Kinetics of Electrode Reactions in Liquid Ammonia Part 2.-Fe1"/FeI1 and C O ~ ~ ' / C O ~ ~ Redox Couples BY OLIVER R. BROWN * AND SEAN A. THORNTON Electrochemistry Research Laboratories, Department of Physical Chemistry, University of Newcastle upon Tyne NEl 7RU Received 19th June, 1973 Polarisation data has been obtained for the CoII/CoIII and FeII/FeIII couples at rotating disc electrodes of platinum, gold and vitreous carbon at - 30°C in ammonia acid solution. In addition, metal wires and carbon rod have been used to examine the low level behaviour of the cobalt system and to record charge measurements in the oxidation of the anode surfaces. The iron couple is reversible on carbon but is inhibited by anodic films on gold and platinum. Apparent standard rate constants for the CoII/CoIII reaction are affected only slightly by the electrode material, but the oxidation branch is noticeably inhibited by anodic film formation. The mechanism of inhibition by anodic films is discussed.FeIIl is unstable in acid ammonia. Electrode kinetic studies of redox reactions have not been made previously in liquid ammonia although polarographic waves for metal ion reductions have been analysed. Similar studies have been carried out in ammoniates at higher tempera- t u r e ~ . ~ ' ~ The reduction of Co"' to ColI in such media was shown to be polaro- graphically irreversible.2* 5 * The chemistry of Fe"' ammines does not appear to have been reported previously. In Part 1 we described the anodic oxidation of a platinum surface in acid am- monia solutions and the effect on the kinetics of hydrogen oxidation.Now we describe the effect of surface oxidation on two redox reactions and extend the study to gold and vitreous carbon anodes. EXPERIMENTAL All measurements were made in gently boiling ammonia solutions at - 30& 1°C. Cells were those used previously.' Working electrodes were of vitreous carbon (Vitreous Carbons Ltd.), gold and platinum (Johnson-Matthey). Rotating disc electrodes (r.d.e.) of these three materials were prepared by machining truncated cones, half angle lo", smaller diameter 4 mm, which were push-fitted into matching teflon holders. Stationary electrodes of glassy carbon were prepared by sealing a rod, diameter 3 mm, into Pyrex whereas those of platinum and gold were made by sealing the metal wires into Pyrex and soda glass respectively.All noble metal electrodes were immersed in warm concentrated HCl prior to use and carbon electrodes were treated with 1 : 1 HN03 : H2S04 cleaning mixture. Gold and platinum r,d.e. were polished with y-alumina but those of carbon were brought to a glassy finish on very fine emery paper lubricated with ethanol because alumina particles were found to embed in the carbon surface. Counter electrodes were platinum sheets, approximately 2.5 cm square. As before,' the reference electrode was Pb/O.l N Pb(NO& and all potentials are quoted relative to it. Bulk conversions of CoIII to CoI1 were carried out at the platinum sheet electrode using as subsidiary electrode a small area of glassy carbon at which, owing to mass transfer control, the prevailing reaction was nitrogen evolution.FelI solutions were obtained in precise 140. R . BROWN AND S. A. THORNTON 15 concentration by the anodic dissolution at controlled current, of pure iron (Johnson- Matthey) in situ. In order that this dissolution should proceed at an acceptable rate without promoting passivation of the iron it was found necessary first to reduce cathodically the iron electrode for one minute by evolving hydrogen from the surface. A large vitreous carbon rod electrode (area 12 cm2) was employed to convert FelI to Fe"I at controlled current density. Tn these bulk oxidations of Fe to FeII and FeII to FelI1 platinum microcathodes (area 4x cm2) were used in order that hydrogen evolution should predominate over- whelmingly over the reverse of the anode process.Except when otherwise stated, the background electrolyte was 1 mol dmF3 NH4N03 prepared as described previously.' However, in the FelI/FelI1 experiments the ammonia was first carefully condensed over sodium and the resulting blue solution redistilled. Hexa- amminocobalt(TI1) trinitrate was prepared in the usual way and was added to the working compartment in preweighed tubes. Potentiostats, function generator and integrator were designed and built in this laboratory. Current against voltage curves were recorded on a Bryans 26,000 X- Y recorder and transients on a TOA EPR-2T chart recorder. 2 0 N I E 3 IC - B .- .= { OO RESULTS THE cO1l'/cO" SYSTEM On rotating disc electrodes of all three materials well-defined waves were obtained for the reduction of CO~~I, using 1 mo1dm-3 NH4N03 as supporting electrolyte.Limiting currents varied linearly with u)* (co = angular rotational velocity of the electrode) although a small intercept current remained at u) = 0, probably attributable to natural convection occurring in the gently boiling solution (fig. 1). Waves were analysed by plotting currents in the form log (id - i ) / i against E (fig. 2). In addition the relation log [(id-i)2/iJ against E was tested for the glassy carbon results; the superior linearity of the former plot suggested that the reaction order was unity rather than two. This order was confirmed by the independence of E+ on the concentration of ColI1. Parameters describing the reduction waves are tabulated 5.0 0.016 ELECTRODE REACTIONS I N LIQUID AMMONIA (table 1).These data were obtained at potential scan rates of 0.01 V s-l on electrodes which, prior to the scan, had been subjected to a hydrogen evolution current of 1 mA cm-2 for some 2 min in order to remove anodic films from the surface. 0.2 0.1 0 -0.1 EIV FIG. 2.-Polarisation data for CoIII reduction at rotating disc electrodes. The points 0, 0 and + refer to a 21 mol m-3 solution at platinum, an 11 mol m-3 solution at gold, and a 13.5 mol m-3 solution at glassy carbon. TABLE 1 .-CoiIJ REDUCTION WAVES iiidc Tafel slope/ E+ mV A cm mol- 1 mV ca tilode (at 0 = 150rads-I)/ (at o = 150rads-l)/ gold + 54 780 92 platinum + 80 804 71 carbon - 75 800 112 After a substantial fraction of the ColI1 in solution had been converted in a bulk reduction to CoII, polarisation data was collected at low current densities on stationary electrodes for both cathodic and anodic branches.In order to minimise hysteresis between scans in the anodic and cathodic directions, low potential sweep rates Mean currents are presented in the form of Butler-Volmer plots (fig. 3) and the reaction parameters are presented in table 2. Apparent standard rate constants are calculated in the usual way according to the relation io = nFkoa&Ea&II. a is the cathodic transfer coefficient, obtained from the Tafel slope (= 2.303 RT/ctnF). Although activity coefficients of ionic species in ammonia deviate even further from unity than in aqueous media,g we have in this work identified activities with con- centrations.V s-l) were used.0. R. BROWN AND S. A . THORNTON 17 The anodic branch of the polarisation curve was measured by cyclic linear potential scans in sections at higher current density values, using for each section scan rates sufficiently small to produce a tolerably small hysteresis. Results for platinum and gold are shown in fig. 4. IOU 1C N I 1 .c Eo % \ 5 h fs 2 0. I 4 v 0.0 0.0 0 I I I -0. I 0 0. I IlV FIG. 3.-Butler-Volmer plots for the CoII/CoIII system at platinum (O), gold (0) and glassy carbon (+) electrodes. The concentrations of CoIII and CoII are as given in table 2. TABLE 2 . - C o I I / C o ~ ~ ~ REACTION AT STATIONARY ELECTRODES Tafel slope ko(apparent)/ i0lAcm-2 at EelrnV cm s-1 [CoIIIl I [CO'II I electrode Ee/mV mol m-3 mol m-3 EolV carbon 362 12.6 5.4 0.345 6 .5 ~ lo-' 65 6 . 8 ~ gold 368 14.4 6.1 0.35 1.03 x 62 1.01 x platinum 344 14.3 6.6 0.33 4 . 9 ~ 10-7 79 4.g5 x 10-7 Oxidation of the platinum electrode surface for two minutes has a corresponding effect upon the E+ value for the subsequent reduction of ColI1. Results in table 3 show how progressively severe surface oxidation diminishes the electrode activity towards the redox process. These reduction waves were recorded at 0.03 V s-l immediately after the surface oxidation treatment.18 ELECTRODE REACTIONS IN LIQUID AMMONIA The effect of ionic strength on the kinetics of ColI1 reduction was investigated by increasing the concentration of ammonium nitrate from 0.01 to 0.1 and then to 1.0 mol dnr3 ; in each case the Co"' concentration was 3 mol m-3.The concentration of NH,NO, in the reference compartment was 1.0 mol dm-3 throughout. On vitreous carbon electrodes the reaction rate appeared to be independent of ionic strength but on gold and platinum electrodes the potential shift values (dE/d log CNH4N03)i of - 0.04 and - 0.05 V respectively were obtained. too N E 0 2 1 0 2 t t o 441 3 I I I L 1.5 0.7 0.9 EIV FIG. 4.-Tafel plots for CoII oxidation at platinum (0) and gold (0) electrodes. The concentrationi of CoII and CoIII are as in table 2. TABLE 3.-EFFECT OF PLATINUM OXIDATION UPON REDUCTION OF CO"' Pretreatment E$at w = 150 rad s-l)/ cathodic potential/V mV Tafel slope/mV 0.6 + 45 75 0.8 - 50 95 1 .o -255 125 THE Fe1I1/Fet1 SYSTEM The addition of ammonia to an aqueous solution of Ferrl precipitates the hydrous oxide.l0 Therefore it was necessary to exclude water rigorously from the liquid ammonia electrolyte when examining the electrochemical behaviour of this system.Iron dissolved quantitatively in ammonium nitrate solution to form Fe" nitrate. This was proved easily by a comparison of the weight loss from the anode and the charge passed during the dissolution. On glassy carbon Fe" gave well-formed oxidation waves only after the anode had been oxidised at +2.0 V for several minutes. Otherwise diffusion limited plateau currents were not obtained. This effect arose presumably from deactivation of large0. R . BROWN AND S . A . THORNTON 19 areas of the carbon surface. The phenomenon was not investigated further. Acti- vated carbon surfaces gave stationary electrode voltammograms (fig.5) characteristic of a simple, essentially reversible one electron transfer system. The rotating disc polarisation curves for a 7.8 mol m-3 solution of Fe" were analysed by plotting log [(id-i)/i] against E (fig. 6). Linear plots were obtained with slopes 51 mV (at o = 150 rad s-l) and 49 mV (w = 37.5 rad s-l) ; the small deviations from the theoretical N I 4 E \ *-l -0.2 - -0.1 - I I I 0.6 0.8 LO EIV FIG. 5.-Cyclic voltammogram at a stationary glassy carbon disc in an 11 mol m-3 FeII solution taken at a scan rate of 0.03 V s-'. EIV FIG. 6.-Polarisation data for Fell oxication at a glassy carbon rotating disc electrode in a 7.5 mol m-3 FeII solution taken at 150 (+) and 37.5 (0) rad s-l.20 6.0 LO 20 0- ELECTRODE RBACTIONS IN LIQUID AMMONIA - - - k 0.7 0.9 I ..I EIV FIG.7.-Cyclic voltammograms taken at 0.03 V s-' at a gold anode rotating at 150 rad s-l int 7.1 mol m-3 FeIl solution. The broken line indicated the return scan when the anodic limit:; restricted to +0.9 V. I i I 0.0 0.9 EIV FIG. 8.-Polarisation data for FeII oxidation at a gold rotating disc electrode in a 7.1 mol m-3 FeII solution. The points (a), (b), (c), and (d) are obtained after holding the electrode at +0.7, +0.95, + 1 .O and + 1.2 V respectively.0. R . BROWN AND S . A . THORNTON 21 value (47 mV at - 30°C) may indicate a small degree of irreversibility but the reaction was considered too rapid to obtain a value for the rate constant on carbon. Assuming equal diffusion and activity coefficients for Fe" and Fe"', the standard potential was equated to the E& (0.833 k0.003 V).A charge of 180 C was passed in an attempt to oxidise an appreciable fraction of the bulk Fe" to Fe"'. Thus a 30 % conversion should have been effected on the 0.55 dm3 solution (initially 1 I mol M - ~ in Fe") but the diffusion limited cathodic current, measured subsequently at a rotating carbon disc electrode, corresponded to a much smaller conversion and this current feu rapidly with time indicating a half life for FeIrL of approximately 3 min. The decay was not simply an oxidation of the solvent to nitrogen with simultaneous regeneration of Fe", because the diffusion limited Fe" oxidation current did in fact decrease as a result of the conversion, although only about half as much as expected.Addition of water to the anhydrous system accelerated the decomposition of Fe"'. FIG. 9.-Polarisation data for FeII oxidation at a platinum rotating disc electrode in a 7.0 mol m-3 FeII solution taken after 3 min at + 0.9 V (0) and + 1.3 V (@). The broken line is an extrapolation of the linear region to the half wave potential. Anodic oxidation of Fe" was examined also on gold and platinum rotating disc electrodes. On these metals it was impossible to scan up to the plateau potentials of the wave without the surface incurring severe oxidation, manifested as a con- siderable hysteresis, the currents on the cathodic (return) scan being much smaller than on the anodic scan (fig. 7). However, if the anodic limit potential for cyclic linear voltammetry was restricted to 0.9 V then the hysteresis on gold was essentially obviated (fig.7) and the polarisation curve, when analysed, showed almost reversible behaviour (E* = 0.848 V, slope 52 mV) illustrated in fig. 8. In the same diagram axe included the polarisation curves obtained with cathodic-going scans on the gold electrode following pre-polarisation for 3 min at various anodic potentials. Reversible22 ELECTRODE REACTIONS I N LIQUID AMMONIA behaviour could be regained by reactivating the gold electrode by a brief excursion below 0 mV. On platinum, inhibition of Felt oxidation by anodisation of the metal surface was more severe than on gold. Thus a platinum anode pretreated by polarisation at 0.9 V gave, on scanning in a cathodic direction, the polarisation curve shown in fig.9, indicating irreversible behaviour. DIRECT MEASUREMENTS OF ELECTRODE SURFACE OXIDATION Previously we described coulometric transient experiments in which smooth platinum surfaces were oxidised by a potential step and subsequently reduced at the end of the pulse. We now report analogous results obtained using gold wire and vitreous carbon disc electrodes (fig. 10). -0.4 V for gold and -0.7 V for carbon. Initial cathodic base potentials were FIG. 10.-Charge measurements on (a) a gold wire and (b) a glassy carbon disc in 1 mol dm-3 ammonium nitrate solution showing anodic (+) and cathodic (0) charges. DISCUSSION The reaction Fe"/Fe"' was found to be almost completely reversible on vitreous carbon and on reduced gold electrodes. This is not surprising, for no major re- arrangement of ammonia ligand molecules is necessary in converting one outer orbital sp3d2 complex ion to another.Fe3+(d5, sp3d2) + e + Fe2+ (d6, sp3d2). The corresponding reaction in aqueous acid solution is well known to be quasi reversible. Any marked deviation from reversibility of the FelI/FelIt couple in ammonia, such as that shown by platinum, must arise from very extensive inhibition of the electron transfer process. A gold surface prepolarised at 0.9 V where the oxidation0. R . BROWN AND S. A . THORNTON 23 charge is 0.4 mC cm-2 (fig. lO(a)) shows reversible behaviour. This oxidation charge corresponds to one or two monolayers of oxidised material. Fe" oxidation is decidedly irreversible when the oxidation charge is 1 mC cm-2 (at 1.2 V) and it exhibits a high Tafel slope associated with an electrode reaction at an inhibited surface.Although a surface oxidation charge of 1 mC cm-2 is reached on vitreous carbon at a much less positive potential (0.7 V), it is unlikely that full surface coverage is approached. The reason is that vitreous carbon is a microporous material l4 presenting a rough surface on an atomic scale; the carbon electrode used in these studies exhibited a low-frequency capacitance value at cathodic potentials of 0.4 mF cm-2 contrasting with 5 0 ~ F c m - ~ of the gold electrode. We conclude that the carbon surface is not extremely oxidised ( < 0.4 mC real cm-2) even at 1.5 V. Hence the reversible behaviour of the FeI1/Fe1I1 couple on vitreous carbon corresponds with the activity expected of an uninhibited surface. A platinum surface acquires an oxidation charge of 1 mC cm-2 before an electrode potential of 0.7 V is reached (ref.(7) fig. 5). As the oxidation of Fe" proceeds only under more anodic conditions, the irreversible behaviour observed with platinum is not unexpected. Even so, the polarisation behaviour at less positive potentials (fig. 9) does not diverge very far from that obtained on carbon. This behaviour contrasts markedly with the severe inhibition observed for hydrogen evolution and oxidation on platinum.' The difference must arise because the redox process is a low overlap reaction whereas the hydrogen electrode reaction proceeds through chemisorbed hydrogen atom intermediates. Previous workers have reported the formation of yellow films and yellow solutions at platinum anodes in ammonia.1s Carbon was not observed to dissolve and gold was heavily oxidised.In the present work we have not attempted to establish the chemical identities of the adsorbed layers on Pt and Au anodes. Herlem et all6 examined the anodic behaviour of silver and mercury in ammonia ; in acid solutions simple ions were formed but in neutral and basic media, amides, imides and nitrides were proposed as oxidation products. Also the salts Hg,NX where X = ClO,, bromide or iodide were found. A recent study of titanium oxidation l7 indicated the formation of an insoluble anodic film composed of Ti(NH2)4 mixed with a small amount of ammonium salt ; these are supposed to arise from the reaction of ammonia with the initially formed titanium halide.It seems probable that the films formed on gold and platinum in the present study are amides, imides or nitrides. Clearly their affect upon the kinetics of redox reactions shows that they are poor electronic con- ductors and the failure of the films to thicken suggests that they are also poor ionic conductors. Inhibition by surface oxide has been reported for the redox system FeIr/FeIrl in aqueous acid so1ution.18 The effect was more severe on platinum than on gold and was ascribed to an increase in the work function of the metal owing to the dipolar nature of the metal-oxygen bond. In terms of the theories of electron transfer reactions at electrodes 19* 2o the presence of anodic films can reduce the electron transfer rate either by (1) affecting Wand W* (respectively the work required to bring the product and reactant ions from the bulk of solution to the reaction site) or (2) reducing K the transmission coefficient which depends upon the magnitude of A the quantum splitting between the upper and lower states arising from the interaction of the degenerate reactant and product systems in the transition state.Hale has discussed 21 the dependence of A upon xp, the distance separating the ion from the metal electrode. He suggested that A decreases by an order of magnitude for each additional 0.7-1.0 A increase in xp. However, the reaction is likely to be completely24 ELECTRODE REACTIONS IN LIQUID AMMONIA adiabatic ( K = 1) until xp increases beyond 15 A. A surface oxidation charge density of 1 mC cm-2, above which platinum manifests inhibition towards the Fe"/Fe"' redox system in ammonia, corresponds to a film thickness somewhat lower than this value.It is possible therefore that effect (1) plays a part in the observed inhibition. This will arise if the point of zero charge (P.z.c.) on the oxidised surface occurs at less anodic potentials than that of the bare metal. Cobalt(II1) reduction takes place at electrode potentials considerably more negative than the Eo for the FeIr/FerIJ system. Even the Eo for the Corl/Corrl reaction is 0.5 V more cathodic than that of iron. Therefore the extent of oxidation of the anode surface, and consequently the effects of inhibition are less under the conditions used to study Co"' reduction.Nevertheless the rate constant of the Corl/Collr couple is very small (< cm s-l) and the Tafel plots are non linear (especially the anodic branches). The low rate of homogeneous exchange between ColI and ColI1 ammines and also the irreversible nature of the CO~~(NH,),/CO~~~(NH,), electrode process in aqueous solutions are well established. The high reorganisation energy for this process is clearly related to the fact that Corl ammine is an outer orbital complex whereas Co"' ammine uses the 3d orbitals for octahedral hybridisation. Thus extensive ligand rearrangement is necessary to achieve the transition state. The cathodic Tafel slope on platinum lies in the range 7 1-8 1 mV at potentials more cathodic than 0.4 V at which point the surface oxidation charge is 0.3 mC cm-2 but the kinetics are retarded at more anodic potentials (fig.3). This result suggests that monolayer coverage by amido radicals (or an equivalent partial coverage by imide or nitride) does not affect the electron transfer rate. Even allowing for the high roughness factor, carbon electrodes commence to oxidise at quite cathodic potentials (fig. 10) and this is manifested by curvature of Tafel plots beyond 0.2 V (fig. 3). On gold, appreciable variation in the transfer coefficient, a, with potential takes place where the surface oxidation is less than 20 pC cm-2. This phenomenon clearly cannot be ascribed to inhibition of the electron transfer process by a surface layer. If, on the other hand, variation in a were to be interpreted in terms of a corresponding variation in p, the symmetry factor for the reaction,22 then similar behaviour should be observed on carbon and platinum electrodes.Short of postulating a mechanism involving ion-electrode overlap, the only explanation for kinetics which differ with change of electrode material lies in terms of the variation of P.Z.C. from one substance to another. The polarisation data recorded for gold (fig. 3) can be explained qualitatively if the P.Z.C. potential for that metal is supposed to be more positive than 0.3 V where the curvature becomes significant. Systematic variation of ionic strength by alteration of the ammonium nitrate supporting electrolyte concentration corro- borates this indication of the position of the P.Z.C. potential of gold. Following this argument, the p.2.c.of platinum should occur at a still more positive potential because curvature in the Tafel plots appears only at potentials positive to 0.4 V and also the shift in the cathodic potential with ionic strength was most pronounced on platinum. Only one previous measurement of the polarisation behaviour of the system CoIr/CorJ1 has been made in liquid ammonia at low temperatures.2 Despite a value of 0.058 V for the slope of the wave analysis plot, which approaches the theoretical value 0.047 V for a reversible one electron transfer process, Laitinen concluded that the polarographic reduction of ColIr nitrate in unbuffered ammonia solution was irreversible because he was unable to observe the reverse process at a dropping mercury electrode. Laitinen determined the E+ as 0.04 V relative to the Pb/O.l N Pb(NO& reference electrode, A comparison of this value with our E+ values obtained at rotating disc electrodes must allow for Our results clearly support that view.0.R . BROWN AND S . A . THORNTON 25 this difference in mass transfer regimes. The mean Nernst diffusion layer thickness in the polarographic experiment can be estimated as approximately cm whereas the value obtained for the r.d.e. from fig. 1 using the Levich equation is 1.08 x cm at cu = 150 rad s-l. The kinematic viscosity has been taken as 3.74 x cm2 s-l at -30°C. Now the variation in Et with 6, the diffusion layer thickness, is given for an irreversible wave by aF(E%- E")/RT = In (nk,S/D). Therefore under the hydrodynamic conditions in our experiments the E4 at mercury would be moved negative by some 0.05 to 0.06 V (accepting Laitinen's value for a) to between - 0.01 and - 0.02 V.This value lies between the half wave potentials obtained at carbon and gold cathodes. Co" oxidation is clearly inhibited by film formation on the metal anodes (fig. 4). Thcre it is seen that the various sections of the polarisation curves do not join and the anodic transfer coefficients are small. The progressive oxidation of the anode surfaces at increasingly positive potentials accounts for these observations. It will be noted that inhibition of CoIr oxidation by gold oxidation is noticeable even at 0.6 V whereas kinetics of the Ferr/FeIrl system on gold were not affected until the potential exceeded 0.9 V. This is no paradox; the iron couple being rapid, requires extensive inhibition before it ceases to be fully reversible.Several workers have studied 23-26 the cobalt ammine couple in aqueous media. This is possible in the presence of excess ammonia because of the large stability constants of these complexes. Laitinen 2 5 found anodic limiting currents much smaller than the expected diffusion controlled values. He explained them in terms of a kinetic wave : slow - e Co" (outer orbital) --+ Co" (inner orbital) + Corr' (inner orbital). Three major objections exist to this proposal. First, the Co(NH,);+ ion cannot exist in the d2sp3 configuration. The reaction order in CoII(NH& was not unity over the whole concentration range. Finally the authors were unable to explain the hysteresis effects whereby currents considerably larger than the steady state limiting value were obtained on the anodic-going scan. Bartelt 26 explained this phenomenon as inhibition of the anode surface by adsorption of hydrolysis products of Co"' which also affected the exchange currents at higher CorIr concentrations.After extrapolation to eliminate these effects the standard rate constant obtained at 25°C (ito = 1.33 x cm s-l) was compared, using Marcus' theory 27 with the homogeneous isotopic exchange rate constants obtained by various workers.28* 29 Reasonable agreement was obtained 2 5 with one of Traube's values.28 If the temperature difference (55°C) between the present work and that of Bartelt alone accounts for the difference in rate constants, then a value for the mean heat of activation can be calculated from the relation Thus the mean activation energy for the cathodic and anodic reactions of the Corr/Colrl couple on platinum is AH: = 86.7 kJ mol-l.Several attempts have been made previously to relate the electrode potential scale in ammonia to that in water. Jolly 30 calculated the displacement between the scales by assuming that the complexing constant for an ion in aqueous solution is a quantita- tive measure of the free energy change for transferring an ion from ammonia to26 ELECTRODE REACTIONS I N LIQUID AMMONIA water. Pleskov assumed that the solvation energy of the rubidium ion is the same in both solvents.31 An dternative to the Rb/Rb+ couple which might be used in a similar manner is the CO~I(NH,),/CO~~~(NH,), redox system.This cannot be carried out accurately with the Eo value determined in the present work because it applies to a different temperature from that used by previous workers in their deternii- nation of the Eo in aqueous solution.32 If the temperature difference is ignored, then AEo = Eo(s.h.e., aq)-E,(s.h.e., NH3) = 0.6V. This value agrees remarkably well with that calculated by Jolly 30 as the free energy of transfer of the proton between the solvents. We express our gratitude to Mr. S . Clarke for his assistance in carrying out several of the experiments. One of us (S. A. T.) thanks S.R.C. for a research studentship. H. A. Laitinen and C. E. Shoemaker, J. Anzer. Chem. SOC., 1950,72,4975. W. B. Schaap, R. F. Conley and F. C.Schmidt, And. Chem., 1961,33,498. W. Hubicki and M. Dabkowska, Anal. Chem., 1961,33,90. T. C. Ichniowski and A. F. Clifford, J. Inorg. Nuclear Chem., 1961,22, 133. G. W. Leonard and D. E. Sellars, J. Electrochem. SOC., 1955, 102,96. J. Bjerrum and J. P. McReynolds, Inorg. Synth., 1946,2,218. J. J. Lagowski and G. A. Moczygemba in Chemistry of Non Aqueous Solvents, ed. J. J. Lagowski (Academic Press, New York, 1967), vol. 11. l o F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry (Interscience, New York, 2nd edn., 1967), p. 860. ref. (10) pp. 856 and 681. ' A. D. McElroy and H. A. Laitinen, J. Phys. Chem., 1953,57,564. ' 0. R. Brown and S. A. Thornton, J. C. S. Furaduy I, 1973, 69,1568. I2 D. H. Angel1 and T. Dickinson, J. Electrounal. Chem., 1972, 35, 55. I3 D. Gilroy and B. E. Conway, J. Phys. Chem., 1965,69,1259. l4 H. W. Davidson, Nuclear Eng., 1962,7, 159. l 5 A. K. Vijh, J. Electrochem. SOC., 1972, 119, 861. l 6 J-J. Minet, M. Herlem and A. Thiibault, J. Electrounal. Chem., 1971, 30,203. l7 A. Comte, M. Renaud, D. Deroo and M. Rigaud, Compt. rend. C, 1971, 273, 1465. J. P. Hoare, Electrochim. Actu, 1972, 17, 1907. R. A. Marcus, Canad. J. Chem., 1959,37,155. 'O V. G. Levich, Adv. Electrochem. Electrochem. Eng., 1966,4, 249. 21 J. M. Hale in Reactions of Molecules at Electrodes, ed. N. S . Hush (Interscience, London, 1971), '' R. A. Marcus, Ann. Rev. Phys. Chem., 1964,15, 155. 23 H. Bartelt, Electrochim. Acta, 1971, 16, 307. 24 L. N. Klatt and W. J. Blaedel, Anal. Chem., 1967,39, 1065. 2 5 H. A. Laitinen and P. Kivalo, J. Amer. Chem. Soc., 1953, 75, 2198. 26 H. Bartelt and S. Landazury, J. Electroarzal. Chem., 1969, 22, 105. 27 R. A. Marcus, J. Chem. Phys., 1965, 43, 679. 2 8 J. F. Endicott and H. Traube, J. Amer. Chem. SOC., 1964,86, 1686. 29 N. S. Birader, D. R. Stranks and M. S. Vaidya, Trans. Faraduy SOC., 1962,58,2421. 30 W. L. Jolly, J. Chem. Ed., 1956, 33, 512. 31 V. A. Pleskov, Fortschr. Chem. U.S.S. R. 1947, 16, 254. 37 J. Bjerrum, Metal Ammine Formation in Aqueous Solution (Haase, Copenhagen, 1941), p. 250 p. 229.