Environ. Sci. Technol. 2000, 34, 3443-3451
Effects of pH and Applied Potential on Photocurrent and Oxidation Rate of Saline Solutions of Formic Acid in a Photoelectrocatalytic Reactor ROBERTO J. CANDAL,† WALTER A. ZELTNER, AND MARC A. ANDERSON* Water Chemistry Program, University of WisconsinsMadison, 660 North Park Street, Madison, Wisconsin 53706
A photoelectrocatalytic reactor containing titania-coated titanium electrodes was employed to degrade solutions of formic acid (2 mmol dm-3) in 0.01 mol dm-3 NaCl. Reaction rates were increased above that observed for a purely photocatalytic experiment by operating at applied potentials of at least +1.0 V (versus saturated calomel electrode). The kinetics of photodegradation at +1.0 V was modeled effectively using a Langmuir-Hinshelwood-Hougen-Watson expression. Unexpected results were obtained when only the background electrolyte was passed through the reactor. During initial recirculation of this solution with no UV illumination and no applied potential, the pH increased from 6.5 to 9, suggesting ion exchange of chloride ions with hydroxyl ions from the catalyst surface. However, when UV illumination was initiated with an applied potential, the pH decreased to 3.5-4.2, depending on the magnitude of the potential. The cause of this behavior is not known, although there are several explanations. Addition of formic acid to this system buffered the pH near 3, producing the highest rate of degradation at an applied potential of +1.0 V. When the formic acid test solution was adjusted to higher initial pH values, the reaction rate was unaffected until the pH increased above 5, at which point the rate decreased.
Introduction Numerous studies of photocatalytic oxidation of organic compounds and reduction of inorganic ions as a method of purifying water have been conducted for more than 10 years (1-3). Most of these studies employed UV-illuminated titania as the photocatalyst in an effort to develop a simple, inexpensive method of treating water. Recent attention has focused on improving the efficiency of the process in order to make it commercially applicable (4). Practical applications of this technology will likely employ supported photocatalysts (i.e., catalysts that are either coated on or incorporated in a robust substrate) rather than suspensions of the catalyst because the separation of suspended particles from the purified water stream may not be convenient (5). Unfortunately, the surface area of active * Corresponding author phone: (608)262-2674; fax: (608)262-0454; e-mail:
[email protected]. † Present address: INQUIMAE, Universidad de Buenos Aires, Ciudad Universitaria, Pabello´n 2, Buenos Aires, CP 1428, Argentina. 10.1021/es991024c CCC: $19.00 Published on Web 06/30/2000
2000 American Chemical Society
catalyst exposed to the solution is lower in supported systems than in photoreactors that employ suspended catalysts, reducing the catalytic activity. A second difficulty is ensuring good adhesion of photocatalytic films to whatever material is employed as the support. Several researchers have reported that the effectiveness of supported photocatalysts can be improved by applying a positive potential (“bias”) across the photoanode (6-10). This improvement results from a decrease in the recombination rate of photogenerated electrons and holes (6, 8). It has also been shown that photoelectrocatalytic reactors can operate under atmospheres with low content of oxygen (6, 10), and that the negative effect of the presence of salts in the solutions (11, 12) is overcome (6). Note that the performance of these biased photoreactors is affected by the methods used to fabricate the supported catalyst (e.g., firing temperature, number of coatings) and by the characteristics of the test solution (e.g., pH, ionic strength, the types of ions present in solution) (7, 13). All of these studies of photoelectrocatalysis employed titania coatings deposited on either Pyrex or conductive glass. While these materials are highly useful for research purposes, they are problematic for constructing practical devices. Consequently, we attempted to prepare titania coatings on various metals that are stronger, lighter, and/or cheaper than glass, in particular copper, aluminum, stainless steel, and titanium (14). When employed in biased photoreactors, however, the first three substrates were easily oxidized with the subsequent detachment of the titania film under a positive bias, even with no UV illumination. Titanium, though, could be employed to prepare stable and active photoelectrodes, probably because titanium oxidizes only at applied potentials higher than 3 V (15). These photoelectrodes have been characterized in a batch reactor system (16). These early studies were performed in batch reactors on test systems consisting of the target organic dissolved in pure water. In these studies, the effect of changing pH or adding an electrolyte on the performance of the reactor was seldom investigated. Our purpose in this study was to extend these earlier studies to a system that might be scaled up for treatment of actual wastes by preparing photoelectrodes on flexible titanium foil and incorporating these electrodes into a bench-scale, flow-through reactor operating in a recirculating mode. Because practical applications of this technology will likely involve waste streams that contain dissolved salts at varying pH values (expected to be in the range 3 to 9), the effects of varying the operating conditions (flow rate, applied potential, initial concentration of formic acid, initial pH) on the degradation of saline solutions of formic acid (typically 2 mmol dm-3 in 0.01 mol dm-3 NaCl) were evaluated as well as the relationship between photocurrent and degradation rate.
Experimental Methods Chemicals. Suspensions of titania were prepared using titanium isopropoxide (Ti(i-OPr)4, Aldrich Chemical, 97%), tert-amyl alcohol (t-AmOH, Aldrich Chemical, 99+%) and nitric acid (ACS reagent grade). Test solutions contained formic acid (HCOOH, Fisher Scientific, 88% certified ACS) and sodium chloride. All solutions were made with ultrapure deionized water obtained from a Barnstead NANOpure UV system. Preparation of Photoelectrodes. Photoelectrodes were prepared by dip coating titanium foil (20 cm × 14 cm, 0.05 mm thick, Goodfellow Cambridge Ltd.) at a withdrawal speed VOL. 34, NO. 16, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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of 21.5 cm min-1 using appropriate suspensions of titania that had been previously synthesized by sol-gel methods. In previous studies (14), homogeneous coatings of titania on titanium metal were obtained by depositing an initial layer of nanosized titania particles on the metallic substrate using an alcohol-based suspension. Two additional layers of titania were deposited on the initial coating from an aqueous-based suspension. The resulting materials were fired at 300 °C for 5 h. A more detailed procedure and the characteristics of these electrodes are given in ref 16. Reaction System and Auxiliary Equipment. The reaction system consisted of a 200-mL flow-through photoelectroreactor connected to a 500-mL reservoir. All connections were made with Teflon tubing. The atmosphere was controlled by bubbling either oxygen or nitrogen into the reservoir after the gas stream was previously cleaned and humidified by passing it through a gas bubbler filled with ultrapure water. Test solution was recirculated through the system at a flow rate of ca. 90 mL min-1 using a gear pump (Micropump, model 900-573, 180-3600 rpm). The flow-through photoreactor was fabricated concentrically around an 8-W fluorescent UV bulb (F8T5BL) using two 26-cm long pieces of Pyrex glass tube (22- and 45-mm OD) as the inner and outer walls, providing a total reactor volume of ca. 200 mL. The photoelectrode was rolled into a cylindrical tube and fitted inside the outer glass wall. Two Teflon end caps held the feed lines and cathodes, which were 15-cm long, 5-mm diameter, reticulated vitreous carbon (RVC) rods that contained 500 pores per inch. Platinum wires attached to the end of the rods provided electrical contact. Three cathodes were employed to obtain a relatively uniform electrical field throughout the reactor. More details concerning reactor fabrication and a diagram of the photoreactor are provided in the Supporting Information. The intensity of UV light that penetrated the inner glass wall was measured at the location of the photoanode with a photometer (International Light Inc., Model IL 1400A, Super-Slim probe). This value was 5.1 mW cm-2 (which is equivalent to 1.56 × 10-8 eins s-1 cm-2). The effect of varying the flow rate on reactor performance was evaluated with a two-electrode configuration (i.e., without reference electrode). Results of this test are available in the Supporting Information. All other experiments were performed with a three-electrode configuration. For these latter tests, the potential on the photoanode was held constant (vs saturated calomel electrode (SCE)) by a potentiostat (EG&G, model 6310) connected to the SCE through a semimicro saline bridge (EG&G model K0065) placed inside the reactor close to the photoanode. The current passing through the photoreactor was measured and recorded during all experiments. External mass transfer limitations did not appear to be present as there was no appreciable change in the rate of degradation of formic acid for flow rates between 63 and 127 mL min-1. Although Reynolds numbers calculated for this system indicate that it operates in a laminar flow regime, the placement of the cathodes near an inlet stream set at a 45° angle to the reactor appears to have introduced enough turbulence to ensure good mixing. Under these conditions, the reaction kinetics followed a modified Langmuir-Hinshelwood-Hougen-Watson expression, as presented in the Supporting Information. Experimental Procedure. The performance of the reactor when operated in both photochemical and photoelectrochemical modes was tested using aqueous solutions of formic acid in 0.01 mol dm-3 NaCl. Formic acid concentrations varied from 0.0 to 5.8 mmol dm-3 (0 to 70 ppmw as C). Formic acid was selected as the target organic because it oxidizes without producing stable intermediates and is not decomposed by near-UV light or directly oxidized by oxygen (17, 18). Although 3444
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chloride ions may adsorb on titania surfaces, possibly reducing their photocatalytic activity, this process is reversible and permanent deterioration of the catalyst has never been reported (11, 12). Because the test solution was recirculated through an external reservoir, there is some time required for test solution coming from the actual photoreactor to mix with the test solution in the reservoir. At the flow rate of 90 mL min-1 employed for this study, it should require ca. 6 min for a complete turnover of the 500 mL of test solution placed in this reactor. Prior to irradiation, the test solution was recirculated through the reactor for ca. 45 min. During this time, the flow rate was adjusted to the desired value, and the solution was saturated with oxygen or nitrogen. Gas bubbling continued throughout the entire experiment. The pH of the test solution was measured with a double-junction combination electrode (Orion Model 81-72BN) placed in the reservoir and connected to a Fisher Scientific Accumet 50 pH meter. Changes in the concentration of formic acid were monitored by measuring the amount of carbon remaining in the test solution with a total organic carbon (TOC) analyzer (Shimadzu Instruments, Model TOC 5000). All samples were withdrawn from the reservoir. The first sample was taken just before illumination, and the remaining samples at desired intervals. Duplicate samples were taken at a given time, except for the first and last samples, which were triplicate. The volume of each sample was 2.0 mL, to which was added 3.0 mL of 1 vol % hydrochloric acid before the analysis. Concentrations of chloride ions were measured using an ion selective electrode (HNU ISE-30-17-00) after selected experiments were completed. This process required withdrawal of 20-mL samples from the reservoir and adjustment of the ionic strength of the samples to 0.1 mol dm-3 by adding 0.4 mL of 5 mol dm-3 NaNO3. The possible presence of peroxides in treated test solutions was evaluated by adding ca. 50 mg of solid NaI to ca. 2 mL of sample placed in a test tube. Formation of a brown-orange color from the production of I3 indicated the presence of oxidants (likely peroxides) in the test solution. Samples for testing of peroxides were taken from the reservoir both immediately after and 12 h after the experiment was completed. The polarization curve for the supporting electrolyte was determined by measuring the steady-state photocurrents at different applied potentials, as a 0.01 mol dm-3 NaCl solution was recirculated through the reactor. Oxygen was bubbled during this experiment. It was not possible to employ this method to determine the polarization curve for the formic acid test solution because the photocurrent continually decreased as the formic acid was oxidized. Instead, the polarization curve of the formic acid solution was determined by measuring the photocurrent at the onset of photodegradation as the applied potential was scanned across the range of interest (-0.5 to +2.5 V). However, the latter approach leads to initial photocurrents at a given voltage that are considerably higher than the steady-state photocurrent. For purposes of this study, measurements of the steady-state photocurrent appeared more appropriate.
Results and Discussion Effect of the Applied Potential on the Degradation Rate of HCOOH. Figure 1 shows the % carbon remaining in solution as a function of time for a purely photocatalytic experiment (no applied potential) and for experiments conducted at three different applied potentials. When a potential of 0.0 V (vs SCE) was applied across the photoanode, during the first 100 minutes the rate of degradation was similar to that observed in the purely photocatalytic experiment. At longer times, however, the rate of degradation in the photocatalytic experiment decreased dramatically while the rate of deg-
FIGURE 1. Effect of applied potential on the degradation of formic acid with the oxygenated test solution (initial concentration of 2.0 × 10-3 mol dm-3 in 0.01 mol dm-3 NaCl) recirculated at 90 mL min-1. The potential on the working electrode was measured against a saturated calomel electrode. (NAP ) No Applied Potential.)
TABLE 1. Faradaic Efficiencies and Apparent Quantum Efficiencies Observed in Tests Conducted for 150 Min during Feb 1999a applied potential (V vs SCE)
mol formic acid degraded
total mol of electrons
FE (%)
Φapp (%)
NAPb 0.0 1.0 2.0
0.536 × 10-3 0.647 × 10-3 0.700 × 10-3 0.711 × 10-3
1.12 × 10-3 1.53 × 10-3 1.89 × 10-3
116 91 75
2.7 3.3 3.6 3.7
a The test solution was 2.0 × 10-3 mol dm-3 formic acid in 0.01 mol dm-3 NaCl with an applied potential of 1.0 V vs SCE and flow rates between 85 and 90 mL min-1. b NAP ) No Applied Potential.
radation in the photoelectrocatalytic experiment was not appreciably affected. The rate of degradation increased noticeably when the applied potential was increased to 1.0 V; however, further increasing the applied potential to 2.0 V caused only a limited increase in the degradation rate. The application of potentials higher than the TiO2 flat band potential across a photoelectrode increases the concentration of photogenerated holes (or hydroxyl radicals formed by subsequent oxidation of water) on the surface by decreasing the rate of recombination of photogenerated holes and electrons (6, 8, 19). As a result, as the potential increases, the rate of oxidation of formic acid increases, until most of the photogenerated electrons are removed either by the electric field or by reaction with dissolved oxygen. Further increasing the applied potential beyond this value does not improve the degradation rate. The potential at which the maximum rate of degradation is achieved depends on the conditions employed in synthesizing the photoelectrode but appears to be no higher than 2.0 V (14 and this work). If the applied potential exceeds 2.5 V, direct electrooxidation of formic acid begins to occur (7, 16). Table 1 shows the number of moles of formic acid degraded, the amount of charge passed through the cell, the Faradaic efficiency defined by eq 1, and the apparent quantum efficiencies of the tests calculated from eq 2.
Faradaic efficiency (%) ) 2 × mol HCOOH degraded at potential V × 100% (1) mol electrons passed through electrode apparent quantum efficiency, Φapp ) 2 × mol HCOOH degraded (2) mol of incident photons The Faradaic efficiency of the process decreases as the applied potential increases, as reported earlier (14). Faradaic
FIGURE 2. Effect of applied potential on the temporal variation of pH. The experimental conditions were identical to those employed for Figure 1. efficiencies higher than 100% indicate the competition between oxygen and the electric field for photogenerated electrons. At low applied potentials, these electrons are scavenged primarily by oxygen before they can reach the counter electrode. Therefore, they do not contribute to the measured photocurrent. However, the number of moles of formic acid degraded depends on the number of photogenerated holes. At low potentials, since many of the photogenerated electrons do not contribute to the photocurrent, large Faradaic efficiencies can be obtained. As the applied potential increases and more photogenerated electrons contribute to the photocurrent, the Faradaic efficiency decreases (14). However, the apparent quantum efficiency increases with increasing applied potential, with a 37% increase in apparent quantum efficiency obtained at the highest applied potential. Faradaic efficiencies in experiments performed with the same electrodes under the same conditions but three months earlier than the experiments shown in Table 1 were on average 25% lower. These results were a consequence of the higher amount of charge circulated through the reactor in the early experiments, although the amounts of degraded formic acid were similar in both sets of experiments. This change in behavior may be related to changes that occurred in the surface of the TiO2 film as it was used. Such changes might affect the interaction of the electrode with oxygen and thus the subsequent removal of electrons by oxygen. However, detailed studies of possible changes in these films during use must still be performed. Changes in pH during Photodegradation. Figure 2 displays the pH changes that occurred during the degradation of formic acid. In the purely photocatalytic experiment, the pH increased with time while, in the other cases, the pH decreased only a few tenths of a pH unit with time. Comparison of Figures 1 and 2 indicates that there is some correlation between the pH of the test solution and the degradation rate of formic acid: the fastest degradation of formic acid was observed at the lowest pH in the system which corresponds, in turn, to the highest applied potential. The strong influence of pH on the degradation rate of charged organic compounds was reported before (13, 20). At this point, it is not clear why the pH increases as much as it does when this system is operated in a purely photocatalytic experiment (ca. pH 7 after 180 min). Another question of interest is why the pH decreases in the photoelectrocatalytic experiments. Additional experiments were conducted to clarify these issues. Effect of the Applied Potential on the Photocurrent and pH. The first two subsections describe experiments performed with only the supporting electrolyte present in order to determine the effect of varying the applied potential on the photocurrent and the system pH, respectively. These baseline data make it easier to interpret the effect of varying VOL. 34, NO. 16, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 3. Polarization curve under UV illumination. Test solution: 0.01 mol dm-3 NaCl bubbled with oxygen. Flow rate of solution: 90 mL min-1. the applied potential in systems that contain both formic acid and the supporting electrolyte, as discussed in the next two subsections. In some cases, additional experiments were conducted to elucidate the effect on the photocurrent of operating under either oxygen or nitrogen atmospheres. Effect of the Applied Potential on the Photocurrent: Systems without HCOOH. Figure 3 shows the polarization curve under UV illumination when an oxygenated 0.01 mol dm-3 NaCl solution is recirculated through the reactor. This curve is similar to those previously reported for titania-based photoelectrodes on different conductive substrates (21-24). The potential at which the current changes from cathodic to anodic is -0.25 V. This potential can be correlated with the onset potential in anaerobic systems. The onset potential, the lowest potential at which photocurrent can be detected, is a characteristic of a given semiconductor and is a function of pH. This value of onset potential is close to values measured at pH 3, as previously reported (25). At potentials lower than the onset potential, electrons accumulate on the photoelectrode. In aerobic conditions, as in this experiment, dissolved oxygen can be reduced by these electrons, generating a cathodic current (8). At potentials higher than the onset potential, the photocurrent increases rapidly until the potential reaches 0.75 V. This increase in the photocurrent occurs because the application of positive potentials across the photoelectrode establishes a potential gradient within the titania film, which produces the same effect as band bending in single-crystal photoelectrodes (8). As the positive potential increases, the resulting gradient separates holes and electrons, decreasing their rate of recombination. As a result, the photocurrent increases with the potential until most photogenerated electrons either react at the surface of the photoelectrode with an electron scavenger, such as dissolved oxygen, or transfer to the cathode under the influence of the electric field. Increasing the potential beyond this value (0.75 V vs SCE in this system) will not further increase the photocurrent. However, at potentials higher than 1.75 V, the current under UV illumination does start to increase again, probably as a result of the onset of other Faradaic reactions. Figure 4 shows the temporal variation of the photocurrent at three different applied potentials, using a 0.01 mol dm-3 NaCl test solution acidified with perchloric acid. The initial pH of these solutions was ca. 3.0-3.2; similar to the pH of the 2.08 mol dm-3 formic acid in 0.01 mol dm-3 NaCl solutions used as test solutions in the degradation experiments. The curves present a characteristic spike at the beginning and, after the first 15 min, remain constant, with only minor erratic changes (similar results were obtained when the test solution was not acidified). This behavior was observed and explained before for TiO2 photoelectrodes supported on conductive glass (26). The much larger increase in steady-state photo3446
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FIGURE 4. Temporal variation of photocurrent as a function of applied potential. Test solution: 0.01 mol dm-3 NaCl bubbled with oxygen and acidified to pH 3 with perchloric acid. Flow rate of solution: 90 mL min-1. current when the potential is increased from 0.0 to 1.0 V (9 mA total) than when the potential is increased from 1.0 to 2.0 V (3 mA total) is a result of the development of a saturation photocurrent as shown in Figure 3. However, note that the solutions employed to obtain Figure 4 were adjusted to pH 3, while the pH was not adjusted for the experiments shown in Figure 3. The apparent independence of these results on pH will be discussed in the next subsection. When the unacidified electrolyte solution was deoxygenated with nitrogen, it was not possible to apply 1.0 V across the photoanode under UV illumination. The highest potential on the photoanode was -0.41 V, corresponding to a potential difference of 31 V between the photoanode and the counter electrode. To compare these results with tests performed under oxygen, a potential difference of 4.4 V was applied between these electrodes. Under oxygen, setting the voltage at 4.4 V fixed the potential on the photoanode at 2.0 V and led to behavior similar to that shown in Figure 4 for both acidified and unacidified electrolyte solutions. Under nitrogen at 4.4 V, however, in the unacidified electrolyte the potential on the photoanode was -0.62 V, the pH decreased slightly (9.1 to 8.9) over a 10-minute period, and the photocurrent (ca. 5.5 mA) was lower than in the experiments performed under oxygen (ca. 15 mA). Processes occurring on the cathode (counter electrode) are important because they can limit the current in the reactor. Electrons accumulating on the cathode should reduce water or oxygen in order to achieve electroneutrality in the system. Because the overpotential for reducing water on a reticulated vitreous carbon (RVC) electrode is high (27), it seems more likely that oxygen is reduced (28). This hypothesis explains why it was not possible to reach the desired potential (1.0 V) on the photoanode in a nitrogen atmosphere. Under these conditions, it is likely that the current is limited by the ability of the cathode to reduce either water or traces of oxygen, consequently limiting the potential that can be achieved on the photoanode. When oxygen is present, it can be reduced to superoxide, which then reacts with protons to generate hydrogen peroxide and oxygen (29). When tested for by adding NaI, peroxides (likely H2O2) were detected in all test solutions, both oxygenated and deoxygenated, immediately after the experiments were finished but were not detected after the test solutions sat for 12 h. Typical mechanisms for peroxide removal include the oxidation of other species present in solution and the autodegradation of the peroxides with evolution of oxygen. Two possible sources of peroxide in this system are recombination of photogenerated hydroxyl radicals (3, 19) and partial reduction of O2 on the counter electrode (29), although it is not known which source is responsible for generating peroxides in this system.
result may be a consequence of the low affinity of perchlorate ions for the titania surface (11). Clearly, both UV illumination and applied potential are needed to lower the pH in an oxygenated solution of NaCl after ion exchange occurs. In this situation the concentration of photogenerated holes on the surface is expected to be higher than when the surface is only illuminated. As a consequence, the following processes are enhanced, although these reactions depend on the particular surface species present:
FIGURE 5. Temporal variation of pH when oxygenated solutions of 0.01 mol dm-3 NaCl were recirculated through the reactor under different conditions. Flow rate of solution: 90 mL min-1. (NAP ) No Applied Potential; No UV, 1.0 V ) 1.0 V applied without UV illumination.) The possibility of photocatalytic oxidation of chloride ions in aqueous suspensions of titania has been reported (30). To determine the extent of oxidation of chloride ions (if any), the concentration of this species was measured at the beginning and the end of several experiments. No change in the concentration of chloride ions was detected. It is likely, though, that small changes in the concentrations of chloride ions (ca. 10-4 - 10-5 mol dm-3) would not be detected given the large amount of chloride ion present in the test solution. Effect of the Applied Potential on the pH: Systems without HCOOH. Several researchers have observed that the pH of the test solution affects the rate of degradation of charged organic species (13, 20). Because the pH of test solutions of the background electrolyte changed as this reactor was operated, these changes were studied in some detail. Figure 5 shows the variation in pH with time as the electrolyte alone is recirculated through the reactor under different conditions. The pH of the 0.01 mol dm-3 NaCl solution, after being bubbled with oxygen for 30 min, was 6.40 (not shown in Figure 5). After recirculation through the reactor for 30 min, the pH was in the range of 8.7-9.2 (points at t)0 in Figure 5). When 1.0 V was applied across the photoanode (at t)0 with no UV illumination in Figure 5), the pH increased slightly. Under UV illumination and no applied potential, the pH rose from 9.2 to 9.4 in 135 min. When the reactor was operated under UV illumination and with applied potential in the range 0.0-2.0 V, the pH decreased from 9.2 to 4.2-3.5 in less than 30 min. The final pH depended on the applied potential: the higher the applied potential, the lower the pH, with only slight differences between 1.0 and 2.0 V. The increase in the pH when a solution of NaCl is recirculated through the reactor (without UV illumination and no applied potential) is likely a consequence of the following ionic exchange:
> Ti-OH + Cl- a > Ti-Cl + OH-
(3)
If we assume a surface exchange capacity of 0.92 micromol cm-2 (20) and consider only the geometric surface of the electrode (280 cm2), the pH will be close to 10, if complete exchange occurs. However, the concentration of chloride ions would change by only ca. 5 × 10-4 mol dm-3 in this extreme case. In the actual system, the change in concentration of chloride ions is so small as to be undetectable with an ion specific electrode. The hypothesis of surface exchange is supported qualitatively by a similar experiment performed with NaClO4 as the supporting electrolyte. When an oxygenated solution of 0.01 mol dm-3 NaClO4 was recirculated through the reactor, the pH increased but only to 8. This
> Ti-OH + h+ f > Ti-O• + H+
(4)
> Ti-OH2+ + h+ f > Ti-OH•+ + H+
(5)
Enough current passes through 500 mL of the test solution to generate the number of protons required to produce the observed changes in pH, as shown in Table 2. The decrease in pH indicates that the H+ released to the solution are not balanced by an equivalent amount of base (OH-). The reason for this imbalance is not completely understood because, if H+ are released to the solution by an oxidation process, an equivalent amount of base should be released by a reduction process. There are several possible explanations for this phenomenon: 1. adsorption of OH- by the positively charged surface of the photoanode; 2. partial reduction of O2 to peroxides on the counter electrode [If this process is kinetically unfavorable, it may lead to an imbalance in the concentration of H+ and OH-, shifting the pH toward lower values.]; 3. removal of hydroxide ions on the RVC cathode by some unknown mechanism; 4. “storage” of hydroxide ions by peroxides adsorbed on the electrodes (19) [These peroxides release base when they are involved in oxidation processes.]; 5. reduction of Ti(IV) to Ti(III) in the coating would remove electrons from the system without generating hydroxides at the cathode; and 6. oxidation of the titanium substrate to TiO2 would release protons to the test solution. Further insight into these processes was obtained by simply turning off the UV light, upon which the pH of the system increased slowly as the surface reequilibrated (see Figure 5). Therefore, the first explanation given above is not likely, as surface adsorption/desorption of hydroxide ions should occur rapidly, even considering the 6 min required for turnover of the test solution. If the last explanation was correct, one would expect some indication of corrosion of the electrode. There was no visual evidence that corrosion occurred. In addition, this last process is not reversible. We do not yet have enough data to determine if any of the other 4 hypotheses (or some combination of these processes) explains these observations. However, if the second hypothesis is valid, then the kinetics of reduction of oxygen on the counter electrode may control the kinetics of the processes that occur in the reactor. This situation may occur at high applied potentials, given that the rate of oxidation of formic acid is essentially identical for either 1.0 or 2.0 V applied potential, as shown in Figure 1. Such behavior would be expected if the rate of reduction of oxygen at the RVC cathode is the rate-limiting step in this system at high potentials. The changes in pH that occur as the supporting electrolyte is passed through the reactor explain the unexpected changes in pH that were noted in the previous subsection. In particular, it was unexpected that the onset potential measured with the oxygenated supporting electrolyte was similar to reported onset potentials for TiO2 measured at pH 3, when the initial pH of the supporting electrolyte was 6.4. However, as shown in Figure 2, the pH of this electrolyte was actually between 3 and 4 (and varied depending on the actual value of the applied potential) under the conditions used to measure the polarization curve. In addition, the high pH VOL. 34, NO. 16, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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TABLE 2. Number of Mol of Protons Needed To Obtain the Measured pH and the Corresponding Number of Mol of Electrons Circulated through the Celle applied potential (V)
steady-state current (mA)
pH after 30 min
no. of protonsa (mol)
no. of electronsb (mol)
pH after 60 min
no. of protonsc (mol)
no. of electronsd (mol)
0.0 1.0 2.0
3.0 11.5 15.0
4.2 3.7 3.5
3.2 × 10-5 1.0 × 10-4 1.6 × 10-5
5.6 × 10-5 2.2 × 10-4 2.8 × 10-4
4.0 3.5 3.4
5.0 × 10-5 1.6 × 10-4 2.0 × 10-4
1.1 × 10-4 4.2 × 10-4 5.6 × 10-4
a Number of protons needed to obtain the pH of column 3. b Number of electrons passing through the solution after 30 min. c Number of protons needed to obtain the pH of column 6. d Number of electrons passing through the solution after 60 min. e The measurements and calculations were made at three different applied potentials and two different times.
(majority carrier) into the electrode. In the case of formic acid, the mechanism is
FIGURE 6. Temporal variation of photocurrent as a function of applied potential. Test solution: 2.08 mmol dm-3 formic acid in 0.01 mol dm-3 NaCl bubbled with oxygen. Flow rate of solution: 90 mL min-1. measured in the unacidified electrolyte solution when nitrogen was bubbled through it strongly suggests that dissolved oxygen plays an important role in the process(es) responsible for the decrease in pH that occurs when the oxygenated electrolyte is illuminated under an applied potential. As a separate point, we note that solutions of formic acid in NaCl could be degraded in this system under either nitrogen or oxygen atmospheres. From the discussion above, the question arises as to what species is reduced at the cathode when formic acid degrades under nitrogen. The amount of residual dissolved oxygen in this system under nitrogen is about 2 orders of magnitude less than the amount needed to degrade the formic acid. Therefore, it appears that H+ from the formic acid is the most likely species being reduced. If this assumption is correct, then it may imply that the degradation of nonacidic organic species in a photoelectrochemical reactor under anaerobic conditions will require the addition of H+ from a separate source. Effect of the Applied Potential on the Photocurrent: Systems with HCOOH. When formic acid is present in the solution (Figure 6), the photocurrent decreases continuously with time but is much higher at the start of illumination than with no formic acid present. When the applied potential was 0.0 V, the initial photocurrent in the presence of formic acid was more than four times higher than the photocurrent measured without formic acid. At applied potentials of 1.0 and 2.0 V, the initial photocurrent in the systems with formic acid was almost two times the photocurrent measured in the absence of formic acid. As formic acid degrades, the photocurrent decreases as the amount of formic acid decreases until the photocurrent approaches the steady-state photocurrent measured in the absence of formic acid (compare Figures 4 and 6). The increase in the photocurrent in the presence of formic acid is related to a phenomenon known as “current-doubling” (31, 32). The mechanism involves capture of a photogenerated hole (minority carrier) by the organic molecule to form a reactive intermediate, followed by injection of an electron 3448
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+ HCOOH + hvb f HCOO• + H+
(6)
HCOO• f CO2 + H+ + eeb
(7)
As a result, two electrons flow through the external circuit for every photon absorbed. This mechanism leads to an unexpectedly low density of intermediate species such as surface radicals (e.g., Tis-O•). Because radicals trapped at the surface are effective sites for recombination reactions, the overall photocurrent is higher than expected (23, 24). Therefore, the increase in photocurrent that occurs when formic acid is present results from both the current doubling phenomenon and a reduction in the rate of electron-hole recombination processes. One would expect the photocurrent to decrease as the concentration of formic acid decreases during photoelectrocatalytic oxidation at all applied potentials. This effect was observed in this study and has been reported before (33). Effect of the Applied Potential on the pH: Systems with HCOOH. The rise in pH observed in the pure photocatalytic experiment (see Figure 2) can be explained in terms of the oxidation of formic acid and its elimination as CO2. As shown in Figure 5, because of the ionic exchange between Cl- and OH- on the TiO2 film, the final pH value of the test solution should be ca. 9 once all the formic acid has photooxidized. As the pH rises, the reaction rate decreases when the pH exceeds 4, until finally there is essentially no degradation of formic acid when the pH exceeds 5. In the photoelectrocatalytic experiments, however, the pH decreases only slightly with time, and the degradation rate of formic acid is essentially constant. At the beginning of these experiments, the buffer capacity of formic acid maintains the pH near 3. As the reaction progresses and the formic acid oxidizes, the pH remains low because of the generation of H+, as discussed above. Effect of pH on the Performance of the Reactor. Given the strong influence of pH on the degradation rate of charged organic compounds and on the photoelectrochemical process, additional experiments were conducted at different pH values in order to determine the behavior of the system under different test conditions. Effect of pH on the Rate of Oxidation of Formic Acid. For this test, the initial pH of the oxygenated test solutions (2.08 mmol dm-3 formic acid in 0.01 mol dm-3 NaCl) was adjusted by adding different amounts of NaOH. All of the experiments were conducted with an applied potential across the photoanode of 1.0 V. Figures 7 and 8 show the variation in the percentage of carbon remaining in solution and the pH, respectively, with reaction time for test solutions with different values of the initial pH, where initial pH refers to the pH of the test solution before it is recirculated through the reactor. These figures indicate a clear relationship between pH and the rate of
FIGURE 7. Temporal variation of the percentage of carbon remaining in an oxygenated solution of 2.08 mmol dm-3 formic acid in 0.01 mol dm-3 NaCl. The initial pH of the test solution was adjusted to the value indicated in the figure. Flow rate of solution: 90 mL min-1.
FIGURE 8. Temporal variation of pH in the experiment described in Figure 7. degradation. In a pure photocatalytic experiment conducted at pH 8.4, the pH remained constant during the experiment with essentially no degradation of formic acid. In the photoelectrocatalytic experiments, however, the pH decreased until it reached a minimum value (in the range of 3.5-5.0). During this period, the rate of degradation was similar in all cases. However, in the systems with initial pH 6.1 and 9.5, the pH increased after ca. 125 and 35 min, respectively, and, in both cases, the rate of degradation of formic acid decreased within 30 min after the pH started to rise. The initial reduction in pH observed in these tests is likely related with the reduction in pH observed in the experiments performed with the background electrolyte (Figure 5). The amount of H+ generated at the beginning of the experiment, estimated from Figure 5, is enough to reduce the pH of the alkalinized formic acid solution from 9.5 to 4.5. At this pH, CO2 formed by oxidizing formic acid can be vented from the reservoir by bubbling oxygen. As a result, an equivalent amount of OH- (strong base) is liberated to the solution and the pH rises:
NaHCOO + 1/2 O2 f CO2 + NaOH
(8)
Kim et al. (13) reported that the degradation rate of formic acid changed sharply with pH. The maximum degradation rate was reached near pH 3.4 and decreased abruptly at higher and lower pH. The dependence of the degradation rate of organics with pH was explained as a consequence of pHmediated changes in the adsorption of organics on the surface of TiO2. Another possible explanation is that the reduction potential of the TiO2 valence band changes enough with pH
FIGURE 9. Effect of pH on photocurrent in oxygenated solutions of 0.01 mol dm-3 NaCl with and without formic acid (formate concentration: 2.08 mmol dm-3). Curve A: Initially contained only supporting electrolyte, with formic acid added after 15 min. Curve B: Initially contained only supporting electrolyte held near pH 11, with sodium formate at pH 12 added after 30 min. Applied potential: 1.0 V (vs SCE). Flow rate of solution: 90 mL min-1. to affect the rate of degradation of formic acid. However, the reduction potential of the valence band decreases 59 mV for each unit increase in pH, from ca. +2.9 V (vs NHE) at pH 3 to ca. +2.8 V at pH 5 (34). These redox potentials are positive enough to oxidize formic acid directly at both pH values (ECO2/HCOOH ) -0.40 V and -0.52 V (vs NHE), respectively, as calculated for a 2.08 × 10-3 mol dm-3 solution of formic acid). Clearly, the decrease in the reduction potential as the pH increases is not the reason for the lack of photocatalytic oxidation of formic acid at pH 5 or higher. In theory, it should be possible to estimate the changes in pH that occur in this system as a function of the concentration of formic acid that remains in the test solution using simple acid/base calculations with suitable approximations. In practice, however, several constraints limit our ability to perform a meaningful calculation. (1) The amount of OH- released to the test solution when it is originally recirculated through the reactor is likely to be pH dependent. This value could be estimated if the pH of the test solution was measured just before the UV light and the potential were turned on, but this measurement was not performed for this experiment. (2) The number of equivalents of acid released to the solution when the reactor is turned on cannot be measured directly. This value can be estimated from Figure 5 if one assumes that this value is independent of pH and is the same in all the experiments performed at a given applied potential. Because the mechanism for the release of acid is not known, it is not clear that this value will be independent of pH. (3) When the pH of the test solution was monitored just before turning on the UV light and the potential, the amount of OH- released during recirculation was observed to vary somewhat for otherwise identical test solutions. This observation indicates the difficulty associated with obtaining absolutely identical conditions at the surface of the photoelectrode from experiment to experiment. Despite the difficulties listed above, efforts were made to calculate the amount of formic acid that would need to be degraded for the test solution initially at pH 9.5 to return to pH 9.5. Oxidation of only 25% of the initial amount of formic acid was required under one set of assumptions. If the assumed number of equivalents of acid released during operation of the reactor was increased by 25%, then oxidation of 40% of the formic acid caused the pH to rise to 9.5. Given that 60% of the formic acid was actually oxidized during the experiment before the pH returned to 9.0 (much less 9.5), it is clear that further studies are required before this system can be effectively modeled. Effect of pH on Photocurrent. Figure 9 shows the close correlation between the degradation rate of formic acid and VOL. 34, NO. 16, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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the photocurrent. Both experiments were performed under UV illumination, with oxygen saturation and with +1.0 V applied across the photoanode. In one experiment (system A, Figure 9), the pH of the supporting electrolyte (0.01 mol dm-3 NaCl) was allowed to change freely during the photoelectrochemical process, and the temporal variation of the photocurrent was measured. After the first 15 min, enough 0.833 mol dm-3 (10,000 ppm) formic acid was injected in the reservoir to make a 2.08 mmol dm-3 solution. The effect of adding formic acid was a sharp increase of the photocurrent because of the concomitant reduction in the rate of electronhole recombination in the system with the onset of photooxidation, as described earlier in the discussion of “current doubling”. In the other experiment (system B, Figure 9), the pH of the supporting electrolyte was adjusted to 11 and held near that value by adding small portions (160 µL in the first 30 min) of 5 mol dm-3 NaOH to the system. Note that the photocurrents in both systems are nearly identical over the first 15 min of testing (i.e., with no formic acid added) even though the pH is quite different in the two systems. After 30 min, enough sodium formate (alkalinized to pH ) 12) was added to system B to make a 2.08 mmol dm-3 solution. In this case only a slight increase in the photocurrent was noticed. This result supports the observation that, at pH values higher than 5, the oxidation rate of formic acid is low. The “current doubling” effect is diminished because only a few formate radicals are produced that can then inject electrons into the conduction band. These results demonstrate the relationship between the photocurrent and the degradation rate of formic acid. Because formic acid does not photocatalytically oxidize in this system at high pH, the photocurrent measured in the presence of formate at high pH is much lower than at low pH. Once again, these data emphasize the importance of pH as a process variable in the operation of photoelectrocatalytic reactors.
Practical Implications Practical applications of photocatalytic oxidation have been limited, in part, by the relatively slow rates of degradation obtained with this technology. Photoelectrocatalysis has been considered as one means of improving the performance of these systems. As in earlier studies, we observed that applying a potential across the photocatalyst improved its performance when compared to a strictly photocatalytic system, although this increase was nowhere near the one or 2 orders of magnitude improvement needed to drive the commercialization of this technology. However, our flow-through biased photoreactor was an initial design that was not optimized, so further improvements in performance are likely. If the cathodic reaction (i.e., dissolved oxygen accepting electrons) is the rate-limiting reaction in this system, then the use of other cathode materials might well improve the performance of this reactor. As we have shown in this paper, however, a number of processes occur in this reactor during operation, some of which are rather unexpected. Because of the complexity of this system, a better understanding of these processes is needed to determine the types of wastes that can be treated using this technology. At its present stage of development, photoelectrocatalysis can probably only be considered for treating relatively dilute wastes (maybe 100 ppm or less) under conditions that provide residence times on the order of 1 h. Note, though, that within these constraints, this technology can be employed to treat rather complex wastes such as acidic anaerobic wastes containing dissolved salt (as discussed earlier) and mixed wastes containing organics as well as 3450
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dissolved metals that can be electroplated. This reactor was employed to treat a saline solution of formic acid and dissolved copper, as will be discussed in a future publication.
Acknowledgments This research was conducted with support from the U.S. Environmental Protection Agency National Risk Management Research Laboratory under Contract No. CR 824125-01-0.
Supporting Information Available More details are available about three aspects of the operating characteristics of the biased photoreactor employed in this study: the reactor design, the effect of varying the flow rate on the rate of degradation of formic acid in the reactor, and the kinetics of degradation. This material is available free of charge via the Internet at http://pubs.acs.org.
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Received for review September 4, 1999. Revised manuscript received May 18, 2000. Accepted May 22, 2000. ES991024C
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