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Anodic behaviour of manganese in alkaline medium
Electrochimica Acta 46 (2001) 2487– 2498 www.elsevier.nl/locate/electacta
Anodic behaviour of manganese in alkaline medium Bouzid Messaoudi 1, Suzanne Joiret 2, Michel Keddam 3, Hisasi Takenouti * Physique des Liquides et Electrochimie, UPR 15 du CNRS, Uni6ersite´ Pierre & Marie Curie, C.P. 133 4 place Jussieu, 75252 Paris, Cedex 05, France Received 28 September 2000; received in revised form 17 January 2001
Abstract Voltammetry, Electrochemical Impedance Spectroscopy (EIS), Rotating Ring-Disk Electrode techniques (RRDE), Electrochemical Quartz Crystal Microbalance (EQCM) measurements, and in-situ Raman Spectroscopy were applied to investigate the anodic behaviour of Mn in 1 M NaOH solution over a wide potential range. Prior to these experiments, for EQCM, an improved plating bath was designed for coating the thin gold electrode of the quartz sensor with Mn. The results obtained revealed clearly that various oxides, depending on the electrode potential, cover the electrode surface. The oxidation–reduction processes between these different oxides and the associated exchange of species with the solution constitute the main characteristic of this electrode. When the Mn electrode is left in 1 M NaOH solution, it becomes spontaneously passive through two consecutive steps. In-situ Raman spectroscopy indicated that the electrode surface is covered by Mn3O4, Mn2O3, and MnO2 as the potential is shifted towards more anodic values. The polarisation curves showed two anodic current peaks, in agreement with the two-step passivation process. EIS spectra exhibited the typical shape of passivation reactions with a large capacitive loop in the low frequency range. The double layer capacitance and the faradaic capacitance determined from EIS data indicate the increase in expanded surface area and bulk volume of the surface oxide with anodic potential. From RRDE measurements, the dissolution of Mn through Mn2 + and Mn3 + species were evaluated. EQCM measurements corroborated the growth of surface oxide species with the potential, and gave valuable information on the nature of the chemical species involved in the oxidation–reduction processes. A reaction mechanism of the Mn electrode in 1 M NaOH in a wide potential range is proposed. © 2001 Elsevier Science Ltd. All rights reserved. Keywords: Electrochemical deposition; Electrochemical impedance spectroscopy; Quartz crystal microbalance; In-situ Raman spectroscopy; Manganese oxides
1. Introduction According to its electrochemical and thermodynamic properties, manganese and its oxidation compounds are * Corresponding author. Tel.: + 33-1-44274158; fax: +331-44274074. E-mail address: [email protected] (H. Takenouti). 1 On leave from Bejaia University, 06000 Bejaia, Algeria; Fax: + 213-3421-4332; e-mail: [email protected] 2 Fax: +33-1-44274074; e-mail: [email protected] 3 Fax: +33-1-44274074; e-mail: [email protected]
good candidates as materials for positive or negative electrodes of primary and secondary batteries. Up to now, application is limited essentially to the positive electrode as manganese dioxide (MnO2) in saline and in alkaline cells of non-rechargeable zinc batteries and as spinels in lithium batteries. To investigate optimal conditions to an eventual use of this metal or its oxides, it is necessary to acquire more information on the electrochemical behaviour in a wide pH domain [1,2]. This study was aimed at gaining new data on the electrochemical behaviour of Mn in alkaline solution by the combination of a number of
0013-4686/01/$ - see front matter © 2001 Elsevier Science Ltd. All rights reserved. PII: S 0 0 1 3 - 4 6 8 6 ( 0 1 ) 0 0 4 4 9 - 2
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electrochemical and spectroscopic techniques: voltammetry, EIS, RRDE, EQCM and, in situ Raman spectroscopy. Among various electrolytes selected, NaOH or KOH solutions in the concentration range from 0.1 to 5 M, the present paper deals only with Mn 1 M NaOH. Electrochemical processes proposed in literature for manganese and manganese oxides are based upon many speculations, leading sometimes to contradictory or ambiguous conclusions. Many of the interpretations call for further experimental work. Effects of experimental conditions onto the electrochemical behaviour of Mn and manganese oxides may constitute a topic of great relevance. Even, the electrochemistry of MnO2, certainly by far the most studied manganese oxide still remains an extremely active field. The use of in situ techniques will be of great help for elucidating reaction mechanisms of manganese oxide in various media. Except for linear or cyclic voltammetries, no experimental technique was applied to metallic manganese for mechanistic purposes. Compared with MnO2 and Mn species in solution, few works are devoted to metallic Mn in the literature. Two references may be quoted for the use of metallic Mn as a negative electrode of a primary battery in alkaline medium [3,4] in the 50’s and 60’s, respectively. Most of the studies of the anodic behaviour of Mn in alkaline media were essentially conducted during these two decades [5–16], and only a few works were conducted outside of this period [17–21]. Many authors pointed out the high complexity of the electrochemical behaviour of manganese in alkaline media. The thermodynamic equilibria of the documented manganese oxides with respect to potential and pH at 25°C are displayed by the Pourbaix’s diagrams in Fig. 1 [22,23].
Fig. 1. Potential pH diagram of manganese in aqueous medium at 25°C from Ref. [23].
It is worth to recall that manganese is an abundant element on earth, it is found as ores, native metallic nodules in deep sea and it plays a major role in the metabolism of living organisms. In addition to MnO2 in battery applications, this element is also largely used in metallurgy as alloying element. It is reported in the literature that in spite of its larger number of oxidation states, (up to 7), the physical and chemical properties of Mn are somehow alike to that of Fe, the neighbour element in the periodic table. In some extent Mn will therefore be compared to Fe since our laboratory has a rather long experience of the latter [24– 26].
2. Experimental Preparation of Mn electrodes, test solution, and experimental set-up will be described below.
2.1. Electrodes Two kinds of manganese electrodes were prepared. One was commercially available electrolytic manganese under plate form, ca. 1 mm in thickness (Prolabo, Rectapur 99.9%), from which disks of 9.83 mm in diameter were cut-out. This electrode is referred to as bulky Mn. The disk thus obtained was pasted with a conducting glue to a brass cylindrical shaft to make a rotating disk electrode (RDE) or a rotating ring-disk electrode (RRDE). The lateral surface of these electrodes was protected by a cataphoretic (cationic electrophoresis) paint to avoid any electrolyte infiltration. For RDE, this electrode was then moulded into an epoxy resin to form a large insulating ring allowing a well-established hydrodynamic flow. The final overall diameter of the assembly was 18 mm. The second type of electrode was made of electrolytic manganese deposited on the thin gold plate electrode of the quartz sensor. This electrode is referred to as EQCM Mn. The plating conditions will be described in detail in the Section 3. RRDE was constituted of a platinum ring, inner and outer diameters were 10 and 11 mm, respectively. Inside this ring, a bulky Mn disk electrode described above was inserted. The gap between these two electrodes was filled-up with silicon sealant allowing an easy replacement of the disk electrode. The collection efficiency of RRDE was calculated as 0.254. The rotation speed of RRDE was 1000 rpm. The reference electrode used was Hg HgO in 1 M KOH or calomel in saturated KCl (SCE). The potential of Hg HgO electrode with respect to SCE is −0.1 V in 1 M NaOH, and all the potentials in this paper will be referred to the Hg HgO reference. The counter-electrode was platinum gauze of large surface area surrounding the cylindrical cell wall.
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Fig. 2. Experimental set-up for EQCM measurements. (a) EQCM electrode with a 6 MHz oscillating quartz crystal. (b) Flow chart of experimental devices.
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ducting glue ensured the connections to both electrodes of the quartz sensor. To avoid any electrolyte leakage to the backside and also to protect the electrical leads, a silicon sealant covered the major non-reactive part of the quartz crystal as illustrated in Fig. 2(a). This electrode was immersed almost at the centre of the electrochemical cell. The experimental set-up for EQCM measurements is shown in Fig. 2(b). The working-electrode output of the potentiostat (Sotelem, model PGStat ZI) was connected to the EQCM electrode. A Frequency Counter (Sefram, model 2721) measured the resonance frequency of the device giving the mass change according to Sauerbrey’s relationship. Other experimental set-ups were similar to that of RRDE experiments. All experiments were monitored by a home-made computer program. In situ Raman spectra were measured with Argon laser (Spectra Physics, model Stabilite 2016) coupled with the CCD based multi-channel spectrometer (Dilor, model Omars 89). The diameter of the laser spot was about 1 mm. Laser power and the integration time applied were 100 mW – 100 s for the spectrum at − 0.65 V and 50 mW –450 s for the spectra at 0.35 and 0.6 V. In these experiments, the solution was opened to the atmosphere. The potential scan was stopped during anodic or cathodic scan at the selected value and held constant during Raman spectra recording.
2.2. Test solution A 1 M NaOH solution was prepared with pro-analisi grade chemical (Merck, Prolabo, or EDS) and doubly ion-exchanged water. Except for in-situ Raman spectroscopy experiments, the solution was bubbled by N2 (Air Liquide, U-quality) before and during the experiments to avoid the formation of carbonate precipitates. The solution for electrochemical plating of Mn will be given below.
3. Results and discussion
2.3. Experimental de6ices
3.1. Electroplating of Mn for EQCM electrode
The EIS measurements were carried out with Solartron equipments (ECI 1286 and FRA 1254) coupled with an analog filter (KEMO, model VBF8) to improve the signal:noise ratio when the impedance is high (typically when the modulus is greater than 10 kV). RRDE experiments were performed with a homemade bi-potentiostat that allows for compensation of the ohmic-drop coupling between the ring and the disk [27]. The disk potential and the disk and the ring currents were digitally recorded by means of a channel scanning voltmeter (Keithley, model 2000 Multimeter). A triangular voltage generator (Amel, model 567) allowed performing potential scan experiments. For EQCM experiments, the Mn-coated quartz blade was flat-fixed on a epoxy-glass circuitry board. A con-
Several compositions of plating bath of Mn are reported in literature [28 –30] but the Mn deposits obtained were found not satisfactory for the present application. This is the reason why new plating conditions were devised. The main objectives were as follows: Similar reactivity to that of bulky Mn electrode. Compatibility with EQCM measurements, i.e. even for a thin film, homogeneity of the deposit, smoothness to avoid a deep polishing, porosity as low as possible. The solution composition was:100 g l − 1 MnSO4·H2O+ 75 g l − 1 (NH4)2SO4 + 60 g l − 1 NH4CNS The solution was buffered spontaneously to pH 3.6.
First, the conditions at which manganese was plated electrochemically onto the thin gold plate electrode of the quartz crystal are described. Then the various experimental results obtained for Mn in 1 M NaOH will be presented and discussed in the perspective of a reaction mechanism.
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Table 1 Mn plating conditions on quartz sensor Current density Solution volume Electrode surface area Temperature
250 mA cm−2 200 ml 0.27 cm2 25°C
Duration Agitation Counter electrode Atmosphere
36–46 s Magnetic bar, 2 Hz Pt grid N2 bubbling
Fig. 3. MEB pictures of Mn electrodeposited on the gold plate electrode of the quartz sensor, nominal frequency 6 MHz. (a) and (b): two different magnifications.
The operating conditions were summarised in Table 1: The EQCM measurements during Mn plating in the above conditions indicated that the current efficiency was ca. 90%, and EDS analysis indicated 99% purity. The electrode reactivity is found similar to that of a bulky Mn. In Fig. 3, the MEB pictures of Mn deposit are shown for two different magnitudes. The surface is smooth enough to perform EQCM experiments after a slight polishing of the electrode surface with alumina powder (4 mm) on a cotton-tip.
state value. This potential change indicates that the pre-passive state is not stable enough and transforms into a more stable one associated to a higher oxidation state of Mn. The kinetics of pre-passivation and the transformation into the final passive state changes from one run to the next one, and depends also on the type (bulky or EQCM) of Mn electrode used. The open-circuit potential at the steady state is − 0.03 V for both electrodes.
3.2. Chronopotentiometry of Mn at open-circuit In Fig. 4, the potential changes at open-circuit conditions of two stationary electrodes, bulky and EQCM, in mildly stirred 1 M NaOH are presented. The potential profiles are similar for both electrodes. When the electrode was immersed in the solution, the potential shifted towards more anodic value, reached a maximum, then decayed in the cathodic direction. This variation may be explained by a spontaneous passivation of the Mn electrode in the highly alkaline medium. This process will be referred to as pre-passivation. However, after exhibiting a minimum ( −0.84 and − 1.18 V, respectively for bulky and EQCM electrodes) the potential increases again to finally attain a steady-
Fig. 4. Chronopotentiogram of stationary Mn electrodes at open-circuit conditions in 1 M NaOH. The solution was mildly stirred and deoxygenated by N2 bubbling.
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Fig. 5. In situ Raman spectra of bulky Mn 1 M NaOH, solution opened to air, stationary electrode
According to Fig. 1, the pre-passive species would correspond to Mn(OH)2 and the passive species to Mn2O3. Surface analysis by Raman spectroscopy gave information on the composition of surface species at various potentials.
3.3. Analysis of surface species by in situ Raman spectroscopy Fig. 5 shows several examples of Raman spectra. These spectra were interpreted on the basis of data reported in Ref. [31]. At potentials smaller than −0.35 V, even very cathodic ones, Mn3O4 (Haussmanite) was observed. Neither MnO nor Mn(OH)2 was found, but this does not indicate that these species were not formed at all, but they may not be sufficiently crys-
Fig. 6. Polarization curves obtained at slow potential scan rates. The arrows indicate the scan direction. Bulky Mn in 1 M NaOH, deoxygenated by N2 bubbling, rotation speed: 1000 rpm.
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tallised to be detectable. Above 0 V, the oxidation of Mn3O4 leading to a phase similar to Mn2O3 (Bixbyite) was observed, and for E \0.5 V, a part of Mn2O3 transforms into MnO2 (Ramsdellite). The steady-state potential obtained for the open-circuit conditions corresponds therefore to the potential range where both Mn3O4 and Mn2O3 are present simultaneously. In this case, Mn is not oxidised up to MnO2. A similar sequence of phase transformations was established for Fe in the same solution [27]. For potentials more anodic than 0.7 V, the oxygen-gas evolution prevents observation of Raman spectra. At 0.8 V the solution becomes violet attesting the formation of MnO− 4 . The main difference is that on an Fe electrode, Fe3O4 is present in the whole potential range, from hydrogen to oxygen evolution. To obtain further information, the phase transformation by potential change was examined by cyclic voltammetry.
3.4. Cyclic 6oltammetry Fig. 6 displays two polarisation curves obtained with slow potential scan rates. During the anodic scan at 1 mV s − 1, two current peaks (anodic peaks 1 and 2 at − 0.5 and 0 V, respectively) can be clearly observed, indicating a two-step passivation process. During the cathodic scan, the reduction of oxidised manganese species can be seen for the potential range up to − 0.6 V as a broad peak of cathodic current then it becomes anodic. This means that a large amount of the surface layer grown during the anodic scan is reduced by the cathodic scan and at the end, the metallic Mn substrate is exposed directly to the solution giving rise to an anodic peak at − 0.64 V. When the scan rate was increased to 5 mV s − 1, a marked anodic current peak was observed during the cathodic potential scan, because the surface film that formed during anodic scan could be easily reduced. During the subsequent scans, the anodic excursion of current during the cathodic scan tends to disappear indicating that the electrode surface would be covered progressively by a remnant anodic layer. In other terms, the reduction of oxide is not complete in the potential range covered in this study. Fig. 7 presents the voltamperograms obtained during the second potential cycling at different scan rates. When the scan rate is increased, the first anodic current peak is observed less, and only the second anodic current peak (at about 0 V) can be clearly seen. Beyond the anodic peak 2, the polarisation curves show a current plateau, the height of which depends definitely on the scan rate. However a careful examination of the polarisation curves indicates clearly an inflexion point near − 0.5 V. The corresponding current density will be assimilated to that of the anodic peak 1.
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Fig. 7. Polarization curves obtained at different scan rates. The arrows indicate the scanning direction. Bulky Mn in 1 M NaOH, deoxygenated by N2 bubbling, rotating speed: 1000 rpm.
For the cathodic scan, a current peak at ca. −0.3 V is observed. This current peak can be associated with the anodic peak 2, therefore it will be referred to as the cathodic peak 2. Another current peak will be defined near − 0.5 V also by an inflexion point. This inflexion point defines the cathodic peak 1. The potentials between the corresponding anodic and cathodic peaks are separated by 0.1 –0.3 V, indicating a rather slow electrode kinetics for these reaction processes. When the scan rate decreases, another current peak can be observed near − 0.9 V. No corresponding anodic peak is observed and this current peak will be referred as cathodic peak 0. Fig. 8 displays the variation of the plateau current density with respect to the scan rate. The plateau current is determined at 0.4 V. A linear relationship is observed. If this current is assimilated to the charging
Fig. 8. Plateau current determined during the anodic sweep at E = 0.4 V with respect to the potential scan rate dE/dt. Bulky Mn 1 M NaOH. Rotation speed: 1000 rpm.
Fig. 9. Effect of potential scan rate to the peak currents. Bulky Mn in 1 M NaOH solution, deoxygenated by N2. Rotation speed: 1000 rpm.
of a capacitance, a value of 9.6 mF cm − 2 is estimated by I= C*(dE/dt). The nearly constant capacitance found on the plateau suggests a physical charge storage (supercapacitor behaviour) rather than a redox response which is expected to produce a peak of the capacitance versus the potential profile. Fig. 9 shows the variation of the peak heights with respect to the scan rate. On this figure it can be seen that all of them can be represented by a linear relationship, with a slope close to 0.5 on the logarithmic scale. A diffusion-controlled process may explain this behaviour. This is why the effect of the rotation speed upon the current peak was examined and the results presented in Fig. 10. The current densities for the anodic and cathodic peak 1 are independent of the rotation speed, whereas a slightly negative dependence is observed for peak 2. The reaction rate is therefore
Fig. 10. Effect of disk rotation speed on the peak currents. Bulky Mn in 1 M NaOH solution, deoxygenated by N2. Potential scan rate: 50 mV s − 1.
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Fig. 11. EIS spectra at different potentials of bulky Mn in 1 M NaOH, deoxygenated by N2 bubbling. Rotation speed: 1000 rpm. (O) Experimental EIS data and ( +) calculated EIS data from the fitted parameters.
not determined by bulk diffusion in solution but may be controlled by solid state diffusion in the surface layer. Higher rotation speeds may erode slightly the surface layer in which oxidation –reduction processes are taking place. As stated above, voltammograms presented in Fig. 7 were obtained during the second potential cycle. The eroding effect decreased, therefore, the total amount of oxide remained at the electrode surface (after the first potential cycle). Further, peak 1 is not dependent on the disk rotation speed. This discrepancy may be interpreted by a double structure of the surface film, the inner film that contributes to peak 1 and the outer layer to peak 2. It is also important to notice that a linear I− V relationship is not a sufficient condition to ensure that the reaction rate is diffusion-limited [32]. EIS measurements may allow verifying if the diffusion process intervenes with the overall reaction rate.
3.5. Electrochemical impedance spectroscopy (EIS) To gain a deeper insight into the electrode kinetics, EIS measurements were performed. In Fig. 11, Nyquist impedance spectra obtained at different dc potentials are presented along with the results of regression calculations. No typical feature of diffusion impedance (Warburg-type impedance) was observed, in contrast to
Fig. 12. Equivalent electrical circuit to represent the EIS data.
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the results displayed in Fig. 9. As commented in the previous section the possibility must be considered that the I− V dependence is not related to diffusion at all. The diagrams shown in Fig. 11 exhibit three different shapes depending on the low frequency limit of the impedance (polarisation resistance Rp): Rp B 0, Rp : , and Rp \ 0. The diagrams exhibiting Rp B 0 are observed below ( −0.9 and − 0.8 V) and around −0.5 V. This behaviour corresponds to a passivation process. The diagrams with Rp : , indicate a blocking electrode response (below − 1.0, −0.7 to − 0.6, and from 0 to 0.2 V). The last type of diagram was observed at remaining potential ranges. The origin of negative polarisation resistance with a large absolute value at very cathodic potentials (E B −0.8 V) can be associated with the passivation process of metallic manganese by the formation of Mn(OH)2 on the basis of the potential –pH diagram (Fig. 1). This passivation process is however not observed on the voltammograms but corresponds well to the pre-passive behaviour observed in Fig. 4. In contrast, −0.5 V corresponds to the passivation process associated with the anodic peak 1 (the oxidation of Mn(OH)2 into Mn3O4). The presence of this negative slope in the current potential curve is the origin of a fast potential jump observed also in Fig. 4 [33]. The second anodic peak does not induce a negative Rp. The impedance spectra of these diagrams were represented by the two hierarchical-R-C circuits shown in Fig. 12 and a Simplex algorithm was used to fit the circuit elements. In this circuit, each element represents, respectively: Rsol: solution resistance between the reference and the working electrode (V cm2), Rt: charge transfer resistance (V cm2), Cd, hd: double layer capacitance, with a depression coefficient hd (0B hd B 1), RF: resistance associated with faradaic processes (V cm2), CF, hF: capacitance associated to charge storage in the surface layer by physical and/or redox processes, with a depression coefficient hF (0B aF B 1). The results of regression calculations are also presented in Fig. 11. Except for the EIS spectra at − 1 V, a fairly good agreement between experimental and calculated data was obtained. For EIS spectra at − 1 V, the addition of a third R-C branch didn’t improve the regression, and it is considered that the system has not reached a complete steady-state. In Fig. 13, the variation of the reciprocal of Rt with respect to potential is presented. This conductance indicates essentially the exchange current density of the oxidation –reduction reactions taking place at the electrode interface. Three maxima located approximately at −0.6, −0.1 and +0.2 V are visible. These potentials correspond roughly to the positions of current peaks observed in
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Fig. 13. The reciprocal of the charge transfer resistance with respect to the electrode potential. Data calculated from EIS spectra displayed in Fig. 11.
closely linked to the expanded surface area of an electronically conducting material, the surface species covering the electrode surface may increase up to the potential of the anodic peak 2. The CF change indicates the ability of the surface film to store the electrical charge, most likely by intercalation. This capacitance is about 1 mF cm − 2 at potentials below the anodic peak 1, and then approaches 6 mF cm − 2 when the potential get close to the anodic peak 2. At 0.4 V, the sum of the two capacitances is slightly smaller than that calculated from Fig. 8 (9.6 mF cm − 2), indicating that the dc current contains some dissolution component that results in an overestimation of the capacitance. To check this hypothesis, the RRDE technique was used to evaluate the formation of soluble species with respect to the electrode potential.
3.6. RRDE experiments Figs. 6 and 7, slight divergence can be explained by the rather poor reversibility of the reactions. The increase of 1/Rt above 0.5 V is due to the contribution of the oxygen evolution reaction. The variation of the capacitances (Cd and CF) with potential are presented in Fig. 14. When the EIS spectra indicates a blocking electrode behaviour (Rp \50 kV cm2), the CF value calculated at 10 mHz was used instead of that determined by regression calculation. Cd increases almost monotonically up to the potential corresponding to the anodic peak 2, then remains essentially constant. The Cd values at ca. 60 mF cm − 2 near − 1 V, are reasonable for the double layer capacitance of a flat electrode. At this potential, the hydrogen evolution reaction begins to take place, filling pores of manganese oxide. Only the top electrode surface is therefore accessible to the electrolyte. The same situation is also taking place for potentials greater than 0.6 V, beyond which the oxygen evolution reaction begins to take place. Between 0 and 0.5 V, the Cd value is close to 2 mF cm − 2. Since the double layer capacitance is
Fig. 14. The change of capacitance with respect to potential. Data calculated from EIS spectra displayed in Fig. 11.
The partial current densities relative to the manganese dissolution through Mn2 + and Mn3 + were determined by RRDE experiments. To collect the former species, the ring potential (ER) was set at 0.5 V. The soluble species was oxidised into Mn3 + . The partial current for Mn3 + dissolution was determined by setting ER at −0.25 V. At this potential, the soluble species is reduced into Mn2 + . If the number of electrons involved in the disk process is nD and that in the ring one is, nR, then one can calculate the partial disk current density JD by: JD = (nD/nR) IR/(N0 SD)
(1)
where IR, N0, and SD denote the ring current, the collection efficiency (here the theoretical value is equal to 0.254) and the disk surface area ( = 0.76 cm2), respectively. If the ring reaction is a reduction, nR is negative. In Fig. 15, the overall current density (J), the partial current densities for Mn2 + (JMn(II)) and Mn3 + (JMn(III)) dissolution are presented. The disk potential scan rate was 50 mV s − 1, and the RRDE electrode was rotated at 1000 rpm. IR showed a bias current. The existence of a bias current on a Pt electrode was reported in the literature [34,35]. To evaluate this bias current, it is postulated that no Mn2 + is released at a potential of 0.5 V, the potential at which ER is set to oxidise this species into Mn3 + . Conversely, no Mn3 + is released at a potential of −0.25 V during the cathodic scan. The partial currents calculated from RRDE data are displayed in Fig. 15. During an anodic scan of the disk, the release of Mn2 + decreases progressively up to the anodic current peak 2, and then remains negligible. During the cathodic scan, even at very anodic potentials, there is an important release of Mn2 + . The emission of this species is maximal at the cathodic peak 2. A further current
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Fig. 15. The partial dissolution current of Mn as Mn2 + and Mn3 + at the disk determined from RRDE measurements. Bulky Mn in 1 M NaOH, deoxygenated by N2 bubbling. Potential scan rate: 50 mV s − 1 (third scan). Rotation speed: 1000 rpm.
increase is observed for a potential lower than −0.6 V, and another partial current maximum was observed at − 0.9 V. A slight current maximum was also observed on the disk at this potential. A high ring-current density observed at very cathodic potentials (E B −1 V) is certainly due to the oxidation of H2 formed at the disk, and not to Mn2 + . As for the release of Mn3 + at the disk, it increases progressively with potential in the whole range examined. A steep ring-current increase above 0.5 V is attributed to the reduction of O2 formed at the disk, and is not related to Mn3 + . A significant Mn3 + dissolution (ca. 0.25 mA cm − 2) is therefore taking place above the anodic current peak 2 for the overall current of 1 mA cm − 2. This explains the discrepancy between the capacitance value determined by the polarisation curve and by the EIS measurements (CF). JMN(III) is negative at a low potential (E B −1 V). This result indicates that there is a parasitic oxidation process taking place at the ring electrode. The origin of this process is unknown but one of the possible explanations is again the oxidation of H2 at the disk. The sign of JMN(III) is opposite to that observed experimentally on the ring by virtue of Eq. (1). Note that JMN(II) and the ring current have the same sign. The results obtained by RRDE are pointing out the following features: Oxides or oxy-hydroxyde such as Mn(OH)2 and Mn2O3 dissolved chemically and release the corresponding cations into the solution, Mn2 + and Mn3 + . The formation of Mn2 + from metallic manganese through the discontinuities of surface oxide layer may take place, because the surface film is thin in this potential range. Reduction of Mn2O3 into Mn3O4 is accompanied by the release of Mn2 + , in contrast the oxidation of Mn3O4 does not produce emission of Mn2 + .
The comparison of voltammetric and RRDE data indicates clearly that there is a formation of multi-layer surface species. This remark support the hypothesis made above to interpret the peak height with respect to the disk rotation speed (§ 3.4). The presence of multilayer surface species is also partly corroborated by EIS data through a high Cd value corresponding to a large expanded surface area. The Cd change suggests also that the amount of surface species increases up to about 0 V, and then remains essentially constant above this potential. To verify this indirect indication, the formation of surface species was followed by EQCM measurements. EQCM measurements may help at identifying the species involved in various oxidation – reduction processes through their gram-equivalent characteristics.
Fig. 16. Mass change determined by EQCM. EQCM Mn in 1 M NaOH, deoxygenated by N2 bubbling. Potential scan rate: 5 mV s − 1 (first to third scans). Stationary electrode.
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where DQ=
&
t2
idt
(3)
t1
where F stands for the Faraday constant (96500 A s mol − 1), and t1 and t2 defined the integration time interval (between two successive data acquisitions). It is important to note that Dm determined experimentally (Fig. 16) depends on two factors; the mass gain by oxide formation (Dmfilm) and the mass loss due to the metal dissolution (Dmdissolution). The amount of metal dissolved (Dm(dissolution)) was evaluated from RRDE experiments (see Fig. 15). Fig. 17. Gram equivalent calculated from data shown in Figs. 15 and 16.
3.7. EQCM measurements In Fig. 16, the mass change during the first three potential scans is shown. The potential scan rate was 5 mV s − 1. At the first scan, the mass remains almost constant up to −0.7 V, that is at the foot of the anodic peak 1, then increases steeply up to the anodic peak 2. Beyond that point the mass continues to increase but at a slower rate. During the cathodic scan, down to the cathodic peak 2, the mass remains essentially constant, then decreases up to −0.5 V. No marked current peak was observed in the voltammogram in contrast to the bulky manganese electrode as presented in Fig. 6, but EQCM data indicates clearly that the electrochemical process is changing at this potential. The mass increases slightly between −0.5 and −0.75 V, then decreases again. During the second anodic sweep, the mass continues to decrease up to the foot of the anodic peak 2. If the mass at 0.6 V (the extreme of the potential scan) is compared for the first and the second scan, a mass loss of 3.5 mg cm − 2 is observed. This difference corresponds approximately to the metal dissolved during one potential cycle. Between the second and the third scan, the electrode mass increases slightly, once the electrode is covered by oxides the anodic dissolution is hindered markedly. The surface oxide layer protects the metal from further dissolution. The determination of species involved in oxidation – reduction processes will be carried out even though indirectly, by evaluating the gram-equivalent (Meq), i.e. the number of grams per Faraday, at different potential ranges. The mass of surface film (Dmfilm) and the charge (DQ) change within a given time interval (here ca. 3 s) were used to calculate Meq according to the following equations. Meq =FDmfilm/(DQ)
(2)
Dm=Dmfilm − Dmdissolution
(4)
Dmdissolution is evaluated from the partial dissolution current density determined by RRDE experiments and Faradays law. That is to say, even if Dm is negative, Dmfilm in Eq. (2) may be positive. Fig. 17 shows the gram-equivalent calculated by means of RRDE and EQCM data obtained during the first anodic scan at a scan rate of 50 mV s − 1. It should be emphasised that in EQCM experiments, Dmfilm does not contain the mass of Mn because this species is only transferred from the substrate (metallic state) to the layer (oxidised state) and remains always part of the electrode. For EB − 0.5 V, Meq is close to −20 g equivalent. That is, there is a mass loss associated to the film formation. This corresponds to positively charged species coming out of the surface during the film formation. The process taking place in this range involves therefore dehydration by expulsion of H3O+ (19 g equivalents). For the potential range corresponding to the anodic peak 2, there is an oxidation process with a gram-equivalent close to 17 g equivalents. The oxidation process is therefore accompanied with an insertion of OH−, as charge compensating species. For E \ 0.2 V, Meq is close to 8 g equivalents indicating that the oxidation process is taking place with the formation of oxide, not hydroxide. In the potential range corresponding to the anodic peak 1, Meq is nearly 0, which is likely in the experimental error for a deprotonation reaction. Results established on the basis of a large set of experimental techniques are summarised below. 1. Metallic manganese immersed in a strongly alkaline medium becomes passive spontaneously following a two-step mechanism. 2. RRDE measurements clearly indicated that Mn dissolves through Mn2 + , the first passivation process taking place at about − 0.9 V. 3. This passivation process would correspond to the formation of Mn(OH)2, according to the Pourbaix diagram.
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4. The dissolution to Mn2 + may follow two reaction paths, direct ionisation of Mn and chemical dissolution of Mn(OH)2. 5. Fig. 17 indicates that the oxidation of this species is taking place by the elimination of H3O+. 6. The second passivation process (Fig. 6) takes place at E: −0.5 V, 7. This second passivation process would be associated with the formation of Mn3O4. 8. The third passivation is taking place at E: 0 V. 9. This reaction is due to the formation of Mn2O3 on the basis of Raman spectra and the Pourbaix diagram. 10. Mn(III) species dissolves chemically. 11. The reduction of Mn2O3 to Mn3O4 gives rise to the formation of soluble Mn2 + species. 12. EQCM measurements indicate that this species is either an hydroxide or an hydrated structure. 13. Above 0.5 V, Mn2O3 transforms into MnO2 (see Fig. 5). This species is however not soluble. 14. Above 0.6 V, together with the O2 evolution reac2− tion, Mn is dissolving through MnO− 4 and MnO4 . From these experimental observations, aggregating data from the Pourbaix diagram, the following reaction scheme is proposed (Scheme 1). The electrochemical reactions that match best the insertion and expulsion of charge and mass determined by RRDE and EQCM experiments are as follows:
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−0.9 V Mn·2H2O+ 2H2O = Mn(OH)2 + 2H3O+ + 2e −
(5)
−0.5 V 3Mn(OH)2 = Mn3O4·2H2O+ 2H+ + 2e − 0 V Mn3O4·2H2O+ OH
(6)
−
= 2MnOOH+ Mn(OH)3 + e − 0.2 V 2(2MnOOH +Mn(OH)3) +3OH = (6MnO2)·5H2O+ 3H+ + 6e −
(7) −
(8)
The passive nature of Mn3O4 is verified by the mass change of the Mn electrode left at the open circuit potential. In Fig. 18, the rate of mass change of an EQCM Mn electrode in 1 M NaOH is presented. The potential change is different from that indicated in Fig. 4: a plateau replaces the first anodic potential peak and the duration of the pre-passive period is longer. This electrode is therefore considered as less active than that presented in Fig. 4. This may be due to an air-grown film during a longer storage of the electrode used in Fig. 18. The film formation is fast during the first 10 min. On the pre-passive plateau, the rate of film formation is almost constant and equals ca. 15 ng cm − 2 s − 1. The electrode potential was −0.6 V corresponding to reactions 5 and 6 above. When the total mass gain reaches about 80 mg cm − 2, the electrode potential jumps to 0 V. The oxide is transformed probably into the mixture of MnOOH and Mn(OH)3.
4. Conclusions
Scheme 1.
Fig. 18. Rate of mass change (dm/dt) and open-circuit potential. EQCM Mn in 1 M NaOH, deoxygenated by N2 bubbling. Magnetic stirring.
Firstly, the optimal conditions to electroplate Mn on quartz crystal microbalance were investigated. The electroplating bath was composed essentially of manganese sulphate. The surface morphology is smooth. The electrochemical behaviour of this Mn is similar to those used for other experiments, i.e. bulky Mn machined into a rotating disk electrode. From in situ Raman spectroscopy, the surface species were characterised. In the low potential range, Mn3O4 covers the surface. When the electrode potential is shifted towards more anodic potentials, this species transforms into Mn2O3 and then MnO2. No Mn(OH)2 was observed, this species would be badly crystallised in the studied medium. The potential borders at which these different species are observed are in agreement with the potential –pH diagram (Pourbaix diagram). In alkaline medium, Mn is covered spontaneously with an oxide layer indicated by a marked pre-passive plateau at ca. − 1 to −0.6 V versus Hg HgO depending on experiments (see Figs. 4 and 18). During this period, Mn(OH)2 and Mn3O4 films are forming progressively. When these oxide covers the electrode surface sufficiently, the mass gain of ca. 80 mg was
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determined by EQCM measurements and the potential jumps to ca. 0 V. The final species becomes Mn2O3 (or hydroxide). The electrode mass no longer increases and the electrode potential becomes constant. Two potential plateaux observed in open-circuit experiments are confirmed by polarisation curves obtained with different potential scan rates. There are actually two current maxima, respectively at −0.5 (anodic peak 1) and 0 V versus Hg HgO (anodic peak 2). The RRDE indicates however, a hidden current peak at a more cathodic potential of −0.9 V versus Hg HgO. The partial disk current associated with Mn2 + release from the disk decreases with potential. The emission of Mn2 + stops when the electrode potential approaches that of the anodic peak 2 for the anodic potential scan. For the cathodic scan, the release of Mn2 + takes place even at very anodic potentials. The reduction of Mn2O3 to Mn3O4 would be accompanied by the formation of soluble Mn2 + . Two or three current peaks, according to the potential scan rate, are observed during the cathodic potential scan. The oxidation process of Mn to Mn(OH)2 will take place with the expulsion of H3O+, whereas the oxidation of Mn(OH)2 to Mn3O4 is accompanied with no significant mass change and thus with the deprotonation reaction. In contrast, the oxidation of the latter into Mn2O3 proceeds with the insertion of OH− ions and that of Mn2O3 into MnO2 with O2 − . The reaction scheme and mechanisms satisfying different experimental observations are proposed. References [1] B. Messaoudi, S. Joiret, M. Keddam, H. Takenouti, 192e`me Meeting of Electrochemical Society, Paris, extended abstracts, vol. 97 –2 (1997) 569. [2] B. Messaoudi, M. Keddam, H. Takenouti, Journe´ es d’Electrochimie, Abstract CA 10 –43, Toulouse (France), June 1999. [3] R.I. Agladze, L.N. Dzhaparidze, Soobshchemiya, Akad. Nauk Gruzin, SSR, 11 (1950) 539; Chem. Abstr. 48 (1954) 467. [4] L.N. Dzhaparidze, D.G. Otiashvili, Trudy Inst. Priklad. Khim. i Elektrokhim., Akad. Nauk Gruzin, SSR, 7 (1960) 73; Chem. Abstr. 56 (1962) 1277. [5] T. Heumann, in: H. Fischer, K. Hauffer, W. Wiederholt (Eds.), ‘Passivierende Film und Deckschichten, Anlaufschichten’, S. 256/88, Berlin-Go¨ ttingen-Heidelberg, 1956, p. 268. [6] S.A. Zaretskii, I.G. Zharnitskii, I.A. Bogdanova, Soviet Electrochem. Proc. 4th conference on Electrochem. Moscow, 1956 (1957); New York, 1961, Bd. III, S. 110/3. [7] R.I. Agladze, G.M. Domanskaya, Elektrokhim. Margantsa, Akad. Nauk. Gruzin, SSR, 1 (1957) 503; Chem. Abstr. 54 (1960) 103. [8] (a) R.I. Agladze, G.M. Domanskaya, Zh. Prikl. Khim. 24 (1951) 787. (b) R.I. Agladze, G.M. Domanskaya, J. Appl. Chem. USSR 24 (1951) 897. .
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[9] R.I. Agladze, G.M. Domanskaya, Soobshchemiya, Akad. Nauk Gruzin, SSR, 18 (1957) 695; Chem. Abstr. 52 (1958) 10759. [10] (a) R.I. Agladze, G.M. Domanskaya, Zh. Prikl. Khim. 24 (1951) 915. (b) R.I. Agladze, G.M. Domanskaya, J. Appl. Chem. USSR 24 (1951) 1041. [11] L. Gerhard, Ann. Univ. Saraviensis Sci., 4 (1955) 35; Chem. Abstr. 52 (1958) 131; Chim. Indust. Paris, 76 (1956) 1081. [12] I.G. Berikashvili, Tr. Inst. Prikl. Khim. i Elektrochim., Akad. Nauk. Gruzin, SSR, 1 (1960) 63; Chem. Abstr. 56 (1962) 1288. [13] I.G. Berikashvili, Soobshchemiya, Akad. Nauk. Gruzin, SSR, 23 (1959) 41; Chem. Abstr. 54 (1960) 7375. [14] V.V. Losev, Diss. Moskau (1955), through Ref. [7] [15] J. Besson, Proc. 11th International Congress on pure Applied Chemistry, London, 1947, Bd. 5, S. 687/93 (1950). [16] J. Besson, L. Gerhard, Ann. Univ. Saraviensis Sci., vol. nos. 4, 25, 28, and 33 (1955); Chem. Abstr. 53 (1959) 2879. [17] G. Grube, H. Metzger, Z. Elektrochem. 29 (1923) 17. [18] G. Grube, Z. Elektrochem. 33 (1927) 389. [19] R. Lorenz, Z. Anorg. Allgem. Chem. 12 (1896) 393. [20] R.I. Agladze, Tr. Inst. Metal i gorn. dela, Akad. Nauk. Gruzin, SSR, vol. nos. 2, 1 and 31 (1949); Chem. Abstr. 50 (1956) 2323. [21] L.N. Dzhaparidze, D.G. Otiashvili, Tr. Inst. Prikl. Khim. i Elektrokhim., Akad. Nauk. Gruzin, SSR, 4 (1963) 9; Chem. Abstr. 61 (1964) 11615. [22] A.M. Moussard, J. Brenet, F. Jolas, M. Pourbaix, J. van Mylder, Atlas d’e´ quilibres Electrochimiques, Gauthier-Villars, Paris, 1963, p. 286. [23] G. Charlot, Les re´ actions chimiques en solution aqueuse et caracte´ risation des ions, Masson, Paris, 1983, p. 227. [24] I. Epelboin, M. Keddam, J. Electrochem. Soc. 117 (1970) 1052. [25] M. Keddam, O.R. Mattos, H. Takenouti, J. Electrochem. Soc. 128 (1981) 257 and 266. [26] N. Benzekri, R. Carranza, M. Keddam, H. Takenouti, Corros. Sci. 31 (1990) 627. [27] M. Keddam, S. Joiret, X.R. No´ voa, M.C. Perez, H. Perrot, H. Takenouti, in: M.B. Ives, J.L. Luo, J. Rodda (Eds.), Proceedings of 7th International Symposium on Passivity of Metals and Semiconductors, The Electrochemical Society, Pennington NJ, 2000, p. 799, PV 99 – 42. [28] E.M. Ungiadze, R.I. Agladze, Elektrokhim., Akad. Nauk. Gruzin, SSR, 27 (1967) 52. [29] K.E. Heusler, J. Electrochem. Soc. 110 (1963) 703. [30] W.E. Bradt, H.H. Oaks, Trans. Electrochem. Soc. 71 (1937) 279. [31] M.C. Bernard, A. Hugot-Le-Goff, B.V. Thi, S. Cordoba de Torresi, J. Electrochem. Soc. 140 (1993) 3065. [32] P. Bernard, M. Keddam, H. Takenouti, J. Electroanal. Chem. 396 (1995) 325. [33] I. Epelboin, C. Gabrielli, M. Keddam, H. Takenouti, in: J.O.M. Bockris, B.E. Conway, E. Yeager, R.E. White (Eds.), Comprehensive Treatise on Electrochemistry, vol. 4, Plenum Press, New York, 1981, p. 151. [34] J. Besson, J. Guitton, Manipulations d’Electrochimie, Masson, Paris, 1972, p. 185. [35] B. Tre´ millon, Electrochimie Analytique et Re´ actions en solution, Tome 2, Masson, Paris, 1993, p. 288.
Anodic behaviour of manganese in alkaline medium Bouzid Messaoudi 1, Suzanne Joiret 2, Michel Keddam 3, Hisasi Takenouti * Physique des Liquides et Electrochimie, UPR 15 du CNRS, Uni6ersite´ Pierre & Marie Curie, C.P. 133 4 place Jussieu, 75252 Paris, Cedex 05, France Received 28 September 2000; received in revised form 17 January 2001
Abstract Voltammetry, Electrochemical Impedance Spectroscopy (EIS), Rotating Ring-Disk Electrode techniques (RRDE), Electrochemical Quartz Crystal Microbalance (EQCM) measurements, and in-situ Raman Spectroscopy were applied to investigate the anodic behaviour of Mn in 1 M NaOH solution over a wide potential range. Prior to these experiments, for EQCM, an improved plating bath was designed for coating the thin gold electrode of the quartz sensor with Mn. The results obtained revealed clearly that various oxides, depending on the electrode potential, cover the electrode surface. The oxidation–reduction processes between these different oxides and the associated exchange of species with the solution constitute the main characteristic of this electrode. When the Mn electrode is left in 1 M NaOH solution, it becomes spontaneously passive through two consecutive steps. In-situ Raman spectroscopy indicated that the electrode surface is covered by Mn3O4, Mn2O3, and MnO2 as the potential is shifted towards more anodic values. The polarisation curves showed two anodic current peaks, in agreement with the two-step passivation process. EIS spectra exhibited the typical shape of passivation reactions with a large capacitive loop in the low frequency range. The double layer capacitance and the faradaic capacitance determined from EIS data indicate the increase in expanded surface area and bulk volume of the surface oxide with anodic potential. From RRDE measurements, the dissolution of Mn through Mn2 + and Mn3 + species were evaluated. EQCM measurements corroborated the growth of surface oxide species with the potential, and gave valuable information on the nature of the chemical species involved in the oxidation–reduction processes. A reaction mechanism of the Mn electrode in 1 M NaOH in a wide potential range is proposed. © 2001 Elsevier Science Ltd. All rights reserved. Keywords: Electrochemical deposition; Electrochemical impedance spectroscopy; Quartz crystal microbalance; In-situ Raman spectroscopy; Manganese oxides
1. Introduction According to its electrochemical and thermodynamic properties, manganese and its oxidation compounds are * Corresponding author. Tel.: + 33-1-44274158; fax: +331-44274074. E-mail address: [email protected] (H. Takenouti). 1 On leave from Bejaia University, 06000 Bejaia, Algeria; Fax: + 213-3421-4332; e-mail: [email protected] 2 Fax: +33-1-44274074; e-mail: [email protected] 3 Fax: +33-1-44274074; e-mail: [email protected]
good candidates as materials for positive or negative electrodes of primary and secondary batteries. Up to now, application is limited essentially to the positive electrode as manganese dioxide (MnO2) in saline and in alkaline cells of non-rechargeable zinc batteries and as spinels in lithium batteries. To investigate optimal conditions to an eventual use of this metal or its oxides, it is necessary to acquire more information on the electrochemical behaviour in a wide pH domain [1,2]. This study was aimed at gaining new data on the electrochemical behaviour of Mn in alkaline solution by the combination of a number of
0013-4686/01/$ - see front matter © 2001 Elsevier Science Ltd. All rights reserved. PII: S 0 0 1 3 - 4 6 8 6 ( 0 1 ) 0 0 4 4 9 - 2
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electrochemical and spectroscopic techniques: voltammetry, EIS, RRDE, EQCM and, in situ Raman spectroscopy. Among various electrolytes selected, NaOH or KOH solutions in the concentration range from 0.1 to 5 M, the present paper deals only with Mn 1 M NaOH. Electrochemical processes proposed in literature for manganese and manganese oxides are based upon many speculations, leading sometimes to contradictory or ambiguous conclusions. Many of the interpretations call for further experimental work. Effects of experimental conditions onto the electrochemical behaviour of Mn and manganese oxides may constitute a topic of great relevance. Even, the electrochemistry of MnO2, certainly by far the most studied manganese oxide still remains an extremely active field. The use of in situ techniques will be of great help for elucidating reaction mechanisms of manganese oxide in various media. Except for linear or cyclic voltammetries, no experimental technique was applied to metallic manganese for mechanistic purposes. Compared with MnO2 and Mn species in solution, few works are devoted to metallic Mn in the literature. Two references may be quoted for the use of metallic Mn as a negative electrode of a primary battery in alkaline medium [3,4] in the 50’s and 60’s, respectively. Most of the studies of the anodic behaviour of Mn in alkaline media were essentially conducted during these two decades [5–16], and only a few works were conducted outside of this period [17–21]. Many authors pointed out the high complexity of the electrochemical behaviour of manganese in alkaline media. The thermodynamic equilibria of the documented manganese oxides with respect to potential and pH at 25°C are displayed by the Pourbaix’s diagrams in Fig. 1 [22,23].
Fig. 1. Potential pH diagram of manganese in aqueous medium at 25°C from Ref. [23].
It is worth to recall that manganese is an abundant element on earth, it is found as ores, native metallic nodules in deep sea and it plays a major role in the metabolism of living organisms. In addition to MnO2 in battery applications, this element is also largely used in metallurgy as alloying element. It is reported in the literature that in spite of its larger number of oxidation states, (up to 7), the physical and chemical properties of Mn are somehow alike to that of Fe, the neighbour element in the periodic table. In some extent Mn will therefore be compared to Fe since our laboratory has a rather long experience of the latter [24– 26].
2. Experimental Preparation of Mn electrodes, test solution, and experimental set-up will be described below.
2.1. Electrodes Two kinds of manganese electrodes were prepared. One was commercially available electrolytic manganese under plate form, ca. 1 mm in thickness (Prolabo, Rectapur 99.9%), from which disks of 9.83 mm in diameter were cut-out. This electrode is referred to as bulky Mn. The disk thus obtained was pasted with a conducting glue to a brass cylindrical shaft to make a rotating disk electrode (RDE) or a rotating ring-disk electrode (RRDE). The lateral surface of these electrodes was protected by a cataphoretic (cationic electrophoresis) paint to avoid any electrolyte infiltration. For RDE, this electrode was then moulded into an epoxy resin to form a large insulating ring allowing a well-established hydrodynamic flow. The final overall diameter of the assembly was 18 mm. The second type of electrode was made of electrolytic manganese deposited on the thin gold plate electrode of the quartz sensor. This electrode is referred to as EQCM Mn. The plating conditions will be described in detail in the Section 3. RRDE was constituted of a platinum ring, inner and outer diameters were 10 and 11 mm, respectively. Inside this ring, a bulky Mn disk electrode described above was inserted. The gap between these two electrodes was filled-up with silicon sealant allowing an easy replacement of the disk electrode. The collection efficiency of RRDE was calculated as 0.254. The rotation speed of RRDE was 1000 rpm. The reference electrode used was Hg HgO in 1 M KOH or calomel in saturated KCl (SCE). The potential of Hg HgO electrode with respect to SCE is −0.1 V in 1 M NaOH, and all the potentials in this paper will be referred to the Hg HgO reference. The counter-electrode was platinum gauze of large surface area surrounding the cylindrical cell wall.
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Fig. 2. Experimental set-up for EQCM measurements. (a) EQCM electrode with a 6 MHz oscillating quartz crystal. (b) Flow chart of experimental devices.
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ducting glue ensured the connections to both electrodes of the quartz sensor. To avoid any electrolyte leakage to the backside and also to protect the electrical leads, a silicon sealant covered the major non-reactive part of the quartz crystal as illustrated in Fig. 2(a). This electrode was immersed almost at the centre of the electrochemical cell. The experimental set-up for EQCM measurements is shown in Fig. 2(b). The working-electrode output of the potentiostat (Sotelem, model PGStat ZI) was connected to the EQCM electrode. A Frequency Counter (Sefram, model 2721) measured the resonance frequency of the device giving the mass change according to Sauerbrey’s relationship. Other experimental set-ups were similar to that of RRDE experiments. All experiments were monitored by a home-made computer program. In situ Raman spectra were measured with Argon laser (Spectra Physics, model Stabilite 2016) coupled with the CCD based multi-channel spectrometer (Dilor, model Omars 89). The diameter of the laser spot was about 1 mm. Laser power and the integration time applied were 100 mW – 100 s for the spectrum at − 0.65 V and 50 mW –450 s for the spectra at 0.35 and 0.6 V. In these experiments, the solution was opened to the atmosphere. The potential scan was stopped during anodic or cathodic scan at the selected value and held constant during Raman spectra recording.
2.2. Test solution A 1 M NaOH solution was prepared with pro-analisi grade chemical (Merck, Prolabo, or EDS) and doubly ion-exchanged water. Except for in-situ Raman spectroscopy experiments, the solution was bubbled by N2 (Air Liquide, U-quality) before and during the experiments to avoid the formation of carbonate precipitates. The solution for electrochemical plating of Mn will be given below.
3. Results and discussion
2.3. Experimental de6ices
3.1. Electroplating of Mn for EQCM electrode
The EIS measurements were carried out with Solartron equipments (ECI 1286 and FRA 1254) coupled with an analog filter (KEMO, model VBF8) to improve the signal:noise ratio when the impedance is high (typically when the modulus is greater than 10 kV). RRDE experiments were performed with a homemade bi-potentiostat that allows for compensation of the ohmic-drop coupling between the ring and the disk [27]. The disk potential and the disk and the ring currents were digitally recorded by means of a channel scanning voltmeter (Keithley, model 2000 Multimeter). A triangular voltage generator (Amel, model 567) allowed performing potential scan experiments. For EQCM experiments, the Mn-coated quartz blade was flat-fixed on a epoxy-glass circuitry board. A con-
Several compositions of plating bath of Mn are reported in literature [28 –30] but the Mn deposits obtained were found not satisfactory for the present application. This is the reason why new plating conditions were devised. The main objectives were as follows: Similar reactivity to that of bulky Mn electrode. Compatibility with EQCM measurements, i.e. even for a thin film, homogeneity of the deposit, smoothness to avoid a deep polishing, porosity as low as possible. The solution composition was:100 g l − 1 MnSO4·H2O+ 75 g l − 1 (NH4)2SO4 + 60 g l − 1 NH4CNS The solution was buffered spontaneously to pH 3.6.
First, the conditions at which manganese was plated electrochemically onto the thin gold plate electrode of the quartz crystal are described. Then the various experimental results obtained for Mn in 1 M NaOH will be presented and discussed in the perspective of a reaction mechanism.
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Table 1 Mn plating conditions on quartz sensor Current density Solution volume Electrode surface area Temperature
250 mA cm−2 200 ml 0.27 cm2 25°C
Duration Agitation Counter electrode Atmosphere
36–46 s Magnetic bar, 2 Hz Pt grid N2 bubbling
Fig. 3. MEB pictures of Mn electrodeposited on the gold plate electrode of the quartz sensor, nominal frequency 6 MHz. (a) and (b): two different magnifications.
The operating conditions were summarised in Table 1: The EQCM measurements during Mn plating in the above conditions indicated that the current efficiency was ca. 90%, and EDS analysis indicated 99% purity. The electrode reactivity is found similar to that of a bulky Mn. In Fig. 3, the MEB pictures of Mn deposit are shown for two different magnitudes. The surface is smooth enough to perform EQCM experiments after a slight polishing of the electrode surface with alumina powder (4 mm) on a cotton-tip.
state value. This potential change indicates that the pre-passive state is not stable enough and transforms into a more stable one associated to a higher oxidation state of Mn. The kinetics of pre-passivation and the transformation into the final passive state changes from one run to the next one, and depends also on the type (bulky or EQCM) of Mn electrode used. The open-circuit potential at the steady state is − 0.03 V for both electrodes.
3.2. Chronopotentiometry of Mn at open-circuit In Fig. 4, the potential changes at open-circuit conditions of two stationary electrodes, bulky and EQCM, in mildly stirred 1 M NaOH are presented. The potential profiles are similar for both electrodes. When the electrode was immersed in the solution, the potential shifted towards more anodic value, reached a maximum, then decayed in the cathodic direction. This variation may be explained by a spontaneous passivation of the Mn electrode in the highly alkaline medium. This process will be referred to as pre-passivation. However, after exhibiting a minimum ( −0.84 and − 1.18 V, respectively for bulky and EQCM electrodes) the potential increases again to finally attain a steady-
Fig. 4. Chronopotentiogram of stationary Mn electrodes at open-circuit conditions in 1 M NaOH. The solution was mildly stirred and deoxygenated by N2 bubbling.
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Fig. 5. In situ Raman spectra of bulky Mn 1 M NaOH, solution opened to air, stationary electrode
According to Fig. 1, the pre-passive species would correspond to Mn(OH)2 and the passive species to Mn2O3. Surface analysis by Raman spectroscopy gave information on the composition of surface species at various potentials.
3.3. Analysis of surface species by in situ Raman spectroscopy Fig. 5 shows several examples of Raman spectra. These spectra were interpreted on the basis of data reported in Ref. [31]. At potentials smaller than −0.35 V, even very cathodic ones, Mn3O4 (Haussmanite) was observed. Neither MnO nor Mn(OH)2 was found, but this does not indicate that these species were not formed at all, but they may not be sufficiently crys-
Fig. 6. Polarization curves obtained at slow potential scan rates. The arrows indicate the scan direction. Bulky Mn in 1 M NaOH, deoxygenated by N2 bubbling, rotation speed: 1000 rpm.
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tallised to be detectable. Above 0 V, the oxidation of Mn3O4 leading to a phase similar to Mn2O3 (Bixbyite) was observed, and for E \0.5 V, a part of Mn2O3 transforms into MnO2 (Ramsdellite). The steady-state potential obtained for the open-circuit conditions corresponds therefore to the potential range where both Mn3O4 and Mn2O3 are present simultaneously. In this case, Mn is not oxidised up to MnO2. A similar sequence of phase transformations was established for Fe in the same solution [27]. For potentials more anodic than 0.7 V, the oxygen-gas evolution prevents observation of Raman spectra. At 0.8 V the solution becomes violet attesting the formation of MnO− 4 . The main difference is that on an Fe electrode, Fe3O4 is present in the whole potential range, from hydrogen to oxygen evolution. To obtain further information, the phase transformation by potential change was examined by cyclic voltammetry.
3.4. Cyclic 6oltammetry Fig. 6 displays two polarisation curves obtained with slow potential scan rates. During the anodic scan at 1 mV s − 1, two current peaks (anodic peaks 1 and 2 at − 0.5 and 0 V, respectively) can be clearly observed, indicating a two-step passivation process. During the cathodic scan, the reduction of oxidised manganese species can be seen for the potential range up to − 0.6 V as a broad peak of cathodic current then it becomes anodic. This means that a large amount of the surface layer grown during the anodic scan is reduced by the cathodic scan and at the end, the metallic Mn substrate is exposed directly to the solution giving rise to an anodic peak at − 0.64 V. When the scan rate was increased to 5 mV s − 1, a marked anodic current peak was observed during the cathodic potential scan, because the surface film that formed during anodic scan could be easily reduced. During the subsequent scans, the anodic excursion of current during the cathodic scan tends to disappear indicating that the electrode surface would be covered progressively by a remnant anodic layer. In other terms, the reduction of oxide is not complete in the potential range covered in this study. Fig. 7 presents the voltamperograms obtained during the second potential cycling at different scan rates. When the scan rate is increased, the first anodic current peak is observed less, and only the second anodic current peak (at about 0 V) can be clearly seen. Beyond the anodic peak 2, the polarisation curves show a current plateau, the height of which depends definitely on the scan rate. However a careful examination of the polarisation curves indicates clearly an inflexion point near − 0.5 V. The corresponding current density will be assimilated to that of the anodic peak 1.
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Fig. 7. Polarization curves obtained at different scan rates. The arrows indicate the scanning direction. Bulky Mn in 1 M NaOH, deoxygenated by N2 bubbling, rotating speed: 1000 rpm.
For the cathodic scan, a current peak at ca. −0.3 V is observed. This current peak can be associated with the anodic peak 2, therefore it will be referred to as the cathodic peak 2. Another current peak will be defined near − 0.5 V also by an inflexion point. This inflexion point defines the cathodic peak 1. The potentials between the corresponding anodic and cathodic peaks are separated by 0.1 –0.3 V, indicating a rather slow electrode kinetics for these reaction processes. When the scan rate decreases, another current peak can be observed near − 0.9 V. No corresponding anodic peak is observed and this current peak will be referred as cathodic peak 0. Fig. 8 displays the variation of the plateau current density with respect to the scan rate. The plateau current is determined at 0.4 V. A linear relationship is observed. If this current is assimilated to the charging
Fig. 8. Plateau current determined during the anodic sweep at E = 0.4 V with respect to the potential scan rate dE/dt. Bulky Mn 1 M NaOH. Rotation speed: 1000 rpm.
Fig. 9. Effect of potential scan rate to the peak currents. Bulky Mn in 1 M NaOH solution, deoxygenated by N2. Rotation speed: 1000 rpm.
of a capacitance, a value of 9.6 mF cm − 2 is estimated by I= C*(dE/dt). The nearly constant capacitance found on the plateau suggests a physical charge storage (supercapacitor behaviour) rather than a redox response which is expected to produce a peak of the capacitance versus the potential profile. Fig. 9 shows the variation of the peak heights with respect to the scan rate. On this figure it can be seen that all of them can be represented by a linear relationship, with a slope close to 0.5 on the logarithmic scale. A diffusion-controlled process may explain this behaviour. This is why the effect of the rotation speed upon the current peak was examined and the results presented in Fig. 10. The current densities for the anodic and cathodic peak 1 are independent of the rotation speed, whereas a slightly negative dependence is observed for peak 2. The reaction rate is therefore
Fig. 10. Effect of disk rotation speed on the peak currents. Bulky Mn in 1 M NaOH solution, deoxygenated by N2. Potential scan rate: 50 mV s − 1.
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Fig. 11. EIS spectra at different potentials of bulky Mn in 1 M NaOH, deoxygenated by N2 bubbling. Rotation speed: 1000 rpm. (O) Experimental EIS data and ( +) calculated EIS data from the fitted parameters.
not determined by bulk diffusion in solution but may be controlled by solid state diffusion in the surface layer. Higher rotation speeds may erode slightly the surface layer in which oxidation –reduction processes are taking place. As stated above, voltammograms presented in Fig. 7 were obtained during the second potential cycle. The eroding effect decreased, therefore, the total amount of oxide remained at the electrode surface (after the first potential cycle). Further, peak 1 is not dependent on the disk rotation speed. This discrepancy may be interpreted by a double structure of the surface film, the inner film that contributes to peak 1 and the outer layer to peak 2. It is also important to notice that a linear I− V relationship is not a sufficient condition to ensure that the reaction rate is diffusion-limited [32]. EIS measurements may allow verifying if the diffusion process intervenes with the overall reaction rate.
3.5. Electrochemical impedance spectroscopy (EIS) To gain a deeper insight into the electrode kinetics, EIS measurements were performed. In Fig. 11, Nyquist impedance spectra obtained at different dc potentials are presented along with the results of regression calculations. No typical feature of diffusion impedance (Warburg-type impedance) was observed, in contrast to
Fig. 12. Equivalent electrical circuit to represent the EIS data.
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the results displayed in Fig. 9. As commented in the previous section the possibility must be considered that the I− V dependence is not related to diffusion at all. The diagrams shown in Fig. 11 exhibit three different shapes depending on the low frequency limit of the impedance (polarisation resistance Rp): Rp B 0, Rp : , and Rp \ 0. The diagrams exhibiting Rp B 0 are observed below ( −0.9 and − 0.8 V) and around −0.5 V. This behaviour corresponds to a passivation process. The diagrams with Rp : , indicate a blocking electrode response (below − 1.0, −0.7 to − 0.6, and from 0 to 0.2 V). The last type of diagram was observed at remaining potential ranges. The origin of negative polarisation resistance with a large absolute value at very cathodic potentials (E B −0.8 V) can be associated with the passivation process of metallic manganese by the formation of Mn(OH)2 on the basis of the potential –pH diagram (Fig. 1). This passivation process is however not observed on the voltammograms but corresponds well to the pre-passive behaviour observed in Fig. 4. In contrast, −0.5 V corresponds to the passivation process associated with the anodic peak 1 (the oxidation of Mn(OH)2 into Mn3O4). The presence of this negative slope in the current potential curve is the origin of a fast potential jump observed also in Fig. 4 [33]. The second anodic peak does not induce a negative Rp. The impedance spectra of these diagrams were represented by the two hierarchical-R-C circuits shown in Fig. 12 and a Simplex algorithm was used to fit the circuit elements. In this circuit, each element represents, respectively: Rsol: solution resistance between the reference and the working electrode (V cm2), Rt: charge transfer resistance (V cm2), Cd, hd: double layer capacitance, with a depression coefficient hd (0B hd B 1), RF: resistance associated with faradaic processes (V cm2), CF, hF: capacitance associated to charge storage in the surface layer by physical and/or redox processes, with a depression coefficient hF (0B aF B 1). The results of regression calculations are also presented in Fig. 11. Except for the EIS spectra at − 1 V, a fairly good agreement between experimental and calculated data was obtained. For EIS spectra at − 1 V, the addition of a third R-C branch didn’t improve the regression, and it is considered that the system has not reached a complete steady-state. In Fig. 13, the variation of the reciprocal of Rt with respect to potential is presented. This conductance indicates essentially the exchange current density of the oxidation –reduction reactions taking place at the electrode interface. Three maxima located approximately at −0.6, −0.1 and +0.2 V are visible. These potentials correspond roughly to the positions of current peaks observed in
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Fig. 13. The reciprocal of the charge transfer resistance with respect to the electrode potential. Data calculated from EIS spectra displayed in Fig. 11.
closely linked to the expanded surface area of an electronically conducting material, the surface species covering the electrode surface may increase up to the potential of the anodic peak 2. The CF change indicates the ability of the surface film to store the electrical charge, most likely by intercalation. This capacitance is about 1 mF cm − 2 at potentials below the anodic peak 1, and then approaches 6 mF cm − 2 when the potential get close to the anodic peak 2. At 0.4 V, the sum of the two capacitances is slightly smaller than that calculated from Fig. 8 (9.6 mF cm − 2), indicating that the dc current contains some dissolution component that results in an overestimation of the capacitance. To check this hypothesis, the RRDE technique was used to evaluate the formation of soluble species with respect to the electrode potential.
3.6. RRDE experiments Figs. 6 and 7, slight divergence can be explained by the rather poor reversibility of the reactions. The increase of 1/Rt above 0.5 V is due to the contribution of the oxygen evolution reaction. The variation of the capacitances (Cd and CF) with potential are presented in Fig. 14. When the EIS spectra indicates a blocking electrode behaviour (Rp \50 kV cm2), the CF value calculated at 10 mHz was used instead of that determined by regression calculation. Cd increases almost monotonically up to the potential corresponding to the anodic peak 2, then remains essentially constant. The Cd values at ca. 60 mF cm − 2 near − 1 V, are reasonable for the double layer capacitance of a flat electrode. At this potential, the hydrogen evolution reaction begins to take place, filling pores of manganese oxide. Only the top electrode surface is therefore accessible to the electrolyte. The same situation is also taking place for potentials greater than 0.6 V, beyond which the oxygen evolution reaction begins to take place. Between 0 and 0.5 V, the Cd value is close to 2 mF cm − 2. Since the double layer capacitance is
Fig. 14. The change of capacitance with respect to potential. Data calculated from EIS spectra displayed in Fig. 11.
The partial current densities relative to the manganese dissolution through Mn2 + and Mn3 + were determined by RRDE experiments. To collect the former species, the ring potential (ER) was set at 0.5 V. The soluble species was oxidised into Mn3 + . The partial current for Mn3 + dissolution was determined by setting ER at −0.25 V. At this potential, the soluble species is reduced into Mn2 + . If the number of electrons involved in the disk process is nD and that in the ring one is, nR, then one can calculate the partial disk current density JD by: JD = (nD/nR) IR/(N0 SD)
(1)
where IR, N0, and SD denote the ring current, the collection efficiency (here the theoretical value is equal to 0.254) and the disk surface area ( = 0.76 cm2), respectively. If the ring reaction is a reduction, nR is negative. In Fig. 15, the overall current density (J), the partial current densities for Mn2 + (JMn(II)) and Mn3 + (JMn(III)) dissolution are presented. The disk potential scan rate was 50 mV s − 1, and the RRDE electrode was rotated at 1000 rpm. IR showed a bias current. The existence of a bias current on a Pt electrode was reported in the literature [34,35]. To evaluate this bias current, it is postulated that no Mn2 + is released at a potential of 0.5 V, the potential at which ER is set to oxidise this species into Mn3 + . Conversely, no Mn3 + is released at a potential of −0.25 V during the cathodic scan. The partial currents calculated from RRDE data are displayed in Fig. 15. During an anodic scan of the disk, the release of Mn2 + decreases progressively up to the anodic current peak 2, and then remains negligible. During the cathodic scan, even at very anodic potentials, there is an important release of Mn2 + . The emission of this species is maximal at the cathodic peak 2. A further current
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Fig. 15. The partial dissolution current of Mn as Mn2 + and Mn3 + at the disk determined from RRDE measurements. Bulky Mn in 1 M NaOH, deoxygenated by N2 bubbling. Potential scan rate: 50 mV s − 1 (third scan). Rotation speed: 1000 rpm.
increase is observed for a potential lower than −0.6 V, and another partial current maximum was observed at − 0.9 V. A slight current maximum was also observed on the disk at this potential. A high ring-current density observed at very cathodic potentials (E B −1 V) is certainly due to the oxidation of H2 formed at the disk, and not to Mn2 + . As for the release of Mn3 + at the disk, it increases progressively with potential in the whole range examined. A steep ring-current increase above 0.5 V is attributed to the reduction of O2 formed at the disk, and is not related to Mn3 + . A significant Mn3 + dissolution (ca. 0.25 mA cm − 2) is therefore taking place above the anodic current peak 2 for the overall current of 1 mA cm − 2. This explains the discrepancy between the capacitance value determined by the polarisation curve and by the EIS measurements (CF). JMN(III) is negative at a low potential (E B −1 V). This result indicates that there is a parasitic oxidation process taking place at the ring electrode. The origin of this process is unknown but one of the possible explanations is again the oxidation of H2 at the disk. The sign of JMN(III) is opposite to that observed experimentally on the ring by virtue of Eq. (1). Note that JMN(II) and the ring current have the same sign. The results obtained by RRDE are pointing out the following features: Oxides or oxy-hydroxyde such as Mn(OH)2 and Mn2O3 dissolved chemically and release the corresponding cations into the solution, Mn2 + and Mn3 + . The formation of Mn2 + from metallic manganese through the discontinuities of surface oxide layer may take place, because the surface film is thin in this potential range. Reduction of Mn2O3 into Mn3O4 is accompanied by the release of Mn2 + , in contrast the oxidation of Mn3O4 does not produce emission of Mn2 + .
The comparison of voltammetric and RRDE data indicates clearly that there is a formation of multi-layer surface species. This remark support the hypothesis made above to interpret the peak height with respect to the disk rotation speed (§ 3.4). The presence of multilayer surface species is also partly corroborated by EIS data through a high Cd value corresponding to a large expanded surface area. The Cd change suggests also that the amount of surface species increases up to about 0 V, and then remains essentially constant above this potential. To verify this indirect indication, the formation of surface species was followed by EQCM measurements. EQCM measurements may help at identifying the species involved in various oxidation – reduction processes through their gram-equivalent characteristics.
Fig. 16. Mass change determined by EQCM. EQCM Mn in 1 M NaOH, deoxygenated by N2 bubbling. Potential scan rate: 5 mV s − 1 (first to third scans). Stationary electrode.
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where DQ=
&
t2
idt
(3)
t1
where F stands for the Faraday constant (96500 A s mol − 1), and t1 and t2 defined the integration time interval (between two successive data acquisitions). It is important to note that Dm determined experimentally (Fig. 16) depends on two factors; the mass gain by oxide formation (Dmfilm) and the mass loss due to the metal dissolution (Dmdissolution). The amount of metal dissolved (Dm(dissolution)) was evaluated from RRDE experiments (see Fig. 15). Fig. 17. Gram equivalent calculated from data shown in Figs. 15 and 16.
3.7. EQCM measurements In Fig. 16, the mass change during the first three potential scans is shown. The potential scan rate was 5 mV s − 1. At the first scan, the mass remains almost constant up to −0.7 V, that is at the foot of the anodic peak 1, then increases steeply up to the anodic peak 2. Beyond that point the mass continues to increase but at a slower rate. During the cathodic scan, down to the cathodic peak 2, the mass remains essentially constant, then decreases up to −0.5 V. No marked current peak was observed in the voltammogram in contrast to the bulky manganese electrode as presented in Fig. 6, but EQCM data indicates clearly that the electrochemical process is changing at this potential. The mass increases slightly between −0.5 and −0.75 V, then decreases again. During the second anodic sweep, the mass continues to decrease up to the foot of the anodic peak 2. If the mass at 0.6 V (the extreme of the potential scan) is compared for the first and the second scan, a mass loss of 3.5 mg cm − 2 is observed. This difference corresponds approximately to the metal dissolved during one potential cycle. Between the second and the third scan, the electrode mass increases slightly, once the electrode is covered by oxides the anodic dissolution is hindered markedly. The surface oxide layer protects the metal from further dissolution. The determination of species involved in oxidation – reduction processes will be carried out even though indirectly, by evaluating the gram-equivalent (Meq), i.e. the number of grams per Faraday, at different potential ranges. The mass of surface film (Dmfilm) and the charge (DQ) change within a given time interval (here ca. 3 s) were used to calculate Meq according to the following equations. Meq =FDmfilm/(DQ)
(2)
Dm=Dmfilm − Dmdissolution
(4)
Dmdissolution is evaluated from the partial dissolution current density determined by RRDE experiments and Faradays law. That is to say, even if Dm is negative, Dmfilm in Eq. (2) may be positive. Fig. 17 shows the gram-equivalent calculated by means of RRDE and EQCM data obtained during the first anodic scan at a scan rate of 50 mV s − 1. It should be emphasised that in EQCM experiments, Dmfilm does not contain the mass of Mn because this species is only transferred from the substrate (metallic state) to the layer (oxidised state) and remains always part of the electrode. For EB − 0.5 V, Meq is close to −20 g equivalent. That is, there is a mass loss associated to the film formation. This corresponds to positively charged species coming out of the surface during the film formation. The process taking place in this range involves therefore dehydration by expulsion of H3O+ (19 g equivalents). For the potential range corresponding to the anodic peak 2, there is an oxidation process with a gram-equivalent close to 17 g equivalents. The oxidation process is therefore accompanied with an insertion of OH−, as charge compensating species. For E \ 0.2 V, Meq is close to 8 g equivalents indicating that the oxidation process is taking place with the formation of oxide, not hydroxide. In the potential range corresponding to the anodic peak 1, Meq is nearly 0, which is likely in the experimental error for a deprotonation reaction. Results established on the basis of a large set of experimental techniques are summarised below. 1. Metallic manganese immersed in a strongly alkaline medium becomes passive spontaneously following a two-step mechanism. 2. RRDE measurements clearly indicated that Mn dissolves through Mn2 + , the first passivation process taking place at about − 0.9 V. 3. This passivation process would correspond to the formation of Mn(OH)2, according to the Pourbaix diagram.
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4. The dissolution to Mn2 + may follow two reaction paths, direct ionisation of Mn and chemical dissolution of Mn(OH)2. 5. Fig. 17 indicates that the oxidation of this species is taking place by the elimination of H3O+. 6. The second passivation process (Fig. 6) takes place at E: −0.5 V, 7. This second passivation process would be associated with the formation of Mn3O4. 8. The third passivation is taking place at E: 0 V. 9. This reaction is due to the formation of Mn2O3 on the basis of Raman spectra and the Pourbaix diagram. 10. Mn(III) species dissolves chemically. 11. The reduction of Mn2O3 to Mn3O4 gives rise to the formation of soluble Mn2 + species. 12. EQCM measurements indicate that this species is either an hydroxide or an hydrated structure. 13. Above 0.5 V, Mn2O3 transforms into MnO2 (see Fig. 5). This species is however not soluble. 14. Above 0.6 V, together with the O2 evolution reac2− tion, Mn is dissolving through MnO− 4 and MnO4 . From these experimental observations, aggregating data from the Pourbaix diagram, the following reaction scheme is proposed (Scheme 1). The electrochemical reactions that match best the insertion and expulsion of charge and mass determined by RRDE and EQCM experiments are as follows:
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−0.9 V Mn·2H2O+ 2H2O = Mn(OH)2 + 2H3O+ + 2e −
(5)
−0.5 V 3Mn(OH)2 = Mn3O4·2H2O+ 2H+ + 2e − 0 V Mn3O4·2H2O+ OH
(6)
−
= 2MnOOH+ Mn(OH)3 + e − 0.2 V 2(2MnOOH +Mn(OH)3) +3OH = (6MnO2)·5H2O+ 3H+ + 6e −
(7) −
(8)
The passive nature of Mn3O4 is verified by the mass change of the Mn electrode left at the open circuit potential. In Fig. 18, the rate of mass change of an EQCM Mn electrode in 1 M NaOH is presented. The potential change is different from that indicated in Fig. 4: a plateau replaces the first anodic potential peak and the duration of the pre-passive period is longer. This electrode is therefore considered as less active than that presented in Fig. 4. This may be due to an air-grown film during a longer storage of the electrode used in Fig. 18. The film formation is fast during the first 10 min. On the pre-passive plateau, the rate of film formation is almost constant and equals ca. 15 ng cm − 2 s − 1. The electrode potential was −0.6 V corresponding to reactions 5 and 6 above. When the total mass gain reaches about 80 mg cm − 2, the electrode potential jumps to 0 V. The oxide is transformed probably into the mixture of MnOOH and Mn(OH)3.
4. Conclusions
Scheme 1.
Fig. 18. Rate of mass change (dm/dt) and open-circuit potential. EQCM Mn in 1 M NaOH, deoxygenated by N2 bubbling. Magnetic stirring.
Firstly, the optimal conditions to electroplate Mn on quartz crystal microbalance were investigated. The electroplating bath was composed essentially of manganese sulphate. The surface morphology is smooth. The electrochemical behaviour of this Mn is similar to those used for other experiments, i.e. bulky Mn machined into a rotating disk electrode. From in situ Raman spectroscopy, the surface species were characterised. In the low potential range, Mn3O4 covers the surface. When the electrode potential is shifted towards more anodic potentials, this species transforms into Mn2O3 and then MnO2. No Mn(OH)2 was observed, this species would be badly crystallised in the studied medium. The potential borders at which these different species are observed are in agreement with the potential –pH diagram (Pourbaix diagram). In alkaline medium, Mn is covered spontaneously with an oxide layer indicated by a marked pre-passive plateau at ca. − 1 to −0.6 V versus Hg HgO depending on experiments (see Figs. 4 and 18). During this period, Mn(OH)2 and Mn3O4 films are forming progressively. When these oxide covers the electrode surface sufficiently, the mass gain of ca. 80 mg was
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determined by EQCM measurements and the potential jumps to ca. 0 V. The final species becomes Mn2O3 (or hydroxide). The electrode mass no longer increases and the electrode potential becomes constant. Two potential plateaux observed in open-circuit experiments are confirmed by polarisation curves obtained with different potential scan rates. There are actually two current maxima, respectively at −0.5 (anodic peak 1) and 0 V versus Hg HgO (anodic peak 2). The RRDE indicates however, a hidden current peak at a more cathodic potential of −0.9 V versus Hg HgO. The partial disk current associated with Mn2 + release from the disk decreases with potential. The emission of Mn2 + stops when the electrode potential approaches that of the anodic peak 2 for the anodic potential scan. For the cathodic scan, the release of Mn2 + takes place even at very anodic potentials. The reduction of Mn2O3 to Mn3O4 would be accompanied by the formation of soluble Mn2 + . Two or three current peaks, according to the potential scan rate, are observed during the cathodic potential scan. The oxidation process of Mn to Mn(OH)2 will take place with the expulsion of H3O+, whereas the oxidation of Mn(OH)2 to Mn3O4 is accompanied with no significant mass change and thus with the deprotonation reaction. In contrast, the oxidation of the latter into Mn2O3 proceeds with the insertion of OH− ions and that of Mn2O3 into MnO2 with O2 − . The reaction scheme and mechanisms satisfying different experimental observations are proposed. References [1] B. Messaoudi, S. Joiret, M. Keddam, H. Takenouti, 192e`me Meeting of Electrochemical Society, Paris, extended abstracts, vol. 97 –2 (1997) 569. [2] B. Messaoudi, M. Keddam, H. Takenouti, Journe´ es d’Electrochimie, Abstract CA 10 –43, Toulouse (France), June 1999. [3] R.I. Agladze, L.N. Dzhaparidze, Soobshchemiya, Akad. Nauk Gruzin, SSR, 11 (1950) 539; Chem. Abstr. 48 (1954) 467. [4] L.N. Dzhaparidze, D.G. Otiashvili, Trudy Inst. Priklad. Khim. i Elektrokhim., Akad. Nauk Gruzin, SSR, 7 (1960) 73; Chem. Abstr. 56 (1962) 1277. [5] T. Heumann, in: H. Fischer, K. Hauffer, W. Wiederholt (Eds.), ‘Passivierende Film und Deckschichten, Anlaufschichten’, S. 256/88, Berlin-Go¨ ttingen-Heidelberg, 1956, p. 268. [6] S.A. Zaretskii, I.G. Zharnitskii, I.A. Bogdanova, Soviet Electrochem. Proc. 4th conference on Electrochem. Moscow, 1956 (1957); New York, 1961, Bd. III, S. 110/3. [7] R.I. Agladze, G.M. Domanskaya, Elektrokhim. Margantsa, Akad. Nauk. Gruzin, SSR, 1 (1957) 503; Chem. Abstr. 54 (1960) 103. [8] (a) R.I. Agladze, G.M. Domanskaya, Zh. Prikl. Khim. 24 (1951) 787. (b) R.I. Agladze, G.M. Domanskaya, J. Appl. Chem. USSR 24 (1951) 897. .
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