Photodegradation of chloroform and urea using Ag-loaded titanium dioxide as catalyst

War. Res. Vol. 25, No. 7, pp. 823-827, 1991 Printed in Great Britain.All rights reserved

0043-1354/91 $3.00+ 0.00 Copyright © 1991 PergamonPress plc

PHOTODEGRADATION OF CHLOROFORM A N D UREA USING Ag-LOADED TITANIUM DIOXIDE AS CATALYST MARCIA M. KONDOand WILSONF. JARDIM* Instituto de Quimica, UNICAMP, CP 6154, 13081 Campinas, SP, Brazil (First received December 1989; accepted in revised form January 1991) AImtraet--The photodegradation of chloroform and of urea in aqueous solutions was investigated under near-u.v, and visible radiation using the anatase form of TiO2 as catalyst. It has been shown that the catalytic activity of the TiO2 can be increased by loading silver, in a proportion of 1% (w/w), onto the oxide surface. In a solution containing 200 mg I - ' of chloroform, the photodegradation achieved using the Ag-loaded oxide was 44% compared to the 35% obtained using the pure (unloaded) oxide. This photoactivity was further enhanced in a solution containing 25 mg 1-~ of chloroform. Upon reuse, a decrease in the photocatalytic activity was observed in both oxides, but to a lesser extent in the Ag-loaded TiO2. An increase in the photocatalytic activity of the silver-loaded oxide was also noted in the degradation of a 100 mg 1-1 aqueous solution of urea. After 12 min of irradiation, the At-loaded oxide destroyed 83% of this compound, compared to a 16% efficiencyobtained for the pure form of anatase. After 6 h of irradiation, 0.5 mg 1-I of silver ions were detected in the At-loaded TiO2 suspension. Key words--photodegradation, chloroform, urea, Ag-loaded TiO2, photocatalyst, semiconductors

INTRODUCTION Water quality control is one of the most important aspects for preserving human lives. Population growth has always been accompanied by a demand for freshwater to attend mainly agricultural, industrial and domestic needs. Considering that in some areas the supply of freshwater may not suffice the growing demand, water reuse has to be considered as an alternative to solve part of this problem. However, as the list of potentially toxic and hazardous substances brought to the aquatic system due mainly to agricultural and industrial activities lengthens, there is a need for effective water treatment processes to assure water quality criteria (Agg et al., 1987). In addition, while attempting to provide drinking water free of pathogenic microorganisms, man can also generate some toxic compounds that were absent in the water before the disinfection step, usually chlorination. During the chlorination, volatile organochloride compounds can be formed due to the reaction of chlorine with naturally occurring organic material such as humic and fulvic acids. Among these compounds, chloroform (CHCI3) is one of the trihalomethanes (THM) that in formed in large amounts (Krasner et al, 1989; Scully et al., 1988; Tan and Wang, 1987; Jodellah and Weber, 1985; Peters et al., 1980). This compound has been proved to be carcinogenic to rats and mice (Cotruvo, 1981). It is believed that the main THM precursors are the polyhydroxi aromatic structures present in humic *Author to whom all correspondence should be addressed.

compounds ubiquitous in natural waters. THMs formed during chlorination have been associated with some parameters such as raw-water dissolved organic carbon (DOC), u.v. absorbance, fluorescence and chlorine demand. However, these are surrogate parameters that do not take into consideration the chemical composition of the naturally occurring organic substances. Bruchet et al. (1990) investigated the relationship between THM formation potential (THMFP) and DOC in natural waters. The authors pointed out that carbohydrates contribute less to T H M F P than proteins and polyphenolic compounds. The World Health Organization (WHO) recommend 0.03 mg 1-~ as a guideline value for CHC13 in drinking water. The U.S.EPA established the maximum contaminant level (MCL) as 0.10mg1-1 for total trihalomethanes (TTHM). in Canada, the maximum acceptable limit for TTHM was established at 0.35 mg 1-1, while the European Economic Community (EEC) recommend the value of 0.001 mg 1-~ for TTHM (Sayre, 1989). Reducing THM formation can be achieved using the following alternatives: (a) replace the traditional chlorination procedure by another disinfectant to guarantee biological standards; (b) carry out a pretreatment to minimize the THM percursores in water: and (c) destroy the THM formed prior to domestic supply (Cotruvo, 1981). Under the chemical, economical and ecotoxicological point of view, the ideal process to destroy any organic contaminant or pollutant is by its total oxidation, since the products are CO2 and H20, non-toxic and ubiquitous substances. Although

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thermodynamically favoured, the complete oxidation of many organic compounds involves kineticaily slow reactions that usually need the use of catalyst to become economically feasible. Heterogeneous photocatalysis is perhaps one of the most promising processes because it can combine total oxidation of organic matter with low cost using solar light as u.v. and visible source of radiation (Matthews, 1986; Ollis, 1985). Semiconductors such as metal oxides (TiO2, ZnO and W O 3) and metal sulphides (CdS) are usually used as the photocatalysts to assist the mineralization of a wide range of organic compounds, including dyes, phenols and organochlorides (Bahnemann et al. 1987; Matthews, 1987; Darwent and Lepre, 1986). When excited by u.v. or visible radiation, semiconductors generate electron/hole pairs. The positive holes can act as strong oxidizing agents (Schiavello, 1987). Although many semiconductors have been shown to assist the photodegradation of organic materials, titanium dioxide, TiO2, has been investigated in detail by many workers (Tan and Wang, 1987; Bahnemann et al., 1987; Matthews, 1986; Pruden and Ollis, 1983; Hsiao et al., 1983; Oliver et al., 1979). In this work, the photodegradation of chloroform (CHCI3) and urea [(NH2)2CO ] were investigated under laboratory conditions using TiO2 as photocatalyst. The photodegradation of both organic compounds was compared in aqueous suspensions of the natural oxide and a Ag-loaded oxide. MATERIALS AND METHODS Reagents Titanium dioxide (anatase 99.9%, gold label, Lot No. 02516HP) was supplied by Aldrich. The determined surface area (BET) was 9.35 m 2g-1. Chloroform (Merck) was used at concentrations of 200 and 25 mg 1-'. Urea (Merck) was used at a concentration of 100 mg 1-t. The Ag-loaded titanium oxide containing 1% (w/w) silver was made according to the following steps. Firstly, to a TiO2 slurry, made by adding 9.2 ml of a 0.1 M solution of AgNO3 to 10g of TiO 2, was added about 10ml of a I% (w/v) solution of Na2CO3. The suspension was dried and then allowed to bake for 6 h at 400°C. The determined surface area (BET) of the Ag-loaded oxide was 8.49m2g-L According to X-ray data, there was no change in the crystal structure of the anatase after the described treatment.

moved was used. The light intensity at the bulk solution varied between 40 and 100 W m -2 (Black-Ray UV-meter, 365 nm). Analytical methods Determination of chloride. Photodegradation of CHC13 was monitored indirectly by measuring the chloride photoproduced according to equation (I), as proposed by Pruden and Ollis (1983a, b) and Hsiao et al. (1983)

CHC13 + H 2 0 + 1/2 O,---,CO2 + 3H + + 3C1-.

RESULTS AND DISCUSSION The comparison between the photocatalytic activity of both catalyst is presented in Fig. 2. When a solution containing 200 mg l -~ of chloroform is irradiated using either the pure or the Ag-loaded oxide, the amount of C H C l 3 degraded, measured as chloride, varies considerably over the time of the experiment (6 h). While the unloaded TiO 2 degraded

N,o-- ~

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Samples Natural and distilled water samples were contaminated with chloroform. For urea, only distilled water was used. The pH was adjusted to 7.0 with a 0.1 M solution of NaOH, and the oxide was added to make a 0.1% (w/v) suspension. For studies of the photodegradation of CHC13 in natural waters, samples were collected from the Taquaral Lake (Campinas, SP) 10 cm below the surface. The samples were filtered using a 0.45/~m membrane (Sartorius). In one of the experiments, part of the organic matter was removed by using Sep-Pak (C~s) columns. The extent of TOC removal after the Sep-Pak treatment was not quantified. Photoreactor The Pyrex photoreactor used in these experiments is shown in Fig. 1. As a near u.v. source a 125W medium pressure mercury lamp (HPL-N) with the glass bulb re-

(1)

At regular intervals, 25 ml aliquots of irradiated sample were collected and filtered in 0.45/~m membrane. The ionic strength was adjusted by addition of 0.5 ml of 5 M NaNO 3 solution and the chloride was measured using a chloride ion-selective electrode (Radelkis OP-CI-071 I-P) in conjunction with a Micronal B375 pH meter. Determination o f CO:. The urea photodegradation was monitored indirectly be measuring the CO2 produced according to a flow injection analysis procedure described elsewhere (Jardim et al., 1989). pH measurements were made using a Ross-8104 electrode (ORION) in a conjunction with a Micronal B375 pH meter. Soluble silver was measured in a 0.1% (w/v) suspension containing only the Ag-loaded TiO 2 using a Perkin-Elmer 5000 AA. Prior to the analysis, 50 ml of the suspension were filtered using 0.45 pm membrane and acidified with HNO 3 to pH 2.0.

A

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Fig. I. The photoreactor used in the experiment showing the water cooled radiation source (A), the sample container (B) and the magnetic stirrer (C).

Photodegradation using Ag-loaded TiO2 80 70 60 "7

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Fig. 2. Photoproduction of chloride in 200 mg 1-l solution of chloroform obtained using fresh Ag-loaded TiO2 (Q) and pure TiO2 (A). The efficiency of both forms after reuse is also shown using the respective open symbols. A typical standard deviation obtained for triplicates is shown. about 35% of the total amount of chloroform originally present in the solution, for the Ag-loaded oxide this amount reached 44%. According to Pruden and Ollis (1983b), chloroform can be completely mineralized using the u.v./TiO: combination. The lower rates of degradation obtained in this work may be due to experimental conditions, specially concerning the design as well as the characteristics and power output of the u.v.

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continuous stirring. Under the same conditions, the adsorption observed in the pure oxide suspension reached 2rag 1-1 of chloride (10%). However, as shown in Fig. 3, this same behaviour was not noticed so markedly in the mineralization of a solution containing 25 mg 1- ~ of chloroform when using the Ag-loaded oxide. Despite different kinetics of chloride photoproduction, after 60 min of irradiation, the activities of both fresh and reused oxide were similar. It was also noted that at lower CHCI 3 levels (25 compared to 200 mg l-t), the photodegradation is more efficient, increasing from 44 to 63% of the total amount of chloroform originally present in the solution. The photodegradation of chloroform in natural waters has also been investigated in this work. According to Zafiriou et al. (1984), natural waters may contain many organic and inorganic compounds that absorb light (chromophores) and consequently contribute to the indirect photolysis of some organochloride compounds. As shown in Fig. 4, the degradation of chloroform is inhibited in the Taquaral Lake water sample when compared to the efficiency obtained using distilled water. Since the process was still inhibited after filtration and even more after the removal of some organic matter (Sep-Pak), the possibility of only adsorbing CHCI3 onto naturally occurring adsorbents as a possible explanation for such behaviour has to be ruled out. It is believed that the naturally occurring organic matter may act as a photosynthesizer in this process. However, when present at higher levels, penetration of light decreases in the bulk of the solution, and consequently the quantum efficiency also decreases.

source.

It is also interesting to note the difference in the kinetic aspects of chloride photoproduction in both solutions. When using the Ag-loaded oxide, the concentration of chloride in the solution was 40 mg 1-1, after 15 min, while in the same time interval chloride was not detected when the irradiated suspension contained the unloaded form of TiO2. The photoactivity of both reused oxides in a 200 mg ! - ' solution of chloroform is also shown in Fig. 2. A decrease in catalytic activity is noticed with both but specially for the unloaded form of the oxide. Compared to the first run, the decrease in the activity of the Ag-loaded oxide was 27%, whereas for the unloaded oxide the decrease was 74%. Since a decrease in the catalytic activity was observed in both cases, one may suggest that it could be due to chloride adsorption at the active sites of the oxide surface. Considering that for TiO2 p H ~ is c. 6.5 (Boehm, 1971), the adsorption of chloride is theoretically favoured as the reaction proceeds due to a pH drop in the reactional medium. Indeed, in a Ag-loaded oxide suspension spiked with 20 mg !- ] of chloride an adsorption of 4 mg 1- t of chloride (20%) was observed after 60min of

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Fig. 3. Photoproduction of chloride in a 25 nag 1- ~solution of chloroform for fresh Ag-loaded TiO2 (O) and the reused oxide (C)). Bars show typical standard deviation obtained for triplicates.

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MARCIAM. KONOOand WILSONF. JARDIM

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Fig. 4. Photoproduction of chloride in distilled water (0), unfiltered lake water (O), filtered lake water (A), and Sep-Pak treated lake water (A). All samples spiked with 25 mgl -I of chloroform and using Ag-loaded TiO2 as catalyst. Bars show typical standard deviation obtained for triplicates. According to equation (1), total mineralization of chloroform (and any other organic compound) can be followed by measuring the CO: produced and trapped in the solution. Matthews (1986) followed the CO2 production in a experiment where 21 different organic compounds including benzene, phenol, catechol and chloroform were photodegraded. The stoichiometry shown in equation (1) was investigated by monitoring the photoproduction of CO2, CI- and H ÷ simultaneously. The results are shown in Table 1. Considering the experimental error associated with the determination of pH ( + 0.05 pH units), there is no significant difference (P = 0.05) between the yield observed for [H + ] compared to [CO:] in Table 1. The discrepances observed for [CI-] compared to [CO2] may be due to either the adsorption of the ionic species onto the oxide or to formation of transient species, as pointed out by Matthews (1986). No attempts were made to identify other chemical species formed in the solution. The pH of the solution at the end of irradiation was 3.32 and 3.39 using Ag-loaded and the pure form of oxide, respectively. According to the Ollis (1985), catalytic activity to TiO2 decreases in acidic solutions, Table I. Photoproduction of C1 , H+ and CO2 after 1h of irradiation* (all valuesin mM) Ag-loadedTiO2 Pure TiO2 [C1 ] [H÷] lEO2] ICl ] [H÷] [CO:] Expected't" 0.63 0.63 0.21 0.63 0.63 0.21 Observed 0.40 0.48 0.17 0.31 0.41 0.15 Yield (%) 63.5 76.2 81.0 49.2 65.1 71.4 *[CHCI3]= 25 mg I-t. t'Accordingto equation(I).

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Fig. 5. Photoproduction of CO2 in a 100 mg 1-l solution of urea using the Ag-loaded TiO2 (O) and the pure TiO2 (©).

pH around 3.5-4.5. However, no changes were observed in the amount of C1- produced when 25 mg 1- ~chloroform solution was irradiated with the initial pH adjusted to 3.0, using either forms of oxide. The data obtained for the photoproduction of CO2 in a 100 mg 1- ~solution of urea using either oxide are presented in Fig. 5. The superior catalytic activity of the Ag-loaded oxide is clearly confirmed by the fact that it takes just 12min to destroy 83% of the original amount of urea added. In the same interval of time, the pure oxide converted only 16% of the urea into CO2. The increase in the activity of the Ag-loaded oxide presented in Figs 2 and 5 could be explained by assuming that transition metals can decrease the semiconductor band gap (3.2 eV for the anatase, according to Zielinski and Sobczynski, 1985), favouring electron transfer from the valence band to the conduction band. Sato et al. (1989) observed an increase in the photoactivity of TiO 2 upon calcination. The authors also noticed a decrease in the photocurrent in a TiO2 film electrode as a function of the increase of the calcination temperature, with a maximum at a temperature near 420°C. In oxides treated at this temperature, the photocurrent action spectra shows a peak current at 335-350 nm. The lamp used in this work as u.v. source shows an emission peak at 366 nm, which is very close to the ideal value pointed out by Sato et al. (1989). When the X-ray data of the pure and the Ag-loaded form of TiO: were compared, no change was observed in the crystalline structure of anatase form, concluding that the metal loading is a typical surface process.

Photodegradation using Ag-loaded TiO2 Silver was detected at 0.5 mg 1 - ~ in solution after 6 h of irradiation in a 0.1% (w/v) suspension of the Ag-loaded oxide without chloroform. Although the Ag ÷ ions may show some photocatalytic properties, no degradation was observed, when 25 m g l - solution of chloroform was irradiated in the presence of 0.5 mg 1-t of silver ions and in the absence of TiO 2. This value is 10 times higher than the maximum contaminant level (MCL) and may require further action when dealing with potable waters (Willey, 1987). CONCLUSIONS The results presented in this paper indicate that the photodegradation of organic compounds using a semiconductor as a catalyst should be considered as a promising alternative to water and wastewater purification. Although the photocatalytic activity of pure TiO2 has been explored for more than a decade (Oliver et al., 1979) for the mineralization of organic compounds such as chloroform, benzene, and polychlorinated biphenyls, and more recently for atrazine and s-triazine herbicides (Pelizzetti et al., 1990), the data obtained in this work show that the TiO2 activity can be further increased when silver is incorporated onto the anatase surface. The Ag-loaded oxide has shown to be more effective than the pure oxide in the degradation of both chloroform and urea. It is also important to emphasize that the experimental conditions used in these experiments are comparable to those found in the environment, specially concerning the u.v. radiation characteristics. Since the modified oxide showed an increase in the photocatalytic activity under these experimental conditions, the possibility of using this semiconductor in conjunction with sunlight as the u.v.-visible source has to be effectively investigated as an alternative low cost process for water and wastewater purification and reuse. Acknowledgements--We would like to thank Dr Carol H. Collins for revising the manuscript. One of us (M.M.K.) was partially funded by FAPESP (Fundaqio de Amparo Pesquisa do Estado de Silo Paulo).

REFERENCES

Agg A. R., Zabel B. A. and Zabel T. F. (1987) Trends in water pollution and control. Wat. Pollut. Control 86, 378-388. Bahnemann D. W., Fisher C. H., Hoffman M. R., Hung A. P., Monig J. and Kormann C. (1987) Mechanistic study of photocatalytic decomposition of organic compounds on semi-conductor particles. 194th Meeting o f American Chemical Society, New Orleans, pp. 528-534. Boehm H. P. (1971 ) Acidic and basic properties of hydroxylated metal oxide surfaces. Discuss. Faraday Soc. 52, 264-275. Bruchet A., Rousseau C. and Mallevialle J. (1990) PyrolysisGC-MS for investigating high-molecualr-weight THM precursors and other refractory organics. J. Am. Wat. Wks Ass. 82, 66-74.

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Cotruvo J. A. (1981) THMs in drinking water. Envir. Sci. Technol. 15, 268-274. Darwent J. R. and Lepre A. (1986) Photo-oxidation of methyl orange sensitized by zinc oxide. J. Chem. Soc., Faraday Trans. 2 82, 1457-1468. Hsiao C. Y., Lee and C. L. and Ollis D. F. (1983) Heterogeneous photocatalysis: degradation of dilute solutions of dichloromethane (CH2C12), chloroform (CHCI3) and carbon tetrachlorine (CC14) with illuminated TiO2 photocatalysts. J. Catal. 82, 418-423. Jardim W. F., Pasquini C., Guimaries J. R. and Faria L. C. (1990) Short-time toxicity test using Escherichia coil: monitoring CO 2 production by flow injection analysis. War. Res. 24, 351-354. Jodellah A. M. and Weber W. J. Jr (1985) Controlling trihalomethane formation potencial by chemical treatment and adsorption. J. Am. Wat. Wks Ass. 77, 95-100. Krasner S. W., McGuire M. J., Jacangelo J. G., Patania N. L., Reagan K. M. and Aieta E. M. (1989) The occurrence of disinfection by-products in US drinking water. J. Am. Wat. Wks Ass. 81, 41-53. Matthews R. W. (1986) Photo-oxidation of organic materials in aqeous suspensions of titanium dioxide. War. Res. 20, 569-578. Matthews R. W. (1987) Solar-electric water purification using photocatalytic oxidation with TiO 2 as a stationary phase. Sol. Energy 38, 405-413. Oliver B. G., Cosgrove E. G. and Carey J. H. (1979) Effect of suspended sediments on the photolysis of organics in water. Envir. Sci. Technol. 13, 1075--1077. Ollis D. F. (1985) Contaminant degradation in water. Envir. Sci. Technol. 19, 480-484. Pelizzetti E., Maurino V., Minero C., Carlin V., Pramauro E., Zerbinati O. and Tosato M. L. (1990) Photocatalytic degradation of atrazine and other s-triazine herbicides. Envir. Sci. Technol. 24, 1559-1565. Peters C. J., Young R. J. and Perry R. (1980) Factors influencing the formation of haloforms in the chlorination of humic materials. Envir. Sci, Technol. 14, 1391-1395. Pruden A. L. and Ollis D. F. (1983a) Photoassisted heterogeneous cataysis: the degradation of trichloroethylene in water. J. Catal. 82, 404-417. Pruden A. L. and Ollis D. F. (1983b) Degradation of chloroform by photoassisted heterogeneous catalysis in dilute aqueous suspension of titanium dioxide. Envir. Sci. TechnoL 17, 628~31. Sato S., Sobczynski A., White J. M., Bard A. J., Campion A., Fox M. A., Maclouk T. E. and Webbcr S. E. (1989) Photochemical properties of ultrathin TiO2 films prepared by chemical vapor deposition. J. Photochem. Photobiol. A: Chem. 50, 283-290. Sayre I. M. (1988) International standards for drinking water. J. Am. War. Wks Ass. 80, 53-60. Schiavello M. (1987) Basic concepts in photocatalysis. In Photocatalysis and Environment--Trends and Applications (Edited by Schiavello M.), pp. 351-360. Kluwer Academic, The Netherlands. Scully F. E. Jr, Howell G. D., Kravitz R. and Jewell J. T. (1988) Proteins in natural waters and their relation to the formation of chlorinated organics during water disinfection. Envir. Sci. Technol. 22, 537-542. Tan C. K. and Wang T. C. (1987) Reduction of trihalomethanes in a water-photolysis system. Envir. Sci. Technol. 21, 508-511. Willey B. R. (1987) Finding treatment options for inorganics. War. Engng Mgmt 137, 2~-31. Zafiriou O. C., Joussot-Dubien J. T., Zepp R. G. and Zika R. G. (1984) Photochemistry of natural waters. Envir. Sci. Technol. 18, 358A-371A. Zielinski S. and Sobczynski A. (1985) Photoevolution of hydrogen from water on modified TiO2, I. Acta chim. Hung. 120, 229-234.