Microscopic Measurement of pH with Iridium Oxide Microelectrodes

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Anal. Chem. 2000, 72, 4921-4927

Microscopic Measurement of pH with Iridium Oxide Microelectrodes David O. Wipf* and Fuyun Ge†

Department of Chemistry, Mississippi State University, Mississippi State, Mississippi 39762 Thomas W. Spaine‡ and John E. Baur*

Department of Chemistry, Illinois State University, Normal, Illinois 61790-4160

Microscopic pH electrodes were produced by deposition of hydrous iridium oxide onto carbon fiber microelectrodes. The electrodes exhibit two linear regions of potentiometric response between pH 2-6 and pH 6-12. The electrodes respond to pH changes within 50 ms, and an equilibrium value is reached within 30 s. By using these electrodes as probes in the scanning electrochemical microscope, dynamic pH changes occurring at or near a surface can be measured and pH maps of the surface can be generated. Vertical pH profiles and images of pH were obtained at substrates where electrochemical (oxidation and reduction of H2O2, hydrogen evolution) or enzymatic (glucose oxidase) reactions involving proton transfers occur. pH plays an important role in many interfacial reactions. The rates of important electrochemical reactions such as corrosion and metal deposition are greatly dependent on proton concentration. Electrochemical reactions often produce or consume protons, thus altering the pH near the electrode interface. Bioanalytical techniques based on reactions of immobilized enzymes, such as ELISA, enzyme electrodes, enzyme-based optical sensors, etc., can also involve the production or consumption of protons at the liquid-solid interface. Microorganisms and individual cells can perturb their local pH during respiration, and pH changes have recently been observed concurrently with the release of neurotransmitters.1,2 In each of these cases, the pH at or near the interface affects the rate of the reaction, influences reaction mechanism, and/or provides information about the process itself. Although pH is readily measured in the bulk of solution, measurement of proton concentration with high spatial resolution at or near the interface can further our understanding of the pH control of these processes. * Corresponding authors. D.O.W.: (e-mail) [email protected]; (fax) (662) 325-1618. J.E.B.: (e-mail) [email protected]; (fax) (309) 438-5538. † Present address: DBS Communications, Inc., 126 W. Center Court, Schaumburg, IL 60195. ‡ Present address: Department of Chemistry, Iowa State University, Ames, IA 50011. (1) Chen, J. C.; Chesler, M. Proc. Natl. Acad. Sci. U.S.A. 1992, 89, 77867790. (2) Michael, D.; Travis, E. R.; Wightman, R. M. Anal. Chem. 1998, 70, 586A592A. 10.1021/ac000383j CCC: $19.00 Published on Web 09/13/2000

© 2000 American Chemical Society

The development of the scanning electrochemical microscope (SECM) has enabled the imaging of chemical species at or near a variety of substrates.3 In this technique, a microscopic electrode (the probe) is scanned across a substrate, and the electrochemical response is plotted as a function of position. In this manner, a chemical and/or topographical map of the surface can be obtained, providing the sensor has a rapid temporal response compared to the scanning velocity. In the generator-collector (GC) mode, the image obtained with the SECM is dependent upon the selectivity of the probe. For example, if the probe is a pH microsensor, the pH of the solution close to the interface can be measured and plotted as a function of tip position. There have been previous reports of spatially resolved pH measurements made with microscopic pH sensors. Bard’s group used an antimony electrode to measure concentration profiles and to image enzymatic activity using the SECM.4 Lewandowski et al.5 and Tanabe et al.6 used oxidized iridium metal microelectrodes to investigate the effect of pH on corrosion at steel surfaces. More recently, local pH near surfaces has been measured using an Sb/ Sb2O3 electrode,7 neutral carrier-based liquid membrane electrodes,8,9 and a two-dimensional semiconductor pH imaging sensor.10 Of the solid-state metal oxides used for pH sensing, the most widely used has been iridium oxide.11 These oxide electrodes are normally prepared by the electrochemical oxidation of iridium metal electrodes,11-13 but they can also be prepared by thermal (3) Bard, A. J.; Fan, F.-R.; Mirkin, M. V. In Electroanalytical Chemistry; Bard, A. J., Ed.; Marcel Dekker: New York, 1993; pp 244-370. (4) Horrocks, B. R.; Mirkin, M. V.; Pierce, D. T.; Bard, A. J.; Nagy, G.; Toth, K. Anal. Chem. 1993, 65, 1213-1224. (5) Lewandowski, Z.; Funk, T.; Roe, F.; Little, B. J. Microbiologically Influenced Corrosion Testing; Kearns, J. R., Little, B. J., Eds.; ASTM STP 1232; American Society for Testing and Materials: Philadelphia, 1994; pp 61-69. (6) Tanabe, H.; Togashi, K.; Misawa, T.; Kamachi Mudali, U. J. Mater. Sci. Lett. 1998, 17, 551-553. (7) Honda, T.; Murase, K.; Hirato, T.; Awakura, Y. J. Appl. Electrochem. 1998, 28, 617-622. (8) Klusmann, E.; Schultze, J. W. Electrochim. Acta 1997, 42, 3123-3134. (9) Park, J. O.; Paik, C. H.; Alkire, R. C. J. Electrochem. Soc. 1996, 143, L174L176. (10) Nomura, S.; Nakao, M.; Nakanishi, T.; Takamatsu, S.; Tomita, K. Anal. Chem. 1997, 69, 977-981. (11) Glab, S.; Hulanicki, A.; Edwall, G.; Ingman, F. Crit. Rev. Anal. Chem. 1989, 21, 29-47. (12) Mozota, J.; Conway, B. E. Electrochim. Acta 1983, 28, 1-8. (13) Conway, B. E.; Mozota, J. Electrochim. Acta 1983, 28, 9-16.

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decomposition,14,15 by sputter deposition,16-20 or by mixing the iridium oxide in an inert matrix.21 More recently, we have reported a new method for preparing pH-sensitive iridium oxide electrodes by electrodeposition of the oxide from alkaline solutions of iridium(III) oxide onto carbon substrates22 These findings suggest a method for preparation of an iridium oxide ultramicroelectrode by coating the exposed end of a carbon fiber ultramicroelectrode. In contrast, other ultramicroelectrode construction methods start with iridium wire. However, iridium wire is brittle and is difficult to fashion into ultramicroelectrodes. In addition, iridium microwires are extremely costly and are not commercially available with diameters of less than 50 µm. Smaller diameters are prepared by etching, which is wasteful and adds an additional preparation step. In contrast, electrochemical deposition permits the formation of an iridium oxide electrode of essentially any shape or size, depending on the substrate. Using iridium oxide-coated carbon fiber microelectrodes has the benefit of employing well-characterized, inexpensive carbon fiber electrodes as a substrate. The coated electrodes also use much less costly iridium than all-metal microelectrodes. In this work, we describe the fabrication and characterization of carbon fiber-based iridium oxide (CF/IrOx) microelectrodes. Additionally, we demonstrate that these electrodes can be used as SECM probes to examine the dynamics of pH changes near surfaces that produce or consume protons. EXPERIMENTAL SECTION Electrodes. Carbon fiber electrodes were prepared by sealing individual carbon fibers (radius, 5 µm) into glass capillaries with epoxy and polishing at a 45° angle using a micropipet beveller.23 The overall probe diameter, including glass and epoxy insulator, was ∼40 µm. Prior to modification with iridium oxide, the carbon fiber electrodes were rinsed in hot toluene. The iridium oxide deposition solution was prepared as described previously.22 Briefly, a solution of 27 mM Na3IrCl6 in 0.1 M HCl was boiled until the solution color changed from olive green to light brown. After sparging with nitrogen, this solution was made basic by addition of 6 M NaOH until the solution pH was ∼12.5. Iridium oxide was deposited onto the electrode surface by the application of a constant potential of between 0.6 and 0.7 V vs Ag/AgCl for 2 min. The presence of the surface iridium oxide was verified by recording background voltammograms in 0.5 M H2SO4 or 0.1 M Na2CO3. All potentials were recorded with respect to a 3 M Ag/ AgCl reference electrode. Instrumentation. A voltammetric analyzer (model BAS 100B, Bioanalytical Systems, West Lafayette, IN) was used for electrodeposition and cyclic voltammetry during the preparation and (14) Ardizzone, S.; Carugati, A.; Trasatti, S. J. Electroanal. Chem. 1981, 126, 287-292. (15) Lin, S. M.; Wen, T. C. Electrochim. Acta 1994, 39, 393-400. (16) Katsube, T.; Lauks, I.; Zemel, J. N. Sens. Actuators 1982, 2, 399-410. (17) Kreider, K. G.; Tarlov, M. J.; Cline, J. P. Sens. Actuators B 1995, B28, 167172. (18) Bardin, M.; Loheac, P.; Petit, M.; Plichon, V.; Richard, N. New J. Chem. 1995, 19, 59-63. (19) Kato, A.; Konno, Y.; Yanagida, Y.; Yamasoto, M.; Taguchi, T.; Motohashi, R.; Katsube, T. Anal. Sci. 1991, 7, 1577-1580. (20) Kreider, K. Sens. Actuators B 1991, B5, 165-169. (21) Fog, A.; Buck, R. P. Sens. Actuators 1984, 5, 137-146. (22) Spaine, T. W.; Baur, J. E. J. Electroanal. Chem. 1998, 443, 208-216. (23) Kelly, R. S.; Wightman, R. M. Anal. Chim. Acta 1986, 187, 79-87.

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testing phase of the electrodes. The potentiometric pH response of the CF/IrOx electrodes was recorded using a high-impedance electrometer (model 6512, Keithley Instruments, Inc., Cleveland, OH). The potential response was measured by either immersing the electrode in a series of standard pH buffers or by comparison of the potential response to a commercial glass pH electrode during an automated titration of H3PO4 with NaOH.22 The SECM was of local construction and employed a bipotentiostat (model EI-400, Cypress Systems, Lawrence, KS) for voltammetric control of the substrate.24 Probe placement was controlled by a three-axis piezoelectric positioner (Burleigh Instruments Inc., Fishers, NY) interfaced to a personal computer. The output of the electrometer was recorded as a function of the iridium oxide probe position, and the data were later converted to pH by use of the calibration curve. To simultaneously perform voltammetry at the substrate electrode and potentiometry at the CF/IrOx electrode, the electrometer was electrically isolated (“floated”) from the potentiostat ground. This required use of a galvanic isolation amplifier (ISO212J, Burr Brown Corp., Tucson, AZ) placed between the electrometer and the data acquisition card of the computer. For SECM experiments, the CF/IrOx electrode tip was positioned ∼10 µm above the test sample with the aid of a video microscope (model K2 Infinity Photo-Optics, Boulder, CO). The exact distance between the tip and sample was set by slowly (0.1 µm/s) moving the tip toward the substrate. Contact was noted by a distinct change in tip signal. Setting the contact distance as zero, the tip was moved away to set the desired tip-sample distance. Reagents. The iridium salts used for preparation of the deposition solutions were obtained from Aldrich (Milwaukee, WI). Biotin-labeled glucose oxidase, 150 units/mg of protein, 5.0 mol of biotin/mol of protein (Sigma, St. Louis, MO), ExtrAvidin (Sigma), sulfosuccinimidyl-6-(biotinamido) hexanoate (NHS-LCbiotin, Immuno Pure, Pierce Chemical Co., Rockford, IL), β-Dglucose (ICN, Costa Mesa, CA), 4-nitrobenzenediazonium tetrafluoroborate (97%, Aldrich Chemical Co.), and tetrabutylammonium hexafluorophosphate (Bu4NPF6, SAChem Inc., Austin TX) were used as received. All other chemicals used here are reagent grade. Ferrocenium cation, FeCp2+, solutions were prepared by oxidation of ferrocene with H2O2. Enzyme Immobilization. Enzyme immobilization used biotin-avidin coupling.25-27 The carbon was first modified by covalent attachment of a monolayer of 4-nitrophenyl. Attachment occurred after several cyclic voltammetric scans from 0.4 to 0 V at a scan rate of 200 mV/s in a solution containing 4-nitrobenzenediazonium tetrafluoroborate in 0.1M Bu4NPF6 CH3CN solution.28 Reduction of the 4-nitrophenyl to 4-aminophenyl was accomplished by several cyclic voltammetric scans from 0.0 to -1.2 V (vs Ag/AgCl ref) in pH 4.3 acetate buffer. After production of the 4-aminophenyl-modified carbon electrode surface, the modified electrode was rinsed with phosphate buffer solution (PBS) and then soaked for 2 h in room-temperature PBS (pH 7.4) (24) Wipf, D. O. Colloids Surf. A 1994, 93, 251-261. (25) Luo, S.; Walt, D. R. Anal. Chem. 1989, 61, 1069-1072. (26) Pantano, P.; Morton, T. H.; Kuhr, W. G. J. Am. Chem. Soc. 1991, 113, 18231833. (27) Pantano, P.; G., K. W. Anal. Chem. 1993, 65, 623-630. (28) Allongue, P.; Delamar, M.; Desbat, B.; Fagebaume, O.; Hitmi, R.; Pinson, J.; Saveant, J. M. J. Am. Chem. Soc. 1997, 119, 201-207.

containing 5 mg/mL sulfo-NHS-LC-biotin. After being rinsed again with PBS, the electrode was soaked in PBS buffer containing 2 mg/mL ExtrAvidin. After formation of a biotin-avidin complex, the electrode was rinsed with PBS and then soaked for 24 h in a 4 °C pH 6.1 PBS containing 1 mg/mL biotinylated glucose oxidase to produce the immobilized surface. RESULTS AND DISCUSSION Characterization of the CF/IrOx Electrode. Hydrous iridium oxide can be electrodeposited onto glassy carbon electrodes of conventional size (radii of several millimeters) by repeated potential cycling between -0.3 and +1.0 V vs Ag/AgCl.22 In this method, a basic, saturated solution of hydrous iridium(III) oxide is first prepared from solutions of IrCl63- or IrCl62- according to the following reactions: 0.1 M HCl

IrCl63- 9 8 Ir(OH2)2Cl4- or 100 °C, 2-3 h 0.1 M HCl, EtOH

IrCl62- 9 8 Ir(OH2)2Cl4- (1) 100 °C, 2-3 h NaOH

Ir(OH2)2Cl4- 98 Ir2O3‚xH2O

(2)

Hydrous iridium(IV) oxide is then anodically deposited onto the carbon surface according to the following reaction:

Ir2O3‚xH2O (aq) + 2OH- f 2IrO2‚xH2O(s) + H2O + 2e- (3)

The resulting oxide film is insoluble but undergoes a reversible redox reaction in the solid state, the potential of which is dependent on the solution pH (see below). Initial experiments using carbon fiber microelectrodes as the deposition substrate, however, showed that this method causes large, dendritic deposits on the carbon fiber surface. In extreme cases, the crystalline entity completely envelops the tip (including the glass insulator) of the microelectrode. Because a small, welldefined probe is required for high-resolution SECM imaging, an alternative deposition procedure was developed. We have found that even, controllable coatings of iridium oxide on carbon fibers can be prepared by applying a constant potential just at the onset of oxygen evolution. The deposition potential is usually in the range of +0.5 to +0.7 V and depends on the pH of the deposition solution and the individual carbon fiber. The exact value is found by recording the cyclic voltammogram of the deposition solution with each carbon fiber and reversing the scan direction at the onset of the large oxygen evolution wave. A constant potential is then applied for a time of 1-2 min, depending on the desired oxide thickness. We have found that coatings formed by depositions of less than 1 min are prone to loss, and therefore, deposition times of 2 min were used for the majority of this work. Figure 1 shows cyclic voltammograms at a CF/IrOx electrode in 0.5 M H2SO4 and in 0.1 M Na2CO3. Peaks characteristic of surface-confined species are observed in both the acidic and basic solutions. The very small peak separation of the waves suggests that the charge-transfer reaction is apparently much faster at the microelectrode than previously observed at larger iridium oxidemodified glassy carbon electrodes.22 The thickness of the oxide

Figure 1. Cyclic voltammograms of a CF/IrOx electrode recorded at 0.1 V s-1 in (A) 0.5 M H2SO4 and (B) 0.1 M Na2CO3 at a CF/IrOx microelectrode.

coating is estimated by integrating the voltammograms and using the surface density for hydrous iridium oxide of 7.8 × 10-7 mol cm-2 µm-1 reported by Pickup and Birss.29 The oxide layer thickness calculated from the data in Figure 1 is 0.05 µm, assuming an effective electrode radius of 5.9 µm (5-µm fiber beveled at 45°). Each of the electrodes used in this work had a similar oxide layer thickness. The coating on the carbon fiber is thinner than previously obtained at a 3-mm-diameter glassy carbon electrode (0.2 µm), which likely accounts for the more rapid redox kinetics observed at the microelectrode compared to a similarly modified glassy carbon electrode.22 Presumably hydroxide ions are transported more rapidly through the thinner oxide on the microelectrode, producing the redox behavior expected for a surface-bound species. Iridium oxide electrodes are known to exhibit super-Nernstian pH response, varying from 59 to ∼80 mV per decade, with the exact response depending on the preparation method.30 The potential of the CF/IrOx electrodes was measured as a function of the reported pH at a commercial glass pH electrode during a titration of phosphoric acid with sodium hydroxide. Table 1 summarizes the results of four different electrodes prepared with the same deposition solution. The data in this table are separated (29) Pickup, P. G.; Birss, V. I. J. Electroanal. Chem. 1987, 220, 83-100. (30) Glab, S.; Hulanicki, A.; Edwall, G.; Ingman, F. Crit. Rev. Anal. Chem. 1989, 21, 29-47.

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Table 1. pH Response of CF/IrOx Electrodes Determined by Titration of H3PO4 with NaOHa electrode

pH range

slope (mV/decade)

intercept (mV)

A

entire range 6 entire range 6 entire range 6 entire range 6

-73.6 (0.99) -58.5 (3.8) -80.5 (2.0) -82.2 (1.5) -79.2 (3.7) -86.6 (2.4) -82.2 (1.9) -77.8 (4.6) -87.7 (3.5) -80.2 (1.2) -78.9 (4.8) -83.4 (2.7)

1097 (25.0) 1040 (35.8) 1165 (16.3) 999 (110) 984 (119) 1042 (101) 972 (59.2) 951 (56.1) 1024 (51.0) 1065 (83.0) 1057 (81.2) 1095 (71.5)

B C D

a Values are averages of five titrations, and numbers in parentheses are standard deviations. The intercept is at pH 0.

into two pH regions because hydrous iridium oxide electrodes are known to show two different linear regions of pH response with the transition point around pH 6.29,31 This behavior has been attributed to two different mechanisms.31 The first (cf. Figure 1A), predominating below pH 6, is

Ir(OH)4(r-4)- + 1.1H+ + e- f Ir(OH)q(q-3)- + 1.1H2O (4) And the second, predominating above pH 6 (cf. Figure 1B), is

Ir(OH)4(r-4)- + e- f Ir(OH)q(q-3)- + 1.4OH-

(5)

The Nernst equations (at 25 °C) for these half-reactions can be written as

ECF/IrOx ) E0 - x0.0592pH

(6)

ECF/IrOx ) E0 - y0.0592pH

(7)

and

where x and y are the stoichiometric coefficients on the H+ and OH- in reactions 4 and 5, respectively. The values of x and y can be calculated from the slopes of the calibration curves (Table 1) over the appropriate pH range. The coefficient on the H+ in reaction 4 and the OH- in reaction 5 are those found for iridium oxide containing some palladium deposited onto a glassy carbon electrode.31 Analysis of the average slopes in Table 1 gives values of 1.25 and 1.43 for the coefficients on H+ and OH- in eqs 4 and 5, respectively, for the CF/IrOx electrodes. The variability in the potentiometric response seen in Table 1 is an indication of how reproducibly we could produce these electrodes. Although these slopes are within the range seen in the literature,30 calibration of individual electrodes is essential for precise results. Despite the slightly different slopes seen in the pH ranges of 2-6 and 6-12, calibration curves acquired over the pH range of 2-12 were linear (r2 of ∼0.99). Our previous work with iridium oxide deposited onto glassy carbon electrodes has shown a slow potentiometric response to 4924

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Figure 2. Temporal response of a CF/IrOx electrode to a change in solution pH accomplished by adding a 0.1 M phosphate buffer (pH 4.0) to a 0.5 mM phosphate buffer (pH 6.2). Inset: enlargement of the response immediately after the pH change. The potential was measured at 25-ms intervals.

pH near the end points of the titration. This is thought to be a result of the rate-limiting hydroxide diffusion through the thick iridium oxide films on these electrodes. In contrast, the potential response of the CF/IrOx electrodes as a function of pH is linear near the equivalence point, indicating that the response time of the CF/IrOx electrode is comparable to that of the glass electrode. Figure 2 shows further evidence that the response of the CF/ IrOx electrodes is very rapid. This figure shows the pH measured by a CF/IrOx electrode as a function of time following the addition of a pH 4.0 phosphate buffer (0.1 M phosphate) to a dilute pH 6.2 phosphate buffer (0.5 mM phosphate). The CF/IrOx electrode shows an immediate large change in potential and a constant value within 30 s. The inset of Figure 2 shows that most of the potential change occurs during the first 50 ms after the addition of the buffer. This temporal response is fast enough to image pH in solution and to follow pH changes associated with cyclic voltammetry. CF/IrOx electrodes were stable whether stored in air or buffer solution during the 1-2-month period of this study. The electrodes could be renewed by a pretreatment consisting of several 1 V/s CVs in 0.5 M H2SO4 at potentials between hydrogen and oxygen evolution followed by a 2-min polarization at +200 mV (vs Ag/ AgCl) in pH 7 buffer. A more in-depth investigation of the stability of carbon-substrate IrOx electrodes has been previously published.22 Measurement of pH Changes During Cyclic Voltammetry. A CF/IrOx electrode was placed within 10 µm of a 300-µmdiameter Pt electrode substrate immersed in 10 mM H2O2 in unbuffered 0.5 M KCl using the SECM. As a cyclic voltammogram of H2O2 is recorded at the substrate, the CF/IrOx electrode is able to follow changes in pH very close to the electrochemical interface. Protons are consumed in the reduction of H2O2 according to the reaction

H2O2+ 2H+ + 2e- f 2H2O

(8)

The result is an increase in solution pH near the substrate electrode as H+ is depleted. Conversely, the oxidation of H2O2

Figure 3. (A) Cyclic voltammogram of the 300-µm-diameter Pt substrate electrode immersed in 10 mM H2O2 in 0.5 M KCl (v ) 0.020 V s-1). (B) Potentiometric response as a function of substrate potential for a CF/IrOx electrode positioned within 10 µm of the 300-µm Pt substrate.

generates protons:

H2O2 f O2 + 2H+ + 2e-

(9)

resulting in a local decrease in solution pH near the substrate electrode as H+ is liberated. Figure 3 shows the voltammogram of H2O2 (part A) and the corresponding potential and pH recorded at the CF/IrOx electrode (part B). Initially (from about -0.6 to -0.1 V), the H2O2 is being reduced and therefore the local solution is basic (pH ∼8.5). As the potential of the substrate is made more positive, the consumption of protons ceases and the pH near the substrate becomes neutral with only slight changes. As the oxidation of H2O2 becomes favorable at a substrate potential of ∼0.5 V, a rapid decrease in pH is observed, corresponding to the production of H+. The local pH remains nearly constant until an applied potential of ∼+1.1 V, where water is oxidized, and an extremely large amount of H+ is generated. In this substrate potential region, the CF/IrOx electrode indicates that the pH becomes large and negative. Since large, negative pH values are unwarranted, it is likely that the response of the CF/IrOx electrode becomes nonlinear in very acidic solutions. The presence of large amounts of O2 and Cl2 from the oxidation of water and Cl- is also likely a cause of the nonlinearity, as iridium oxide electrodes are known to exhibit a potentiometric response to changes in O2. This fact, in conjunction with the slow dissolution of electrodeposited iridium oxide from carbon in strong acids,22 discourages use of the CF/IrOx elec-

trodes in strongly acidic solutions.During the cathodic scan, the response of the CF/IrOx electrode does not closely match that of the anodic scan until the potential becomes less than -0.1 V, where H+ is again being consumed. This indicates that the large amount of H+ generated upon oxidation of the peroxide and water lingers near the electrode surface. Note that the presence of the CF/IrOx tip will hinder diffusion of H+ away from the electrode. Upon consumption of H+ by the reduction of H2O2 (and possibly direct reduction of the H+ to H2), the pH recorded by the CF/ IrOx electrode returns to that of the initial KCl solution at the initial potential. Concentration Profiles at Microelectrode Substrates. Because the CF/IrOx electrode is used as a potentiometric electrode, approach curves do not resemble those encountered when the feedback mode of the SECM is employed. However, if the substrate is creating a hydrogen ion gradient via the production or consumption of H+ at the surface, the pH should be a function of distance of the CF/IrOx probe from the substrate. To investigate this pH gradient, a 300-µm-diameter substrate electrode was immersed in a pH 6.0 phosphate buffer and the potential of the substrate was poised at increasingly negative potentials. As the potential of the substrate becomes more negative, protons in solution are reduced to H2, effectively lowering the concentration of protons near the surface and increasing the pH. Figure 4 shows the dependence of pH and H+ on distance of the CF/IrOx electrode from the Pt substrate electrode immersed in buffer for several different substrate potentials. At large distances (>250 µm), the solution pH is practically unaffected by the electrode reaction. When the substrate potential is set to 0 V, there is no voltammetric reaction and therefore the pH remains constant throughout the region close to the electrode. However, at negative applied potentials, the H+ near the electrode is depleted due to hydrogen evolution and the pH is observed to increase as the substrate surface is approached. This increase is most dramatic at the more negative applied potentials where the rate of H+ reduction is greatest. Figure 4B shows the H+ concentration profile calculated from the data in Figure 4A. At applied potentials between -0.8 and -1.2 V, a nearly linear relation between H+ concentration and distance exists. However, at applied potentials more negative than -1.2 V, deviation from linearity is observed. This deviation may arise from the fact that bubbles of H2 gas are generated at the substrate at these potentials. The bubble formation introduces a significant convection component that disrupts the diffusional profile. Imaging pH with the SECM. pH images of a 10-µm-diameter Pt electrode reducing H+ recorded with a CF/IrOx electrode are shown in Figure 5. These images show the distribution of pH near the substrate electrode immersed in unbuffered (0.5 M KCl, Figure 5A-E) and buffered (pH 6.1 phosphate, Figure 5F-J) solutions. The potential of the Pt substrate was varied from -0.2 (Figure 5A and F) to -1.2 V (Figure 5E and J). To obtain the images, a CF/IrOx electrode was lowered to within 5 µm of the electrode surface and then was rastered across the region near the Pt substrate. Scan rates of 2-10 µm/s were used to acquire these images. In part, the response time of the CF/IrOx electrode limits the fastest practical scan rates. For 1-µm resolution at a scan rate of 10 µm/s, the response of the electrode must settle to its final value within 0.1 s. As observed in Figure 2, the electrode response is marginal for tip scan rates of 10 µm/s. For best results, scan rates of less than 2.5 µm/s are recommended. The potenAnalytical Chemistry, Vol. 72, No. 20, October 15, 2000

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Figure 4. Effect of substrate potential on the approach curves for a CF/IrOx electrode near a 300-µm-diameter Pt electrode. Tip approach and withdraw speed is 5 µm/s. (A) Dependence of pH on probe-substrate distance at several substrate potentials. (B) Dependence of hydrogen ion activity (calculated from the data in part A) on probe-substrate distance at several substrate potentials. The electrolyte was a pH 6.0 phosphate buffer with no chloride or citric acid.

tiometric response was then converted to a pH using data from a precalibration. As the potential of the electrode becomes more negative, proton reduction occurs at an increased rate and therefore H+ near the electrode surface is depleted, causing an increase in local pH. When the applied potential is -0.2 V, the rate of proton reduction should be essentially zero. However, a small perturbation in the pH can be observed in the unbuffered solution (Figure 5A). This change, ∼0.07 pH unit, could be due to a small underpotential reduction of H+ and/or the electrochemical adsorption of protons to the platinum surface without subsequent hydrogen evolution:

Pt + H+ + e- f Pt-H

(10)

When the potential of the platinum substrate is adjusted to more negative values, changes in pH near the electrode surface are clearly visible. At an applied substrate potential of -0.4 V (Figure 5B), a change of 0.5 pH unit is observed within ∼30 µm of the Pt 4926 Analytical Chemistry, Vol. 72, No. 20, October 15, 2000

Figure 5. pH images of a 10-µm Pt substrate at various applied potentials in unbuffered 0.5 M KCl (A-E) and in pH 6.1 phosphate buffer (F-J). The substrate potential was -0.2 (A, F), -0.4 (B, G), -0.6 (C, H), -1.0 (D, I), and -1.2 V (E, J). Images were acquired at scan rates of between 2 and 10 µm/s at a tip-substrate distance of 5 µm.

substrate. As the potential is made more negative (Figure 5CE), the perturbation grows to well over 60 µm from the substrate electrode. Note that in Figure 5E, where the applied substrate potential is -1.2 V, the pH reaches 11 near the surface and drops off only to ∼8.5 at the edge of the frame. In buffered solution, the same experiments produce quite different results. Remarkably, pH changes of less than 0.02 pH unit can be visualized through applied potentials of -0.6 V even though appreciable noise is evident (Figure 5H). When the applied potential reaches -1.0 V, the pH perturbation is 0.08 unit, approximately that observed in the unbuffered solution at an applied potential of -0.2 V. Even at applied potentials of -1.2 V, the pH change is less than 0.6 unit, and the perturbation disappears at distances greater than ∼20 µm from the electrode center. Thus, the overall effect of the buffer is to restrain the pH changes to a smaller region and a smaller magnitude. Imaging pH at an Immobilized Enzyme. The CF/IrOx is also useful for studying reactions of immobilized enzymes. Many biological reactions are mediated by enzymes such as oxidases and dehydrogenases and involve a proton transfer. Although the result of the previous section suggests that production of small amounts of H+ can be masked by buffers, reactions occurring in

Figure 6. pH image of glucose oxidase immobilized onto a carbon microelectrode substrate (33-µm diameter). (A) Image obtained with FeCp2+ in solution and the substrate at open circuit. (B) Same as (A), but plotted on the same pH scale as (C) for comparison purposes. (C) Image obtained with the substrate biased at +0.6 V to regenerate FeCp2+ from FeCp2 produced by the enzymatic reaction. Images were acquired at scan rates of 5 µm/s at a tip-substrate distance of 10 µm.

unbuffered or weakly buffered solutions should result in a measurable change in local pH. In an attempt to visualize this perturbation, glucose oxidase (GOx) was immobilized onto a 33µm-diameter carbon using a biotin-avidin linkage. In the presence of ferrocenium ion (FeCp2+) as a redox mediator, the glucose oxidase converts glucose to gluconic acid and generates protons:

glucose + H2O + 2FeCp2+ f gluconic acid + 2FeCp2 + 2H+ (11)

Because the rate of the enzymatic reaction depends on pH, this reaction must be carried out in a buffer in order to maintain enzyme activity. However, by using a buffer with a low buffer capacity, the pH distribution at the modified electrode can be recorded. Figure 6A shows the pH image obtained at the GOx-modified carbon microelectrode in a solution containing 50 mM glucose and 1 mM FeCp2+ in a 0.5 mM phosphate buffer. This image clearly shows the production of protons by enzymatic reaction of the glucose to produce a decrease in pH of 0.4. For this image, the GOx-modified carbon microelectrode was left at open circuit. Under these conditions, the rate of the enzymatic reaction is limited by the concentration of FeCp2+ near the enzyme-modified electrode. In a second experiment, a saturated solution of FeCp2 was used as a mediator solution. The GOx-modified carbon electrode is biased to +0.6 V, producing FeCp2+ at the carbon surface. Although the concentration of FeCp2 in aqueous solution is very low, the pH signal in Figure 6C is significantly larger than that in Figure 6A. The biased electrode produces a small amount of FeCp2+ from the solution of FeCp2, but the predominant increase is due to a feedback mechanism operating between the immobilized GOx and the underlying carbon electrode. The FeCp2 produced during the enzymatic reaction is converted back to FeCp2+ at the electrode by the following reaction.

FeCp2 f FeCp2+ + e-

(12)

Under these conditions, therefore, the overall reaction is the

oxidation of glucose

glucose + H2O f gluconic acid + 2H+ + 2e-

(13)

catalyzed by the biased electrode. The effect of this redox catalysis is shown by the image in Figure 6C. The application of 0.6 V to the carbon substrate results in a much larger 0.8 pH unit change. This effect can be dramatized by comparing Figure 6C to Figure 6B. Figure 6B is the data from Figure 6A (no substrate potential applied) plotted on the same scale as the data in Figure 6C. CONCLUSIONS The experiments described in this paper demonstrate that the CF/IrOx electrode probe is well suited as a SECM probe for experiments in which changes of pH near a surface are to be measured. This type of electrode has the advantage over previous pH-sensitive electrodes used in SECM because it is prepared by electrodeposition of the oxide onto the surface of inexpensive carbon fiber electrodes. In this way, smaller pH sensors can be fabricated by further reduction of the carbon fiber substrate size. Thus, electrochemically32,33 or flame34 etched carbon fiber electrodes, carbon ring electrodes formed by pyrolysis,35 or other reduced-size carbon electrodes should be suitable substrates for smaller pH sensors that will be capable of higher resolution imaging. ACKNOWLEDGMENT This work was supported by an award from the Research Corp. (J.E.B). D.O.W. acknowledges support for this work by a grant from the National Science Foundation (CHE-94144101). Received for review April 3, 2000. Accepted July 27, 2000. AC000383J (31) Jaworski, R. K.; Cox, J. A.; Strohmeier, B. R. J. Electroanal. Chem. 1992, 325, 111-123. (32) Kawagoe, K. T.; Jankowski, J. A.; Wightman, R. M. Anal. Chem. 1991, 63, 1589-1594. (33) Schulte, A.; Chow, R. H. Anal. Chem. 1998, 70, 985-990. (34) Strein, T. G.; Ewing, A. G. Anal. Chem. 1992, 64, 1368-1373. (35) Kim, Y.-T.; Scarnulis, D. M.; Ewing, A. G. Anal. Chem. 1986, 58, 17821786.

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