The problem arises when we wish to relate single electrode potentials in real electrochemical cells to the standard hydrogen scale. An obvious way of doing this would be to make any one of the standard half-cells whose potentials have been determined on the standard hydrogen scale by the above procedures. Unfortunately this is not possible, since, as explained ahove, such standard halfcells are hypothetical and we cannot actually prepare a solution that contains all the components in their standard states. In real systems, single ion activities cannot be determined nor can they be adjusted independently. For example, a Hamed cell of the type mentioned ahove could in practice he set u p having an emf equal to its standard value (0.22234 V at 25°C); the hydrochloric acid concentration would be around 1.18 m. Neither half of this cell would constitute a standard half-cell. Although the mean ionic actiuity of the solution in this cell is unitv. the individual ions a r e not in their standard states an-d the halfcells do not have their standard potentials. The practical solution to this problem of necessity involves extra-thermodynamic assumptions. Contrary to the implications of the various texts discussed above, it does not embody an attempt to construct a half-cell with properties approaching the hypothetical S.H.E. By definition, this half-cell consists of a hydrogen electrode (unit fugaci-
ty of Hz) in a solution in which the hydrogen ion activity is unity. On the hydrogen ion activity scale of p H measurement this solution has a p H of zero, and one might therefore imagine that a hydrogen electrode in a solution of pH = 0 constitutes a practical S.H.E. Indeed, this is a commonly quoted "operational" definition of the S.H.E. I t fails in practice because the operational definition of pH does not claim to define uniquely a solution of such low pH. The National Bureau of Standards p H scale (5) has been established by assigning p H values to certain standard solutions, the values being chosen so as to give the maximum possible consistency between precise thermodynamic information, such as dissociation constants of weak acids, and the relationship p H = -log aH+. The scale only has this significance for dilute solutions and for the p H range about 2-12, because it is only under these conditions that the liquid junction potential between the test solution and the electrolyte (usually concentrated KC1) in the reference half-cell required for the measurement remains constant. This condition does not hold for solutions of measured p H as low as zero, and moreover the liquid junction potential with such solutions will vary with the nature of the solution. Thus, if one were to prepare solutions of say HC1, HzSOl or HClOl in which the measured p H was zero, it would be quite incorrect to conclude that in all, or any, of these solutions aH+ = 1. This is simply another way of saying that the potential of a hydrogen electrode in such solutions is undefined on the standard hydrogen scale. It is anoarent from the above discussion that the DH .. scale can provide the proper means for referring potentials to the standard hvdro~en . - scale. Bv definition, the potential of a hydrogen electrode increases by RT/F volt for each ten-fold increase in hydrogen ion activity. The potential of the hypothetical S.H.E. would therefore he (RT/F) X p H above that of the hydrogen electrode in the p H standard solution. Such electrodes have, for accessible electrodes. the most ~reciselydefined potentials on the standard hydrogen scale a n d should be regarded as the basic standards for operational definition of the S.H.E. If a hydrogen electrode in a p H standard solution is a standard on the standard hydrogen scale, why don't we use it as a reference electrode in practice? In fact such an electrode, in combination with a saturated potassium chloride bridge to minimize the junction potential hetween the p H standard solution and the dissimilar solution in the other half-cell, would yield an emf that allows the potential of the latter electrode to he fixed on the standard hydrogen scale. It is experimental convenience which determines that we employ a "secondary standard," rather than the hydrogen electrode itself which requires a source of pure hydrogen and a catalytically active electrode (e.g., Pt) which is subject to poisoning. Calomel electrodes in concentrated potassium chloride are preferred for this purpose, particularly the saturated calomel electrode (S.C.E.) The potential of the S.C.E. on the standard hydrogen scale has been established by exactly the kind of measurement that would he expected from the above discussion, i.e., by setting up cells consisting of the S.C.E. and a hydrogen electrode in p H standard solutions. The accepted value for the potential of the S.C.E. on the standard hydrogen scale is 0.2444 V a t 25°C ( 6 ) .
When precise (better than 1 mV) measurements of potential on the standard hydrogen scale are required, it is advisable to calibrate the particular S.C.E. to be used by the above procedure. This will take into account small variations between individual calomel electrodes and will allow accurate placement of the electrode potential on the standard hydrogen scale, prouided that the solution in the half-cell is dilute and has a p H in the range 2-12. In other conditions the potential of the S.C.E. is uncertain owing to variations in the liquid junction potential. These uncertainties are greatest in concentrated solutions of strong acids and alkalis' and are, of course, of the same nature as the uncertainties in definingpH in such solutions. The majority of reports of single electrode potentials on the standard hydrogen scale seem to he based on measurements with an S.C.E. In many cases details of the method of arriving a t these potentials are not given, and, as we discussed earlier, this could lead the reader to helieve that an S.H.E. was actually used. One final point that should he mentioned is the added confusion brought about by the occasional use, even in the most recent literature, of potentials quoted with respect to a normal hydrogen electrode (N.H.E.). It is not a t all clear whether authors are (a) using this synonymously with the S.H.E., ( b ) employing the zero of the potential scale suggested originally by Nernst and named the N.H.E. at that time, viz. a hydrogen electrode (1atm HZ) in a normal solution of hydrogen ions, or (c) using a hydrogen electrode in a normal solution of an unspecified acid. The question of precisely which solution to use, raised by ( b ) and (c) above, is generally avoided in electrochemistry texts that mention the N.H.E. This is simply because the N.H.E. as originally defined was a hypothetical electrode and no recipe accompanied the definition. There were several attempts to set up electrodes that corresponded to the definition (e.g., the use of 1 M H2S04 has been ascribed to Nernst), hut no universally accepted operational definition developed before the introduction of activities rendered the concept of normal solutions of hydrogen ions, and hence the N.H.E., obsolete. To summarize, we have tried to show that the way in which the S.H.E. is portrayed is likely to mislead both students and researchers, and that there is no real difficulty in precisely defining the zero of the potential scale used in electrochemistry, viz. the standard potential of the hydrogen electrode. Where single electrode potentials are being considered, procedures are available for relating such potentials on the standard hydrogen scale, or in other words to quote them against the S.H.E. The S.H.E. itself is hypothetical and does not enter into these procedures as an operational half-cell. We feel that more attention should he given to stressing this point. The N.H.E. has neither fundamental nor operational significance and should he regarded as of historical interest only. Literature Cited (1) Bockis, J. O'M, and M d y , A. K. N., ,'Modem Elenmchamistnl," Plenum has, New York. 1970. p. 655. (2) Vetter, K.J.,"Electrachemical Kimtiea." Aesdcmie Press, New York, 1967, p. 22. (3) IWB. D. J. 0..and Jan., G. d.. "Reference Electrodes." Academic Presa, New Ymk,
, w., .."..
S.,and Ehlem, R. W..J Arne,. Chprn. Soc., 55,2179 (1933). (5) Bates, R. G.. "Determination of pH." John Wiley & Sons. bc.. New Yolk. 1964. P. 75. (6) Reference (5).p. n8. (4) Harned, H.
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