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Bromide oxidation and bromine reduction in propylene carbonate
Journal of Electroanalytical Chemistry 547 (2003) 109 /113 www.elsevier.com/locate/jelechem
Short Communication
Bromide oxidation and bromine reduction in propylene carbonate V. Vojinovic a, S. Mentus b,*, V. Komnenic a a
b
Faculty of Metallurgy and Technology, Podgorica University, 81000 Podgorica, Yugoslavia Faculty of Physical Chemistry, Belgrade University, Studentski trg 16, 11000 Belgrade, Serbia and Montenegro, Yugoslavia Received 13 November 2002; received in revised form 12 February 2003; accepted 28 February 2003
Abstract Bromide oxidation and bromine reduction in propylene carbonate (PC) /1 M LiClO4 was investigated voltammetrically using polycrystalline platinum rotating disc electrode. The corresponding voltammograms were compared to those of iodide oxidation and HCl decomposition. The oxidation 2Br 0/Br2/2e was found to proceed through two steps, which are caused by the formation of a stable intermediate Br3 ion. The stability constant of tri-bromine ion was estimated to be 105.5 M. The diffusion coefficient of Br ion in this solution was determined to be 3.41 /10 6 cm2 s 1. In the case of bromine reduction, a time dependence of the shape of the voltammogram was observed, which showed that one is dealing with a direct bromination of PC. # 2003 Elsevier Science B.V. All rights reserved. Keywords: Bromide oxidation; Platinum; Propylene carbonate; Rotating disc; Tri-bromide ion
1. Introduction Propylene carbonate (PC) is an aprotic solvent, which became commercially available in the early sixties. This solvent possesses a high relative dielectric permittivity of 64. Its boiling point close to 250 8C and a broad voltage window of electrochemical inertness, which, according to some authors, amounts to more than 8 V [1], make it advantageous in comparison to water as a medium for electrochemical investigations. Being relatively stable in contact with metallic lithium, and having simultaneously a high solubility power for many lithium salts, it was widely investigated as a potential solvent for lithium batteries [1,2]. Because PC has been used exclusively with lithium electrodes, other electrochemical processes have not been investigated. For instance, for a classical hydrogen electrode in this solvent, only a part of a stationary current /potential curve recorded in a mixture of HCl/ SbCl5 by L’Her and Courtot-Coupez [3], and an acido-
* Corresponding author. E-mail address: [email protected] (S. Mentus).
base titration curve by Jaksic et al. [4] are found in the literature. Since a PC/Li-salt solution was shown to be electrochemically stable up to 4 V versus a Li/Li reference electrode [1], one might expect that such a solution would present a suitable medium for all halogen/halogenide electrodes. However, only the iodine/iodide redox pair has been investigated in a PCbased solution [5,6]. The need to investigate other halogen/halogenide redox pairs is supported by the fact that recently, a patent application appeared, relating to a lithium battery, in which a bromine/graphite intercalate as the cathode [7] and a PC based solution as the electrolyte were proposed, but the authors said nothing about the redox properties and chemical stability of the cathode. A knowledge of the anodic processes in aprotic solvents may also have other practical implications: for instance, by the method of the anodic dissolution of metals in the presence of inorganic acids or complexing agents, non-aqueous salts [8], and organometallic complexes [9], respectively, have been synthesized. To contribute to the understanding of the electrochemical behaviour of halogen/halogenide electrodes in
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PC based solutions, in this work, bromide oxidation on a platinum rotating disc in a LiBr solution in PC was examined. The reactivity of anodically evolved bromine with the solvent was demonstrated separately, by recording the voltammograms of bromine reduction.
2. Experimental HPLC grade (99.7%, Aldrich) PC (with an initial maximal water content of 0.02%) was additionally dried ˚ dry zeolite, and distilled under in contact with 4 A reduced pressure. Lithium perchlorate was dried at 120 8C, first in air and then under reduced pressure. The supporting electrolyte solution was 0.5 M LiClO4 in PC. The HCl solution was obtained by blowing dry HCl through the supporting solution. A sample of this solution was diluted with water, and its concentration was determined by titration with standard aqueous NaOH solution. Bromide and iodide ions were added as the LiBr and NaBr salts, which were dried at 120 8C before use. Bromine was dropped directly into the supporting solution, and its concentration was controlled by mass difference before and after addition. Since all the chemicals used are hygroscopic, the solutions were prepared in a dry box. A glass electrochemical cell with double walls was used. The temperature was maintained at 25 8C by circulating the thermostated liquid from an ultrathermostat. The working electrode was a rotating platinum disc in a cylindrical teflon holder. The counter-electrode was a wide platinum foil. An aqueous saturated calomel electrode was used as the reference electrode. To avoid the wetting of the electrolyte bulk, the reference electrode was placed in a tube-like compartment, separated by a glass sinter. During measurements, the cell was maintained under an argon atmosphere. The electrochemical measurements were carried out by an EG&G PAR Model 273 potentiostat/galvanostat. The viscosity of the solutions was measured by an Ostwald capillary viscosimeter.
3. Results and discussion 3.1. Bromide oxidation The conductivity of the PC/LiClO4 system was measured elsewhere [10 /12] and for a 0.5 M LiClO4 solution at 25 8C it was 0.0052 V 1 cm 1 [12]. Contrary to the value in aqueous solutions of similar concentration, this is a rather low conductivity, which may cause a considerable IR drop through the electrolyte. Therefore, the electrolyte resistance between the working and the reference electrodes was determined by an a.c. impe-
Fig. 1. The voltammograms recorded on a rotating Pt disc in the PC/ 0.5 M LiClO4 solution at a scan rate of 20 mV s 1 and at a rotation frequency of 5 rps, after the addition of the following compounds: 1HCl, 2-NaI, 3-LiBr, 4-Br2, freshly prepared solution, 5-Br2, 18 h aged solution. The concentrations are marked in the inserted legend.
dance method. It was estimated that, with the current densities attained in this work, the electrolyte resistance may cause an IR drop of a maximum of 15 mV. Fig. 1 shows the voltammograms recorded in the presence of HCl, LiBr, NaI, and Br2 in the PC/0.5 M LiClO4 solution. From Fig. 1, curve 1, the decomposition voltage of HCl may be estimated as 1.38 V. This figure is very close to 1.358 V, presenting the difference of the standard electrode potentials of hydrogen and chlorine electrodes in aqueous solution. This indicates that the solubility Gibbs energies of HCl in water and in PC are practically equal. Curve 3 relates to the oxidation of bromide ions. It is seen that the oxidation proceeds through two steps, each of which produces a limiting current. By changing the rotation rate, it was concluded that the limiting current of each step is of a diffusion nature, and the height ratio of the higher against the lower current plateau is 3:2, independent of the rotation rate. This is in complete agreement with the oxidation of iodide ions in the same solution, presented by curve 2. This behaviour was also seen previously for iodide oxidation in nitromethane [13], dimethyl formamide [14 /16], acetonitrile [17,18] and dimethyl sulphoxide [19 /21]. The two-step nature of the iodide oxidation was explained as a consequence of the formation of a stable intermediate tri-iodide ion in the solution [5,13 /21]. Therefore, an analogous twostep oxidation scheme may be proposed for the bromide ions studied here, involving the formation of a stable complex tri-bromide ion as an intermediate species: 3Br Br 3 2e Br 3 3=2Br2 e
(1) (2)
Studies of bromide oxidation in aprotic media have
V. Vojinovic et al. / Journal of Electroanalytical Chemistry 547 (2003) 109 /113
been carried out to date only in acetonitrile [22,23], and a similar two-step process was found in this medium. The limiting current density (jlim), concentration (c), rotation frequency (v ) and kinematic viscosity (n ) are correlated to the diffusion coefficient (D ) by the Levich equation [24]: jlim 0:62nFD2=3 n 1=6 v1=2 c
(3)
Introducing the data from Fig. 1, curve 3, in this equation, together with a value of the kinematic viscosity of 0.038 cm2 s 1, which was determined by a capillary viscosimeter, the diffusion coefficient of bromide ion was determined to be 3.41 /10 6 cm2 s 1. The stability constant of tri-iodide ion in PC was found previously to be 107.7 M [25]. The voltammetric curve itself is unsuitable for thermodynamic calculations due to polarisation effects. Nevertheless, the voltammograms recorded under similar circumstances reflect the relative influence of thermodynamic quantities. Actually, the difference between the first and the second halfwave potentials in the voltammograms of both iodide and bromide oxidation in Fig. 1 are proportional to the logarithm of the stability constants of tri-iodide and tribromide complex ions [26]. On the basis of both this proportionality and the previously found value of the stability constant of tri-iodide ion [25], the stability constant of tri-bromide ion in PC was estimated to be 105.5 M.
111
That is, owing to the high stability of tri-bromide ions, this mixture is actually Br2/Br3 . The oxidation of tribromide ions (reaction Eq. (2)) is responsible for the anodic current of voltammograms 4 and 5, while the reduction of bromine excess to tri-bromide (reverse reaction Eq. (2)) is responsible for the first anodic plateau (voltage region 0.3 /0.7 V) of voltammograms 4 and 5. With further cathodic polarisation, this current continues to the current of tri-bromide ion reduction (reverse reaction Eq. (1)). The transition of voltammogram 4 to voltammogram 5, accompanying the solution ageing, indicates the decrease in free bromine excess, i.e. the tendency of the concentrations of Br2 and Br to become equal. The other interesting feature in the system under investigation is the appearance of hydrogen ions. Thus, in voltammograms 4 and 5, one may observe that, in the voltage region that lies more negative than 0.0 V, an additional reduction wave is superposed to the limiting current of Br3 reduction. This new wave appears in the same potential region in which the reduction of proton originating from HCl occurs, as voltammogram 1 shows. The occurrence of the acid was proven additionally by means of methyl orange indicator, in a sample of the solution investigated diluted by water. The features shown by voltammograms 4 and 5 suggest that the following relatively slow chemical reaction between PC and bromine proceeds in the solution.
3.2. Bromine reduction
C4 H6 O3 Br2 C4 H5 BrO3 H Br
In the voltammogram of bromide oxidation (Fig. 1, curve 3) there is no indication that the final anodic product, elemental bromine, reacts chemically with the solvent. However, the voltammograms of elemental bromine reduction offer more information on this subject. Fig. 1, curve 4, presents the voltammogram of a freshly prepared (only aged for /30 min) solution of bromine in the basic PC/0.5 M LiClO4 solution. Contrary to the expectations for the case when only the oxidized species is present in the solution (in which only cathodic processes are to be expected), in the bromine solution under investigation, the voltammetric curve shows diffusion-limited processes in both the positive and negative directions with respect to the open circuit potential. In addition, the system displays a relatively stable open circuit potential (0.94 V for a freshly prepared solution). With increasing time, one may observe an additional shift of the voltammogram toward the region of the anodic currents, i.e. with increasing time, curve 4 shifts to curve 5. In analogy with the voltammograms of the I2/I system in organic solvents [18 /22,26], and based on the similarity to curve 3 in Fig. 1, one may conclude that both voltammograms 4 and 5 indicate the presence of a mixture of Br2/Br , with elemental bromine in excess.
This is accompanied by a fast tri-bromide ion formation involving bromine excess: Br2 Br XBr 3
(4)
(5)
One may exclude the assumption that any impurity existing in PC, rather than the solvent itself, is responsible for the reactions with bromine, because, not only was very pure solvent used, but also, the concentration of water, being the main impurity (/0.02%, or /0.01 M), was additionally reduced by drying with dry zeolite. On the other hand, unlike the reaction with impurities, the voltammetric waves increased proportionally to the bromine concentration when it was increased up to 0.029 M. It is noteworthy that the system under investigation, in both the chemical and electrochemical sense, behaves similarly to the dimethylsulphoxide/bromine system. Martin et al. [27] have shown by the methods of organic chemistry, that dimethylsulphoxide (DMSO) undergoes direct bromination; however, the reaction is not simple and, as a consequence, a variety of products, such as trimethylsulphonium bromide and HBr, appear. To explain the mechanism of bromination, the authors assumed that the following reaction occurred in the initial stage:
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V. Vojinovic et al. / Journal of Electroanalytical Chemistry 547 (2003) 109 /113
(CH3 )2 SO2Br2 [(CH3 )2 SOBr ]Br 3
(6)
The bromination of DMSO may proceed to completion if a compound able to eliminate acid is added to the solution. In addition, we found that the voltammetric curves of the DMSO/Br2 system, immediately after solution preparation, reflect the presence of a high concentration of bromide ions and strong acid ([28], and unpublished results). Although Fig. 1 shows the reactions between PC and Br2, voltammetric curves 4 and 5 do not allow us to conclude more than is expressed by Eq. (4), and the reaction mechanism may be more complex. For instance, either disproportionation of bromine to positively and negatively charged ions, or adducts similar to that defined by Eq. (6) may be assumed to appear in the initial stage of the reaction. The disproportionation of halogens in organic solvents has already been reported in the literature. Iodine disproportionation was proposed elsewhere in order to explain the shape of the voltammetric curves in pyridine [13,29] as well as to explain the appearance of tri-iodide ion spectra in the UV /Vis spectral region in dimethylsulphoxide [30,31]. A pronounced disproportionation was evident in the halogen/solvent system involving mixed halogens, for instance Cl2/I2 [32]. Further to this, the final product of PC bromination may not be that proposed in Eq. (4), and its exact composition remains also a matter for further investigation. The progressive bromine consumption produces no observable change in the distance between the anodic and the final cathodic current plateau in voltammograms 4 and 5, since bromine becomes replaced by an equivalent amount of other electroactive species. This investigation shows that bromide oxidation in PC presents, in fact, an electrochemical reaction of an EC type. With regard to chloride oxidation, shown in the anodic part of voltammogram 1, one may see that, in comparison to other halogenide oxidations, chloride oxidation requires the highest positive potential. Therefore, anodic chlorine evolution is accompanied most probably by a direct anodic substitution of hydrogen in PC molecules. As is obvious in voltammogram 1, the reaction products adsorb on the electrode surface, causing passivation, and therefore diffusion limitation of chloride oxidation appears to be unattainable. Also, there is no evidence for the formation of intermediate complex ions, which is in accordance with the lowest values of the stability constants of polychlorine species in organic solutions, in comparison to other polyhalogens [32]. Among the studies in aprotic solvents, chloride oxidation was studied also in acetonitrile [33,34], nitromethane [35] and dimethyl formamide [36,37], and in each solvent the limiting diffusion current of
chloride oxidation was reported to be attained. However, similarly to the case studied here, almost always, the oxidation was found to occur as a one-step process, i.e. without intermediate complex ion formation.
Acknowledgements One of the authors (SM) is grateful to the Ministry of Science, Technology and Development of the Republic of Serbia, for supporting this work through the contract entitled ‘Structure, thermodynamic and electrochemical properties of modern materials for energy conversion and components in electronics’. The authors are indebted to Prof. H. Lund, University of Aarhus, Denmark, for directing their attention to Ref. [27]. The authors are grateful also to Dr C. Amatore, Ecole Normale Superieure, Paris, France, for helpful discussion.
References [1] A.M. Christie, C.A. Vincent, J. Appl. Electrochem. 26 (1996) 255. [2] R. Koksbang, I.I. Olsen, D. Shackle, Solid State Ionics 69 (1994) 320. [3] M. L’Her, J. Courtot-Coupez, J. Electroanal. Chem. 48 (1973) 265. [4] Lj.N. Jaksic, R.M. Dzudovic, R.P. Mikajlovic, Z.D. Stanic, J. Serb. Chem. Soc. 65 (2000) 587. [5] K.J. Hanson, C.W. Tobias, J. Electrochem. Soc. 134 (1987) 2204. [6] K.J. Hanson, M.J. Matlosz, C.W. Tobias, J. Newman, J. Electrochem. Soc. 134 (1987) 2210. [7] P.K. Sharma, S. Narayanan, G.S. Hickey, U.S. Patent No. 6,042,964. [8] J.J. Habeeb, F.F. Said, D.G. Tuck, J. Chem. Soc. 1 (1981) 118. [9] H. Lehmkuhl, W. Eisenbach, Liebigs. Ann. Chem. 4 (1975) 672. [10] J. Barthel, R. Walcher, H.-J. Gores, in: B.E. Conway, J.O.M. Bockris (Eds.), Modern Aspects of Electrochemistry, vol. 13, Plenum Press, New York, 1979, p. 54. [11] D. Batisti, G.A. Nazri, B. Klassen, R. Aroca, J. Phys. Chem. 97 (1993) 5826. [12] N. Cvjeticanin, S. Mentus, PCCP 1 (1999) 5157. [13] R.T. Iwamoto, Anal. Chem. 31 (1959) 955. [14] H. Breant, C. Sinicki, CR Acad. Sci. (Paris) 260 (1965) 5016. [15] C. Sinicki, Bull. Soc. Chim. (France) (1966) 194. [16] L.N. Ivanovskaya, S.V. Gorbachev, DANN SSSR 118 (1958) 530. [17] J. Desbarres, Bull. Soc. Chim. (France) (1961) 502. [18] V.A. Macagno, M.C. Giordano, A.J. Arvia, Electrochim. Acta 14 (1969) 335. [19] M.C. Giordano, J.C. Bazan, A.J. Arvia, Electrochim. Acta 11 (1966) 741. [20] M.C. Giordano, J.C. Bazan, A.J. Arvia, Electrochim. Acta 11 (1966) 1533. [21] A.J. Arvia, M.C. Giordano, A.J. Podesta, Electrochim. Acta 14 (1969) 389. [22] T. Iwasita, M.C. Giordano, Electrochim. Acta 14 (1969) 1045. [23] A.I. Popov, D.H. Geske, J. Am. Chem. Soc. 80 (1958) 5347. [24] B.G. Levich, Physico-chemical Hydrodynamics (chapter 2), Prentice-Hall, Englewood Cliffs, NJ, 1962. [25] M. L’Her, D. Morin-Bozec, J. Courtot-Coupez, J. Electroanal. Chem. 61 (1975) 99.
V. Vojinovic et al. / Journal of Electroanalytical Chemistry 547 (2003) 109 /113 [26] G. Ciric-Marjanovic, S. Mentus, J. Serb. Chem. Soc. 59 (1994) 639. [27] D. Martin, A. Berger, R. Pescuel, J. Prakt. Chem. 312 (1970) 683. [28] V. Vojinovic, S. Mentus, V. Komnenic, M. Pjesˇcic, XV Jugoslav Symposium on Electrochemistry, Palic, Yugoslavia, Jun 2001, Abstracts, p. 23. [29] R.A. Zingaro, C.A. Van der Werf, J. Kleinberg, J. Am. Chem. Soc. 73 (1951) 88. [30] P. Klaeboe, Acta. Chem. Scand. 18 (1964) 27. [31] M.C. Giordano, J.C. Bazan, A.J. Arvia, J. Inorg. Nucl. Chem. 28 (1966) 1211.
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[32] A. Popov, in: V. Gutman (Ed.), Halogen Chemistry, vol. 1, Elsevier, 1974, pp. 225 /264. [33] J.M. Kolthoff, J.F. Coetzee, J. Am. Chem. Soc. 79 (1957) 1852. [34] L. Sereno, V.A. Macagno, M.C. Giordano, Electrochim. Acta 17 (1972) 561. [35] J.V. Nelson, R.T. Iwamoto, J. Electroanal. Chem. 7 (1961) 218. [36] G. Sinicki, P. Desportes, M. Breant, R. Rosset, Bull. Soc. Chim. 2 (1968) 829 (France). [37] R.L. Benoit, M. Guay, J. Desbarres, Can. J. Chem. 46 (1968) 1261.
Short Communication
Bromide oxidation and bromine reduction in propylene carbonate V. Vojinovic a, S. Mentus b,*, V. Komnenic a a
b
Faculty of Metallurgy and Technology, Podgorica University, 81000 Podgorica, Yugoslavia Faculty of Physical Chemistry, Belgrade University, Studentski trg 16, 11000 Belgrade, Serbia and Montenegro, Yugoslavia Received 13 November 2002; received in revised form 12 February 2003; accepted 28 February 2003
Abstract Bromide oxidation and bromine reduction in propylene carbonate (PC) /1 M LiClO4 was investigated voltammetrically using polycrystalline platinum rotating disc electrode. The corresponding voltammograms were compared to those of iodide oxidation and HCl decomposition. The oxidation 2Br 0/Br2/2e was found to proceed through two steps, which are caused by the formation of a stable intermediate Br3 ion. The stability constant of tri-bromine ion was estimated to be 105.5 M. The diffusion coefficient of Br ion in this solution was determined to be 3.41 /10 6 cm2 s 1. In the case of bromine reduction, a time dependence of the shape of the voltammogram was observed, which showed that one is dealing with a direct bromination of PC. # 2003 Elsevier Science B.V. All rights reserved. Keywords: Bromide oxidation; Platinum; Propylene carbonate; Rotating disc; Tri-bromide ion
1. Introduction Propylene carbonate (PC) is an aprotic solvent, which became commercially available in the early sixties. This solvent possesses a high relative dielectric permittivity of 64. Its boiling point close to 250 8C and a broad voltage window of electrochemical inertness, which, according to some authors, amounts to more than 8 V [1], make it advantageous in comparison to water as a medium for electrochemical investigations. Being relatively stable in contact with metallic lithium, and having simultaneously a high solubility power for many lithium salts, it was widely investigated as a potential solvent for lithium batteries [1,2]. Because PC has been used exclusively with lithium electrodes, other electrochemical processes have not been investigated. For instance, for a classical hydrogen electrode in this solvent, only a part of a stationary current /potential curve recorded in a mixture of HCl/ SbCl5 by L’Her and Courtot-Coupez [3], and an acido-
* Corresponding author. E-mail address: [email protected] (S. Mentus).
base titration curve by Jaksic et al. [4] are found in the literature. Since a PC/Li-salt solution was shown to be electrochemically stable up to 4 V versus a Li/Li reference electrode [1], one might expect that such a solution would present a suitable medium for all halogen/halogenide electrodes. However, only the iodine/iodide redox pair has been investigated in a PCbased solution [5,6]. The need to investigate other halogen/halogenide redox pairs is supported by the fact that recently, a patent application appeared, relating to a lithium battery, in which a bromine/graphite intercalate as the cathode [7] and a PC based solution as the electrolyte were proposed, but the authors said nothing about the redox properties and chemical stability of the cathode. A knowledge of the anodic processes in aprotic solvents may also have other practical implications: for instance, by the method of the anodic dissolution of metals in the presence of inorganic acids or complexing agents, non-aqueous salts [8], and organometallic complexes [9], respectively, have been synthesized. To contribute to the understanding of the electrochemical behaviour of halogen/halogenide electrodes in
0022-0728/03/$ - see front matter # 2003 Elsevier Science B.V. All rights reserved. doi:10.1016/S0022-0728(03)00174-8
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PC based solutions, in this work, bromide oxidation on a platinum rotating disc in a LiBr solution in PC was examined. The reactivity of anodically evolved bromine with the solvent was demonstrated separately, by recording the voltammograms of bromine reduction.
2. Experimental HPLC grade (99.7%, Aldrich) PC (with an initial maximal water content of 0.02%) was additionally dried ˚ dry zeolite, and distilled under in contact with 4 A reduced pressure. Lithium perchlorate was dried at 120 8C, first in air and then under reduced pressure. The supporting electrolyte solution was 0.5 M LiClO4 in PC. The HCl solution was obtained by blowing dry HCl through the supporting solution. A sample of this solution was diluted with water, and its concentration was determined by titration with standard aqueous NaOH solution. Bromide and iodide ions were added as the LiBr and NaBr salts, which were dried at 120 8C before use. Bromine was dropped directly into the supporting solution, and its concentration was controlled by mass difference before and after addition. Since all the chemicals used are hygroscopic, the solutions were prepared in a dry box. A glass electrochemical cell with double walls was used. The temperature was maintained at 25 8C by circulating the thermostated liquid from an ultrathermostat. The working electrode was a rotating platinum disc in a cylindrical teflon holder. The counter-electrode was a wide platinum foil. An aqueous saturated calomel electrode was used as the reference electrode. To avoid the wetting of the electrolyte bulk, the reference electrode was placed in a tube-like compartment, separated by a glass sinter. During measurements, the cell was maintained under an argon atmosphere. The electrochemical measurements were carried out by an EG&G PAR Model 273 potentiostat/galvanostat. The viscosity of the solutions was measured by an Ostwald capillary viscosimeter.
3. Results and discussion 3.1. Bromide oxidation The conductivity of the PC/LiClO4 system was measured elsewhere [10 /12] and for a 0.5 M LiClO4 solution at 25 8C it was 0.0052 V 1 cm 1 [12]. Contrary to the value in aqueous solutions of similar concentration, this is a rather low conductivity, which may cause a considerable IR drop through the electrolyte. Therefore, the electrolyte resistance between the working and the reference electrodes was determined by an a.c. impe-
Fig. 1. The voltammograms recorded on a rotating Pt disc in the PC/ 0.5 M LiClO4 solution at a scan rate of 20 mV s 1 and at a rotation frequency of 5 rps, after the addition of the following compounds: 1HCl, 2-NaI, 3-LiBr, 4-Br2, freshly prepared solution, 5-Br2, 18 h aged solution. The concentrations are marked in the inserted legend.
dance method. It was estimated that, with the current densities attained in this work, the electrolyte resistance may cause an IR drop of a maximum of 15 mV. Fig. 1 shows the voltammograms recorded in the presence of HCl, LiBr, NaI, and Br2 in the PC/0.5 M LiClO4 solution. From Fig. 1, curve 1, the decomposition voltage of HCl may be estimated as 1.38 V. This figure is very close to 1.358 V, presenting the difference of the standard electrode potentials of hydrogen and chlorine electrodes in aqueous solution. This indicates that the solubility Gibbs energies of HCl in water and in PC are practically equal. Curve 3 relates to the oxidation of bromide ions. It is seen that the oxidation proceeds through two steps, each of which produces a limiting current. By changing the rotation rate, it was concluded that the limiting current of each step is of a diffusion nature, and the height ratio of the higher against the lower current plateau is 3:2, independent of the rotation rate. This is in complete agreement with the oxidation of iodide ions in the same solution, presented by curve 2. This behaviour was also seen previously for iodide oxidation in nitromethane [13], dimethyl formamide [14 /16], acetonitrile [17,18] and dimethyl sulphoxide [19 /21]. The two-step nature of the iodide oxidation was explained as a consequence of the formation of a stable intermediate tri-iodide ion in the solution [5,13 /21]. Therefore, an analogous twostep oxidation scheme may be proposed for the bromide ions studied here, involving the formation of a stable complex tri-bromide ion as an intermediate species: 3Br Br 3 2e Br 3 3=2Br2 e
(1) (2)
Studies of bromide oxidation in aprotic media have
V. Vojinovic et al. / Journal of Electroanalytical Chemistry 547 (2003) 109 /113
been carried out to date only in acetonitrile [22,23], and a similar two-step process was found in this medium. The limiting current density (jlim), concentration (c), rotation frequency (v ) and kinematic viscosity (n ) are correlated to the diffusion coefficient (D ) by the Levich equation [24]: jlim 0:62nFD2=3 n 1=6 v1=2 c
(3)
Introducing the data from Fig. 1, curve 3, in this equation, together with a value of the kinematic viscosity of 0.038 cm2 s 1, which was determined by a capillary viscosimeter, the diffusion coefficient of bromide ion was determined to be 3.41 /10 6 cm2 s 1. The stability constant of tri-iodide ion in PC was found previously to be 107.7 M [25]. The voltammetric curve itself is unsuitable for thermodynamic calculations due to polarisation effects. Nevertheless, the voltammograms recorded under similar circumstances reflect the relative influence of thermodynamic quantities. Actually, the difference between the first and the second halfwave potentials in the voltammograms of both iodide and bromide oxidation in Fig. 1 are proportional to the logarithm of the stability constants of tri-iodide and tribromide complex ions [26]. On the basis of both this proportionality and the previously found value of the stability constant of tri-iodide ion [25], the stability constant of tri-bromide ion in PC was estimated to be 105.5 M.
111
That is, owing to the high stability of tri-bromide ions, this mixture is actually Br2/Br3 . The oxidation of tribromide ions (reaction Eq. (2)) is responsible for the anodic current of voltammograms 4 and 5, while the reduction of bromine excess to tri-bromide (reverse reaction Eq. (2)) is responsible for the first anodic plateau (voltage region 0.3 /0.7 V) of voltammograms 4 and 5. With further cathodic polarisation, this current continues to the current of tri-bromide ion reduction (reverse reaction Eq. (1)). The transition of voltammogram 4 to voltammogram 5, accompanying the solution ageing, indicates the decrease in free bromine excess, i.e. the tendency of the concentrations of Br2 and Br to become equal. The other interesting feature in the system under investigation is the appearance of hydrogen ions. Thus, in voltammograms 4 and 5, one may observe that, in the voltage region that lies more negative than 0.0 V, an additional reduction wave is superposed to the limiting current of Br3 reduction. This new wave appears in the same potential region in which the reduction of proton originating from HCl occurs, as voltammogram 1 shows. The occurrence of the acid was proven additionally by means of methyl orange indicator, in a sample of the solution investigated diluted by water. The features shown by voltammograms 4 and 5 suggest that the following relatively slow chemical reaction between PC and bromine proceeds in the solution.
3.2. Bromine reduction
C4 H6 O3 Br2 C4 H5 BrO3 H Br
In the voltammogram of bromide oxidation (Fig. 1, curve 3) there is no indication that the final anodic product, elemental bromine, reacts chemically with the solvent. However, the voltammograms of elemental bromine reduction offer more information on this subject. Fig. 1, curve 4, presents the voltammogram of a freshly prepared (only aged for /30 min) solution of bromine in the basic PC/0.5 M LiClO4 solution. Contrary to the expectations for the case when only the oxidized species is present in the solution (in which only cathodic processes are to be expected), in the bromine solution under investigation, the voltammetric curve shows diffusion-limited processes in both the positive and negative directions with respect to the open circuit potential. In addition, the system displays a relatively stable open circuit potential (0.94 V for a freshly prepared solution). With increasing time, one may observe an additional shift of the voltammogram toward the region of the anodic currents, i.e. with increasing time, curve 4 shifts to curve 5. In analogy with the voltammograms of the I2/I system in organic solvents [18 /22,26], and based on the similarity to curve 3 in Fig. 1, one may conclude that both voltammograms 4 and 5 indicate the presence of a mixture of Br2/Br , with elemental bromine in excess.
This is accompanied by a fast tri-bromide ion formation involving bromine excess: Br2 Br XBr 3
(4)
(5)
One may exclude the assumption that any impurity existing in PC, rather than the solvent itself, is responsible for the reactions with bromine, because, not only was very pure solvent used, but also, the concentration of water, being the main impurity (/0.02%, or /0.01 M), was additionally reduced by drying with dry zeolite. On the other hand, unlike the reaction with impurities, the voltammetric waves increased proportionally to the bromine concentration when it was increased up to 0.029 M. It is noteworthy that the system under investigation, in both the chemical and electrochemical sense, behaves similarly to the dimethylsulphoxide/bromine system. Martin et al. [27] have shown by the methods of organic chemistry, that dimethylsulphoxide (DMSO) undergoes direct bromination; however, the reaction is not simple and, as a consequence, a variety of products, such as trimethylsulphonium bromide and HBr, appear. To explain the mechanism of bromination, the authors assumed that the following reaction occurred in the initial stage:
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(CH3 )2 SO2Br2 [(CH3 )2 SOBr ]Br 3
(6)
The bromination of DMSO may proceed to completion if a compound able to eliminate acid is added to the solution. In addition, we found that the voltammetric curves of the DMSO/Br2 system, immediately after solution preparation, reflect the presence of a high concentration of bromide ions and strong acid ([28], and unpublished results). Although Fig. 1 shows the reactions between PC and Br2, voltammetric curves 4 and 5 do not allow us to conclude more than is expressed by Eq. (4), and the reaction mechanism may be more complex. For instance, either disproportionation of bromine to positively and negatively charged ions, or adducts similar to that defined by Eq. (6) may be assumed to appear in the initial stage of the reaction. The disproportionation of halogens in organic solvents has already been reported in the literature. Iodine disproportionation was proposed elsewhere in order to explain the shape of the voltammetric curves in pyridine [13,29] as well as to explain the appearance of tri-iodide ion spectra in the UV /Vis spectral region in dimethylsulphoxide [30,31]. A pronounced disproportionation was evident in the halogen/solvent system involving mixed halogens, for instance Cl2/I2 [32]. Further to this, the final product of PC bromination may not be that proposed in Eq. (4), and its exact composition remains also a matter for further investigation. The progressive bromine consumption produces no observable change in the distance between the anodic and the final cathodic current plateau in voltammograms 4 and 5, since bromine becomes replaced by an equivalent amount of other electroactive species. This investigation shows that bromide oxidation in PC presents, in fact, an electrochemical reaction of an EC type. With regard to chloride oxidation, shown in the anodic part of voltammogram 1, one may see that, in comparison to other halogenide oxidations, chloride oxidation requires the highest positive potential. Therefore, anodic chlorine evolution is accompanied most probably by a direct anodic substitution of hydrogen in PC molecules. As is obvious in voltammogram 1, the reaction products adsorb on the electrode surface, causing passivation, and therefore diffusion limitation of chloride oxidation appears to be unattainable. Also, there is no evidence for the formation of intermediate complex ions, which is in accordance with the lowest values of the stability constants of polychlorine species in organic solutions, in comparison to other polyhalogens [32]. Among the studies in aprotic solvents, chloride oxidation was studied also in acetonitrile [33,34], nitromethane [35] and dimethyl formamide [36,37], and in each solvent the limiting diffusion current of
chloride oxidation was reported to be attained. However, similarly to the case studied here, almost always, the oxidation was found to occur as a one-step process, i.e. without intermediate complex ion formation.
Acknowledgements One of the authors (SM) is grateful to the Ministry of Science, Technology and Development of the Republic of Serbia, for supporting this work through the contract entitled ‘Structure, thermodynamic and electrochemical properties of modern materials for energy conversion and components in electronics’. The authors are indebted to Prof. H. Lund, University of Aarhus, Denmark, for directing their attention to Ref. [27]. The authors are grateful also to Dr C. Amatore, Ecole Normale Superieure, Paris, France, for helpful discussion.
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