Hydroxyl radical involvement in the decomposition of hydrogen peroxide by ferrous and ferric-nitrilotriacetate complexes at neutral pH

w a t e r r e s e a r c h 4 5 ( 2 0 1 1 ) 3 3 0 9 e3 3 1 7

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Hydroxyl radical involvement in the decomposition of hydrogen peroxide by ferrous and ferric-nitrilotriacetate complexes at neutral pH Yen Hai Dao, Joseph De Laat* Universite´ de Poitiers, Laboratoire de Chimie et Microbiologie de l’Eau (CNRS UMR 6008), Ecole Nationale Supe´rieure d’Inge´nieurs de Poitiers, 40, Avenue du Recteur Pineau, 86 022 Poitiers Cedex, France

article info

abstract

Article history:

The relative rates of degradation of three hydroxyl radical probe compounds (atrazine,

Received 12 December 2010

fenuron and parachlorobenzoic acid (pCBA)) by FeIII/H2O2 (pH ¼ 2.85), FeIIINTA/H2O2

Received in revised form

(neutral pH), FeII/O2, FeIINTA/O2, FeII/H2O2 and FeIINTA/H2O2 (neutral pH) have been

16 March 2011

investigated using the competitive kinetic method. Experiments were carried out in batch

Accepted 21 March 2011

and in semi-batch reactors, in the dark, at 25  C. The data showed that the three probe

Available online 29 March 2011

compounds could be degraded by all the systems studied, and in particular by FeIINTA/ H2O2 and FeIIINTA/H2O2 at neutral pH. The relative rate constants of degradation of the

Keywords:

three probe compounds obtained for all the systems tested were identical and equal to

Oxidation

1.45  0.03 and 0.47  0.02 for kAtrazine/kpCBA and kFenuron/kpCBA, respectively. These values

Fenton-like reactions

as well as the decrease of the rates of degradation of the probe compounds upon the

Competitive kinetics

addition of hydroxyl radical scavengers (tert-butanol, bicarbonate ions) suggest that the

Atrazine

degradation of atrazine, fenuron and pCBA by FeIINTA/O2, FeIINTA/H2O2 and FeIIINTA/H2O2

Fenuron

is initiated by hydroxyl radicals. ª 2011 Elsevier Ltd. All rights reserved.

Parachlorobenzoic acid Tert-butanol Bicarbonate ion

1.

Introduction

Advanced Oxidation Processes (AOPs) can be successfully used in the field of wastewater treatment to reduce the chemical oxygen demand and the toxicity of industrial wastewaters, to convert toxic and biorecalcitrant contaminants into biodegradable by-products, to remove colour or to obtain a complete mineralization of organic pollutants (Pera-Titus et al., 2004). Among the AOPs, the Fenton and the Fenton-like oxidation processes (FeII/H2O2, FeIII/H2O2) have been applied to many industrial wastewaters such as chemical, petrochemical, pharmaceutical and textile effluents (Pignatello et al., 2006; Bautista

et al., 2008). These oxidation processes are also popular methods for the remediation of contaminated soils and groundwater (Bergendahl et al., 2003; Watts and Teel, 2005). The main benefits of the Fenton and Fenton-like reactions are the use of environmentally friendly and low cost reagents. However, these oxidation processes have also some limitations. The FeII/H2O2 and FeIII/H2O2 systems must be operated at low pH values (pH z 3) to prevent the precipitation of ferric oxyhydroxides and to produce hydroxyl radicals. In addition, the efficiency of the Fenton reaction can be markedly decreased in the presence of high concentrations of inorganic salts such as chloride and sulfate ions because of the formation of inactive

* Corresponding author. Tel.: þ33 5 49 45 39 21; fax: þ33 5 49 45 37 68. E-mail address: [email protected] (J. De Laat). 0043-1354/$ e see front matter ª 2011 Elsevier Ltd. All rights reserved. doi:10.1016/j.watres.2011.03.043

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w a t e r r e s e a r c h 4 5 ( 2 0 1 1 ) 3 3 0 9 e3 3 1 7

iron(III)-complexes (De Laat and Le, 2005, 2006). To overcome these drawbacks, the use of organic and inorganic ironchelating agents has been investigated by several authors in order to enhance the efficiency of the homogenous Fenton-like oxidation processes at neutral pH (Sun and Pignatello, 1992, 1993; Li et al., 2007; Lee and Sedlak, 2009; Rastogi et al., 2009) or of heterogenous systems involving ion bearing minerals (Xue et al., 2009) or granular zero-valent iron (Keenan and Sedlak, 2008; Lee et al., 2008). In their pioneering works, Sun and Pignatello (1992) have tested several classes of iron(III)-chelates for the decomposition of H2O2 and of 2,4-dichlorophenoxyacetic at pH 6. Of the 50 compounds tested, nitrilotriacetic acid (NTA) was one of the most active iron(III)-chelate. The iron-NTA/H2O2 system was also found to be effective for the degradation of tetrachloroethene in contaminated soils (Howsawkeng et al., 2001; Ndjou’ou et al., 2006). Kim and Kong (2001) showed that the FeIIINTA/H2O2 system degraded more efficiently 1-hexanol and carbon tetrachloride at pH 9 than at pH 3 and that the degradations of 1-hexanol and carbon tetrachloride are initiated by hydroxyl radical (HO) and by superoxide anion radical (HO2/O2), respectively. The nature of the reactive oxidant species (free and bound HO radicals, high-valent-oxoiron species) generated by the reaction of H2O2 with free and complexed Fe(II) and Fe(III) species in aqueous solution remains a controversial issue (Walling et al., 1975; Yamazaki and Piette, 1991; Bossmann et al., 1998; Pignatello et al., 2006). In the absence of chelating agents, it is generally accepted that the reaction of H2O2 with Fe2þ at low pH yields hydroxyl radicals whereas other oxidants such as the ferryl ion may be produced at circumneutral pH (Gallard et al., 1998; Rivas et al., 2001; Hug and Leupin, 2003; Keenan and Sedlak, 2008; Katsoyiannis et al., 2008). In all these studies, proofs of the formation of one or another oxidizing species are based on EPR spin-trapping techniques, distribution of oxidation by-products, effects of added hydroxyl radical scavengers on reaction rates or on kinetic modeling studies. In this work, the competitive kinetic method has been applied to the degradation of three hydroxyl radical probes (atrazine, fenuron and parachlorobenzoic acid) in order to highlight the formation of hydroxyl radicals by the FeIINTA/ H2O2 and FeIIINTA/H2O2 systems at neutral pH. The degradation of probe compounds by the FeII/O2 and FeIINTA/O2 systems has also been investigated because dissolved oxygen readily oxidizes ferrous ion (King et al., 1995; Rose and Waite, 2002) and FeIINTA (Harris and Aisen, 1973; Welch et al., 2002) at neutral pH. The effects of the addition of HO radical scavengers (tert-butanol or bicarbonate ions) on the degradation rates of probe compounds and of H2O2 will also be examined. However, the effects of experimental parameters on the rate of decomposition of H2O2 by FeIINTA and FeIIINTA will be presented in a next paper and will not be discussed here.

2.

Materials and methods

2.1.

Reagents and preparation of solutions

All chemicals used in this work were of reagent grade or higher and were used without further purification. All aqueous solutions were prepared with Milli-Q water.

The stock solution of pCBA (100 mM) was prepared at pH 10.5. The stock solution of fenuron was prepared with a precise concentration (0.5 mM). The stock solution of atrazine was prepared by dissolving the required amount of atrazine in water. After a mixing time of about 48 h in the darkness, the solution was filtered through a 0.45 mm filter and the concentration of dissolved atrazine in the filtrate was determined by HPLC. Stock solutions of Fe(III) (10 mM) were prepared by dissolving the required amount of iron(III) perchlorate (Fe(ClO4)3, 9 H2O) in HClO4 0.1 M. Since iron(III) perchlorate has a great tendency to hydrate, its degree of hydration was checked regularly by spectrophotometric titration using 1,10-orthophenanthroline. Solutions of FeIIINTA were prepared by mixing appropriate volumes of stock solutions of ferric perchlorate (10 mM) and of sodium nitrilotriacetate (5e7 mM). The pH was then adjusted with NaOH 1 M. All FeIIINTA solutions were freshly prepared each time before use. Aqueous solutions of FeII and FeIINTA at neutral pH are extremely sensitive to dissolved oxygen and must be prepared in oxygen-free water and handled under a protective nitrogen atmosphere. For the FeII/O2 and FeII/H2O2 oxidation experiments, a stock solution of Fe(II) (50.0  0.5 mM) was prepared by dissolving the appropriate weight of iron(II) perchlorate hydrate (Fe(ClO4)2, 7 H2O) in oxygen-free pure water. For the stock solution of FeIINTA (50.0  0.5 mM; FeII:NTA molar ratio ¼ 1:3), 1.435 g of NTA powder (7.5 mmol) was dissolved into 48.0  0.5 mL of water. After dissolution, the NTA solution was deoxygenated by nitrogen bubbling for at least 20 min. Then, 0.98 g (2.5 mM) of iron(II) perchlorate hydrate (Fe(ClO4)2, 7 H2O) was added to the solution and a pale green solution was obtained. A nitrogen flux was then kept over the solution during the experiments.

2.2.

Experimental conditions

All the experiments were carried with initial concentrations of atrazine, fenuron and pCBA of 5 mM and were conducted in the dark and at 25.0  0.5  C. All the oxidation experiments conducted with FeIINTA or FeIIINTA complexes have been performed by using [NTA]0:[FeII]0 and [NTA]0:[FeIII]0 molar ratios of 3:1. Under these conditions, FeII and FeIII were present only as FeIINTA and FeIIINTA complexes at the beginning of the reaction (Motekaitis and Martell, 1994). It should also be noted that uncomplexed NTA is therefore always present in the solutions and shall compete with atrazine, fenuron and pCBA for the reaction with reactive species generated in solution. Oxidation experiments with the FeIII/H2O2 (pH ¼ 2.85) and the III Fe NTA/H2O2 (7 < pH < 9) were carried out using a batch reactor. The batch reactor consisted of a 1.3-L cylindrical double-wall jacketed reactor in order to circulate thermostated water with an external circulating pump connected to a thermostated water bath (25.0  0.5  C) (Fig. S1 in Supplementary material). The reactor was covered by a black plastic film to protect the aqueous solution from ambient light. The reactor was filled with 1.25 L of solution containing FeIII or FeIIINTA and the three organic solutes. The solution was mixed using a magnetic stirrer at nearly 800 rpm during all the course of the reaction. Initial pH was adjusted to the desired value using HClO4 or NaOH. During the course of the reaction, the pH was kept constant using a pH

w a t e r r e s e a r c h 4 5 ( 2 0 1 1 ) 3 3 0 9 e3 3 1 7

transmitter (OPM 223, Endress þ Hauser) and a peristaltic pump (Gilson Minipuls 3) for the injection of NaOH 1 M (liquid flow rate : 10e15 mL/h). The reaction was initiated by adding H2O2 into the reactor. At regular time intervals, samples were withdrawn from the reactor and immediately analyzed for H2O2 (1 or 2 mL samples) or immediately quenched with 1 mL of methanol (for 1 mL sample) for the analyses of HO probe compounds. Because of the very fast reactions between O2 or H2O2 with free FeII or FeIINTA at neutral pH, concentrationetime profiles for HO probe compounds could not be obtained using a batch reactor for the FeII/(O2 or H2O2) and the FeIINTA/(O2 or H2O2) systems. A series of experiments were carried out using a semi-batch reactor with a continuous introduction of iron(II) perchlorate or of iron(II)-NTA. The reactor was equipped with a pH electrode connected to a pH transmitter (OPM 223, Endress þ Hauser) and a WTW CellOx 325 dissolved oxygen sensor connected to a calibrated oxygen meter (WTW Oxi 340i-A/SET). A volume of 1.25 L of a solution containing the three HO probe compounds (pH ¼ 7.0) was introduced into the reactor. For the FeII/O2 and the FeIINTA/O2, the solution was oxygenated by bubbling air at a flow rate of nearly 1 L/min. When the equilibrium concentration of dissolved oxygen in water was reached (z8.3 mg/L), the gas bubbling was stopped and all the valves were closed to avoid external oxygen contact during the reaction. The stock solution of FeII(ClO4)2 or of FeIINTA (50 mM) was then injected at time t ¼ 0, at a constant flow rate (z25 or z50 mL/h) with a Gilson Minipuls 3 peristaltic pump. During the course of the reaction, the pH was kept constant at pH 7.0  0.1 using the pH transmitter and a peristaltic pump for the injection of NaOH (0.1e1 M) or of HClO4 (0.1e1 M) (liquid flow rate  10e15 mL/h). For the FeII/H2O2 and the FeIIINTA/H2O2 oxidation experiments, the solution containing the probe compounds (5 mM each) and H2O2 (0.5 mM) was either deoxygenated by bubbling nitrogen gas ([O2] in water < 0.1 mg O2/L) or aerated with air ([O2] z 8.3 mg O2/L) during at least 30 min before injecting the stock solution of FeII(ClO4)2 or FeII/NTA (50 mM) at a constant flow rate (z25 or 50 mL/h). During the course of the reaction, the pH and the concentrations of dissolved oxygen were noted every 30 or 60 s. Samples were taken from the reactor at various reaction times and immediately analyzed without filtration for H2O2, Fe(II) and for total Fe concentrations. Samples were also withdrawn from the reactor and immediately quenched with methanol (1 mL methanol for 1 mL sample), filtered through 0.45 mm pore size membrane filters and then analyzed by HPLC for the determination of the concentrations of HO probes. Batch experiments have also been carried out with the FeII/ H2O2 and FeIINTA/H2O2 systems. For these experiments, 100 mL of a stock solution of FeII(ClO4)2 (50 mM) or FeIINTA (50 mM, FeII:NTA molar ratio ¼ 1:3) were introduced into a series of flasks containing 100 mL of solution of probe compounds (5 mM each) and H2O2 at concentrations ranging from 0 to 1 mM ([H2O2]0 ¼ 0, 10, 25, 50, 80, 100, 150, 200, 500 and 1000 mM). The initial pH before the introduction of FeII or FeIINTA was equal to 7.0. The addition of the FeII or FeIINTA solution was done under vigorous mixing. After a reaction time of 3 h (residual H2O2 ¼ 0 mM, and complete oxidation of FeII into FeIII), the pH of the solutions was determined and the solutions quenched with methanol (100 mL/mL of sample) to stop the degradation of the probe compounds.

2.3.

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Analytical methods

The concentration of hydrogen peroxide in stock solutions of H2O2 was determined by iodometric titration. Concentrations of hydrogen peroxide in solutions containing Fe(II) or Fe(III) were determined spectrophotometrically using the TiCl4 method (Eisenberg, 1943) and a molar absorption coefficient of 724 M1 cm1 for the titanium peroxocomplex. The concentrations of Fe(II) or of Fe(III) (after reduction of Fe(III) by hydroxylamine hydrochloride) were measured by the o-phenanthroline colorimetric method and by using a molar extinction coefficient of 1.105 104 M1 cm1 at 510 nm for the Fe(II)ephenanthroline complex (Tamura et al., 1974). pH measurements were made with a Meter Lab PHM 240 pH meter calibrated with standard buffers. Dissolved oxygen was analyzed using a calibrated oxygen meter (WTW Oxi 340i-A/SET). Atrazine, fenuron and pCBA were analyzed by HPLC using a Waters HPLC system equipped with a Waters 717 plus auto sampler, a Waters 996 photodiode array detector, a Waters Millennium software, a Waters Spherisorb C8 column (5 mm, 4  250 mm). A mobile phase consisting of 50% methanol and 50% H2O (pH 2, acidification with CF3COOH) was used at a flow rate of 0.8 mL/min. Absorbance was measured continuously in the range of 200e400 nm. Fenuron and pCBA, peaks were normally quantified at 240 nm and atrazine at 220 nm.

2.4.

Competitive kinetic expression

Competitive kinetic experiments using three probe compounds for HO radicals have been conducted to ascertain the formation of hydroxyl radicals. By assuming that the degradation of the model compounds are only due to an attack by the hydroxyl radical and by using pCBA as the reference compound, the relative degradation rates of atrazine and fenuron should be described by the following competitive kinetic expressions: ½Atrazinet kAtrazine ½pCBAt ¼ ln ln ½Atrazine0 kpCBA ½pCBA0

(1)

½Fenuront kFenuron ½pCBAt ln ¼ ln ½Fenuron0 kpCBA ½pCBA0

(2)

where kAtrazine, kFenuron and kpCBA are the absolute rate constants for the reaction of HO radicals with atrazine, fenuron and pCBA respectively. Previous competitive kinetic studies carried out with the Fenton reaction (Fe2þ/H2O2, pH ¼ 3) showed that kFenuron/kpCBA ¼ 1.47  0.04 (Acero et al., 2002). A value of 2.92  0.04 has been determined for kFenuron/katrazine from H2O2/UV experiments (254 nm, pH 7.2) (Mazellier et al., 2007). From these values, a value of 0.50  0.03 can be calculated for kAtrazine/kpCBA. By using a value of 5.0 109 M1 s1 for the reaction rate constant of HO radicals with parachlorobenzoate anion (Elovitz and von Gunten, 1999) and by assuming that the rate constants for the reaction of HO radicals with the acid and the basic forms of pCBA (pKa of pCBA z 4) are identical (kpCBA ¼ 5.0 109 M1 s1), the absolute second-order rate constants kAtrazine and kFenuron should be equal to 2.4 109 M1 s1 and 7.0 109 M1 s1, respectively. If the plots of ln([A]t/[A]0) against ln([B]t/[B]0) (A ¼ atrazine or fenuron, B ¼ pCBA) obtained from different experimental conditions or from various oxidation processes yield

3312

w a t e r r e s e a r c h 4 5 ( 2 0 1 1 ) 3 3 0 9 e3 3 1 7

superimposable straight lines with a slope equal to the expected kA/kB value, one can therefore say that the degradation of the model compounds can be attributed to the hydroxyl radical.

a

1 III

Fe /H2 O 2 ; pH 2.85 Tert -Butanol : 50 mM (Exp.A4)

Results

For the concentrations of reactants used in the present work, control experiments showed that, in the dark, the HO probe compounds were not degraded by H2O2 alone, by FeIII alone (acidic pH), by FeIIINTA alone (acidic and neutral pH), by FeII or by FeIINTA alone in deoxygenated solutions at neutral pH (data not shown). In addition, coagulation experiments carried out with iron(III) perchlorate (0e1 mM) and pH showed that the maximum losses of HO probe compounds by adsorption on the iron(III) hydroxide surface were less than 3% and 5% for a total doses of FeIII equal to 0.25 mM and 1 mM, respectively. Degradation of probe compounds was only observed in the presence of H2O2 for solutions containing iron species (FeII, FeIII, FeIINTA, FeIIINTA) or in the presence dissolved oxygen for solutions containing ferrous species (FeII and FeIINTA).

3.1. Oxidation by FeIII/H2O2 (pH 2.85) and FeIIINTA/ H2O2 (neutral pH)

0.6

Atrazine pCBA III

0.4

Fenuron

Fe /H2 O 2 ; pH 2.85 Tert -Butanol : 0 mM (Exp. A1-A3)

0.2

0

5000

10000

15000

Time (s)

b

1 III

Fe /H2 O2 ; pH 2.85 Tert -Butanol : 0 mM (Exp. A1-A3)

0.8

- ln (Ct / C 0 )

3.

Ct / C0

0.8

Fenuron y = 1.481x

0.6 0.4

Atrazine

0.2

y = 0.462x

Figs. 1 and 2 present data obtained for the degradation of the three HO radical probes by FeIII/H2O2 at pH 2.85 and by FeIIINTA/H2O2 at neutral pH. The experimental conditions and the relative rate constants kAtrazine/kpCBA and kFenuron/kpCBA obtained for all experiments carried with these two oxidation processes have been reported in Table 1. For each experiment, detailed data can be found in the Supplementary material. Identical initial concentrations of FeIII (0.2 mM) and H2O2 (0.5 mM) were used in order to compare the rates of degradation of the organic compounds by FeIII/H2O2 and FeIIINTA/H2O2. As shown the data in Fig. 1b, competitive oxidation experiments carried out with the classical Fenton-like reaction (FeIII/ H2O2, pH ¼ 2.85, Experiments A1eA3 in Table 1) yielded relative rate constant values (kAtrazine/kpCBA ¼ 0.47  0.01 and kFenuron/ kpCBA ¼ 1.48  0.01; mean value of three experiments) which are consistent with the expected values for a HO radical-initiated mechanism of degradation of organic compounds by the Fenton-like reaction. In addition, the degradations of the three organic solutes were totally inhibited in the presence of 50 mM of tert-butanol (Fig. 1a). If HO radicals are generated in solutions, they can react with the probe compounds with secondorder rate constants (ki) roughly equal to 5 109 M1 s1, H2O2 (ki z 3 107 M1 s1), tert-butanol (ki z 6 108 M1 s1, Buxton et al., 1988), free NTA (ki ¼ 6.1 107, 5.5 108 and 4.2 109 M1 s1 at pH 2, 6 and 10, respectively (Sahul and Sharma, 1987)) and with FeIINTA (minor reaction, ki unknown). A comparison of the values of the scavenging term ki Ci (ki : rate constant for the reaction of HO with a solute i, and Ci concentration of the solute i) calculated for tert-butanol (6 108  5 102 ¼ 3 107 s1), a probe compound (z5 109  5 106 z 2.5 104 s1), H2O2 (z3 107  5 104 z 1.5 104 s1) and free NTA (<5 108  2 103 ¼ < 1 106 s1 at pH < 3) indicate that more than 97% of HO radicals would be scavenged by tert-butanol under our conditions. All these data confirm that the hydroxyl radical is the main

0

0

0.2

0.4

0.6

0.8

- ln([pCBA] t / [pCBA]0 ) Fig. 1 e Oxidation of probe compounds by FeIII/H2O2 at pH 2.85 in the absence and in the presence of tert-butanol. (a) Normalized concentrationetime profiles of atrazine, fenuron and pCBA in the absence and in the presence of 50 mM tert-butanol; (b) Plots of the competitive kinetic equation for the determination of the relative rate constants (Experiments A1eA4 in Table 1).

oxidant responsible for the degradation of organic compounds by the FeIII/H2O2 at acidic pH. Fig. 2 presents typical data obtained for the FeIIINTA/H2O2 system in the pH range 7e9 and with a molar FeIII:NTA ratio of 1:3 ([FeIII]0 ¼ 0.2 mM, [NTA]0 ¼ 0.6 mM). At neutral pH, the FeIII/ H2O2 is not efficient for the decomposition of H2O2 and for the oxidation of organic compounds because of the formation of non-soluble and inactive iron(III) hydroxides. As shown the data in Fig. 2, the addition of NTA leads to the formation of soluble FeIIINTA complexes which can decompose H2O2 (Fig. 2a) and degrade the model organic compounds (Fig. 2b) at neutral pH. The decomposition rate of H2O2 by FeIIINTA was faster at pH 8 than at pH 7 or 9, in agreement with previous data obtained by Tachiev et al. (2000) for the study of the effects of pH on the initial rates of decomposition of H2O2 by several iron(III)-complexes with aminopolycarboxylic acids including NTA. For an iron(III) complex such as FeIIINTA, Tachiev et al. (2000) showed that the decomposition rate of H2O2 is strongly dependent on the stability constants for the formation of the various FeIIINTA complexes and on the reactivity of the various FeIIINTA complexes toward H2O2.

w a t e r r e s e a r c h 4 5 ( 2 0 1 1 ) 3 3 0 9 e3 3 1 7

a

1

[H 2 O 2 ]t / [H2O 2 ]0

pH = 7.0 0.8

pH = 8.0 pH = 9.0

0.6 0.4 0.2 0 0

1000

2000

3000

4000

5000

Time (s)

b

1

Tert -Butanol: 10 mM ; pH 8.3

[C] t / [C] 0 (Fenuron)

-

HCO3 : 50 mM ; pH8.3

0.8 -

HCO 3 : 25 mM ; pH8.3

0.6

-

pH 7.0

0.4

HCO 3 : 10 mM ; pH 8.3

pH 9.0

0.2 pH 8.0

0

0

1000

2000

3000

Time (s)

- ln ([C] t / [C] 0 )

c

III

Fe -NTA/H2O 2 7 < pH < 9 NaHCO3 : 0 - 50 mM

2 1.6

Fenuron y = 1.433 x

1.2 0.8

Atrazine y = 0.460 x

0.4 0

0

0.5

1

1.5

2

2.5

-ln([pCBA]t / [pCBA] 0 ) Fig. 2 e Oxidation of probe compounds by FeIIINTA/H2O2 at neutral pH. (a) Normalized concentrationetime profiles of H2O2 at pH 7, 8 and 9; (b) Normalized concentrationetime profiles of fenuron in the absence and in the presence of tertbutanol (10 mM) or of bicarbonate ion (10e50 mM); (c) Plots of the competitive kinetic equation for the determination of the relative rate constants (Experimental conditions are given in Table 1, Experiments B1eB11 and C1eC4).

The degradation rates of model organic compounds were also faster at pH 8 than at pH 7 or 9 and decreased in the presence of tert-butanol (10 mM) or of bicarbonate ions (10e50 mM) (Fig. 2b). Under the conditions used, the inhibition of the degradation rates of the HO probes was greater in the

3313

presence of tert-butanol than of bicarbonate ion, in agreement with the values of the scavenging terms for tert-butanol (ki Ci ¼ 6.0 106 s1) and for bicarbonate ions (ki Ci ¼ 1.2 105, 3.0 105 and 6.0 105 s1 for Ci ¼ 10, 25 and 50 mM and pH ¼ 8.3). The apparent second-order rate constant for the reaction of HO radicals with bicarbonate ion (ki) at pH 8.3 is equal to 1.2 107 M1 s1 and has been calculated by using pKa values of 6.35 and 10.33 for H2CO3 and second-order rate constants of 0, 8.6 106 and 3.9 108 M1 s1 for the reaction of HO radicals with H2CO3, HCO3 and CO32, respectively (Buxton et al., 1988). The plots of ln([Atrazine]0/[Atrazine]t) and of ln([Fenuron]0/ [Fenuron]t) against ln([pCBA]0/[pCBA]t) yield straight lines and the slopes of the straight lines were not affected by the pH of the solutions and by the concentration of bicarbonate (Fig. 2c). The mean values of the ratios kAtrazine/kpCBA and kFenuron/kpCBA determined from the 15 experimental values obtained with the FeIIINTA/H2O2 system were equal to 0.46 (s.d. 0.02) and 1.44 (s.d. 0.03), respectively (Experiments B1eB11, C1eC4, Table 1). These values of relative rate constants as well as the inhibiting effect of bicarbonate ion demonstrate that the degradation of model organic compounds by the FeIIINTA/H2O2 system can be attributed to the hydroxyl radical. It should also be noted that the identical relative rate constants obtained for the degradation of the HO probes by FeIII/H2O2 at pH 2.85 and by FeIIINTA/ H2O2 at pH 7 indicate that the absolute rate constants of HO radicals with the acid and the basic forms of pCBA are quite identical confirming the assumption made above.

3.2. pH

Oxidation by FeII/O2 and by FeIINTA/O2 at neutral

In the classical Fenton-like reaction (Fe2þ/H2O2, acidic pH), the reaction of H2O2 with Fe2þ represents the unique source for the generation of HO radicals and the reduction reaction of iron(III) into iron(II) represents the rate-limiting step for the overall rate of decomposition of H2O2 (De Laat and Gallard, 1999). At neutral pH, iron(II) species (FeOHþ and Fe(OH)2) are readily oxidized by H2O2 and can also be rapidly oxidized by dissolved oxygen at pH >7 (King et al., 1995; Santana-Casiano et al., 2005). It has also been found that the presence of NTA significantly enhances the rate of oxidation of iron(II) by oxygen (Kurimura et al., 1968) because of the formation of more reactive FeIINTA complexes. As dissolved oxygen oxidizes ferrous ion and FeIINTA at neutral pH, a series of experiments was carried out without adding H2O2 to check if the reactions of dissolved oxygen with free FeII species or with FeIINTA lead to the formation of intermediates that can degrade the three model compounds used in the present work. These experiments were carried out in a semibatch reactor with a continuous introduction of FeII(ClO4)2 (50 mM) or FeIINTA (50 mM; FeII:NTA molar ratio ¼ 1:3) at a flow rate of about 25 or 50 ml L1 h1. For each experiment, the feed doses of FeII or of FeIINTA (0.72e2.3 mM L1 h1, Table 2) were calculated from the flow rate delivered by the peristaltic pump and the exact concentration of Fe(II) in the stock solutions of Fe(II). The feed doses were confirmed by the slope of the straight lines obtained for the concentrationetime profiles of total iron in the reactor (Fig. 3a). Data in Fig. 3a show a decrease of the concentration of dissolved oxygen with increasing reaction time and a nearly

3314

w a t e r r e s e a r c h 4 5 ( 2 0 1 1 ) 3 3 0 9 e3 3 1 7

Table 1 e Relative rate constants of degradation of model organic solutes obtained for the FeIII/H2O2 system (experiments A1eA4) and the FeIIINTA/H2O2 system in the absence (experiments B1eB9) and in the presence of sodium hydrogenocarbonate (experiments C1eC4). (Batch reactor; [Fe(III)]0 [ 0.206 mM; [NTA]0 [ 0.6 mM; [H2O2]0 [ 0.5 mM; [Atrazine]0 [ [Fenuron]0 [ [pCBA]0 [ 5 mM; 25  C). Exp. A1 A2 A3 A4

Process tested III

Fe /H2O2 FeIII/H2O2 FeIII/H2O2 FeIII/H2O2 III

pH 2.85 2.85 2.85 2.85

   

0.05 0.05 0.05 0.05

B1 B2 B3 B4 B5 B6 B7 B8 B9 B10 B11

Fe NTA/H2O2 FeIIINTA/H2O2 FeIIINTA/H2O2 FeIIINTA/H2O2 FeIIINTA/H2O2 FeIIINTA/H2O2 FeIIINTA/H2O2 FeIIINTA/H2O2 FeIIINTA/H2O2 FeIIINTA/H2O2 FeIIINTA/H2O2

7.0 7.0 7.0 7.0 8.0 8.0 8.0 8.3 8.3 9.0 9.0

          

0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1

C1 C2 C3 C4 C5

FeIIINTA/H2O2 FeIIINTA/H2O2 FeIIINTA/H2O2 FeIIINTA/H2O2 FeIIINTA/H2O2

8.3 8.3 8.3 8.3 8.3

    

0.1 0.1 0.1 0.1 0.1

[NaHCO3]0 (mM)

Tert-Butanol (mM)

kAtrazine/kpCBA

kFenuron/kpCBA

50

0.471 0.473 0.441 e

1.482 1.476 1.486 e

0.482 0.494 0.458 0.453 0.436 0.451 0.444 0.420 0.441 0.464 0.463

1.437 1.486 1.428 1.404 1.453 1.440 1.439 1.443 1.389 1.403 1.430

0.466 0.488 0.456 0.480 e

1.412 1.417 1.468 1.448 e

0.460 0.020

1.441 0.030

5 10 25 50 10

Mean value Standard deviation

complete conversion of ferrous species into ferric species (z95% for FeII; z97% for FeIINTA at a reaction time of 20 min). The consumptions of dissolved oxygen by free FeII and FeIINTA were equal to 0.243  0.005 mol of O2/mole of FeII or FeIINTA (Table 2), in agreement with the expected overall stoichiometry for the oxidation of ferrous species by O2. The rates of consumption of dissolved oxygen by FeII and FeIINTA were identical because the introduction doses of FeII and FeIINTA were identical and represented the rate-limiting step in the overall rate of consumption of dissolved oxygen at pH 7.0(Fig. 3a). The data also demonstrate that atrazine, fenuron and pCBA could be degraded by FeII and FeIINTA at neutral pH in the presence of dissolved oxygen and that the degradation rates were faster with FeIINTA than with FeII (Fig. 3b). In addition, the

values of the relative rate constants determined from the competitive kinetic expression (kAtrazine/kpCBA ¼ 0.46  0.01 and kFenuron/kpCBA ¼ 1.45  0.02, Fig. 3c, Table 2) showed that the degradation of the organic solutes can be attributed to HO radicals and provide evidence for the formation of HO radicals during the oxidation of FeII and FeIINTA by dissolved oxygen. These results are consistent with other data showing the involvement of hydroxyl radicals in the degradation of organic compounds during the oxidation of iron(II) (Burns et al., 2010) or of FeIINTA (Kachur et al., 1998; Keenan and Sedlak, 2008) by dissolved oxygen at circumneutral pH. In the FeII/O2 and FeIINTA/O2 systems, HO radicals can be formed via the Fenton reaction because H2O2 can be produced from the two-step reduction of dissolved oxygen into HO2/O2 and then into H2O2.

Table 2 e Dissolved oxygen consumption (D[O2]/D[FeII]) and relative rate constants of degradation of probe compounds obtained with the FeII/O2 and FeIINTA/O2 systems (Semi-batch reactor, [O2]0 [ 8.3 ± 0.1 mg O2 LL1; [Atrazine]0 [ [Fenuron]0 [ [pCBA]0 [ 5 mM; Feed dose of FeII or of FeIINTA (FeII:NTA molar ratio [ 1:3) [ 0.72e2.33 mM LL1 hL1; pH [ 7.0 ± 0.1; 25  C). Exp.

Process tested II

D1 D2 D3

Fe /O2 FeII/O2 FeII/O2

E1 E2 E3 E4 E5

FeIINTA FeIINTA FeIINTA FeIINTA FeIINTA

Mean value Standard deviation

(1:3)/O2 (1:3)/O2 (1:3)/O2 (1:3)/O2 (1:3)/O2

FeII introduced (mM/L1 h1)

DO2/D FeII (mol/mol)

kFenuron/kpCBA

kAtrazine/kpCBA

2.295 2.255 1.874

0.243 0.246 0.245

1.427 1.436 1.445

0.466 0.459 0.447

2.350 2.316 2.331 0.727 0.872

0.244 0.241 0.245 0.239 0.24

1.469 1.476 1.472 1.445 1.453

0.478 0.468 0.462 0.458 0.467

0.243 0.003

1.453 0.018

0.463 0.009

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w a t e r r e s e a r c h 4 5 ( 2 0 1 1 ) 3 3 0 9 e3 3 1 7

0.3

a

1

II

Open symbols : Fe /O2 Dark symbols Fe NTA/O2

0.8 Total iron

0.2

0.6

Dissolved oxygen

0.4 0.1

0.2 Ferrous iron

0

0

0

300

600

900

1200

0,2 Total iron

H2O2

0,3 0,2

0,1

0,1

Ferrous iron

0

200

400 600 Time (s)

800

0 1000

1

1 0.8

0.8

II

II

Fe /O2

Atrazine pCBA Fenuron

0.6

[C] t / [C] 0

[C] t / [C] 0

0,4

0

b

0,3

II

Open symbols : Fe /H2O2 (Exp. F1) II Dark symbols : Fe NTA/H2O2 (Exp. G4)

0,5

1500

Time (s)

b

0,6

[Iron species] (mM)

[Iron species] (mM)

[Dissolved oxygen] (mM)

II

[Hydrogen peroxide] (mM)

a

0.4

Fe /H2O2 (Exp. F1)

0.6

Atrazine pCBA Fenuron

0.4 II

II

0.2

Fe -NTA/H2O2 (Exp. G4)

0.2

Fe NTA/O 2

0

0

0

300

600

900

1200

1500

1800

0

200

400

600

800

1000

Time (s)

Time (s)

c

Fenuron

II

Fe NTA / O2 FeII-NTA/O2 II Fe NTA / O2 FeII-NTA/O2 II Fe / O2 FeII/O2 II FeII/O2 Fe / O2

2.50 2.00

- ln ([C]t / [C]0

Fig. 4 e Oxidation of probe compounds by FeII/H2O2 and FeIINTA/H2O2 at neutral pH. (a) Concentrationetime profiles of H2O2 and of iron species (unfiltered samples), (b) Normalized concentrationetime profiles of probe compounds, (Experimental conditions are given in Table 3, Experiments F1 and G4).

3.00

1.50

y = 1.453x

Atrazine

1.00 0.50

y = 0.463x

0.00

0

0.5

1

1.5

2

- ln([pCBA] t / [pCBA]0 ) Fig. 3 e Oxidation of probe compounds by FeII/O2 and FeII/ NTA/O2 at pH [ 7.0. (a) Concentrationetime profiles of O2 and of iron species(unfiltered samples); (b) Normalized concentrationetime profiles of probe compounds; (c) Plots of the competitive kinetic equation for the determination of the relative rate constants (Experimental conditions are given in Table 2).

3.3. Oxidation by FeII/H2O2 and by FeIINTA/H2O2 at neutral pH Oxidation experiments with the FeII/H2O2 ([O2]0 < 0.1 mg/L) and FeIINTA/H2O2 ([O2]0 < 0.1 and 8.5 mg/L) systems were also carried

out in the semi-batch reactor with a continuous introduction of a solution of FeII(ClO4)2 (45 mM) or of FeIINTA (45 mM, FeII:NTA molar ratio ¼ 1:3) into the reactor (experiments F1 and G1eG4 in Table 3 and Fig. 4). The initial concentration of H2O2 was 0.5 mM and the pH was regulated at pH ¼ 7.0  0.1. Contrary to the experiments carried out with FeII/O2, FeIINTA/O2 and FeII/H2O2 systems, ferrous ion was never detected in the experiments conducted with FeIINTA/H2O2 (Fig. 4a) because of the very fast reaction of H2O2 with FeIINTA at neutral pH. The data showed that the hydrogen peroxide consumptions were nearly equal to 1.3 and 1.8  0.1 mol H2O2/mol FeII for the FeII/H2O2 and the FeIINTA/H2O2 systems, respectively. These data will be discussed in a next paper. The data in Fig. 4b also showed that the degradation rates of HO probes were faster with FeIINTA/H2O2 than with FeII/H2O2, probably because the reaction rates by the FeII/H2O2 are likely to be limited by the precipitation of iron(III) hydroxides (De Laat and Gallard, 1999). Similar conclusions could be drawn from data obtained from batch experiments (0.5 mM; [NTA]0 ¼ 0 or 1 mM), initial (pH ¼ 7.0) and with initial concentrations of H2O2 ranging from 0 to 1 mM (Table S4). The plots of competitive kinetic expression to data obtained from batch and semi-batch experiments yielded straight lines with

3316

w a t e r r e s e a r c h 4 5 ( 2 0 1 1 ) 3 3 0 9 e3 3 1 7

Table 3 e Hydrogen peroxide consumption (D[H2O2]/D[FeII]) and relative rate constants of degradation of probe compounds obtained with the FeII/H2O2 and FeIINTA/H2O2 systems in the absence and in the presence of dissolved oxygen. (Semi-batch reactor, [H2O2]0 [ 0.5 mM; [Atrazine]0 [ [Fenuron]0 [ [pCBA]0 [ 5 mM; Feed dose of FeII or of FeIINTA (FeII:NTA molar ratio [ 1:3) [ 1.0e1.2 mM LL1 hL1, pH [ 7.0 ± 0.1; 25  C). Exp. F1 G1 G2 G3 G4

Process tested FeII/H2O2 FeIINTA (1:3)/H2O2 FeIINTA (1:3)/H2O2 FeIINTA (1:3)/H2O2 FeIINTA (1:3)/H2O2

FeII introduced (mM L1 h1) [O2]0 (mg L1) DH2O2/DFeII (mol mol1) kFenuron/kpCBA kAtrazine/kpCBA 1.038 1.193 1.152 1.135 1.033

<0.1 <0.1 <0.1 8.5 8.5

slopes equal to 0.48  0.02 and 1.46  0.02 for kAtrazine/kpCBA and kFenuron/kpCBA, respectively (Table 3). These relative rate constant values as well as the inhibition of the degradation of the three probes molecules in the presence of tert-butanol (50 mM) demonstrate the formation of HO radicals during the catalytic decomposition of FeIINTA by H2O2.

4.

Conclusions

The data obtained in the present work showed that the catalyzed decomposition of oxygen and/or hydrogen peroxide by iron(II) and iron(III)-nitrilotriacetate complexes (FeIINTA/O2, FeIINTA/H2O2 and FeIIINTA/H2O2 systems) can degrade atrazine, fenuron and parachlorobenzoic acid at neutral pH whereas the classical Fenton and Fenton-like oxidation processes (FeII/H2O2 and FeIII/H2O2) were not efficient at pH 7. The competitive kinetic study and the observed decrease of the reaction rates in the presence of hydroxyl radical scavengers (tert-butanol and bicarbonate ions) clearly demonstrate that the hydroxyl radical is the unique oxidant involved in the initial step of the degradation of the three model compounds at neutral pH. These data do not exclude the formation of ferryl ion but they demonstrate that ferryl ion is not the primary oxidant of the model compounds studied in the present work. The data also showed that the rates of decomposition of H2O2 and of the model compounds observed with the FeIIINTA/ H2O2 system were at a maximum at pH z 8. The study of the effects of pH as well as of other parameters on the rates of decomposition of H2O2 by FeIIINTA complexes is currently investigated in our lab and will be presented in a next paper.

Appendix. Supplementary material Supplementary data related to this article can be found online at doi:10.1016/j.watres.2011.03.043.

references

Acero, J.L., Benitez, F.J., Gonzalez, M., Benitez, R., 2002. Kinetics of fenuron decomposition by single-chemical oxidants and combined systems. Ind. Eng. Chem. Res. 41, 4225e4232. Bautista, P., Mohedano, A.F., Casas, J.A., Zazo, J.A., Rodriguez, J.J., 2008. An overview of the application of Fenton oxidation to

1.315 1.832 1.763 1.789 1.955

1.48 1.443 1.410 1.435 1.456

0.501 0.485 0.479 0.481 0.471

industrial wastewaters treatment. J. Chem. Technol. Biot. 83, 1323e1338. Bergendahl, J., Hubbard, S., Grasso, D., 2003. Pilot-scale Fenton’s oxidation of organic contaminants in groundwater using autochthonous iron. J. Hazard. Mater. 99, 43e56. Bossmann, S.H., Oliveros, E., Go¨b, S., Siegwart, S., Dahlen, E.P., Payawan, A.L., Straub, M., Wo¨rner, M., Braun, A.M., 1998. A new evidence against hydroxyl radicals as reactive intermediates in the thermal and photochemically enhanced Fenton reactions. J. Phys. Chem. A 102, 5542e5550. Burns, J.M., Craig, P.S., Shaw, T.J., Ferry, J.L., 2010. Multivariate examination of Fe(II)/Fe(III) cycling and consequent hydroxyl radical generation. Environ. Sci. Technol. 44, 7226e7231. Buxton, G.U., Greenstock, C.L., Helman, W.P., Ross, A.B., 1988. Critical review of rate constants for reactions of hydrated electrons, hydrogen atoms and hydroxyl radicals (OH/O) in aqueous solution. J. Phys. Chem. Ref. Data 17 (2), 513e886. De Laat, J., Gallard, H., 1999. Catalytic decomposition of hydrogen peroxide by Fe(II) in homogenous aqueous solution: mechanism and kinetic modelling. Environ. Sci. Technol. 33, 2726e2732. De Laat, J., Le, T.G., 2005. Kinetics and modelling of the Fentonlike reaction in the presence of sulfate in acidic aqueous solutions. Environ. Sci. Technol. 39, 1811e1818. De Laat, J., Le, T.G., 2006. Effects of chloride ions on the iron(III)catalyzed decomposition of hydrogen peroxide and the efficiency of the Fenton-like oxidation process. Appl. Catal. B. Environ. 66, 137e146. Eisenberg, G.M., 1943. Colorimetric determination of hydrogen peroxide. Ind. Eng. Chem. 15, 327e328. Elovitz, M., von Gunten, U., 1999. Hydroxyl radical/ozone ratios during the ozonation processes. I. The Rct concept. Ozone Sci. Eng. 21, 239e260. Gallard, H., De Laat, J., Legube, B., 1998. Influence du pH sur la vitesse d’oxydation de compose´s organiques par FeII/H2O2. Me´canismes re´actionnels et mode´lisation. New J. Chem., 263e268. Harris, D.C., Aisen, P., 1973. Facilitation of Fe(II) autooxidation by Fe(III) complexing agents. Biochim. Biophys. Acta 329, 156e158. Howsawkeng, J., Watts, R.J., Washington, D.L., Teel, A.L., Hess, T.F., Crawford, R.L., 2001. Evidence for simultaneous abiotic-biotic oxidations in a microbial-Fenton’s system. Environ. Sci. Technol. 35, 2961e2966. Hug, S.J., Leupin, O., 2003. Iron-catalyzed oxidation of arsenic(III) by oxygen and by hydrogen peroxide: pH-dependent formation of oxidants in the Fenton reaction. Environ. Sci. Technol. 37, 2734e2742. Kachur, A.V., Tuttle, S.W., Biaglow, J.E., 1998. Autooxidation of ferrous ion complexes: a method for the generation of hydroxyl radicals. Radiat. Res. 150, 475e482. Katsoyiannis, I.A., Ruettimann, T., Hug, S.J., 2008. pH dependence of Fenton reagent generation and As(III) oxidation and removal by corrosion of zero valent iron in aerated water. Environ. Sci. Technol. 42, 7424e7430.

w a t e r r e s e a r c h 4 5 ( 2 0 1 1 ) 3 3 0 9 e3 3 1 7

Keenan, C.R., Sedlak, D.L., 2008. Ligand-enhanced reactive oxidant generation by nanoparticulate zero-valent iron and oxygen. Environ. Sci. Technol. 42, 6936e6941. Kim, Y.-S., Kong, S.-H., 2001. Treatment of 1-hexanol and carbon tetrachloride in modified Fenton’s reagent using iron(III)-NTA complex. J. Korean Ind. Eng. Chem. 12, 787e792. King, D.W., Lounsbury, H.A., Millero, F.J., 1995. Rates and mechanism of Fe(II) oxidation at nanomolar total iron concentrations. Environ. Sci. Technol. 29, 818e824. Kurimura, Y., Ochiai, R., Matsuura, N., 1968. Oxygen oxidation of ferrous ions induced by chelation. Bull. Chem. Soc. Jpn. 41, 2234e2239. Lee, C., Keenan, C.R., Sedlak, D.L., 2008. Polyoxometalate oxidation of organic compounds by nanoparticulate zerovalent iron and ferrous ion in the presence of oxygen. Environ. Sci. Technol. 42, 4921e4926. Lee, C., Sedlak, D.L., 2009. A novel homogeneous Fenton-like system with Fe(III)-phosphotungstate for oxidation of organic compounds at neutral pH values. J. Mol. Catal. A: Chem. 311, 1e6. Li, Y.C., Bachas, L.G., Bhattacharyya, D., 2007. Selected chloroorganic detoxifications by polychelate (poly(acrylic acid)) and citrate-based Fenton reaction at neutral pH environment. Ind. Eng. Chem. Res. 46, 7984e7992. Mazellier, P., Busset, C., Delmont, A., De Laat, J., 2007. A comparison of fenuron degradation by hydroxyl and carbonate radicals in aqueous solution. Water Res. 41, 4585e4594. Motekaitis, R.J., Martell, A.E., 1994. The iron(III) and iron(II) complexes of nitrilotriacetic acid. J. Coord. Chem. 31, 67e78. Ndjou’ou, A.-C., Bou-Nasr, J., Cassidy, D., 2006. Effect of Fenton reaction dose on coexisting chemical and microbial oxidation in soil. Environ. Sci. Technol. 40, 2778e2783. Pera-Titus, M., Garcıa-Molina, V., Ban˜os, Miguel A., Gime´nez, J., Esplugas, S., 2004. Degradation of chlorophenols by means of advanced oxidation processes: a general review. Appl. Catal. B. Environ. 47, 219e256. Pignatello, J., Oliveros, E., MacKay, A., 2006. Advanced oxidation process for organic contaminant destruction based on the Fenton reaction and related chemistry. Crit. Rev. Environ. Sci. Technol. 36, 1e84. Rastogi, A., Al-Abed, S.R., Dionysiou, D.D., 2009. Effect of inorganic, synthetic and naturally chelating agents of Fe(II) mediated advanced oxidation of chlorophenols. Water Res. 43, 684e694.

3317

Rivas, F.J., Beltran, F.J., Frades, J., Buxeda, P., 2001. Oxidation of p-hydroxybenzoic acid by Fenton’s reagent. Water Res. 35, 387e396. Rose, A.L., Waite, T.D., 2002. Kinetic model for Fe(II) oxidation in seawater in the absence and presence of natural organic matter. Environ. Sci. Technol. 364, 433e444. Santana-Casiano, J.M., Gonza´lez-Da´vila, M., Millero, F.J., 2005. Oxidation of nanomolar levels of Fe(II) with oxygen in natural waters. Environ. Sci. Technol. 39, 2073e2079. Sahul, K., Sharma, B.K., 1987. Gamma radiolysis of nitrilotriacetic acid (NTA) in aqueous solutions. J. Radioanal. Nucl. Ch. 109, 321e327. Sun, Y., Pignatello, J.J., 1992. Chemical treatment of pesticide wastes. Evaluation of Fe(III) chelates for catalytic hydrogen peroxide oxidation of 2,4-D at circumneutral pH. J. Agric. Food Chem. 40, 322e327. Sun, Y., Pignatello, J.J., 1993. Activation of hydrogen peroxide by iron(III) chelates for abiotic degradation of herbicides and insecticides in water. J. Agric. Food Chem. 41 (2), 308e312. Tachiev, G., Roth, J.A., Bowers, A.R., 2000. Kinetics of hydrogen peroxide decomposition with complexed and “free” iron catalysts. Int. J. Chem. Kinet. 32, 24e35. Tamura, H., Goto, K., Yotsuyanagi, T., Nagayama, M., 1974. Spectrophotometric determination of iron(II) with 1. 10-phenanthroline in the presence of large amounts of iron(III). Talanta 21, 314e318. Walling, C., Partch, R.E., Weil, T., 1975. Kinetics of the decomposition of hydrogen peroxide catalyzed by ferric ethylenediaminetetraacetate complex. Proc. Nat. Acad. Sci. USA 72 (1), 140e142. Watts, R.J., Teel, A.L., 2005. Chemistry of modified Fenton’s reagent (catalyzed H2O2 propagations - CHP) for in situ soil and groundwater remediation. J. Environ. Eng. 131, 612e622. Welch, K.D., Davis, T.Z., Aust, S.D., 2002. Iron autoxidation and free radical generation: effects of buffers, ligands and chelators. Arch. Biochem. Biophys. 397, 360e369. Xue, X., Hanna, K., Despas, C., Wu, F., Deng, N., 2009. Effect of chelating agent on the oxidation of PCP in the magnetite/H2O2 system at neutral pH. J. Mol. Catal. A: Chem. 311, 29e35. Yamazaki, I., Piette, L., 1991. EPR spin-trapping study on the oxidizing species formed in the reaction of the ferrous ion with hydrogen peroxide. J. Am. Chem. Soc. 113, 7588e7593.