Environ. Sci. Technol. 1994, 28, 479-403
Photoelectrocatalytic Degradation of Formic Acid Using a Porous Ti02 Thin- Film Electrode Dong Hyun Kim and Marc A. Anderson.
Water Chemistry Program, University of Wisconsin, 660 North Park Street, Madison, Wisconsin 53706 The degradation of formic acid (HCOOH) using titanium dioxide (TiOz)in photocatalytic and photoelectrocatalytic reactions was investigated in order to determine whether electrical biasing could improve the efficiency of photocatalytic reactions. This study addressed the effects of film thickness, biasing potential, presence of oxygen, and added inorganic electrolytes on the photocatalytic degradation of HCOOH. The results of these experiments showed that the degradation of HCOOH in this system was due only to the photocatalytic as opposed to homogeneous photolysis reactions. Degradation efficiency of the photocatalytic reaction was roughly proportional to the Ti02 film thickness. In the photoelectrocatalytic reaction, positive potentials (vssaturated calomel electrode, SCE) improved the degradation efficiency and +O.O V (vs SCE) was enough to obtain a maximum efficiency. The supply of oxygen was essential in the photocatalytic reaction, while the photoelectrocatalytic reaction was not significantly affected by the removal of oxygen. The presence of inorganic electrolytes lowered the efficiency of the photocatalytic degradation of HCOOH. However, the efficiency of photoelectrocatalytic degradation was not affected by inorganic electrolytes. Overall, when used with the bias, the system showed efficient degradation over a wide range of conditions.
Introduction Since the oil crises of the early seventies, there have been many attempts to make solar energya practical energy source. For that purpose, research has examined photocatalytic reactions using UV-illuminated Ti02 in a variety of systems as a means to split water in order to produce HZ (1-3). As environmental problems became more serious, scientists have begun to use these same Ti02 systems to degrade organic contaminants in water. Scientific interest toward UV-illuminated Ti02 increased because it provided an easy and inexpensive way to mineralize toxic organics to unharmful C02 and water. A complete discussion on semiconductor photochemistry of these systems can be found in a recent review (4). Previous research has proven the ability to degrade organic compounds with Ti02 (5-161, including surfactants, phenols, herbicides, chloroform, pesticides, and carboxylic acids. Many of these compounds can be totally and efficiently mineralized. The Ti02 photocatalytic reaction is now generating commercial interest because of its low cost, simplicity, robustness, and ability to achieve extremely low residual organic contaminant levels. However, several improvements must be made before these systems can become economicallyattractive. One improvement is to increase the contact area between the activated Ti02 and the
* Address correspondence to this author; e-mail address: andersona engr.wisc.edu. 0013-936X/94/0928-0479$04.50/0
0 1994 American Chemical Society
solution, which can be done by using porous Ti02 films made by the sol-gel method (17). The second improvement needs to be made in increasing the lifetime of the shortlived holes created by light-induced electronic excitation which are the oxidizing agents in this reaction. Increasing the hole lifetime would increase the reaction rate. In our system, we address both of these issues by applying a biasing potential to a porous Ti02 film in order to withdraw the excited electrons to a counter-electrode before they could recombine with the holes (18). In this study, formic acid (HCOOH) was employed as a target chemical because HCOOH oxidizes to CO2 without forming any intermediates. Furthermore no homogeneous reaction occurs in the range of UV light employed in these studies (300-400 nm).
Experimental Methods Chemicals. All chemicals (Fisher and Aldrich) were used without any further purification. The water used in this experiment was deionized using a Milli-Q water purification system (Millipore Corp.). Ti02 Sol Preparation and Characterization. The Ti02 sol was synthesized according to the sol-gel method (17). A solution with a ratio of 150 mL of Hz0:15 mL of Ti(isoPro)d:l mL of "03 was refluxed at 80 "C for 3 days. The porosity and surface area of unsupported films were characterized by BET analysis using Nz as the adsorbing gas. Particle size and particle size distribution in the sol were measured by laser light scattering (Brookhaven Instruments Corp.). The crystalline form of Ti02 was determined by X-ray diffraction. Ti02 Thin-Film Preparation. A PWM spinner system (Headway Research Inc.) was used to produce a uniform Ti02 coating on tin(1V) oxide-covered glass (F. J. Gray & Co.). The glass slide measured 100 mm X 50 mm X 2.5 mm. The upper 1cm of the glass was taped off before coating in order to provide a place for ohmic contact. Each layer was made by spinning 500 p L of Ti02 sol at 2000 rpm for 30 s. After spinning, the supported Ti02 film was dried at room temperature. The film was fired in an oven using a temperature ramping rate of 3 "C/min. At 400 "C, the film was allowed to dwell for 2 h and was cooled using a ramping rate of 3 "C/min until it obtained room temperature. The process was repeated up to 10 times for the multilayered films. The absorbance of Ti02 on the glass was measured by UV-visible spectroscopy. The thickness of the film was measured by SEM. Reactor System and Auxiliary Equipment. The reactor and supplementary equipment are shown in Figure 1. The main components are the glass reactor, un ultraviolet light source, and a potentiostat. The reactor and the UV lamp were placed in a black acrylic box in order to avoid extraneous illumination. The rectangular (75 mm X 100 mm X 125 mm with a 2.5-mm wall) reactor (Vitro Dynamics, Inc.) was made from borosilicate glass. A plastic cap was placed on the reactor in order to seal it Environ. Sci. Technoi., Vol. 28, No. 3, 1994
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Figure 2. X-ray diffraction measurement fw the unsupported TiOI flred at 400 "C.
UV-Light
Reactor
Flgure 1. Diagram for the reactor system for the photoeleceocataiytic
reaction: A, platinum wire (counter-electrode,cathode):6, saturated calomel electrode (reference electrode): C, tltanium dioxide film electrode (working electrode, anode); D. oxygen supply. and install three electrodes. The reactor, which contained a 800-mL sample solution allowing 7 cm of the supported Ti02 film to be immersed into the solution, was placed 7 em in front of the UV lamp (General Electric, F15T8.blb. 15W). The intensityofliiht,asmeasuredbyanoptometer (United Detector Technology System S380), was 2 mW/ cmz a t 7 cm into the reactor, the position where the Ti02 film was placed. The photoelectrocatalytic reaction employed a potentiostat (IBM, EC/225) connected with a counter-electrode (Pt wire, 3 in. in length with a 0.4-mm diameter), a working electrode (a supported T i 0 2 film), and a reference electrode (a saturated calomel electrode). The computer (IBM, PC/AT) was connected to the potentiostat via an AID converter (PCS Inc.) in order to store current and potential signals. Blank Experiments. The first experiments were designed to ensure that the disappearance of HCOOH occurred as a result of a photocatalytic reaction between UVlightandTiOz. Weneededtodetermine whetherthere was any purely photochemical degradation of HCOOH and if the disappearance of HCOOH could he due to specific adsorption of HCOOH on the reactor, SnOz, or Ti02 film. Also, we needed to determine the effect of SnOz and whether any spontaneous evaporation of HCOOH could occur. Photocatalytic and Photoelectrocatalytic Reactions. The photocatalytic degradationreaction of HCOOH was carried out in the reactor system described above. Given potentials were applied with the three-electrode system to perform photoelectrocatalyticreactions.Oxygen was bubbled through the solution and it was continuously stirred during both the photocatalytic and photoelectrocatalytic reactions. The dissolved oxygen (DO) was measured by DO meter (Great Lakes Instruments, Model 867). The initial concentration of HCOOH was around 60 ppm, and the temperature of the solution was 23 f 3 "C. The amount of degradation was determined with a total organic carbon analyzer (TOC 5000, Shimadzu), which measured the concentration of dissolved organic carbon. 400
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We tested the reactor under a variety of conditions that might occur in practical water treatment. In one experiment, we bubbled nitrogen through the solution instead of oxygen in order to determine if any reaction would occur under anaerobic conditions. In another experiment, we added various inorganic electrolytes which would likely be present in real waste waters. Results and Discussion Properties of Ti02 Sol andunsupported Film. The particles in the TiOz sol were 60 nm in diameter. The surface area and pore radius of an unsupported Ti02 film, fired at 400 OC for 2 h, were 123.9 mz/gm and 18.2 A, respectively. The surface area per gram of substrate was larger than that ofpreviously published studies using Ti02 (5-16). However, the absolute area of exposed TiOz was limited by the area of the supporting glass substrate (35 cm2). The Ti02 was in the anatase form, as determined by X-ray diffraction (Figure 2). While the exact results of the sequential coating process are unclear, absorbance measurements taken afterdepositingseverallayers of film suggest a limitation in the amount of Ti02 which could he coated on the glass (Figure 3). Some spectra indicate a band edge around 370 nm rather than at 390 nm as expected for anatase. This may be due to the different thickness of SnOz in the "as received" conducting glass which is used as a reference. The thickness of a five-layer film as measured by SEM was about 0.5 hm (Figure 4). BlankTests. Specific Adsorption ofHCOOH. There was no change in the solution concentration of HCOOH due to the specific adsorption on the surface of the reactor, Ti02 film, or pure SnOz-coveredglass during a 6-h period. This does not imply that HCOOH is not adsorbed by these surfaces, but rather that there is not sufficient surface areatoproduceany changein theHCOOHconcentrations over this time period. Effect of Tin Oxide. SinceourTiOzisaporousmaterial, HCOOH can reach the supporting conductive SnOz glass. However, no change of HCOOH was detected during the photocatalyticandthe photoelectrocatslytic (0.3Vvs SCE) reactions using a pure SnOz-covered glass instead of a Ti02 film. This demonstrates that either photoelectrocatalytic reactions do not occur on SnOz or that there is not enough surface area to cause any detectable reaction. Euaporation of HCOOH. No spontaneous evaporation of HCOOH from the sample solution was detected while oxygen was bubbled through the solution for 3 h. HomogeneousReaction. HCOOH does not absorb light in the spectral range of 320-400 nm, and no significant
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Absorbance Peaks lor the various number of T102 layers coaled on tin oxide glass (number shown represents number of layers).
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Flgura 4. SEM picture for the five-layel' film of Ti02 on tin oxide-coated glass substrate
homogeneous photodegradation reactions of HCOOH occurred. Therefore, we conclude that the disappearance of HCOOH in our subsequent studies is due to the photocatalytic reaction whichoccursat or near thesurface of the TiOn. Photocatalytic Reaction. Ti02 Film Thickness. Figure 3 shows that the Ti02 film absorbance is roughly proportional to the number of layers deposited on to the conducting glass substrate. As expected, the amount of
degradation is also proportional to the number of layers which is shown in Figure 5. Oxygen. The dissolvedoxygen was 36 ppm when oxygen was supplied. When nitrogen was supplied instead of oxygen, dissolved oxygen was 0.7 ppm, and the amount of HCOOH degraded is shown in Figure 6. Because oxygen is theessentialoxidizingagentinthis reaction, weexpected the degradation of HCOOH to be very slow as oxygen tensions were diminished. There was probably a small Envlron. Sd.Twhnol.. Vol. 28. NO. 3. 1994 481
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Flgure 5. Effect of number of llOElayers in photocatalytic and 0 . 3 4 (vs SCE) photoelectrocatalytic reaction over a 4-h time period.
PureHCOOH
KCI
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Flgure 7. Effectof i M Inorganic ions in phoiocatalytlc and biased (0.3 V vs SCE) reaction over a 100-min time period.
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Figure 8. Effectof oxygen in photocatalyticand photoelectrocatalytic (0.3 V vs SCE) reactions over a 4-h time period.
Figure 8. Effect of NaCi in HCOOH photocatalytic degradation over a 4-h time period.
amount of oxygen adsorbed in the pores of the film before it was inserted into the reactor or a low concentration of oxygen in the solution since the reactor is not sealed, which allowed the readion to take place at all. Beside the reduction of oxygen which occurs at the TiOdaqueous solution interface from photogenerated electrons, Ti02 itself can be reduced. In this case, Ti02 becomes a blue/ purple color. Unfortunately, the change of color in these Ti02 films could not be easily distinguished from their original color. This may be due to the interference caused by SnO2 and also to the thickness of the Ti02 film. Dissolued Inorganic Ions. Upon the addition of 1 M inorganic ions to the HCOOH solution in the form of KCI, NaC104, NaN03, and NanSOa, the degradation efficiencies decreased (Figure 7). The pH of our HCOOH solutions is 3.3. The isoelectric point of titanium dioxide anatase is around 6 (17). Therefore, under the conditions of our experiments, the surface of Ti02 is positively charged, and some HCOOH (pK. = 3.74 a t 25 "C) exists as formate ions. A t this time, we are not certain which formic acid species is more favorably adsorbed on the surface of TiO2. Some species such as arsenite show maximum adsorption a t their pK (19). However, lowered efficiencies in pho-
tocatalytic reactions seem to be attributable to the competition between the formate ion and electrostatically adsorbed ionic species at the TiOz surface. This effect hecomes very noticeable when the electrolyte concentration is higher than the HCOOH concentration. Figure 8 shows how degradation efficiency depends on electrolyte concentration. In an earlier paper (ZO), the effects of perchlorate and nitrate in photocatalytic reactions (salicyclic acid, aniline, and ethanol) were shown to be trivial. The maximum concentration of perchlorate or nitrate employed in these previous studies was 0.1 M. The concentration of those ions in this study was 1M. Higher concentrations used in this study may decrease photocatalytic reactions significantly. Further study of the effects of these inorganic ions is in progress. Photoelectrocatalytio Reaction. ExternalPotential. The amount of degradation increased when an external potential was applied to the Ti02 electrode. We believe this to he due mostly to a decrease in the electron-hole recombination rate. Variable Applied Potential. When various potentials (from 0 to +2.0 V vs SCE) are applied to a five-layer film, the absolute amount of increase in degradation for each
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By biasing these thin, porous Ti02 electrodes, we hope to take the photocatalytic investigation of organics one step closer to practical use. Not only does this system increase degradation efficiency under normal conditions, but it maintains its high efficiency under adverse conditions, such as in the absence of oxygen or in the presence of interfering ions. Further experiments should help to clarify the mechanism of the reaction in this system, the reasons for its effectiveness, and the modifications which should help to make it even better. Acknowledgments
Photo
0.0 Applied
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potential ( V vs. SCE)
Effect of applied potential (vs SCE) using a five-layer Ti02 film over a 4-h time period.
We thank the United States Environmental Protection Agency for funding this research under Grant R81711501. Literature Cited
Flgure g,
applied potential was about 5 ppm (Figure 9). A 0-V potential (vs SCE) was enough to establish the maximum effect of an applied potential in this system. These results suggest that a 0-V potential (vs SCE) provides enough band bending to withdraw electrons to the counterelectrode and implies that the flat-band potential of Ti02 in this experimental condition is below 0 V (vs SCE). In photoelectrocatalytic reactions, the reduction reaction on the counter-electrode may be the reduction of oxygen. The reaction rate may be controlled by this reaction. By enlarging the surface area of the Pt wire by a factor of 2, we hope to examine the effect of surface area. We could see no difference between the two different size Pt wires in photoelectrocatalytic reactions. Ti02Film Thickness. The reaction efficiency is roughly proportional to the number of layers in the film, as for the photocatalytic reaction. However, the results in Figure 5 show some possible limitations of the system. The 10layer film showed no improvement over the photocatalytic reaction alone, which may be because of its relatively high resistance. Also, the efficiency fluctuated for film thicknesses between seven and ten layers, which might be due to the experimental variation. Diffusion limitations within these films have been studied by Sabate et al. and are not likely rate limiting in these systems (21). Oxygen. When we bubbled nitrogen through the solution instead of oxygen, the efficiency decreased, but only by about 20%,as illustrated in Figure 6. This means the photoelectrocatalytic system can also be used in anaerobic conditions, unlike the photocatalytic reaction. This also means that some other species, instead of oxygen, can accept electrons. As a result of the reduction reaction occurring on the counter-electrode,HZevolution may occur. Dissolved Inorganic Zons. The addition of inorganic electrolytes did not affect the efficiency in photoelectrocatalytic reactions as shown in Figure 7. In this case, the presence of two reaction surfaces, the two electrodes instead of one may reduce the competition for reactive sites. Alternatively, a change in the electrical double layers near the semiconductor (electrolyte interface) may be responsible for the increased performance of the biased system.
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Abstract published in Advance ACS Abstracts, December 15,
1993. Envlron. Scl. Technol., Vol. 28, No. 3, 1994
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