Anal. Chem. 1995,67, 976-980
Electrochemical Pretreatment of Carbon Electrodes as a Function of Potential, pH, and Time Alvin L. Beilby,* Tania A. -saki,+ and Howard Y. Stem*
Department of Chemistry, Pomona College, 645 North College Avenue, Claremont, Califomia 91711-6333
The electrochemical pretreatment of carbon electrodes through the oxidation and reduction of the electrode surface is a widely used procedure to improve electrode response; however, there still remain many unanswered questions about the nature of the pretreatment process. The main variables of the pretreatment process are the oxidation and reduction potentials, the composition and pH of the electrolyte solution, and the length of time of oxidation and reduction. This study focuses on the effects of these variables on the primary redox reactions that are responsible for activation of the electrodes. In the first part of this study, an anodic process starting at +1.6 V (SCE) and a cathodic process with a peak potential of approximately - 1.O V were observed with a glassy carbon electrode in solutions of pH 12 or greater. These processes are similar to those reported for pretreatment in acidic and neutral solutions and are distinct from the previously reported anodic process in 1 M NaOH which occurs between +0.9 and +1.5 V and gives a cathodic process with a peak potential at +0.2 V. In the second part of this study, the pH dependence of the anodic and cathodic processes in the pH range from 2 to 12 was examined by changing the pH of the solution during the course of oxidizing and reducing a glassy carbon electrode surface while keeping the electrode under potential control at all times. The results support the idea of the formation of Merent species on the electrode surface when anodization is done at different pH’s. During the course of the study, the length of time of anodization also was found to have an effect on the observed reduction potentials. Electrochemical pretreament was the first technique used to activate carbon electrodes to improve electrode response. In the fist paper on the use of graphite electrodes for voltammetric studies, Lord and Rogers’ reported that anodizing the electrode greatly improved the shape of a slow-scan voltammetric curve for the reduction of Fe(1ID in 0.1 M KCl. Since the time of this first use of electrochemical pretreatment, the technique has become widely used to activate carbon electrodes, and understanding the nature of electrochemical pretreatment on the carbon electrode surfaces has become a major area of study. There still remain, however, many unanswered questions about the nature of the + Present address: Department of Chemistry, University of California, Riverside, Riverside, CA 92521-0403. * Present address: Department of Biochemistry, SJ-70, University of Washington, Seattle, WA 98195. (1) Lord, S. S., Jr.; Rogers, L. B.Anal. Chem. 1954,26, 284-295.
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anodization and cathodization processes involved in electrochemical treatment. The reponse of carbon electrodes to pretreatment depends upon the nature of the pretreatment. The main variables of the electrochemical pretreatment process are the following: (1) the positive (anodic) and negative (cathodic) potential limits to which the electrode is exposed, Le., the pretreatment potentials, (2) the composition of the electrolyte solution, which includes both the pH of the solution and the nature and concentration of the electrolyte ions besides H+, and (3) the length of time at which the electrode is held at the pretreatment potentials (dc pretreatment) or how rapidly the electrode is switched between the potentials (ac pretreatment). A working hypothesis for the surface effects of electrochemical pretreatment, along with the effects of vacuum heat treatment and laser activation on polished glassy carbon, has been given by McCreery2in Figure 46 of his review on carbon electrodes. The main feature of electrochemicalpretreatment shown in the figure is the formation of an electrochemical graphitic oxide film that appears when the electrode is subjected to high positive potentials in acidic or neutral solutions. Treatment with base or electrochemical pretreatment in base (1 M NaOH) is shown to remove the film but to leave many surface oxides. In the study of Beilby and Carlsson3 on electrochemical pretreatment in base (1 M NaOH), an anodic process that takes place between +0.9 and +1.5 V (Ag/AgCl) and that gives rise to a cathodic process with a voltammetric peak potential of a p proximately +0.2 V was shown to activate glassy carbon and pyrolytic carbon film electrodes. This pretreatment process was used by Anjo et al.4 in the analysis of dopamine, where the optimum pretreatment conditions were a solution pH of 1 and a potential of f 1 . 2 V (SCE) for a duration of 5 min. In our study, we show that an additional anodic process takes place beyond +1.6 V. This process appears to be similar to the process reported by Engstrom5for neutral solutions and by Kepley and Bard6for acidic solutions. Because the surface oxide species may include H atoms, it is logical to assume that the type of oxides or the composition of the electrochemicalgraphitic oxide will change, depending upon the pH of the electrolytic solution used for the electrochemical pretreatment. In previous studies where the pH of the electrolyte (2) McCreery, R L. In Electroanalytical Chemisty, Bard, A J., Ed.; Marcel DeWter: New York, 1991; Vol. 17, pp 221-374. (3) Beilby, A L.; Carlsson, A j. Electrounal. Chem. 1988,248, 283-304. (4) Anjo, D. M.; Kahr, M.; Khodabakhsh, M. M.; Nowinski, S.; Wanger, M. Anal. Chem. 1989,61, 2603-2608. (5) Engstrom, R C. Anal. Chem. 1982,54, 2310-2314. (6)Kepley, L. J.; Bard, A J. Anal. Chem. 1988,60, 1459-1467.
0003-2700/95/0367-0976$9.00/0 0 1995 American Chemical Society
solution has been varied, a separate measurement was made for each different pH. In the process, the electrode was moved from one solution to another, causing the electrode to be released from potential control for a period of time and causing the surface of the electrode to be exposed to the air. The loss of potential control and the exposure to air could modify the electrode surface, thus affecting the kinetics of redox reactions. To eliminate these two possible complications, in our study the pH of the electrolyte solution was changed during pauses in the voltage scans during which the electrode remained under full potential control at all times and was not exposed to the air. Most studies on electrochemical pretreatment of carbon electrodes have been concerned with the state of the surface after the primary redox reactions have taken place, whereas our study investigates the primary reactions themselves in order to better understand how the nature of these reactions leads to activation. EXPERIMENTALSECTION Apparatus. Voltammetric curves were obtained with an IBM EC/225 voltammetric analyzer coupled to a Mosely 2D-2AM X-Y recorder. pH measurements were made with an Orion SA 250 pH meter or a Coming 220 pH meter. For the results reported in this paper, a standard Bioanalytical Systems (BAS) 3 mm diameter glassy carbon electrode was used. The electrode was polished between experiments with BAS PK-3 polishing kit materials (6 and 1pm diameter particle diamond polish on cloth followed by 0.05 pm diameter particle alumina on a glass plate). A standard H cell was used with one side containing an auxiliary platinum foil electrode and the other side containing the glassy carbon electrode, a saturated calomel electrode (SCE) ,a combination pH electrode and automatic temperature probe, and a bubbler tube for nitrogen. All experiments were done at room temperature. Reagents. All solutions were prepared from standard reagent grade chemicals. Laboratory grade nitrogen was used to purge the solutions of dissolved oxygen. Ultrapure reagents were not used since only the primary redox reactions were studied. To obtain an initial solution of a given pH, a 0.1 M NaOH solution was neutralized with 1 M H3PO4; 1 M H3P04 was used also to change the pH during a scan. Procedure. Normally a voltammetric cycle was started at 0 V (SCE) scanning at 100 mV/s to the positive scan limit. The scan either was held at the limit for a specified length of time, usually 15 s, or reversed in direction immediately. The scan was stopped at $0.4 V so that the solution could be bubbled with nitrogen for 5 min to remove dissolved 02. Any pH changes were also made at +0.4 V. The scan was continued in either direction, depending upon the nature of the experiment. Normally, the scan was continued in the negative direction until the negative limit was reached. The scan was reversed immediately at the negative limit and returned to 0 V, where, in most cases, the scan was continued for additional cycles without stopping. RESULTS AND DISCUSSION Anodization Potentials in Basic Solutions. The study of Beilby and Carlsson3 on the electrochemical pretreatment of carbon electrodes in 1 M NaOH did not address the issue of whether or not the oxidation reaction taking place between +0.9 and f 1 . 5 V, which gives a reduction reaction peaking at $0.2 V, was the same type of reaction shown by Engstrom5for neutral
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Figure 1. Cathodic portion of cyclic voltammograms for a glassy carbon electrode in pH 12 solution with increasing positive limits: (1) 1.6, (2) 1.8, (3)2.0, and (4)2.2 V limit; 15 s pause at the positive potential limit.
solutions. Two factors suggest that the process reported by Beilby and Carlsson is not the same type as that associated with pretreatment in neutral to acidic solutions. First, the $0.2 V potential for the reduction reaction in 1M NaOH is considerably more positive than the potentials for the reduction reactions in neutral and acidic solutions. Second, the amount of charge associated with the reduction reaction in 1 M NaOH appears to be considerably less than the charge associated with the reduction reaction shown by Eng~trom.~ Since the amount of this charge was not reported by Beilby and Carl~son,~ we did a rough calculation of the charge associated with the cathodic peak for an electrode that had been anodized at +1.5 V in a pH 12 solution for 60 s and obtained a charge of 1.3 x C/cm2. Using the constants of Bowling et al.,7 this charge would correspond to approximately 1.3 monolayers of carbon oxides, whereas the values given for the reduction reaction in neutral solutions correspond to many monolayers.5 To test whether the anodic process described by Engstroms for neutral solutions is seen when an electrode is anodized at potentials greater than +1.5 V in basic solutions, we increased the positive potential limit in 0.2 V increments starting at $1.6 V. The results of this experiment in a pH 12 solution are shown in Figure 1, which focuses on the cathodic peaks at -1.0 V. This reduction peak appears clearly only when the positive potential limit is greater than $1.6 V. The sizes and potentials of these cathodic peaks are similar to the cathodic peaks seen after anodization in neutral and acidic solutions, suggesting that the anodic processes taking place beyond +1.6 V in basic solutions are similar to the anodic processes taking place in the same potential region in neutral and acidic solutions. pH 12 was chosen for the most basic solution for the reported data because the IBM voltammetric analyzer tended to overload at the higher potentials with 1 M NaOH solutions. Preliminary experiments showed that the cathodic peak at +0.2 V associated (7) Bowling, R; Packard, R.T.; McCreery, R L. Langmuir 1989,5, 683-688.
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with the anodic process from +0.9 to +1.5 V in 1 M NaOH appeared also in pH 12 solutions; however, at pH 12 this peak was smaller in size and was shifted approximately 0.1 V positive from the 1M NaOH solution peak. In pH 10 solutions, a cathodic peak could not be detected. The data of Anjo et aL4support our data, suggesting that the anodic process from f0.9 to +1.5 V only takes place above pH 10. The results thus show that there are two anodic processes that must be considered when electrochemical pretreatment is utiliied in basic solutions of pH 12 or greater. One characteristic of the electrochemical graphitic oxide film is the change in color of the electrode due to interference effects produced by the nearly transparent Because no colors were observed on the electrode surface with short times of anodization, we performed an extended anodization in 1M NaOH using a large current capacity potentiostat to see whether colors would be produced on the electrode surface. When a glassy carbon electrode was held at +2.1 V (SCE) for 5 min, a brownishyellow color appeared on the surface of the electrode in a speckled fashion. After washing, the electrode appeared very dull with a speckled appearance but with no color remaining, suggestingthat the electrode surface developed an unstable graphitic oxide film that was physically removed by the washing. The results of Anjo et ale4suggest that, even if electrochemical graphitic oxide is formed in base, it is unstable because a 1 M NaOH solution apparently removed an electrochemical graphitic oxide film formed in an acidic solution. Variation of Oxidation and Reduction Reactions as a Function of pH. A technique that apparently has not been used previously in studies examining pH effects on the pretreatment of carbon electrodes is to keep the electrode under potential control while the pH is changed. This technique would allow the oxidation of an electrode in a solution of one pH and the reduction in a solution of a different pH. Furthermore, an electrode could be oxidizied multiple times at d ~ e r e n pHs t prior to reduction. Because phosphate buffers have been used widely as the solutions for electrochemical pretreatment? we made our pH changes by neutralizing 0.1 M NaOH with 1M H3P04. Even though the formation of intercalation (lamellar) compounds probably does not occur in glassy ~ a r b o n ,the ~ , ~choice of phosphate solutions further removes this process as a complication. Alsmeyer and McCreery12have shown no intercalation formation even in highly oriented pyrolytic graphite with 1 M H3P04. A potential of +0.4 V was selected as the potential for making pH changes because no apparent redox reactions take place at this potential. This potential of +0.4 V is also the point in the voltammetric cycle where the solution was bubbbled with nitrogen for 5 min to remove dissolved oxygen produced during the anodization. In order to interpret the voltammetric curves obtained when the pH is changed, it is necessary first to examine the voltammetric current-voltage curves for the cases where the pH is not changed during the course of the pretreatment voltammetric cycle. Figures 2 and 3 show the voltammetric current-voltage curves for a pH 12 solution and a pH 2 solution, respectively. Voltammetric current-voltage curves were also obtained for pH 4,pH (8) Cabaniss, G. E.; Diamantis, A. A.; Murphy, W. R., Jr.; Linton. R W.; Meyer,
T. J.J. Am. Chem. SOC.1985,107,1845-1853. (9) Engstrom, R C.; Strasser, V. A. Anal. Chem. 1984,56, 136-141. (10) Rojo, A.; Rosenstratten, A.; Anjo, D. Anal. Chem. 1986,58, 2988-2991. (11) Bjelica, L.; Parsons, R; Reeves, R. M. Croat. Chem. Acta 1980,53, 211231. (12) Alsmeyer, D. C.; McCreery, R L. Anal. Chem. 1992,64, 1528-1533. 978 Analytical Chemistry, Vol. 67, No. 5, March 1, 1995
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Figure 2. Cyclic voltammogram for a glassy carbon electrode in pH 12 solution; 15 s pause at the positive potential limit of +2.1 V.
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Figure 3. Cyclic voltammogram for a glassy carbon electrode in pH 2 solution; 15 s pause at the positive potential limit of $2.1 V.
8, and pH 10 solutions. In general, these curves appear very similar to the pH 2 curves, with the exception of the shiftof the
cathodic peak potential with respect to the pH of the solution. This pH-dependent shift is shown in Figure 4,where the potential of the cathodic peak is plotted versus pH. This shift is not a linear function, suggesting that Merent species are produced at different pHs or that the nature of the electrochemical graphitic oxide film changes with the pH at the time of formation. Our voltammetric curves for pH 2 solutions are similar to the curves shown by other gr0ups~3-~5 for several types of carbon electrodes in 0.5 and 1 M H2S04. Their reported cathodic peak potentials, which are (13) Barbero, C.; Silber, J. J.; Sereno, L. J. Electroanal. Chem. 1992,248, 321340. (14) Epstein, B. D.; DalleMolle, E.; Mattson, J. S. Carbon 1971,9, 609-615. (15) Proctor, A.; Shenvood, P.M. A. Carbon 1983,21,53-59.
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Figure 4. Cathodic peak potential as a function of solution pH; 15 s pause at the positive potential limit of +2.1 V.
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0 0.8 1.6 Potential, Volts vs SCE Figure 6. Cyclic voltammogram for a glassy carbon electrode anodized in both pH 12 and pH 2 solution before electrode cathodized in pH 2 solution. pH 12 solution (- - -); pH 2 solution (-); 15 s pauses at the positive potential limit of f 2 . 1 V; during second scan, pH of solution changed from 12 to 2 at +0.4 V and scanned back to positive potential potential limit for anodization in pH 2 solution. -0.8
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0 0.8 1.6 Potential,Volts vs SCE Figure 5. Cyclic voltammogram for a glassy carbon electrode anodized in a pH 12 solution and reduced in pH 2 solution: pH 12 solution (- - -); pH 2 solution (-); 15 s pauses at the positive potential limit of +2.1 V; during second scan, pH of solution changed from 12 to 2 at +0.4 V prior to N2 bubbling. -0.8
between -0.1 and 0 V, are close to the extrapolated potential for pH 0 in Figure 4. Another observation from Figures 2 and 3 is the difference in the anodic curves after the first cycle is completed. At both pH's a clearly d e h e d peak or shoulder at approximately+1.6 Vis seen after the first cycle. This peak was also seen for several types of carbon electrodes in 0.5 and 1M HzS04 solutions and 1 M HClOd ~ o l u t i o n s . ~ ~ JSince 3 - ~ ~ the peak does not appear to shift a p preciably with pH, H+ must not be involved in the reaction. A possible explanation for the anodic processes appearing at lower potentials after the first cycle is that the electrode is now activated for these processes. The additional peak around + L O V which is unique to the pH 12 solution, supports the idea of a difference between the reactions in pH 12 and those in less than pH 12. The first case to examine where the pH is changed during a scan is the experiment in which anodization is done in a solution of one pH followed by cathodization at a different pH. Figure 5 shows the voltammetric current-voltage curve obtained when the anodization is done at pH 12 and the pH of the solution is changed
to 2 at f 0 . 4 V before scanning to the negative potential limit. For comparison purposes the first cycle was done entirely in pH 12 solution. The main observation from Figure 5 is that the cathodic peak has shifted to -0.64 V in pH 2 from -0.96 V in pH 12, whereas the cathodic peak for a voltammetric scan at a constant pH of 2 is at -0.54 V (Figure 3). Although the heights of the peaks varied when the experiment was repeated, the peak potentials were constant within a few millivolts. The 32 mV/pH unit shift suggests that the potential shift represents the change in potential seen when H' is involved in the redox reaction. Taking into account the potential shift with pH, there is still a 0.1 V difference between the pH 2 cathodic peaks in Figures 3 and 5. This difference in peak potential suggests that different species may be formed when the anodizations are done at different pHs. As a control, an additional cycle was done in the now pH 2 solution (data not shown); the voltammetric curve matched well with the curve in Figure 3, indicating that the effects of anodization in pH 12 were removed from the electrode during the 15 s anodization in the pH 2 solution. Another peak to consider is the broad anodic peak that appears on the return scan after the primary reductions. This broad peak with an approximate peak potential of f0.3 V is seen in Figure 3 for a pH 2 solution and also at the same potential in Figure 5 for the case where the primary oxidation is done in a pH 12 solution and the reduction is done in a pH 2 solution. A broad peak centered around -0.3 V is also visible in Figure 5 for the initial cycle done entirely in a pH 12 solution. Cyclic voltammograms for pH 4 and pH 8 solutions also showed this broad peak, with peak potentials of approximately f0.15 and -0.1 V, respectively. A plot of these potentials versus pH was reasonably linear with a Analytical Chemistry, Vol. 67, No. 5, March 1, 1995
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Figure 7. Cathodic peak potential for a pH 2 solution as a function of time of pause at the positive potential limit of +2.1 V.
slope of approximately 60 mV/pH unit. The voltammograms shown by other g r o ~ p s ~for ~ J1~MJ HClO4 ~ and 1M HzS04 have similar peaks with peak potentials of approximately +0.4 V, a value that fits reasonably well on our plot. These data suggest two conclusions. First, the slope of approximately 60 mV/pH unit suggests an oxidation reaction involving one H+. Second, because the peak appears in the same potential region for a cycle done entirely in a pH 2 solution (Figure 3) and for a cycle done in both pH 12 and pH 2 solutions (Figure 5), it is probable that the same species is produced on the electrode surface regardless of the pH of the solution during the primary redox processes. The second case to examine is the case where anodization is done in two different pH solutions prior to cathodization. Figure 6 shows the voltammetric current-voltage curve obtained when this double anodization is performed; Le., after the pH has been changed from 12 to 2, the electrode is reanodized in the now pH 2 solution before scanning to the negative potential limit. Again, for comparison purposes, the first cycle was done entirely in pH 12 solution. After reanodizing at pH 2, the scan to the negative potential limit shows a new peak with a peak potential of -0.8 V and with a shoulder around -0.65 V. Superimposing Figure 6 on Figure 5 suggests that the shoulder represents the peak in Figure 5, Le., the reduction of a species produced in pH 12 solution but reduced in pH 2 solution. The remaining question is whether or not the peak at -0.8 V represents a new species created by the reaction of the species produced on the electrode in pH 12 solution with the species produced on the electrode in pH 2 solution. Again, as a control, an additional cycle was done in the now pH 2 solution (data not shown) giving a voltammetric curve that matched well with the curve shown in Figure 3. Similar results were obtained when going from a pH 12 solution to a pH 4 solution except that the normal cathodic peak from anodization in pH 4 was evident as a shoulder on the new cathodic peak corresponding to the Figure 6 cathodic peak, along with a shoulder representing the pH potential shift of the pH 12 species. Effects of the Length of T i e of Anodization. During the initial phase of this study, we observed that the peak potential for the cathodic process varied with the length of time of anodization for all solutions, regardless of the pH. The variation in peak potential for a pH 2 solution versus pause time at the positive potential limit is shown in Figure 7. It should be noted that zero time on the graph does not represent zero time for reaction since (16) Thorp, H. H. J. Chem. Educ. 1992,69, 250-252. (17)Allred, C. D.; McCreery, R L. Anal. Chem. 1992,64, 444-448.
980 Analytical Chemistry, Vol. 67, No. 5, March 7, 1995
the reaction is taking place from the potential where the reaction starts to the potential limit and then back to the reaction beginning potential. Hence, scan rate would also affect observed potentials because of the effect of scan rate on the length of anodization time. The peak potentials are reproducible for a given pause time and remain constant regardless of the size of the peak. The variation of peak potential with anodization time was first reported by Engstrom5for extended times of anodization. The negative potential shift with increased anodization time was attributed to the additional time required to electrolyze the product when more product is present. An alternative explanation for the much shorter times of anodization used in our study is that the potential shifts represent the formation of different species within the electrochemical graphitic oxide film as it is being created; i.e., there is a continuous change in the structure of the film as the layers of oxide film form. Results similar to ours were shown by Bjelica et for 20 s anodizations at +2.0 V versus cycling with a +2.0 V limit. The potential shift with respect to length of anodization time points out the problem of comparing results from different studies when anodization times and scan rates differ. The time effect may also explain why ac and dc pretreatment procedures may activate electrodes d~erently. CONCLUSIONS The results obtained in this study provide additional data to r e h e the McCreery2 model of glassy carbon surface effects during electrochemical pretreatment. The first addition to the model is that electrochemical graphitic oxide can be formed in basic solutions, i.e., in solutions of a pH greater than 7; however, the graphitic oxide is unstable in highly basic solutions. The second addition is that the lower potential anodic process in highly basic solutions, reported by Beilby and Carlsson? is a different process which apparently does not involve the formation of eletrochemical graphitic oxide. Both potential regions should be considered when electrochemical pretreatment in basic solution for the activation of carbon electrodes is examined. Our studies reinforce the importance of controlling all variables closely when examining the nature of electrochemical pretreatment of carbon electrodes. For redox reactions taking place at the electrode surface, where the nature of the surface oxides may be imp~rtant,'~J~ the fine tuning of the anodization processes by changing the pH during the pretreatment process may provide clues to help understand the nature of electron transfer at carbon electrodes. Changing the pH while keeping the electrode under potential control would also be a useful technique for studying the pH effects of the surface complexes that have been formed by the primary pretreatment processes. ACKNOWLEWMENT This work was supported in part by a grant from the Howard Hughes Medical Institute to Pomona College. Portions of this paper were presented at the 1993 Paci6c Conference on Chemistry and Spectroscopy, Pasadena, CA, October 1993. Received for review December 15, 1994.e
October
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Accepted
AC9410326 @
Abstract published in Advance ACS Abstracts, January 15, 1995.
