Ethanol electro-oxidation: Cyclic voltammetry, electrochemical impedance spectroscopy and galvanostatic oscillation

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Ethanol electro-oxidation: Cyclic voltammetry, electrochemical impedance spectroscopy and galvanostatic oscillation Lijuan Han a, Hua Ju b, Yanhui Xu a,* a b

Institute of Chemical Power Sources, Soochow University, Moye Rd 688, Suzhou 215006, China School of Urban Rail Transportation, Soochow University, 215006 Suzhou, China

article info

abstract

Article history:

In the present work, ethanol electro-oxidation reaction on Pt electrode has been studied in

Received 12 June 2012

details by measuring and analyzing the cyclic voltammetries (CV), the dependence of the

Received in revised form

electrochemical impedance spectroscopies (EIS) on the applied potentials and the galva-

24 July 2012

nostatic potential oscillation. The CV measurement has exhibited a bistable characteristic.

Accepted 9 August 2012

The origin of all the oxidation and reduction peaks has been analyzed. The origin of the

Available online 28 August 2012

bistable characteristic has also been studied by measuring the EIS at different potentials. In addition, the dependence of the galvanostatic oscillation on the applied potential and the

Keywords:

ethanol concentration has been reported.

Galvanostatic oscillation

Copyright ª 2012, Hydrogen Energy Publications, LLC. Published by Elsevier Ltd. All rights

Negative real impedance

reserved.

Cyclic voltammetry Ethanol electro-oxidation

1.

Introduction

In the past decades, direct alcohol fuel cells (DAFC), as one of the most promising chemical power sources, have attracted more and more attention due to their high energy-conversion efficiency, proper operation temperature and simple handling of fuel, as well as convenience to use [1e7]. The main challenge faced by DAFC includes the solution of poor electrooxidation reaction kinetics, the poisoning tendency of the anode catalysts by some intermediates, such as adsorbed CO [8e14], as well as the large over-potential for oxygen reduction reaction. Much attention has been paid on the Pt-based electrocatalysts that are well known to have a good catalytic behavior for methanol, ethanol, propanol and ethylene glycol etc. electro-oxidation reaction [2,5,12e20]. To improve the

electroecatalytic activity of Pt-based electro-catalysts, it is of very importance to understand the electro-oxidation reaction mechanism [2,8e11,21e23], but until now the detail of the electro-oxidation reaction mechanism remain controversial, especially for long-chain alcohols. In the ethanol molecule there are two kinds of protons. One is bonded to the oxygen atom. If it is firstly oxidized, then the first reaction step is CH3 CH2 OH / ½CH3 CH2 Oad þ Hþ þ e The other is bonded to the carbon atom. If it is firstly oxidized, then the first reaction step should be CH3 CH2 OH / ½CH3 CHOHad þ Hþ þ e It can be not still concluded which one of the two reaction pathways is right.

* Corresponding author. Tel.: þ86 512 67261337; fax: þ86 512 67261575. E-mail addresses: [email protected], [email protected] (Y. Xu). 0360-3199/$ e see front matter Copyright ª 2012, Hydrogen Energy Publications, LLC. Published by Elsevier Ltd. All rights reserved. http://dx.doi.org/10.1016/j.ijhydene.2012.08.034

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Some electro-oxidation systems are able to exhibit oscillatory behavior. The appearance of the oscillatory behavior is related with the co-existence of the positive and negative feedback. For the alcohol electro-oxidation, it is commonly accepted that the negative feedback is just the poisoning process. Therefore, the analysis and study of the oscillatory behavior may provide some useful information on the poisoning reaction. It is plausibly concluded that the longchain alcohol has a larger poisoning ability. Therefore we are planning to compare the electrochemistry of methanol, ethanol, propanol, ethylene glycol, propylene glycol and propanetriol etc by comparing their oscillatory response. Our purpose is to explore the reaction mechanism and to find the ways to improve the electro-oxidation reaction kinetics of the alcohols, as results, the fuel cell performance. In the present work, the ethanol electro-oxidation was firstly reported.

2.

Experimental

The solutions were prepared by dissolving the chemicals (CH3CH2OH p.a., H2SO4 p.a.) in ultra-pure water (18.2 MU, Hitech Ultra-pure water system). The working electrode is a Pt metal electrode with an area of 24 mm2. The purity of Pt metal is 99.99%. The counter electrode is a Pt electrode with much larger area, and the reference electrode is Hg/Hg2SO4. The distance between the working electrode and reference electrode was set up at 20 mm. The corresponding electrochemical tests were performed with PE Parc 2273 or CHI 660C electrochemical workstations. To remove some possible organic impurities existing on Pt electrode surface, before electrochemical measurement, the working electrode was pre-treated by sweeping its electrode potential between-in 0.65 V and 0.80 V in 0.5 mol L1 H2SO4 solution until the standard CV was obtained, as showed in Fig. 1.

3.

Results and conclusion

The CV pattern in sulfuric acid electrolyte is useful to analyze the electro-oxidation mechanism of ethanol at Pt electrode.

Fig. 1 e The CV pattern of Pt electrode in 0.5 mol L solution, the scanning rate is 50 mV sL1.

L1

H2SO4

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Fig. 1 shows the CV pattern of Pt working electrode in 0.5 mol L1 sulfuric acid solution. For the two couples of redox peaks positioned at 0.44 V and 0.59 V, the oxidation peak potential is the same as the reduction peak potential, which means that they represent the under-potential electro-deposition and adsorption/desorption of the hydrogen. The oxidation peaks at 0.25 V and 0.36 V should correspond to the two-step formation of Pt oxides/hydroxides. At the potentials of more than 0.8 V the oxygen evolution reaction becomes more obvious. The reduction peak at 0.02 V is believed to relate with the reduction of Pt oxides/hydroxides. The phenomenon that this reduction peak is sharper than the corresponding oxidation peaks may imply that the reduction reaction of Pt oxides/hydroxides is a fast step. The obtained CV pattern of Pt electrode in 0.5 mol L1 H2SO4 solution is same as that reported in literature [3,24,25], which has confirmed our experimental quality. There is no impurity in the solution and on the electrode surface. Fig. 2 shows the CV pattern in 0.1 mol L1 H2SO4 þ x mol L1 CH3CH2OH (x ¼ 0.05, 0.5 and 3.0) solutions at the scanning rate of 1 mV s1. The first main oxidation peak pa1 is found to be at 0.2 V for the solution containing 3.0 mol L1 CH3CH2OH. One shoulder peak pas can be found at the left side of the first main oxidation peak, at about 0.01 V. In some literature [26e29] there is no shoulder peak being reported at the low-potential side of the main oxidation peak. The position of the shoulder peak is not influenced by the methanol concentration and the scanning rates. It is reasonable to conclude that this shoulder peak pas should originate from the non-faradiac adsorption/desorption process of some organics. The formation of the first main oxidation peak is believed to relate with the competitive adsorption between OH and oxygen-containing intermediates. Here the OH adsorbent is formed via the water oxidation decomposition. Usually, the oxygen-containing intermediate is assigned to CO [8e14]. Fig. 3 shows the dependence of the CV pattern on the scanning rates. The potential of the oxidation peak pa2 is about 0.45 V when the scanning rate is 1 mV s1. The decrease in ethanol concentration does make the peak potential slightly more positive, as showed in Fig. 2. In addition, the increase in the scanning rate does also lead to more positive peak potential. These results may imply that the oxidation peak pa2

Fig. 2 e The CV pattern in 0.1 mol LL1 H2SO4 D x mol LL1 CH3CH2OH (x [ 0.05, 0.5 and 3.0) solutions, the scanning rate is 1 mV sL1.

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Fig. 3 e The CV patterns of Pt electrode in 3.0 M CH3CH2OH D 0.1 M H2SO4 solution.

should represent an electrochemical process accompanied by charge-transfer process. The oxidation peak pa3 can not be found when the ethanol concentration is 0.05 mol L1, as seen in Fig. 2, while it appears at 0.96 V in the solution containing 3.0 mol L1 CH3CH2OH. The increase in the scanning rate makes the peak potential more positive. The fact that the oxidation peak position is influenced by the scanning rate may imply that this oxidation peak represents an electrochemical process in which a charge transfer process should be involved. In the negative sweeping potential direction, the oxidation peak par appears at 0.03 V when the ethanol concentration is 3.0 mol L1. And, its peak potential becomes more negative as the ethanol concentration decreased (Fig. 2) or the scanning rate increased (Fig. 3). At lower scanning rates, the oxidation peak par is left-to-right asymmetric. It is believed that there exists an oxidation peak hidden in the low-potential side of the oxidation peak par and the peak potential of this hidden oxidation peak is not influenced by the scanning rates and the ethanol concentrations. It should have a same origin with the oxidation peak pas. The oxidation peak par is believed to originate from the competitive adsorption between OH and CO, and to represent the ethanol electro-oxidation at metallic Pt electrode surface. It has a same origin with the oxidation peak pa1. One small reduction peak can be found at 0.015 V from the Fig. 2, when the ethanol concentration is 0.05 mol L1. The reduction peak is hardly to see when the ethanol concentration increased, but becomes clearer if the scanning rate increased. These phenomena have at least conformed that the reduction peak represents one electrochemical reaction process and the corresponding reduction reaction is of a fast step. In Fig. 1, the reduction peak represents the electroreduction of Pt oxides/hydroxides. We think, the reduction peak in the solution containing ethanol should also represent the electro-reduction of Pt oxides/hydroxides. The ethanol or some intermediates can react with the Pt oxides/hydroxides, which leads to the disappearance of reduction peak in higher ethanol concentration. Fig. 4 shows the dependence of the peak potential and current on the scanning rate for the oxidation peaks pa2 and

Fig. 4 e The dependence of the peak current (Ipa2, (A); Ipa3, (C)) and potential (Ppa2, (B); Ppa3, (D)) on the scanning rate in 3 M CH3CH2OH D 0.1 M H2SO4 solution.

pa3 in 3 M CH3CH2OH þ 0.1 M H2SO4 solution. The relation between the peak current and the square root of the scanning rate is linear for the reversible and irreversible electrochemical reactions. The peak potential is independent of the scanning rates for the reversible reaction, while it has a linear dependence upon the logarithm of the square root of the scanning rate if the reaction is irreversible. It can be seen from Fig. 4 that the peak current of the oxidation peak pa2 has an approximate linearization with the square root of the scanning rate. The peak potential does also

Fig. 5 e The electrochemical impedance spectroscopies of Pt electrode in 3 M CH3CH2OH D 0.1 M H2SO4 solution at different applied potentials.

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Fig. 6 e The electrochemical impedance spectroscopies of Pt electrode in 0.5 M CH3CH2OH D 0.1 M H2SO4 solution at different applied potentials.

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Fig. 8 e The time evolution of the electrode potential at different applied currents in 3.0 M CH3CH2OH D 0.1 M H2SO4 solution.

linearly vary with the logarithm of the square root of the scanning rate. It could be concluded from the Fig. 10 that the electrochemical reaction that is represented by the oxidation peak pa2 is a fully irreversible. For the oxidation peak pa3, the linearity between the peak current and the square root of the scanning rate is very bad, while linearity between peak potential and the logarithm of the square root of the scanning rate is better. At the scanning rate of 50 mV/s, the peak current is much smaller than that expected from the linear relation, which may be caused by some unexpected reasons. However, the linearity becomes better if the CV obtained at 50 mV/s is excluded. Based upon the above-mentioned analysis, it could be concluded that the oxidation peak pa3 is an irreversible reaction peak. The two oxidation peaks pa2 and pa3 should represent the electro-oxidation process of the ethanol on the oxidized electrode surface. Fig. 5 shows the potentiostatic EIS pattern in 0.1 mol L1 H2SO4 þ 3.0 mol L1 CH3CH2OH solution. For 0.06 V, the impedance data in Nyquist plot emerge in a clockwise mode as the applied frequency decreases. Such a kind of EIS pattern is commonly encountered [30]. The impedance spectrum has a qualitative change when the applied potential is increased up to 0.09 V. As the frequency decreases, the real impedance becomes negative for the frequency of about 0.264 Hz. At the

lowest frequency (0.001 Hz) the real impedance becomes positive again. The appearance of the negative real impedance for non-zero finite frequencies may imply that such an electrochemical system can exhibit unstable behavior [31]. At higher potential than 0.36 V, the impedance spectra becomes “normal” as that obtained at 0.06 V. Fig. 6 shows the potentiostatic EIS pattern in 0.1 mol L1 H2SO4 þ 0.5 mol L1 CH3CH2OH solution. At 0.01 V, the impedance data emerge in a clockwise mode as the frequency decreases. When the frequency decreased from 0.333 Hz to 0.264 Hz, the imaginary impedance becomes positive and an inductive arc appears. In general, it is commonly accepted that the inductive arc in the low-frequency domain in Nyquist plot should represent the adsorption/desorption process. The fast transition from the capacitive behavior to the inductive behavior may imply a fast reaction kinetic for adsorption/ desorption process. The EIS pattern obtained at 0.05 V is similar as that obtained at 0.09 V in the solution containing 3.0 mol L1 CH3CH2OH. When the electrode potential was increased up to more than 0.37 V, the impedance spectra becomes “normal”. The EIS pattern that is obtained at 0.15 V in the solution containing 0.05 mol L1 CH3CH2OH is somewhat different, as

Fig. 7 e The electrochemical impedance spectroscopies of Pt electrode in 0.05 M CH3CH2OH D 0.1 M H2SO4 solution at different applied potentials.

Fig. 9 e The time evolution of the potential at the applied currents in 0.1 mol LL1 H2SO4 D 0.5 mol LL1 CH3CH2OH solution.

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Fig. 10 e The time evolution of the electrode potential at the applied currents in 0.1 mol LL1 H2SO4 D 0.05 mol LL1 CH3CH2OH solution.

shown in Fig. 7. As the frequency decreases, the impedance data emerge in a counterclockwise mode. The real impedance becomes negative for the lowest frequency, not for the nonzero finite frequencies. Based on such a kind of impedance spectrum, it can not be concluded whether the electrochemical system is able to exhibit a nonlinear behavior. To our knowledge, there is no literature concerning the unstable behavior of an electrochemical system having such a kind of impedance spectrum. In the following paragraph, it can be found that such an electrochemical system is not able to exhibit oscillatory behavior. Fig. 8 shows the galvanostatic potential oscillation curves in 0.1 mol L1 H2SO4 þ 3.0 mol L1 CH3CH2OH solution. For the 0.05 mA current, the oscillation starts at about 2.7 h. The oscillation amplitude can reach 0.314 V (from 0.18 V to 0.134 V). The average potential is about 0.0 V. The oscillation belongs to the relax-type oscillation. The potential jumps very quickly from the high-potential state to the low-potential state, and then slowly wanders back to the high-potential state. As the time goes, the oscillation becomes faster. The situation is similar for the 0.1 mA current but the oscillation amplitude becomes larger and reached 0.344 V (from 0.21 V to 0.134 V). As the current increases, the induction duration becomes shorter, from 2.7 h for 0.05 mA, to 1.8 h for 0.1 mA and to 0.46 h for 0.5 mA, before the oscillation behavior starts to occur. For 0.5 mA current, the oscillation amplitude is 0.33 V. There is no induction duration being found in the timeeevolution curve of the potential when the applied current is 0.9 mA. The oscillation has a relatively high frequency but the oscillation amplitude is relatively small for 0.9 mA. As the current increased up to 1.3 mA, the oscillation becomes slow again. At 1.4 mA, there is no oscillatory behavior being found. Fig. 9 shows the galvanostatic potential oscillation curves in 0.1 mol L1 H2SO4 þ 0.5 mol L1 CH3CH2OH solution. For 0.1 mA current, the oscillation amplitude is about 0.29 V. And the induction duration is 4 h. Our experiment has showed that the induction duration can reach more than 12 h if the applied current decreased furthermore. The oscillation type remains unchanged when the applied current is changed for the

0.1 mol L1 H2SO4 þ 3 mol L1 CH3CH2OH solution, as showed in Fig. 8, i.e., the oscillation is always the relax-type oscillation. But when the ethanol concentration decreased down to 0.5 M, the oscillation transited from the relax-type oscillation for the low currents to the nail-type one for the high currents. In addition, the parameter space in which the system is able to exhibit the oscillatory behavior becomes thinner. When the current is increased up to 0.6 mA, the oscillatory behavior disappears. Fig. 10 shows the time evolution curves of the electrode potential under a galvanostatic operation mode in 0.1 mol L1 H2SO4 þ 0.05 mol L1 CH3CH2OH solution. In our experimental conditions there is no oscillation being found. According to the literature [8e11,21e23], the ethanol electro-oxidation mechanism can be summarized as follows, nPt þ CH3 CH2 OH ðaqÞ / Ptn  CH3 CH2 OH

(1)

Ptn  CH3 CH2 OH / Ptn  CH3 CHOH þ Hþ þ e

(2)

Ptn  CH3 CHOH / Ptn  CH3 COH þ Hþ þ e

(3)

Ptn  CH3 COH / Ptn  CH3 CO þ Hþ þ e

(4)

Pt þ Ptn  CH3 CO / Ptn  CO þ Pt  CH3

(5)

In common, the reaction pathway expressed by these five reactions is more possible in the acidic medium. In the alkaline medium, the most possible reaction pathway could be expressed by the following reactions. Ptn  CH3 CH2 OH þ OH / Ptn  CH3 CH2 O þ H2 O þ e

(6)

Ptn  CH3 CH2 O þ OH / Ptn  CH3 CHO þ H2 O þ e

(7)

Ptn  CH3 CHO þ OH / Ptn  CH3 CO þ H2 O þ e

(8)

It should be noted that some intermediates are possible to leave the electrode surface to the solution as the incomplete by-products, such as CH3COOH, CH3CHO, CH4 and CH3CH3 etc. The adsorbed CH3 group has two possible fates. It can combine with the adsorbed proton to form methane, or two CH3 groups combine to form ethane, Pt  CH3 þ Pt  H/2Pt þ CH4 Pt  CH3 þ Pt  CH3 / 2Pt þ CH3 CH3

(9) (10)

There exist two possible explanations on the oxidation peak pa1 formation. In the first explanation, the peak current is contributed to the direct oxidation reaction pathway of ethanol. As the potential increases some formed intermediates deactivate the electrode and then the oxidation current declined. In this explanation, the electrode surface is covered by some adsorbed intermediates if the electrode potential is kept at a potential more than the pa1, the peak potential of the main oxidation peak, for instance, at 0.45 V in Fig. 2. In fact, we could directly say that the direct oxidation reaction is of very slow rate, although there is no direct evidence. Usually, the anode in direct ethanol fuel cell operates at the potential near to pa1, which does also imply that the first explanation must be incorrect. The reason is simple. The

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main intermediate that deactivates the Pt electrode is the adsorbed CO. According to literature [32], the CO has a stronger blocking effect for the entrance of other species into the Pt electrode surface. The electrode surface is full of CO if the electrode was kept at Ppa1 for long time and if the first explanation is correct. In such a condition, the potentiostatic current will decay down to zero. In the second explanation, the electrode surface is covered by OH-type species if its potential is kept at 0.45 V in Fig. 2. The competitive adsorption reaction between OH and CO species leads to the formation of the oxidation peak pa1. Once the Pt electrode is immersed into the acidic solution containing ethanol, the non-faradiac chemical reaction will take place on the electrode surface. As results, some intermediates formed on the surface and adsorbed on it. One of the main intermediates is CO. As the potential increases, the indirect electro-oxidation occurs, at same time the OH species starts to form via the decomposition of the water and remove the adsorbed CO via the following reactions Ptn  CO þ Pt e OH / Ptm  COOH þ ðn þ 1  mÞPt

(11)

Ptm  COOH/mPt þ CO2 þ Hþ þ e

(12)

The CO adsorption on the Pt surface is dependent of the electrode potential [33]. As the potential increases and the reactions (11) and (12) proceed, at last, all the CO-type species existing on the electrode surface are removed and the electrode surface is covered fully by the OH. As mentioned in the literature [32], CO has a stronger blocking effect for the entrance of other species. But the OH-covered surface is open for the entrance of other species. It can be expected that the OH-covered electrode should have smaller impedance in comparison with the CO-covered electrode. The impedance measurement has partially proved this suggestion. For the 0.1 mol L1 H2SO4þ3.0 mol L1 CH3CH2OH solution, the impedance arc obtained at 0.06 V (the CO-covered surface) is larger than that obtained at 0.36 V (the OH-covered surface), as showed in Fig. 5. The situation in 0.1 mol L1 H2SO4þ0.5 mol L1 CH3CH2OH solution is similar, as showed in Fig. 6. The impedance arc obtained at 0.01 V has the biggest diameter. However, the situation in 0.1 mol L1 H2SO4 þ 0.05 mol L1 CH3CH2OH solution is different. The impedance spectrum in this solution does not support our suggestion. As mentioned above, the shoulder peak in the positive sweeping potential direction (PSPD) and the hidden oxidation peak in the negative sweeping potential direction (NSPD) have a same origin and are believed to be due to the non-faradiac adsorption/desorption process of some organics. In principle, the peak position should be not influenced by the potential-scanning direction for a non-faradiac adsorption process. From the CVs showed in Figs. 2 and 3, it can be found that their position is being affected by the ethanol concentration and the potential-scanning direction as well as the scanning rates. This phenomenon should originate from the disturbance of the main oxidation peaks pa1 and par in our experimental condition. All the possible reaction pathways for ethanol electrooxidation were summarized in Fig. 11. The CV measurement

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Fig. 11 e Multi-pathway reaction model of ethanol electrooxidation.

has confirmed that the electrochemical system could exhibit unstable behavior. And the EIS results have evidenced that the parameter space in which the electrochemical system can exhibit nonlinear behavior falls into the potential range where the first main oxidation peak appears. The average potential of the oscillatory potentials is about 0.0 V. All the phenomena imply that the oscillation relates with the electrode processes represented by the shoulder and the main oxidation peak pa1 and the pas. At the potential maximum, the electrode is covered by OH, while at the potential minimum it is covered by oxygen-containing incomplete oxidation products, such as CH3CHO, CH3CO or CO. If the direct electro-oxidation reaction pathway can be omitted, then the positive feedback includes the reactions (1)e(4), while the negative feedback is related with the reaction (5). In practice, the reaction represented by the oxidation peak pa1 is employed. There should be two strategies to improve the electroecatalytic activity depending on the electro-oxidation reaction rate. When the electrode works in the low-potential region (LPR), as indicated in Fig. 3, the electro-oxidation reaction rate can be improved by using some additives/ supplementary catalysts to promote the formation of adsorbed OH species, because in this potential region the removal of the adsorbed CO is crucial. If it is hoped that the electrode works in the high-potential side (HPS) of the main oxidation, as indicated in Fig. 3 (in this case a larger operation current is expected), then the improvement in the electro-oxidation reaction rate should be realized by using such a kind of additives or supplementary catalysts that are able to delay or inhibit the formation of adsorbed OH species, because in this case the detrimental factor to increase the operation current is the poisoning effect of the adsorbed OH.

4.

Conclusion

In the article, the ethanol electro-oxidation reaction has been studied in details. The cyclic voltammetry measurements have shown that in higher ethanol concentration there exist three oxidation peaks whose position is dependent of the scanning rate. At the low-potential side of the first oxidation peak there is a shoulder peak. In the negative sweeping potential direction one asymmetric oxidation peak can be

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found. In his low-potential side there should be a hidden oxidation peak that is belied to corresponds to the nonfaradiac adsorption/desorption process. The two oxidation peaks at more positive potential should correspond to the electro-oxidation at oxidized electrode surface. It is found that the potentio-static impedance measurement exhibits a qualitative change as the applied potential changes. In general, in EIS Nyquist plot, the EIS data emerge in the clockwise mode, but in some special potentials the EIS data emerge in the counterclockwise mode. The real impedance becomes negative for non-zero finite frequencies and positive for the lowest frequencies in higher ethanol concentrations (3.0 mol L1 and 0.5 mol L1). In 0.1 mol L1 H2SO4 þ 0.05 mol L1 CH3CH2OH solution, the real impedance becomes negative for the lowest frequencies, not for non-zero finite frequencies. The galvanostatic measurements have showed that the potential oscillation can be found for the higher-concentration ethanol solutions. There is no oscillatory behavior being found in 0.1 mol L1 H2SO4 þ 0.05 mol L1 CH3CH2OH solution. The galvanostatic oscillation amplitude can reach more than 0.34 V, and the oscillation frequency is in an order of magnitude of about 104e101 Hz, depending on the applied current and the ethanol concentration.

[9]

[10]

[11]

[12]

[13]

[14]

[15]

Acknowledgment

[16]

This work was supported by the Foundations No. SRF for ROCS, SEM and BK2009110.

[17]

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