Photocatalytic removal of fenitrothion in pure and natural waters by photo-Fenton reaction

Chemosphere 57 (2004) 635–644 www.elsevier.com/locate/chemosphere

Photocatalytic removal of fenitrothion in pure and natural waters by photo-Fenton reaction Aly S. Derbalah, Nobutake Nakatani, Hiroshi Sakugawa

*

Graduate School of Biosphere Science, Hiroshima University, 1-7-1 Kagamiyama, Higashi-Hiroshima 739-8521, Japan Received 30 September 2003; received in revised form 29 July 2004; accepted 10 August 2004

Abstract The photocatalytic removal kinetics of fenitrothion at a concentration of 0.5 mg l
1 in pure and natural waters were investigated in Fe(III)/H2O2/UV–Vis, Fe(III)/UV–Vis and H2O2/UV–Vis oxidation systems, with respect to decreases in fenitrothion concentrations with irradiation time using a solar simulator. Fenitrothion concentrations were determined by HPLC analysis. Furthermore, total mineralization of fenitrothion in these systems was evaluated by monitoring the decreases in DOC concentrations with solar simulator irradiation time by TOC analysis. It was shown that the degradation rate of fenitrothion was much faster in the Fe(III)/H2O2/UV–Vis system than the Fe(III)/UV–Vis and H2O2/ UV–Vis systems in both pure and river waters. Consequently, the mineralization rate of fenitrothion was much faster in the Fe(III)/H2O2/UV–Vis system than in the other two systems. The high OH generation rate measured in the Fe(III)/H2O2/UV–Vis system was the key to faster degradation of fenitrothion. Increases in the concentrations of H2O2 and Fe led to better final degradation of fenitrothion. These results suggest that the photo-Fenton reaction (Fe(III)/H2O2/UV–Vis) system is likely to be an effective method for removing fenitrothion from contaminated natural waters.
2004 Elsevier Ltd. All rights reserved. Keywords: Pesticides; Photo-Fenton reaction; Contaminated water; Iron; Hydrogen peroxide; Hydroxyl radical

1. Introduction Organophosphorus compounds have recently been used as an alternative to organochlorine compounds for pest control. Widespread use and disposal of these organophosphorus compounds resulted in the release of their residues into natural water, thus inducing an environmental problem. Such chemicals are included in several priority lists of pollutants in Japan and other

*

Corresponding author. Tel./fax: +81 824 24 6504. E-mail address: [email protected] (H. Sakugawa).

countries, owing to their widespread use and high toxicity. Fenitrothion (Fig. 1), an organophosphorus pesticide, is considered to be a common river water pollutant, and is detected with high frequency in river water all over Japan (Numabe et al., 1992; Okumura and Nishikawa, 1995; Itagaki et al., 2000; Kondoh et al., 2001; Tanabe et al., 2001; Sudo et al., 2002; Derbalah et al., 2003). Furthermore, fenitrothion residues in natural water undergo photodegradation, resulting in the release of many toxic metabolites, some more toxic than the parent compound, to aquatic organisms (Eto, 1974; Amoros et al., 2000; Derbalah et al., 2004). Moreover, fenitrothion and its photoproducts are suspected

0045-6535/$ - see front matter
2004 Elsevier Ltd. All rights reserved. doi:10.1016/j.chemosphere.2004.08.025

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A.S. Derbalah et al. / Chemosphere 57 (2004) 635–644

Fig. 1. Chemical structure of fenitrothion.

endocrine disruptors (Tamura et al., 2001; Turner et al., 2002; Derbalah et al., 2004). Subsequently, the presence of fenitrothion and its metabolites in natural water has led to a search for highly effective methods for their removal or decomposition into environmentally compatible compounds. Advanced oxidation processes (AOP) are potentially useful for treating pesticides in wastes because they generate hydroxyl radical (OH), a powerful oxidant. One way to generate OH is by well-known reactions of hydrogen peroxide (H2O2) with Fe2+ and Fe3+ salts (Pignatello, 1992; Arnold et al., 1995). The applications of Fe2+/H2O2 (Fenton reaction) and Fe3+/H2O2 (Fentonlike reaction) systems to hazardous pesticides have been reviewed (Murphy et al., 1989; Pignatello, 1992; Sun and Pignatello, 1993). The advantages of these methods are that they are cheaper than TiO2 particles or O3 generators, and they are straightforward; no special equipment such as complex reaction vessels or an O3 generator is required (Lunar et al., 2000). However, in many cases, the use of H2O2 as the oxidant source requires a stoichiometric amount of Fe2+, thereby requiring a high concentration of these inorganic species. Moreover, under dark (thermal) Fe3+/H2O2 conditions, mineralization is markedly slower and does not go to completion, even with high concentrations of H2O2 (Pignatello, 1992). However, it has recently been discovered that the oxidizing power of Fenton-type systems is greatly enhanced by irradiation with UV or UV–Vis light, and this technique has proved to be very powerful in destroying persistent pesticides such as methylparathion, endosulfan, methamidofos, oxamyl and alachlor (Haag and Yao, 1992; Sun and Pignatello, 1993; Pignatello and Sun, 1995; Penuela and Barcelo, 1996, 1998; Fallmann et al., 1999). The majority of sunlight is visible light; therefore it must be of importance in the environmental field utilizing visible light to degrade pollutants (Xie et al., 2000). Although the Fe3+ ion is known to absorb weakly in the solar UV region (290–400 nm), the absorption spectra of its hydrated and other complexed ion species are shifted toward the visible region (Knight and Sylva, 1975), which might make their use in sunlight possible (Larson et al., 1991). Moreover, in the Fe3+/H2O2 system, the rate of free radical production and, consequently, the rate of contaminant removal, are significantly accelerated by the effective reduction of Fe3+ to Fe2+. In addition, in other cases, formation of organic Fe3+ complexes also accelerates H2O2 decomposition

(Rivas et al., 2002). This means that the use of Fe3+ instead of Fe2+ in the Fenton reaction contributes to the economy of the process since the reduction step of the sludges produced is avoided. In this sense, direct reutilization of the catalyst, Fe3+, (after acid re-dissolution) can be accomplished with the subsequent savings in chemicals (reducing agent) (Rivas et al., 2002). However, as regards pesticide elimination using the Fenton reaction, many of the studies conducted up till now employed pesticide concentrations at mg l
1 levels in water, which is not close to the actual environmental concentrations of the pesticides in water (Pignatello, 1992; Pignatello and Sun, 1995; Huston and Pignatello, 1999; Benitez et al., 2002). Moreover, most of these studies were conducted using pure water; and pure water is not able to closely reflect the efficiency of these methods under natural contaminated water conditions. Thus, application of these methods using a lower concentration of pesticides, which is much closer to the actual environmental levels, and in natural river water, is of special interest. Also, an important measure of the success of AOP is the monitoring of total mineralization of target compounds to CO2, because this signifies the overall destruction of possibly toxic organic intermediates in addition to the destruction of the active ingredient itself (Huston and Pignatello, 1999). Therefore, in this study, we attempted to find effective methods to destruct fenitrothion by using a photocatalytic Fenton reagent, to evaluate the efficiency of this method under a low concentration close to that present in the environment so that the behavior of fenitrothion can more closely reflect what could happen in actual environmental situations, and to evaluate the efficiency of this method in river water under natural conditions.

2. Materials and methods 2.1. Chemicals Fenitrothion (98.5% purity) was obtained from Kanto Chemicals Company, Japan. Ferric chloride (FeCl3 Æ 6H2O) (99% purity) was obtained from Katayama Chemicals, and 30% H2O2 was obtained from Kanto Chemicals Company, Japan. Benzene, acetonitrile and phenol (99.5% purity) were obtained from Nacalai Tesque, Japan, and methanol was obtained from Wako Chemical Company, Japan. Due to the low solubility of fenitrothion in water, a stock solution of 100 mg l
1 fenitrothion was prepared by making appropriate dilutions in methanol (a photochemically inert organic solvent) and stored in a refrigerator at 4 C. Working standard solutions of fenitrothion were prepared by making appropriate dilutions in MilliQ water, and were

A.S. Derbalah et al. / Chemosphere 57 (2004) 635–644

stored in a refrigerator at 4 C. For TOC analysis, fenitrothion was dissolved directly in MilliQ water to avoid the influence of methanol on dissolved organic carbon (DOC) concentration. 2.2. Photodegradation experiments In the photodegradation experiments, a solar simulator (Oriel, Model 81160-1000) unit equipped with a 300 W Xenon lamp (O3 free) and special glass filters (Oriel, AM0 and AM1.0) restricting the transmission of wavelengths below 300 nm was used. River water collected from the Kurose River at the Hinotsume site (Derbalah et al., 2003) in May 2003, of which the chemical composition is given in Table 1, was filtered through a glass fiber filter (GC-50, diameter: 47 mm, nominal rating: 0.5 lm, Advantec) before use. Ferric chloride was used as a source of Fe catalyst because this salt remains unchanged before and after oxidation and this made the study of the reaction and the future engineering scale-up simpler, as the system remains homogeneous (Murphy et al., 1989). The solution was prepared by addition of the desired amount of fenitrothion (0.5 mg l
1) to either filtered river or MilliQ water, and after that the mixture was carefully agitated. Then, freshly prepared FeCl3 at Fe concentration of 50 mg l
1 was added, and finally, H2O2 to a final concentration of 0.05% was added. The initial pH of the solution was adjusted to 2.8 with 1 M hydrochloric acid (HCl) for all experiments (Pignatello, 1992; Huston and Pignatello, 1999). During the irradiation, the solution in the quartz cell (60 ml) was mixed well with a stirring bar and the temperature was kept constant at 20 C. Each experiment was replicated three times. Solutions from the irradiated samples were removed at regular intervals for HPLC analysis. To account for the catalytic reaction, one experiment was carried out in the absence of H2O2 (Fe(III)/UV–Vis). In addition, another experiment was carried out in the absence of catalyst, to account for the degradation of fenitrothion with H2O2 (H2O2/UV–Vis). To account for any possible degradation of fenitrothion under dark conditions, three experiments were carried out in the absence of light, with Fe(III)/H2O2 and with H2O2 and Fe3+ separately. We also carried out experiments to check the Table 1 Chemical composition of the Kurose river water collected for photocatalytic degradation of fenitrothion Analytical item

Concentration level

NO
3 NO
2


66 lM 9 lM 328 lM 160 lM 3.87 mg C l
1 7.0

Cl SO2
4 DOC pH

637

degradation rate of 0.5 mg l
1 fenitrothion in the presence and absence of methanol (0.5 mg l
1), which was used to dissolve the fenitrothion standard in water. We did not find much difference between the degradation rates in the presence and absence of methanol (within ±10%) (data not shown). 2.3. Analytical methods The irradiated samples were analyzed directly by an HPLC system equipped with a pump (LC-10Ai, Shimadzu), a sample injector (Rheodyne Model 9725, sample loop size 50 ll) and a UV–Vis detector (SPD-10A, Shimadzu). A Supelcosil LC-18 (Supelco, particle size 5 lm) 250 mm · 4.6 mm ID column was used. A guard column (Supelcosil LC-18, 5 lm, 20 mm · 4.0 mm ID) was fitted before the analytical column. A mixture of acetonitrile and MilliQ water (60:40) was used as the mobile phase with isocratic elution. The flow rate was maintained at 1.0 ml min
1 and the UV detector wavelength for fenitrothion was 220 nm (Kiso et al., 1996). The losses in DOC with irradiation time were estimated by TOC 5000-A (Shimadzu Company, Japan) for an initial fenitrothion concentration of 3 mg C l
1. After irradiation, water samples were acidified with HCl and then injected directly into the TOC analyzer, which was calibrated with standard solutions of hydrogen potassium phthalate. 2.4. Measurement of OH formation rate Hydroxyl radical were quantified by trapping with added benzene, and measuring the product phenol by high-performance liquid chromatography (HPLC) (Faust and Allen, 1993), using a mobile phase of 60:40 CH3CN:H2O, a Supelcosil LC-18 column (5 lm, 250 mm · 4.6 mm) and a fluorescence detector operating at excitation and emission wavelengths of 270 and 298 nm, respectively. The flow rate was maintained at 1.0 ml min
1. At the beginning of each experiment, an aliquot of aqueous benzene stock solution (20 mM) was added to either pure or filtered river water to give an initial benzene concentration, [benzene]0, of 1.2 mM. The solution was then spiked with Fe3+ and H2O2, agitated to mix the solution, transferred to a quartz cell (60 ml) and irradiated using a solar simulator. Due to the high concentration of Fe3+ (50 mg l
1) and H2O2 (0.05%) under Fenton-like reaction conditions, diluted solution (100 times) was employed to measure the  OH formation rates in both pure and river waters. The initial pH of the diluted solution was adjusted to 2.8 in all experiments (Pignatello, 1992; Huston and Pignatello, 1999). During the irradiation, the solution in the quartz cell (60 ml) was mixed well with a stirring bar and the temperature was kept at 20 C. Solutions were

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removed at regular intervals from the irradiated samples for HPLC analysis. The total concentration of OH produced by the photoreaction in a given time period, [OH]total, was determined as follows: ½ OH ¼ ½Phenol=ðY ÞðF Þ ð1Þ

phenol from OH mediated oxidation of benzene. Thus, significant loss of phenol through mineralization is highly unlikely.

where the phenol concentration [Phenol] is measured as function of irradiation time, Fbenzene,OH is the fraction of  OH that reacts with benzene (in competition with other  OH scavengers), and Yphenol is the yield of phenol formed per benzene molecule oxidized by OH. Fbenzene,OH was assumed to be 1.00 in this study. We note that actual Fbenzene,OH may be a little lower than 1.00 due to possible reactions of OH with other scavengers such as H2O2 and Fe used for the Fenton agents, and dissolved chemical species present in the river water sample used for the photo-Fenton reaction. In addition, as phenol is not the only product of benzene photo-oxidation, the reactions of OH with other photo-oxidized products such as polyhydroxybenzene derivatives (i.e. catechol), and also with phenol oxidation products, possibly scavenge OH. Thus, the formation rate of OH calculated using Eq. (1) might be lower than the real formation rate. However, under our experimental conditions, the molar ratios of H2O2 and Fe to benzene (1.2 mM) are 13% and 0.7%, respectively, while the reaction rate constants of OH with benzene, H2O2 and Fe2+ ion are 7.8 · 109, 2.7 · 107 and 4.3 · 108 M
1 s
1, respectively (Buxton et al., 1988). Moreover, the conversion of benzene to phenol is always limited to <0.2% under the experimental conditions employed in this study. Furthermore, the scavenging effect of dissolved chemical species occurring in river water on OH was examined by Takeda et al. (2004) and Nakatani (2004) and they showed that benzene (1.2 mM) was sufficient to scavenge almost all generated OH (92%) in the presence of dis
solved organic matter, NO
2 , and HCO3 present in the Kurose river water. Therefore, the effect of other chemical species (occurring in the experimental aquatic systems for the photo-Fenton reaction in the present study) on OH scavenging is very limited compared with that of benzene and, for practical purposes, it can be assumed that Fbenzene,OH = 1.00. Arakaki and Faust (1998) experimentally determined the yield of phenol from the benzene–OH reaction in air-saturated aqueous solution at 20 C and showed a phenol yield of approximately 0.75. In our study, [OH]total in Eq. (1) was corrected by a factor of 0.75 (as Yphenol). Losses of phenol under Fenton reaction conditions through mineralization may be expected, however, it has been reported that the destruction rate of phenol by direct photolysis and through oxidation by peroxyl radicals and other oxidants is expected to be slow (Faust and Allen, 1992) in comparison to the formation rate of

3.1. Fenitrothion degradation by photo-Fenton reaction

phenol

3. Results and discussion

benzene;OH

3.1.1. (A)-MilliQ water The first parameter considered in this study was the decrease in fenitrothion concentration with irradiation time, with respect to the half-life and complete degradation. The results shown in Fig. 2 show that the degradation rate of fenitrothion was greatly enhanced by irradiation in a Fe(III)/H2O2/UV–Vis system relative to in systems with Fe(III)/UV–Vis or H2O2/UV–Vis. The half-lives of fenitrothion were 10, 265 and 420 min in Fe(III)/H2O2/UV–Vis, Fe(III)/UV–Vis and H2O2/ UV–Vis systems, respectively. Furthermore, near complete degradation of fenitrothion (92%) was achieved in the Fe(III)/UV–Vis/H2O2 system within 120 min, in comparison to only 34% and 24% for the Fe(III)/UV– Vis and H2O2/UV–Vis systems, respectively, with the same irradiation time (Fig. 2). This great enhancement in the fenitrothion degradation rate in the photo-Fenton reaction system Fe(III)/H2O2/UV–Vis compared to that in other irradiation systems is due to the higher rate of OH generation in this system (105 lM min
1, compared to 10.7 lM min
1 with Fe(III)/UV–Vis and 2.19 lM min
1 with H2O2/UV–Vis) (Table 2). The OH generation rates in these three aquatic systems were calculated from the initial rate measurements. This high rate of OH generation under photo-Fenton like reaction conditions is due to many reasons. Firstly, the photolysis of Fe(OH)2+, which is the predominant species of Fe(III) (Larson et al., 1991), directly generates OH, (Eq. (2)) (Sun and Pignatello, 1993), while the Fe2+ produced from this reaction in the presence of H2O2 may go

100

% Conversion

total

Fe(III)/H2O2/UV-Vis H2O2/UV-Vis Fe(III)/UV-Vis

80 60 40 20 0

0

60

120

180

240

300

360

420

480

540

Irradiation time (min) Fig. 2. Degradation of fenitrothion at initial concentration of 0.5 mg l
1 under Fe(III)/H2O2/UV–Vis, Fe(III)/UV–Vis and H2O2/UV–Vis systems in pure water. H2O2 = 0.05%, Fe = 50 mg l
1 and pH = 2.8.

A.S. Derbalah et al. / Chemosphere 57 (2004) 635–644 Table 2 Calculated hydroxyl radicals generation rate under different oxidation systems in both pure and river waters in the presence and absence of light: H2O2 = 0.05%, Fe = 50 mg l
1, pH = 2.8

Fe(III)/H2O2/UV–Vis Fe(III)/UV–Vis H2O2/UV–Vis Fe(III)/H2O2/Dark Fe(III)/Dark H2O2/Dark

Generation rate of OH (lM min
1) MilliQ water

River water

105 ± 2 10.7 ± 0.20 2.19 ± 0.19 1.20 ± 0.02 ND ND

142 ± 8.5 4.40 ± 0.36 2.02 ± 0.02 4.60 ± 0.09 ND 0.022 ± 0.002

Each experiment was replicated three times (n = 3). ND = not detected.

on to generate additional OH via a Fenton reaction (Eq. (3)) (Pignatello, 1992; Catastini et al., 2002) ð2Þ FeðOHÞ2þ þ hm ! Fe2þ þ  OH þ OH
Fe2þ þ H2 O2 ! Fe3þ þ  OH þ OH


ð3Þ

Secondly, photolysis of H2O2 itself occurs, which leads to the formation of OH (Eq. (4)) (Benitez et al., 2002) H O þ hm ! 2 OH ð4Þ 2

2

Moreover, photodecomposition of complexes of Fe3+ with carboxylates such as oxalate, which are ubiquitously present in natural water, generates Fe2+ (Eq. (5)), which can in turn react with H2O2 to generate additional OH (Eq. (3)) (Pignatello and Sun, 1995) Fe3þ
ðRCO
Þ þ hm ! Fe2þ þ  RCO
ð5Þ 2

3.1.2. Effect of Fe concentration on the degradation rate of fenitrothion The photodegradation kinetics of fenitrothion at various concentrations of Fe were investigated, to estimate their effect on the degradation rate of fenitrothion (0.5 mg l
1). As shown in Fig. 3a, a high concentration of Fe (100 mg l
1) leads to slightly better final degradation of the last 20% of fenitrothion than the use of Fe at a concentration of 50 mg l
1. However, for the first 80% of fenitrothion, the degradation rates at these two concentrations (100 and 50 mg l
1) of Fe were almost identical. In contrast, the lowest concentration of Fe tested (25 mg l
1) leads to a slower rate of fenitrothion degradation than that achieved with Fe concentrations of 50 and 100 mg l
1. This indicates that an increase in Fe concentration resulted in a faster fenitrothion

2

Fig. 2 shows that, after 20 min of irradiation, the degradation rate of remaining fenitrothion was much slower than during the first 20 min in the Fe(III)/H2O2/UV–Vis system. This is due to the low concentration of remaining fenitrothion after 20 min of irradiation (20% of initial concentration), which leads to a lower reaction rate with OH, and scavenging reactions of OH such as photodegradation products of fenitrothion and the reactions represented by Eqs. (6) and (7) exceed the rate of reaction of OH with fenitrothion (Sun and Pignatello, 1993; Catastini et al., 2002; El-Morsi et al., 2002; Wang and Lemley, 2002) Fe2þ þ  OH ! Fe3þ þ OH
ð6Þ  OH þ H O ! HO þ H O 2 2 2 2

et al., 2002). On the other hand, in the Fe(III)/UV–Vis system, there are different sources that generate OH. Firstly, the Fe(OH)2+ complex is the predominant species for generating OH (Eq. (2)) (Larson et al., 1991). In addition, Fe3+ and its hydrated complex (Fe(OH)2+) absorbs light at wavelengths up to 410–500 nm; therefore, the reaction can be carried out efficiently at longer wavelengths than those required for H2O2 (Huston and Pignatello, 1999; Malato et al., 2002). It is concluded that Fe(III) in combination with UV– Vis light and H2O2 is more effective than Fe(III)/UV–Vis or H2O2/UV–Vis systems at degrading fenitrothion. This is consistent with the results obtained by both Haag and Yao (1992) and Doong and Chang (1998).

ð7Þ

The degradation rate of fenitrothion in a H2O2/UV– Vis system was slightly slower than in an Fe(III)/UV–Vis system, due to the lower formation rate of OH in the H2O2/UV–Vis system (Table 2). This low rate of OH generation can be attributed to the fact that OH are generated only through direct photolysis of H2O2, as H2O2 only absorbs weakly above 300 nm (Benitez

100

(a) 80 60 40 % Conversion

Degradation system

639

Fe (25 mg l-1) Fe (50 mg l-1) Fe (100 mg l-1)

20

0

(b) 80 60 40 H2O2 (0.025%) H2O2 (0.05%) H2O2 (0.1%)

20

0 0

20

40

60

80

100

Irradiation time (min) Fig. 3. Effect of Fe concentration (a) and H2O2 concentration (b) on the degradation of fenitrothion (0.5 mg l
1) under Fe(III)/ H2O2/UV–Vis system. Experimental conditions: (a) H2O2 = 0.05% and pH = 2.8, (b) Fe = 50 mg l
1 and pH = 2.8.

A.S. Derbalah et al. / Chemosphere 57 (2004) 635–644

degradation rate, and this finding agrees with that reported previously by Fallmann et al. (1999). All experiments in this study were carried out using initial Fe concentrations of 50 mg l
1, due to the lack of any significant difference in the fenitrothion degradation rate between 50 and 100 mg l
1 Fe. Moreover, the Fe needs to be removed after treatment, thus the use of a much higher Fe concentration (100 mg l
1) is not recommended. Also, the use of a much higher concentration of Fe leads to the self-scavenging of OH by Fe2+ (Eq. (6)) (Catastini et al., 2002; El-Morsi et al., 2002; Wang and Lemley, 2002). 3.1.3. Effect of H2O2 concentration on the degradation rate of fenitrothion The photodegradation kinetics of fenitrothion at various H2O2 concentrations were also studied, to evaluate the influence of H2O2 on the degradation rate of 0.5 mg l
1 fenitrothion. Fig. 3b shows that 0.1% H2O2 leads to slightly better final degradation of fenitrothion than 0.05% H2O2, when degradation of the last 20% of fenitrothion is considered. However, the degradation rate of the first 80% of fenitrothion with 1% H2O2 was slower than that with 0.05% H2O2, due to self-scavenging of OH by H2O2 (Eq. (7)) (El-Morsi et al., 2002; Wang and Lemley, 2002). The rate of fenitrothion degradation at the lowest level of H2O2 tested (0.025%) was slower than those of H2O2 at levels of 0.05% and 0.1%. This result reflects the positive correlation between the concentration of H2O2 and the degradation rate of organic pollutants (Kang et al., 1999; El-Morsi et al., 2002). Therefore, the initial concentration of H2O2 used in this study was 0.05%. This is due to no significant difference between results obtained using this concentration and the 0.1% concentration. Secondly, the use of a much higher concentration of H2O2 sometimes leads to a reduction in the degradation rate of organic pollutants through self-scavenging of OH by H2O2 (El-Morsi et al., 2002). 3.1.4. (B)-River water Photocatalytic removal of fenitrothion dissolved in natural water collected from the Kurose River in May 2003 was evaluated using Fe(III)/H2O2/UV–Vis, Fe(III)/UV–Vis and H2O2/UV–Vis systems. As shown in Fig. 4a, the degradation rate of fenitrothion in river water in an Fe(III)/H2O2/UV–Vis system was much faster than in Fe(III)/UV–Vis and H2O2/UV–Vis systems. The fenitrothion half-lives were 16, 310, and 465 min for Fe(III)/H2O2/UV–Vis, Fe(III)/UV–Vis and H2O2/ UV–Vis systems, respectively. Moreover, fenitrothion was almost completely degraded (93%) within 120 min of irradiation in the Fe(III)/H2O2/UV–Vis system only. The faster degradation rate of fenitrothion with the Fe(III)/H2O2/UV–Vis system than with the Fe(III)/

100

(a)

Fe(III)/H2O2/UV-Vis H2O2/UV-Vis Fe(III)/UV-Vis

80 60 40 20 0

% Conversion

640

(b)

Fe(III)/H2O2 H 2O 2 Fe(III)

80 60 40 20

0

(c)

Fe(III)/H2O2 H 2O 2 Fe(III)

80 60 40 20 0 0

60

120

180

240

300

360

420

480

540

Reaction time (min) Fig. 4. Photocatalytic degradation of fenitrothion (0.5 mg l
1) under UV–Vis light in river water (a) and degradation of fenitrothion (0.5 mg l
1) under dark condition in pure water (b) and river water (c). Experimental conditions: H2O2 = 0.05%, Fe = 50 mg l
1, pH = 2.8.

UV–Vis and H2O2/UV–Vis systems is due to the higher rate of OH generation in the Fe(III)/H2 O2/UV–Vis system (142 lM min
1, compared to 4.40 lM min
1 with Fe(III)/UV–Vis and 2.02 lM min
1 with H2O2/UV–Vis) (Table 2). This high rate of OH generation under photo-Fenton reaction conditions is due to the reasons already described elsewhere in this study. Furthermore, the results show that the degradation rate of fenitrothion with Fe(III)/UV–Vis was slightly faster than for the H2O2/UV–Vis system, which is consistent with the results for the OH formation rate given in Table 2. The results presented in Table 2 indicate that the rate of OH formation in river water was higher than that in pure water for the Fe(III)/H2O2/UV–Vis system, in contrast to the rate of OH formation in other systems. This high rate of OH formation in the Fe(III)/H2O2/UV–Vis system in river water can firstly be attributed to the presence of extra Fe (97 lM) as a natural component of the Kurose River water at the Hinotsume site (Mostofa and Sakugawa, unpublished data). It can also be attributed to the presence of
NO
3 (60 lM), NO2 (9 lM), and dissolved organic

A.S. Derbalah et al. / Chemosphere 57 (2004) 635–644

matter in river water samples (Table 1). This extra
Fe, NO
3 (Eq. (8)) (Mack and Bolton, 1999), NO2 (Eq. (9)) (Mack and Bolton, 1999), and dissolved organic matter (Eq. (10)) (Fukushima et al., 1998) in river water samples leads to the generation of additional OH under light conditions (Skurlatov et al., 1983; Southworth and Voelker, 2003; Derbalah et al., 2004)
 NO
3 þ H2 O þ hm ! NO2 þ OH þ OH

ð8Þ


 NO
2 þ H2 O þ hm ! NO þ OH þ OH

ð9Þ

DOM þ hm ! DOM þ  OH

ð10Þ

In addition, the retarding effects of Cl
and SO2
4 present in the river water at lM levels (Table 1) (Pignatello, 1992) on the rate of OH formation are considered to be negligible due to the very high amounts of OH generated in the Fe(III)/H2O2/UV–Vis system (Table 2). We note that dissolved organic matter such as humic substances is known to scavenge OH, but is also able to generate OH in the presence of light (Mansour et al., 1999; Sakkas et al., 2002). Moreover, it is known that the photochemical properties of humic substance result from complex phenomena and depend on many factors, such as the type of structure (fulvic acid or humic acid) and the origin of each substance (Sakkas et al., 2002). Therefore, the type and origin of dissolved organic matter can have different influences on the photoprocess (Sakkas et al., 2002). Although the rate of OH formation in river water was higher than that in pure water (Table 2), the rate of fenitrothion degradation in the Fe(III)/H2O2/UV– Vis system in natural water was slightly slower than in pure water, especially early on during irradiation (first 30 min) (Figs. 2 and 4a). This may be attributed to the intense competition between the degradable portion of dissolved organic matter in the water body and fenitrothion for OH to react with during the early stage of the reaction (Amon and Benner, 1996; Epling and Lin, 2002; Sakkas et al., 2002). This subsequently resulted in a slightly slower rate of fenitrothion degradation in river water than in pure water. In addition, there is competition between humic substances (that co-exist in river waters) and fenitrothion to absorb light, since humic substances are known to absorb sunlight energy, which reduces the direct photolysis rate of fenitrothion (Epling and Lin, 2002; Sakkas et al., 2002). Furthermore, the degradation rates of fenitrothion in Fe(III)/UV–Vis and H2O2/UV–Vis systems in river water were slightly slower than those in pure water (Figs. 2 and 4a). This finding is consistent with that reported by Larson et al. (1991), due to the reasons described earlier for the Fe(III)/H2O2/UV–Vis system in pure and natural waters.

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3.2. Fenitrothion degradation in the dark 3.2.1. (A)-MilliQ water Degradation of fenitrothion at a concentration of 0.5 mg l
1 was carried out in the absence of light in Fe(III)/H2O2, Fe(III) and H2O2 systems, to determine the dark reactions of these systems. The results presented in Fig. 4b show that the degradation rate of fenitrothion in the Fe(III)/H2O2 system was faster than those of the H2O2 and Fe(III) systems. This is because the Fe(III)/H2O2 system generates OH in the absence of light (Table 2) (Sun and Pignatello, 1993; Bandara et al., 1996; Wu et al., 1999) (Eqs. (3) and (11)), which accelerates the degradation rate of fenitrothion in this system compared to in the H2O2 and Fe(III) systems Fe3þ þ H O ! Fe2þ þ HO  þ H ð11Þ 2

2

2

3.2.2. (B)-River water The degradation rate of fenitrothion at a concentration of 0.5 mg l
1 was carried out in the absence of light for the Fe(III)/H2O2 system as well as the Fe(III) and H2O2 systems, to account for the dark reaction in natural water. The results presented in Fig. 4c show that the degradation rate of fenitrothion in the Fe(III)/H2O2 system was faster than that in the Fe(III) and H2O2 systems. This is attributed to the higher rate of OH formation in the Fe(III)/H2O2 system as opposed to the H2O2 and Fe(III) systems (Table 2). This accelerates the rate of fenitrothion degradation in the Fe(III)/H2O2 system compared to that in the Fe(III) and H2O2 systems. The high rate of OH generation in the Fe(III)/H2O2 system in comparison to that of the H2O2 system can be explained by following reasons: firstly, the Fe(III)/H2O2 system is able to generate  OH through a Fenton-like reaction (Sun and Pignatello, 1993; Bandara et al., 1996; Wu et al., 1999) (Eqs. (3) and (11)). Secondly, addition of a large amount of Fe3+ (a spike) to the Fe(III)/H2O2 system would generate a much higher amount of OH than the H2O2 system alone, even though the H2O2 system may generate OH by reacting with Fe3+ occurring in the Kurose river water. Moreover, Fig. 4b and c showed that the degradation rates of fenitrothion in river water were slightly faster than in pure water for the Fe(III)/H2O2 and H2O2 systems. In the case of the H2O2 system, this is due to the presence of Fe3+ with an average concentration of 97 lM in the Kurose River water in addition to the spiked amount of H2O2. For the Fe(III)/H2O2 system, the presence of the spiked Fe, in addition to the natural level of Fe in river water and the spiked H2O2, is responsible for the higher rate of OH generation through a Fenton-like reaction than in MilliQ water (Table 2).

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% Conversion

100 80

Fe(III)/H2O2/UV-Vis H2O2/UV-Vis Fe(III)/UV-Vis

60 40 20 0

0

100

200

300

400

500

600

700

Irradiation time (min) Fig. 5. Loss of DOC concentration in photocatalytic mineralization of fenitrothion at initial DOC concentration of 3 mg C l
1 under Fe(III)/H2O2/UV–Vis, Fe(III)/UV–Vis and H2O2/UV–Vis systems in pure water. H2O2 = 0.05%, Fe = 50 mg l
1 and pH = 2.8.

3.3. Total mineralization of fenitrothion To confirm the total mineralization of fenitrothion in the Fe(III)/UV–Vis/H2O2, Fe(III)/UV–Vis and H2O2/ UV–Vis systems in pure water, decreases in the DOC concentration with increasing irradiation time were estimated over 10 h of irradiation, from an initial DOC concentration of 3 mg C l
1. Fig. 5 shows that the mineralization rate of fenitrothion in the Fe(III)/H2O2/UV– Vis system was much faster than in the Fe(III)/UV–Vis and H2O2/UV–Vis systems. Furthermore, the fenitrothion mineralization rate was slightly faster in the Fe(III)/ UV–Vis system than the H2O2/UV–Vis system. These results for the fenitrothion mineralization rate obtained with the three different irradiation systems are consistent with the OH formation rate data presented in Table 2. After 2 h of irradiation, the mineralization rate of the remaining fenitrothion in the Fe(III)/UV–Vis/H2O2 system was, however, much slower than that achieved during the first 2 h. This is due to the low concentration of fenitrothion remaining after 2 h of irradiation, which leads to competition between fenitrothion and the Fenton reagent (and fenitrothion degradation products) for the OH (Eqs. (6) and (7)) (El-Morsi et al., 2002; Wang and Lemley, 2002). Therefore, near-complete mineralization of fenitrothion (93%) was achieved after 10 h of irradiation in the Fe(III)/H2O2/UV–Vis system, while 60% of the compound was mineralized within the first 2 h of irradiation. These results reflect the superiority of the Fe(III)/H2O2/UV–Vis system over the other systems for the photocatalytic removal of fenitrothion in contaminated water.

4. Conclusions The results obtained in this study clearly indicate that the photo-Fenton reaction is very promising for the

complete removal of fenitrothion from contaminated agricultural waters, especially considering the much lower concentration of fenitrothion present in natural water in comparison to the concentration used in this study. The advantage of this technique is that it provides the possibly of another approach, using natural sunlight, to the treatment of water contaminated with fenitrothion. Moreover, the use of sunlight instead of artificial light for the photo-Fenton reaction would dramatically lower the costs of the process, and thus provides a major step towards industrial application. However, the disadvantage of this method is its sensitivity to pH and the need to remove the Fe after treatment. This can be overcome by immobilization of Fe ions on Nafion membranes (Maletzky et al., 1999), so it is now possible to carry out the photo-Fenton reaction over a wider pH range without additional separation of Fe after the treatment process.

Acknowledgment The authors are grateful to Mr. Khan Mostofa for his help with the DOC analysis.

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