EnWon. &I. T&d.
1999, 27. 1776-1782
An Investigation of Ti02 Photocatalysis for the Treatment of Water Contaminated with Metals and Organic Chemicals Michael R. Prairie,' Lindsey R. Evans, Bertha M. Stange, and Sheryi L. Marlinez
Center for Advanced Energy Technology, Sandia National Laboratories, Albuquerque, New Mexlco 87185 Laboratory experiments were performed to investigate Ti02 photocatalysis for treating water contaminated with dissolved metals (Ag,Au, Cd, Cr, Cu, Hg, Ni, and Pt) and a variety of organics (e.g., methanol, formic acid, salicylic acid, EDTA, phenol, andnitrobenzene). It was found that only those metals with half-reaction standard reduction potentials more positive than 0.3 V (vs normal hydrogen electrode) can be treated using Ti02 as the photocatalyst. Kinetic data illustrating the synergism between oxidation and reduction are presented. Experiments using singly substituted benzenes as electron donors show that the rate of reduction of Cr(VI) is correlated with Hammett u constants. Photoefficiencies approaching 85% were measured for the conversion of Cr(V1) to Cr(II1) using citric acid as the reductant. In contrast, photoefficiency was only 4% when oxidizing salicylic acid using 02 as the oxidant. It is concluded that efficient designs of photocatalyticsystemsfor wastewater treatment must take into account both oxidation and reduction Drocesses. Introduction Semicondudor photocatalysishas received considerable attention in recent years as an alternative for treating water polluted with hazardous organic chemicals (1-14). To a lesser extent, it has also been studied for application towatercontainingmetals (fordecontaminationandmetal recovery) (15-26) and for making highly dispersed supported metal catalysts (23,27,28). There is also a good deal of literature on the fundamental processes of photocatalysis (29-32) including the effects of metal electron scnvengers for enhancing water splitting and the production of 02 (33-36). These reports make it apparent that photocatalysishas many potential applicationsfor treating water containing organics, for removing metals from water, and for splitting water. It is also quite clear that these processes are intrinsically interrelated. The energetics and charge-transfer processes involved in a typical photocatalytic process are illustrated in Figure 1. Bandgap illumination (hu)of a semiconductorparticle suspended in water causes electronic transitions from the valence band to the conduction band, leaving holes in the former. These electrons and holes then either migrate to the particle surfaceand become involved in redox reactions or they recombine and simply liberate heat. Conductionband electrons are consumed in reactions that reduce oxidants (Ox O x w ) while holes are filled via oxidation reactions (Red Red&). In this paper, we describe an investigation of the kinetic and electrochemical relationships among these two processes. Hydroxyl radicals are generated by the oxidation of water at the valence hand of Ti02 (2, 7,9,11,37). This occursatastandardpotentialof2.8V(%),whichdecreases with increasing pH. All values for redox potentials in this report are referred to the normal hydrogen electrode. As
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Author to whom correspondence should ba addnssed. 177e Env*on. Sci. Techm., Vd. 27, NO. 9, 1993
Flguv 1. Energetlcs and reactbans Involved In semlcocductw photocatalysis. The values In pamttwses are for TlO, (anatase VII nwmal hydrogwn elecbode, pH 7).
water is oxidized, reduction must occur at the conduction hand to maintain electroneutrality. During typical photocatalytic oxidation of organic components, oxygen is reduced to become superoxide (-0.56 V) and/or perhydroxyl(-0.13 V) radicals (%), depending on pH. These radicals eventually form hydroxyl radicals which enter into the oxidation cycle (2, 7,9,11,37). Besides oxygen, any dissolved species with a reduction potential more positive than the conduction hand of the photocatalyst can,inprinciple,consume electronsandcompletethe redox cycle. Hg(II), Ag(I), and Cr(VI) are examples of oxidants that are reduced by Ti02 photocatalysis. Hg(I1) and Ag(I)are converted into metallic crystallites on the photocatalyst (21,22,24,30,31). Cr(VI)isconvertedintoCr(II1) (20, 25, 26) which, in general, remains in solution (this depends on p H at high pH, Cr(II1) is hydrolyzed and precipitates). It should he noted that the process is not truly catalytic when metals are plated out onto the photocatalyst because the catalyst surface is being consumed. However, the destruction of organics is catalytic. In this paper, we present work that explores in detail the kinetic interactions between reduction and oxidation processes from the point of view of obtaining efficient destruction of organics and efficient reduction of toxic metals. We show how the type and concentration of the oxidant affects oxidation and vice versa for reduction. We also look at how the standard potential of the reduction process affects the rate of oxidation of organics, and we correlate the rates of reduction with the Hammett u constants for singly Substituted benzenes. In addition, we show that it is p w i h l e to obtain photoefficiencies approaching unity by optimizing both half-reactions. Experimental Section The following metal salts (reagent grade) were used AuCk, Cd(N03)~4HzO,Ni(NOd~6Hz0,CuS01.5H20, KZCrzO,, HgC12, HzPtCla.xHz0, AgNOz, and &NOS. Inductively coupled plasma/mass spectroscopy was used to analyze for Cd (0.003), Cu (0.003),Ni (0.005). and Pt (0.002), and atomic absorption was used for Hg (4), Au (10),andAg (2). Cr(V1) (25)wasanalyzedcolorimetrically 0015936X/93/0927-1778$04.0010
t 3 1993 Amehan Chmkal Socbty
using diphenylcarbazide (39). Detection limits in microgram per liter are indicated in parentheses. In several instances, we verified that all the Cr(V1)that was consumed was in fact converted to Cr(II1). This was achieved using potassium permanganate to convert the Cr(II1) to Cr(V1) followed by colorimetric determination of the resulting Cr(V1)using diphenylcarbazide (39). In all cases, the mass balance held, and no material was unaccounted for. The organics were chosen for their differing electrochemical characteristics; they were salicylic acid (SA), mandelic acid, citric acid, disodium EDTA, methanol, ethanol, acetic acid, formaldehyde, formic acid, aniline, benzaldehyde, benzonitrile, benzoic acid, nitrobenzene, and phenol. All were at least reagent grade. UV/visible spectroscopy at 296 nm was used to determine SA concentrations, Of the metal salts used, KzCr207 prevents determination of SA; HzPtC16 and AuC4 interfere slightly; and CuSO4 distorts the 296-nm absorption band indicating some type of chemical interaction, possibly the formation of copper salicylate. A Shimadzu 5000total organic carbon (TOC) analyzer was acquired toward the end of this study, so some TOC data are also discussed. Otherwise, we did not analyze for any other organics. Batch experiments were performed in a glass-jacketed reactor vessel resembling a pot. It is open on top (13-cm i.d., 1.5 L) and equipped with two sample ports. A quartz lid is sealed in place with a Viton O-ring. Oxygen-free experiments were carried out by sparging nitrogen (900 cm3/min, 20 OC, 1atm) through a stainless steel frit into the reaction solution for 15 min prior to and throughout the experiment. Nitrogen leaked out of the system through the sample port. The contents of the reactor were magnetically stirred throughout each experiment and cooled with tap water to maintain constant temperature (20 f 5 "C). Each experiment used a 300-mL solution illuminated with ultraviolet light for up to 60 min. The photocatalyst was anatase titanium dioxide powder (Tioxide Tilcom HACS, 270 m2/g) suspended in solution at 0.1 wt %. Variables included the type and amount of organic, the type and amount of metal, the absence or presence of dissolved oxygen, and pH. Metals and organics were used between 0 and 5 mM. Initial pH was 6 in all cases except for studies on the reduction of Cr(V1) when it was varied between 1and 6. pH adjustments were made using dilute "03 and NaOH. All water used was carbon filtered and deionized (0.5-3 p U at 1 cm). To start each experiment, a 100-W UV spot lamp (Spectroline Model MB-100 with a Sylvania Par 38 mercury bulb; 7000 pW/cm2 UV366 at 40 cm) without a visible filter was placed on the quartz lid shining into the reactor solution. Integral lamp output (12 cm away) between 300 and 390 nm was found to be 38.7 mW/cm2 using a spectroradiometer. The principal emission lines are at 365 nm (4.2 mW cm-2 nm-l), 406 nm (2.3 mW cm-2 nm-l), and 437 nm (4.4 mW cm-2nm-l). Periodic control experiments (0.1 wt % TiO2,30 mg/L SA, pH 6, no metal) verified no deterioration of the lamp during the entire course of the experiments. Two series of experiments were performed to determine photocatalytic reaction rates as functions of light intensity. These tests were carried out in a 200-mL glass-jacketed reactor vessel filled with 100 mL of reaction mixture 10.1 wt % TiO2,50 mg/L Cr(VI), 0.72 mM citric acid, pH 2 or 0.1 wt % TiO2, no metal, 0.22 mM SA, pH 61 open to the atmosphere and wrapped with aluminum foil. Square
Table I. Metals Removal from Water by Adsorption and Photocatalytic Reduction*
metal AgW Au(1II)b CU(IIP Cd(IUb CrlVIY Hg(1Ijb Ni(II)b Pt(1V)b
fraction removed after 40-min UV (ambient Od,%
fraction removed after 40-min UV (no 021, %
36
85
42
>94d 99.97
13
lOOe 17
I 20 42 I 16
I 84 98 I 58
fraction adsorbed (dark), %
17
I 84
loor I 92
a I = insignificant difference in concentration before and after treatment. Parameters: 0.1 wt % TiOz; 30 mg/L salicylicacid. pH 6, metals are reduced to metallic state onto the photocatalyst. pH 2, Cr(V1) is reduced to Cr(III), 50% of which stays in solution, the rest precipitates. d Residual silver was below the calibration limit of 4 mg/L. e Residual gold was below the detection limit of 0.009 mg/L. f Residual mercury was below the detection limit of 0.004 mg/L.
neutral density filters (5 cm X 5 cm) were used to attenuate light from the spot lamp. Also,the spot lamp was equipped with a colored UV band-pass filter that only passes UV light between 290 and 420 nm (peaking at 72% transmittance at 365 nm). Light intensity for each experiment was determined using a hand-held radiometer (ColeParmer Series 9811) calibrated at 5 points between 0 and 20 mW/cm2by potassium ferrioxalate actinometry (40) in the small reactor vessel. For all experiments, the starting solution with the catalyst was allowed to equilibrate in the absence of UV illumination for 15min (sometimes under nitrogen purge). At this time (t = 0), a 1-mL sample was removed, and the lamp was placed on the reactor. Other samples were later removed at known times. The samples were filtered through 0.2-pm Teflon syringe filters to remove the catalyst. When necessary,samples for metals analysis were stabilized in 2% "03 (Fisher, Optima Grade) or 2% HC1 (Fisher, Ultrex Grade).
Results and Discussion Effect of Metals on the Oxidation of Salicylic Acid. Table I summarizes results for the eight metals using 30 mg/L (0.22 mM) salicylic acid as the reductant. Experimental parameters are listed in the caption. Adsorption is the first thing that happens when suspended Ti02 is contacted with dissolved metal ions. This process is reversible, does not require light, and is very sensitive to pH because (1)pH affects the surface charge of TiOz, thereby affecting electrostatic interactions between it and the substrate, and (2) pH affects the extent of hydrolysis and species distribution of the metal ions (41). Upon exposure to UV light for 40 min, gold(III), silver(1) (from the nitrate), chromium (VI), mercury(II), and platinum(IV) are reduced as indicated in the table. The tabulated data also show that oxygen inhibits the photocatalytic reduction of silver, mercury, and platinum, probably by competing with the metals for conduction-band electrons. Cr(V1)and Au(II1)are reduced so rapidly that competition by 02 is negligible. As illustrated in Figure 2, SA disappears at different rates depending on what oxidant is used; note that oxidation of SA occurs only for those metals that experience reduction (Table I). When neither 0 2 nor metals are present, SA remains essentially unreacted. If the reactor is open to atmospheric air, fairly rapid disappearance of SA is observed. In the absence of 0 2 , rapid destruction Environ. Sci. Technol., Vol. 27, No. 9, 1993
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0
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20 30 TIME (min)
Figure 2. Effect of varlous oxidants on the photocatalytic Oxidation of salicylic acid, pH 6, 0.1 wt % Ti02, 30 mg/L salicylic acid initially (curves start at about 15 mg/L as a result of adsorption). (0)with ambient air, no metal. The following were purged with Nz: (0)no metal; (V)Hg(I1); ('I AgN02; ) ( 0 )Cu(I1); (m) Ni(I1); (A)Cd(I1); (A) R(IV); (e) AgNOs; (0) Au(III). I
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0.0 0.5 1.o 1.5 STANDARD ELECTRODE POTENTIAL (V) vs. NHE
Flgure 3. Initial Oxidation rates of SA from Figure 2 versus standard reductlon potential (NHE) for each metal. The dashed line is the Tafel equation, In(d = -4.89 - 3.98 (-0.5- P ) ,where r is Initial Oxidation rate and F' is standard potential. The parameters were obtained by linear regresslon. O2 reductlon Is shown at the potential for formation of perhydroxyl radlcal. The oxidation rate for O2was not included in the regression analysis.
of SA occurs for Pt(IV), Hg(II), Au(III), and Ag(1) and to a lesser extent for AgNOZ. Salicylic acid destruction ends after about 25 min for Hg(I1) and Ag(I), possibly corresponding to the depletion of the dissolved metal. In contrast, Pt(1V) reduction only reaches 92% after 40 min (Table I), probably as a result of exhaustion of SA. Only Au(II1) appears to have been present at an initial concentration that balanced the amount of SA initially present. Figure 3 shows initial oxidation rates from the data displayed in Figure 2 (obtained by fitting the data to exponential functions and analytically determining slopes at t = 0) plotted against standard reduction potentials for the corresponding metal reduction reaction (standard electrode potentials from refs 38, 39, 42, and 43). The 1778
rate of SA oxidation for Cd2+could not be included in Figure 3 because it was zero within experimental error. Theory suggests that reduction processes follow the Tafel formulation (44), which states that current (Le., rate) at an electrode depends exponentially on the difference between the electrode potential (-conduction band) and the reaction potential. The resemblance of the Tafel line in Figure 3 and the experimental data therefore suggests that the oxidation of SA is controlled by the rate of reduction. When reduction is slow, most of the electron/ hole pairs simply recombine. The data in Figure 3 may also serve as a guide for predicting which metals can be reduced using Ti02 photocatalysis. Plotting the data in rectilinear coordinates (not shown) suggests that metals reduction requires an overpotential of about 0.8 V since the conduction band of Ti02 is at about -0.5 V (pH 6) and SA oxidation only becomes significant at about +0.3 V. A different catalyst would probably be required to carry out reduction reactions at potentials more negative than this value. Interestingly, oxygen (present initially at about 0.3 mM at our conditions) does not follow the pattern seen for the metals. One explanation is that in addition to consuming conduction-band electrons, the products of oxygen reduction, superoxide, and perhydroxyl radicals feed into and accelerate the oxidative pathways. The error bars in Figure 3 for Pt(IV), Hg(II), and Au(111) were added to account for the complex aqueous chemistry of these species and uncertainty in reduction potential due to the formation of chloride complexes. For Pt: PtC16'PtC14'2Cl- [E" = 0.74V (42)], PtC1& PtO 4C1- [E" = 0.73V (42)1, and Pt2+ PtO [(E" = 1.2V (4211; for Hg: HgClz HgO + 2C1- [E" = 0.41V (2411, HgC1d2- HgO + 4C1- [E" = 0.46V (24)1,and Hg2+ HgO [E" = 0.8OV (42)l; for Au: AuC14- AuO + 4C1- [E" = 1.OV (38)I,andAu3+-Auo [E" = 1.5V (38)l. Furthermore, the reduction rzactions taking place in the pot reactor are influenced by pH, products of metals hydrolysis, surface chemistry, and the presence of salicylicacid. Nevertheless, the ranges indicated in the figure are thought to encompass most possibilities. Also, AgNOz was considered separate from Ag+ derived from AgN03 because of its low dissociation constant [1.2 X lo4 (3811 and independent Ag + NOz-; E" = 0.59V reduction reaction [AgN02 (42)l. It is conceivable that the enhancements in oxidation rate seen for Hg, Ag, Au, and especially Pt arise from catalytic action on the metal crystallites formed on the Ti02 and not from the synergetic effect of simultaneous reduction. To discount this hypothesis, we removed the Pt-doped Ti02 catalyst from one run, rinsed it with water, and then reused it under ambient air without aqueous Pt. Despite this catalyst's black color, it produced a destruction rate very similar to that seen for the fresh Ti02 in ambient 0 2 and lacking any dissolved metal. Similar behavior is expected for the other metals. Another complication arises because SA is known to form stable metal complexes that may be photoactive. However, experiments in the absence of Ti02 verified that SA/metal complexes, if they exist, do not contribute to the observed behavior. Mercury was chosen for further study because of its toxicity and prevalence as a water contaminant. Figure 4 shows the removal of mercury compared with the simultaneous destruction of SA. Oxygen inhibits mercury reduction. In the presence of oxygen, mercury accelerates the destruction of SA (cf.,Figure 2). When no air is present
Environ. Scl. Technoi., Vol. 27, No. 9, 1993
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