H2O2 System in the Presence of

Emmanuel Mousset , Luigi Frunzo , Giovanni Esposito , Eric D. van Hullebusch , Nihal Oturan , Mehmet A. Oturan. Applied Catalysis B: Environmental 201...
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Environ. Sci. Technol. 2005, 39, 1811-1818

Kinetics and Modeling of the Fe(III)/H2O2 System in the Presence of Sulfate in Acidic Aqueous Solutions

H2O2) indicates that each active intermediate formed (FeO2+ or •OH) oxidizes another ferrous ion. The overall rate of oxidation of Fe(II) by H2O2 obeys the following pseudosecond-order kinetics

JOSEPH DE LAAT* AND TRUONG GIANG LE Laboratoire de Chimie de l'Eau et de l'Environnement, CNRS UMR 6008, Ecole Supe´rieure d'Inge´nieurs de Poitiers, Universite´ de Poitiers, 40, avenue du Recteur Pineau, 86 022 Poitiers Cedex, France

where k represents the pseudo-second-order rate constant. k is not a true second-order constant since i) the overall reaction of decomposition of H2O2 involves several elementary reactions and (ii) the rate constant k depends on the speciation of Fe(II). In a recent study of the rate of oxidation of Fe(II) by H2O2 conducted with Fe(II) in excess ([Fe(II)]0/[H2O2]0 g 2 mol/ mol, 25 °C, 1 < pH < 3, ionic strength: 0-1 M), we found that k was not affected by the presence of perchlorate, nitrate or chloride (k ) 55 M-1 s-1) but increased in the presence of sulfate (18). k increased from 55 to 78 M-1 s-1 when the fraction of total Fe(II) that was complexed with SO42increased from 0 to ≈ 100%. For the homogeneous Fe(III)/H2O2 system (Fenton-like reaction), the spontaneous reaction of H2O2 with Fe(III) primarily leads to the formation of iron(III)-peroxo complexes (7, 9, 16, 19). A recent spectrophotometric study conducted in NaClO4/HClO4 solutions demonstrates the formation of two iron(III)- peroxo complexes at pH < 3 and provides the equilibrium constants and the UV/visible spectra of these two complexes (20):

This work examined the effect of sulfate ions on the rate of decomposition of H2O2 by Fe(III) in homogeneous aqueous solutions. Experiments were carried out at 25 °C, pH e 3 and the concentrations of sulfate ranged from 0 to 200 mM ([Fe(III)]0 ) 0.2 or 1 mM, [H2O2]0 ) 10 or 50 mM). The spectrophometric study shows that addition of sulfate decreased the formation of iron(III)-peroxo complexes and that H2O2 does not form complexes with iron(III)sulfato complexes. The rates of decomposition of H2O2 markedly decreased in the presence of sulfate. The measured rates were accurately predicted by a kinetic model based on reactions previously validated in NaClO4/HClO4 solutions and on additional reactions involving sulfate ions and sulfate radicals. At a fixed pH, the pseudo-firstorder rate constants were found to decrease linearly with the molar fraction of Fe(III) complexed with sulfate. The model was also able to predict the rate of oxidation of a probe compound (atrazine) by Fe(III)/H2O2. Computer simulations indicate that the decrease of the rate of oxidation of organic solutes by Fe(III)/H2O2 can be mainly attributed to the complexation of Fe(III) by sulfate ions, while sulfate radicals play a minor role on the overall reaction rates.

Introduction Fenton’s reaction (Fe(II)/H2O2), Fenton-like reagents (Fe(III)/H2O2) and photo-Fenton processes (Fe(II) or Fe(III)/ H2O2/UV) can be used to oxidize organic pollutants present in industrial wastewaters. Optimum pH for the application of these advanced oxidation processes is pH ≈ 3 (1-3). The mechanisms and the kinetics of the reactions of H2O2 with Fe(II) and Fe(III) have been the subject of numerous investigations over the last century and are still far from being fully understood (4-16). It is now widely accepted that the first step for the reaction of H2O2 with Fe(II) is the formation of a transient iron(II)H2O2 complex. Depending on the nature of the ligands and on the pH, the transient complex may decompose to give Fe(IV) species (i.e. FeO2+) or free hydroxyl radicals (•OH) (14, 17-18). In organic-free solution and in the presence an excess of Fe(II) over H2O2 ([Fe(II)]0/[H2O2]0 g 2 mol/mole), the stoichiometry of the overall reaction (2 mol of Fe(II)/mol of * Corresponding author phone: +33 5 49 45 39 21; fax: +33 5 49 45 37 68; e-mail: [email protected]. 10.1021/es0493648 CCC: $30.25 Published on Web 01/22/2005

 2005 American Chemical Society

-

d[H2O2] d[Fe(II)] ) -2 ) 2k[H2O2][Fe(II)] dt dt

(1)

Fe3+ + H2O2 H FeIII(HO2)2+ + H+

(2a)

FeOH2+ + H2O2 H FeIII(OH)(HO2)+ + H+

(2b)

Once formed, the iron(III)-peroxo complexes are assumed to decompose in a unimolecular way to yield Fe2+ and HO2• (10):

FeIII (HO2)2+ f Fe2+ + HO2•

(3a)

FeIII (OH)(HO2)+ f Fe2+ + HO2• + OH-

(3b)

The rates of decomposition of H2O2 by Fe(III) in HClO4/ NaClO4 solutions could be predicted very accurately under a wide range of experimental conditions (1 e pH e 3; 0.2 mM e [H2O2]0 e 1 M; 50 µM e [Fe(III)]0 e 1 mM; 1 e [H2O2]0/ [Fe(III)]0 e 5000) (21). The kinetic model (Table 1) takes into account the formation and decomposition of Fe(III)-peroxo complexes, the •OH-mechanism of decomposition of H2O2 by Fe(II) and other known reactions in the Fenton chemistry. This model was also able to predict the rates of decomposition of H2O2 by Fe(II) at higher pH (pH values up to 4) (22) and the rates of oxidation of a probe organic solute (atrazine) by the Fenton and the Fenton-like reagents (23). Besides organic solutes, industrial wastewaters also contain inorganic compounds, in particular sulfate and chloride anions at concentrations levels which can vary from 0.1 to 100 mM. These anions which are also added as reagents for the Fenton and Fenton-like oxidation processes (FeSO4, FeCl3, H2SO4, HCl) may have a significant effect on the overall reaction rates. The effects of inorganic salts on the rates of decomposition of H2O2 and organic compounds are ignored by most of the authors. VOL. 39, NO. 6, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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TABLE 1. Kinetic Model for the Fenton’s Reaction in NaClO4/HClO4 Solutionsa reactions

rate constants (s-1 or M-1 s-1)

R1 R2 R3 R4 R5 R6 R7 R8

Equilibria (I ) 0.1 M) Fe3+ + H2O H FeOH2+ + H+ k1 ) 2.34 × 107 s-1/k-1 ) 1 × 1010 (K1 ) 2.34 × 10-3 M) Fe3+ + 2 H2O H Fe(OH)2+ + 2H+ k2 ) 4.68 × 103 s-1/k-2 ) 1 × 1010 (K2 ) 4.68 × 10-7 M) 2 Fe3+ + 2 H2O H Fe2(OH)24+ + 2H+ k3 ) 1.12 × 107/k-3 ) 1 × 1010 (K3 ) 1.12 × 10-3 M-1) Fe2+ + H2O H FeOH+ + H+ k4 ) 1.90 s-1/k-4 ) 1 × 1010 (K4 ) 1.90 × 10-10 M) H+ + OH- H H2O k5 ) 1 1020/k-5 ) 1 106 s-1 (K5 )1 1014 M-1) H2O2 H HO2- + H+ k6 ) 1.26 × 10-2 s-1/k-6 ) 1 × 1010 (K6 )1.26 × 10-12) O22- + H+ f HO2k7 ) 1.0 × 1010 HO2• H O2•- + H+ k8 ) 1.58 × 105 s-1/k-8 ) 1 1010 (K8 ) 1.58 × 10-5 M)

R9 R10 R11 R12 R13 R14 R15 R16 R17 R18 R19 R20

Reactions of Fe(II) and Fe(III) Fe2+ + H2O2 f Fe3+ + •OH + HOk9 ) 55 + 3+ • FeOH + H2O2 f Fe + OH + 2 HOk10 ) 5.9 × 106 Fe3+ + H2O2 H Ia + H+ k11 ) 3.1 × 107/k-11 ) 1 1010 (KIa ) 3.1 × 10-3) FeOH2+ + H2O2 H Ib + H+ k12 ) 2 × 106/k-12 ) 1 × 1010 (KIb ) 2 × 10-4) Ia f Fe2+ + HO2• k13 ) k14 ) 2.3 × 10-3 s-1 Ib f Fe2+ + HO2•+ OHk14 ) k13 ) 2.3 × 10-3 s-1 Fe2+ + •OH f Fe3+ + OHk15 ) 2.7 × 108 FeOH+ + •OHf Fe3+ + 2OHk16 ) 2.7 × 108 Fe(II) + HO2• f Fe3+ + HO2k17 ) 1.2 × 106 Fe(II) + O2•- f Fe3+ + O22k18 ) 1.0 × 107 Fe(III)+ HO2• f Fe2+ + O2 + H+ k19 ) 2 × 104 Fe(III)+ O2•- f Fe2+ + O2 k20 ) 5 × 107

R21 R22 R23 R24 R25 R26 R27

Reactive Oxygen Radical Reactions H2O2 + •OH f HO2•+ H2O k21 ) 3.3 × 107 • • HO2 + OH f O2 + H2O k22 ) 0.71 × 1010 O2•- + •OH f O2 + OHk23 ) 1.01 × 1010 OH•+ •OH f H2O2 k24 ) 5.2 × 109 H2O2 + HO2•f OH•+ O2+H2O k25 ) 0.5 HO2• + HO2•f H2O2 + O2 k26 ) 8.3 × 105 HO2• + O2•- f HO2- + O2 k27 ) 9.7 × 107

a Fe(II) represents Fe2+ and FeOH+, Fe(III) represents free Fe(III) species (Fe3+, FeOH2+, Fe(OH) + and Fe (OH) 2+). In reactions R11-R14, I and 2 2 4 a Ib represent FeHO22+ and Fe(OH)(HO2)+, respectively.

In the case of sulfate ions, Pignatello (24) showed that the rates of degradation of 2-4-dichlorophenoxyacetic acid by Fe(III)/H2O2 and the first-order rate constant for the Fe(III)catalyzed decomposition of H2O2 at pH 2.7-2.8 were slower in the presence of sulfate than in the presence of perchlorate. The addition of sulfate was also found to decrease the rates of oxidation of 4-chlorophenol (25), direct dyes (26), acetic acid, atrazine and nitro-4 phenol (27) by Fe(II)/H2O2 or Fe(III)/H2O2. Sulfate ions may affect the overall reaction rates because i) sulfate ions form complexes with Fe(II) and Fe(III), ii) the chemical and photochemical properties of sulfato complexes of iron(II) and iron(III) and of free iron species can be different, and iii) the sulfate radicals (SO4•-) formed by the reaction of •OH with HSO4- can be less or much less reactive than •OH. Pulse radiolysis experiments conducted in acidic pH showed that the rate constants for the reaction of •OH (k ≈ 3 × 108 M-1 s-1) or of HO2• (k ≈ 1.2 × 106 M-1 s-1) with Fe2+ ion and the FeSO4 complex are identical and the sulfate ion displaces HO2- from the iron(III)-peroxo complex by an outer-sphere ion-pair formation mechanism. The rates of oxidation of organic solutes by •OH can also be affected by the presence of sulfate ions in acidic pH because •OH radicals react with HSO4- to give sulfate radicals (Table 2). The sulfate radical is a strong oxidant (E° ≈ 2.43 V) (29). As compared to •OH, SO4•- is more selective and a little less reactive than •OH (30). To date, complexation reactions of Fe(II) and Fe(III) by sulfate ions and reactions involving sulfate radicals have never been considered in kinetic models used for predicting the overall rates of decomposition of H2O2 and of organic solutes by the Fenton and Fenton-like reactions. Therefore, this work is designed to better understand the effects of sulfate on the overall reaction rates and to develop a generalized kinetic 1812

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model for the H2O2/Fe(III)/SO42- system at pH e 3. In this paper, we have examined the effects of sulfate on the distribution of Fe(III) in the absence and in the presence of H2O2, on the rate of decomposition of H2O2 in organic-free water and on the rate of oxidation of a probe organic compound (atrazine).

Model Reaction and Kinetic Expressions The kinetic model used in the present work is based on the •OH-radical mechanism of decomposition of H O by Fe(II). 2 2 To our kinetic model of the Fe(III)/H2O2 system in NaClO4/ HClO4 solutions (Table 1), we have incorporated all the reactions listed in Table 2 in order to take into account: (1) The complexation reactions of Fe(II) and Fe(III) by sulfate ions: Equilibrium constants for Fe(III) speciation were corrected for ionic strength from literature values (Table 3). The constants used in the present study are those given by the MINEQL+ software (31) and are identical or very close to published values (32, 33). (2) The reactions involving the formation and the decay of sulfate radicals: These reactions are well-known in radiation chemistry (34).

Experimental Section All chemicals used were reagent grade and were used without further purification. Ferric perchlorate was purchased from Aldrich, hydrogen peroxide (30%, unstabilized) from Fluka and atrazine (> 99% purity) from Cluzeau Info Labo. All the solutions were prepared in Milli-Q purified water (Millipore). Great care was taken to make ferric solutions to prevent precipitation of ferric hydroxide: an appropriate weight of Fe(ClO4)3‚9H2O was diluted in a few milliliters of perchloric acid (0.1 N) and then added to an appropriate volume of ultrapure water to give the desired pH and concentration of Fe(III). The ionic strength was adjusted to the desired value

TABLE 2. Additionnal Reactions for the Fenton’s Reaction in the Presence of Sulfate rate constants (s-1 or M-1 s-1)

reactions

Equilibria (I ) 0.1 M) k28 ) 3.47 × 1011/k-28 ) 1 × 1010 (K28 ) 3.47 × 101 M-1) K29 ) 3.89 × 1012/k-29 ) 1 × 1010 (K29 ) 3.89 × 102 M-1) K30 ) 4.47 × 1013/k-30) 1 × 1010 (K30 ) 4.47 × 103 M-2) K31 ) 2.29 × 1011/k-31 ) 1 × 1010 (K31 ) 2.29 × 101 M-1)

R28 R29 R30 R31

H+ + SO42- H HSO4Fe3+ + SO42- H FeSO4+ Fe3+ + 2 SO42- H Fe(SO4)2Fe2+ + SO42- H FeSO4

R32 R33 R34 R35 R36 R37 R38

HSO4- + •OH f SO4•- + H2O SO4•- + H2O f H+ + SO42- + •OH SO4•- + OH- f SO42- + •OH SO4•- + SO4•- f S2O82SO4•- + H2O2 f SO42- + H+ + HO2• SO4•- + HO2• f SO42- + H+ + O2 SO4•- + O2•- f SO42- + O2

R39 R40 R41 R42 R43 R44 R45 R46 R47 R48 R49

Reactions Involving Fe(II) and Fe(III) FeSO4 + H2O2 f Fe3+ + SO42- + •OH + OHk39 ) 78 FeSO4 + •OH f Fe3+ + SO42- + OHK40 ) 2.7 × 108 FeSO4 + HO2• f Fe3+ + SO42- +HO2k41 ) 1.2 × 106 FeSO4 + O2•- f Fe3+ + SO42- + O22k42 ) 5 × 108 Fe2+ + SO4•- f Fe3+ + SO42k43 ) 3.0 × 108 M-1 s-1 FeOH+ + SO4•- + f Fe3+ + SO42k44 ) 3.0 × 108 M-1 s-1 FeSO4 + SO4•- f Fe3+ + 2 SO42k45 ) 3.0 × 108 M-1 s-1 FeSO4+ + HO2• f Fe2+ + SO42- + O2 + H+ k46 < 1 × 103 FeSO4+ + O2•- f Fe2+ + SO42- + O2 k47 < 1 × 103 Fe(SO4)2- + HO2• f Fe2+ + 2 SO42- + O2 + H+ k48 < 1 × 103 Fe(SO4)2- + O2•- f Fe2+ + SO42- + O2 k49 < 1 × 103

Reactions Involving SO4•k32 ) 3.5 × 105 M-1 s-1 k33 ) 6.6 × 102 s-1 k34 ) 1.4 × 107 M-1 s-1 k35 ) 2.7 × 108 k36 ) 1.2 × 107 M-1 s-1 k37 ) 3.5 × 109 M-1 s-1 k38 ) 0

TABLE 3. Equilibrium Constants as a Function of Ionic Strength ionic strength (M) reaction H+

+ H2 O H + Fe3+ + 2 H2O H [Fe(OH)2]+ + 2H+ 2 Fe3+ + 2 H2O H [Fe2(OH)2]4+ Fe2+ + H2O H [FeOH]+ + H+ H+ + SO42- H HSO4Fe3+ + SO42- H FeSO4+ Fe3+ + 2 SO42- H Fe(SO4)2Fe2+ + SO42- H FeSO4 Fe3+

FeOH2+

log K1 log K2 log K3 log K4 log K28 log K29 log K30 log K31

0

0.1

0.2

0.5

0.6

1

-2.19 -5.67 -2.95 -9.5 1.99 3.92 5.42 2.25

-2.63 -6.33 -2.95 -9.72 1.54 2.59 3.65 1.36

-2.72 -6.47 -2.95 -9.77 1.45 2.32 3.28 1.18

-2.79 -6.57 -2.95 -9.8 1.38 2.11 3.01 1.05

-2.79 -6.57 -2.95 -9.8 1.39 2.12 3.03 1.05

-2.72 -6.47 -2.95 -9.77 1.45 2.32 3.29 1.19

with sodium perchlorate and sodium sulfate. Solutions of ferric salts were prepared daily and kinetic experiments rapidly carried out to prevent the effect of the maturation of monomeric Fe(III) species to polymeric Fe(III) species. All experiments were conducted in a thermostated batch reactor (25.0 ( 0.5 °C) and in the absence of light. Kinetic experiments of the decomposition of H2O2 were initiated by adding H2O2 under vigorous magnetic-stirring to the solution containing ferric salt and atrazine. At various intervals, samples of solution in the batch reactor were withdrawn and analyzed. Absorption spectra of the solutions were measured with a SAFAS DES 190 double beam spectrophotometer and pH measurements were made with a Meter Lab PHM 240 pH meter calibrated with acidic standard buffers between pH 1.0 and 3.0 (25.0 °C). Fe(III) concentrations were measured by the o-phenanthroline method after reduction of Fe(III) with hydroxylamine hydrochlorate and by using a value of  ) 11,000 M-1 cm-1 for the Fe(II)-phenanthroline complex at 510 nm. Hydrogen peroxide was determined iodometrically ([H2O2] > 10-3 M) or spectrophotometrically using TiCl4 method ([H2O2] e 10-3 M;  ) 730 M-1 cm-1) (35). Atrazine was analyzed by HPLC (Resolve C18 column, 5 µm, length: 150 mm, internal diameter: 3.9 mm) and UV detection at 220 nm. The mobile phase composition was water/methanol (40/60, v/v). The flow rate was 0.8 mL/min and the injection volume was 200 µL.

Distribution Calculations. Equilibrium Fe(II) and Fe(III) speciations were calculated using the software package MINEQL+ (31). Table 3 lists the equilibrium expressions and stability constants used for model input. Ionic strength corrections were made using the Davies equation. Kinetic Modeling. Kinetic modeling was performed with GEPASI version 3.30 (36, 37). All reactions and kinetic constants listed in Tables 1 and 2 were used as inputs of the program. Complexation reactions of Fe(III) and acid-base equilibrium were considered to be very fast equilibrium processes. Therefore, equilibrium reactions were written as separate forward and reverse reactions, with fast, nonrate determining forward/reverse rate constants. The reaction parameters ([H+]0, [H2O2]0, [Fe(III)]0, [SO42-]0, [Atrazine]0) were specified as inputs to the program. The concentration-time profiles for H2O2 and atrazine were calculated by the program and compared to the experimental measurements.

Results and Discussion Spectrophotometric Study. UV/visible absorption spectra of Fe(III) solutions have been measured under various experimental conditions ([Fe(III)] ) 5.5 × 10-4 M; pH 1-3) in the absence and in the presence of H2O2 in order to show the effects of sulfate concentrations on the speciation of Fe(III). The measured spectra were compared to calculated spectra in order to verify the accuracy of both equilibrium VOL. 39, NO. 6, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 1. Distribution curves of Fe(III) species as percentage of total Fe(III) versus pH (MINEQL+ calculations with [Fe(III]T ) 1 mM. [SO42-]T ) 33.33 mM. I ) 0.1 M. 25.0 °C). (Fe(OH)2+ and Fe2(OH)24+ have not been reported because they represent less than 1% of total Fe(III)).

FIGURE 4. UV/visible spectra of FeSO4+ and Fe(SO4)2- obtained in the present work. The spectrum of FeSO4+ obtained by Benkelberg and Warneck (40) (symbol) is included for comparison. where l is the path length of the cell, and i and Ri are the molar decadic absorption coefficient at the wavelength λ and the molar fraction of the Fe(III) species, respectively. In the absence of both H2O2 and sulfate, the measured absorbances of the solutions at all pH and λ were identical to the absorbances calculated by using the hydrolysis constants for Fe(III) and the molar absorption coefficients of Fe3+, FeOH2+, Fe(OH)2+ and Fe2(OH)24+ (20, 38). It should be noted that the two latter species represent minor species under our conditions (RFe(OH)2+ + RFe2(OH)24+ , 1%).

FIGURE 2. Effect of the concentration of sodium sulfate on the pH of aqueous solutions of Fe(III): comparison between the measured values (symbols) and the values calculated with MINEQL+ (line).

In the presence of sulfate, the absorption band close to 305 nm is attributed to the formation of iron(III)-sulfato complexes (39). Distribution curves in Figure 1 show that FeSO4+ and Fe(SO4)2- complexes represent the major iron(III) species for sulfate concentration higher than 5-10 mM. By using the speciation constants of Fe(III) listed in Table 3, the UV/visible spectra of the two sulfato-ferric complexes could be calculated with eq 4 from the experimental spectra of Fe(III)/SO42- solutions prepared over a wide range of pH and concentrations (Figure 4). The spectra obtained for FeSO4+ was consistent with those obtained by Jayson et al. (28) (305 nm ) 2200 M-1 cm-1 by using K29 ) 180 M-1 at I ) 0.5 M) and by Benkelberg and Warneck (40) (Figure 4). For the Fe(SO4)2- complex, the measured extinction values (305 nm ) 3600 M-1 cm-1) were higher than the value obtained by Whiteker and Davidson (39) or by Jayson et al. (305 nm ) 3100 M-1 cm-1 by using K30) 1700 M-2 at I ) 0.5 M) (28). In addition, the measured values for the pH (Figure 2) and for the absorbances of Fe(III)/SO42- solutions (Figure 3) were in good agreement with the calculated values.

FIGURE 3. UV/visible absorption spectra of Fe(III) solutions in the presence of increasing concentrations of sulfate. Comparison between measured (line) and computed spectra (symbol) (l ) 0.5 cm. [Fe(III)]T ) 5.5 × 10-4 M. pH ) 2). speciation constants and molar absorptivities of iron(III)complexes. In the absence of H2O2, the addition of Na2SO4 (0-40 mM) to solutions of Fe(III) ([Fe(III)] ) 1 mM, pH 2) leads to an increase of the pH of the solutions (Figure 2) and to the formation of species which present an absorption band close to 305 nm (Figure 3). The absorbance of the solutions can be explained in terms of the weighted sum of the absorbance of individual Fe(III) species

Aλ ) Σi‚l‚Ri‚[Fe(III)]T 1814

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As previously shown, the addition of H2O2 into acidic solutions of Fe(III) immediately leads to the formation of at least two Fe(III)-peroxo complexes (Ia and Ib; reactions R11 and R12 in Table 1) which absorb UV/visible light and in particular in the region 400-600 nm where the other Fe(III) species do not absorb (21). In the presence of sulfate (0-40 mM), the UV/visible absorbance at λ > 380 nm decreases drastically as the concentration of sulfate increases because SO42- ions compete with H2O2 for the formation Fe(III) complexes (Figure 5). By considering the formation of iron(III)-hydroxo, peroxo and sulfato complexes, Figure 5 shows that the calculated absorbances of the solutions in the region 400-600 nm were found to be in good agreement with measured absorbances. From these data, one can assume that H2O2 does not form complexes with iron(III)-sulfato complexes. Effects of Sulfate Concentration on the Rate of Decomposition of H2O2. For all the experiments performed in the present work ([Fe(III)]0 ) 0.2 or 1.0 mM, [H2O2]0 ) 10 or 50

TABLE 4. Measured Pseudo-First-Order Kinetic Constants (kobs) for the Initial Rate of Decomposition of H2O2 by Fe(III) under Various Experimental Conditionsa no.

I (M)

[SO42-]0 (mM)

[H+]0 (mM)

[Fe(III)]0 (mM)

[H2O2]0 (mM)

pH

r(Sulfato-Fe(III)) (%)

r(Ia+Ib) (%)

kobs. (10-5 s-1)

A1 A2 A3 A4 A5 A6 A7 A8 B1 B2 B3 B4 B5 C1 C2 C3 C4 C5 C6 D1 D2 D3 E1b E2b

0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.6 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.089 0.1 0.095 0.1 0.2 0.6 0.6 0.1 0.1

0 2 5 10 20 40 66.67 200 0 2 5 10 33.33 0 2 5 8 20 33.33 33.33 200 200 0 33.33

10 10 11 11.5 13.5 19 25 40 5.1 5.4 6.1 7.1 10.1 1.1 1.1 1.3 1.4 1.9 2.6 100.1 10 10 1.0 2.5

1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.2 1.0 1.0 1.0 0.2249 0.207

49.9 50.0 50.1 49.7 49.4 49.3 48.9 48.5 9.8 9.6 9.5 9.6 9.5 9.7 9.4 9.3 9.4 9.3 9.4 47.6 50.0 49.4 10.66 10.3

2.00 2.01 2.00 2.01 2.01 2.00 2.01 2.02 2.32 2.31 2.29 2.28 2.33 2.93 2.95 2.94 2.95 2.96 2.95 1.18 2.67 2.67 3.01 2.98

0 19.6 38.8 57.3 74.6 87.0 92.8 97.7 0 30.5 53.4 70.8 90.7 0 17.8 39.3 48.5 73.6 84.0 69.4 96.8 96.8 0 83.7

1.267 1.052 0.782 0.571 0.339 0.163 0.089 0.037 0.407 0.278 0.179 0.111 0.037 0.823 0.675 0.507 0.425 0.224 0.130

10.46 8.04 5.78 4.02 2.41 1.25 0.79 0.53 3.83 2.38 1.53 0.99 0.37 8.07 6.01 4.60 3.63 2.05 1.29 0.24 2.07 1.98 8.78 1.15

a

0.141 0.139 0.913 0.148

([•OH]/[SO4•-]) 4418 1688 855 419 193 114 51 1069 426 208 66 4293 1619 1053 414 235 73 205 203 211

T ) 25.0 ( 0.5 °C; ionic strength adjusted with NaClO4. Experiments with solutions containing atrazine ([Atrazine]0 ) 0.88 µM). b

or very close to those used in our previous papers (21). In a recent study of the rate of oxidation of Fe2+ by H2O2 conducted (25 °C, HClO4/NaClO4 solutions pH < 3), the mean value for k9 was found to be equal to 55.0 ( 1.0 M-1 s-1 (18). This value is lower than the one used in our previous works (63 M-1 s-1) (21, 23). Therefore, some rate constants of our model had to be changed. The best prediction of the overall rates of decomposition of H2O2 in NaClO4/HClO4 solutions obtained in the present work as well as in our previous studies was obtained with the set of rate constants listed in Table 1.

FIGURE 5. Absorption spectra of Fe(III)/H2O2 solutions in the presence of increasing concentrations of sulfate. Comparison between measured (lines) and calculated spectra (symbols). ([Fe(III)]T ) 1.13 mM, (H2O2]T ) 1.67 M, pH ) 2). mM) (Table 4), the rates of decomposition of H2O2 follow a pseudo-first-order kinetic law with respect to H2O2 concentration (Figure 5a,b)

ln([H2O2]/[H2O2]0) ) -kobs‚t

(5)

where kobs is the pseudo first-order rate constant for the overall decomposition of H2O2. The measured values for kobs and the experimental conditions are reported in Table 4. As previously observed by Pignatello (24), the presence of sulfate ions drastically decreases the rate of decomposition of H2O2 by Fe(III). Under our experimental conditions ([Fe(III)]T ) 0.2 or 1 mM, pH 2-3), concentrations of sulfate in the range 3-6 mM can lead to a 50%-reduction of kobs (Figure 6a-c). Kinetic Modeling of the Rate of Decomposition of H2O2. First, our kinetic model has been used to simulate the rates of decomposition of H2O2 by Fe(III) in the absence of sulfate in order to check the accuracy of the model presented in Table 1. As compared to our previous model, the rate and equilibrium constants used in the present work are identical

In the presence of sulfate, the rates of decomposition of H2O2 by Fe(III) could be predicted very well by a model including all the reactions listed in Tables 1 and 2 and by assuming that sulfato complexes of iron(III) cannot be reduced by HO2• and by O2•-. As illustrated in Figure 7, the rates of decomposition of H2O2 were overestimated when the rate constants for the reaction of HO2•/O2•- with FeSO4, FeSO4+ and Fe(SO4)2- (k41-k42, k46-k49 in Table 2) were assumed to be equal to the rate constants of HO2•/O2•- with free iron(II) (k17-k18 in Table 1) and free iron(III) species (Fe3+, FeOH2+, Fe(OH)2+, Fe2(OH)24+) (k19-k20 in Table 1). Computer simulations indicate that a good prediction of the rate of decomposition of H2O2 was obtained for k46-k49 < 103 s-1 and for k42 ) 5 × 108 M-1 s-1. By using these values, our kinetic model for the Fe(III)/H2O2/SO42- system (Tables 1 and 2) was found to predict very well the effects of sulfate concentration on the rates of decomposition of H2O2 obtained in the present work (Figure 6a-c). Furthermore, a pseudofirst-order kinetic law with respect to the concentration of H2O2 was also obtained for the predicted rates of decomposition of H2O2 by Fe(III) in the presence of sulfate. As shown in the data in Figure 8, the pseudo-first-order rate constants (kobs) calculated from the slopes of simulated plots of ln([H2O2]t/[H2O2]0) versus reaction time were in good agreement with the measured kobs values for all the experiments and were found to decrease linearly with the total molar fraction of iron(III)-sulfato species (RFe(SO4)+ + RFe(SO4)2-). VOL. 39, NO. 6, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 7. Model fit of the experimental rate of decomposition of H2O2 by Fe(III) ([H2O2]0 ) 9.4 mM; [Fe(III)]0 ) 0.2 mM, [SO42-]T ) 8 mM, pH ) 2.95). Experimental data points (symbols) and model fitting curves for the following rate constants for the reaction of HO2•/O2• with sulfatoferric complexes. Curves a-c: k41 ) k17 ) 1.2 × 106 M-1 s-1, k42 ) k18 ) 1 × 107 M-1 s-1; a: k46 ) k48 ) k19 )2 × 104 M-1 s-1 and k47 ) k49 ) k20 ) 5 × 107 M-1 s-1; b: k46 - k49 ) 2 × 104 M-1 s-1; c: k46 - k49 < 103 M-1 s-1. Curves d-g: k41 ) k17 ) 1.2 × 106 M-1 s-1and k46 - k49 < 103 M-1 s-1, k42 ) 1 × 108, 3 × 108, 5 × 108 or 7 × 108 M-1 s-1.

FIGURE 8. Pseudo-first-order rate constants (kobs) for the decomposition of H2O2 by Fe(III) as a function of [SO42-]0 at pH 2.0 and 3.0. Comparison between measured values (Table 4, solid symbols) and calculated values from our kinetic model (Tables 1 and 2, open symbols) or from our simplified kinetic model (Table 5) by using Eq 12 (+, see Table S4 in Supporting Information). FIGURE 6. First-order plots for the decomposition of H2O2 by Fe(III) for the experiments listed in Table 4 at pH 2.00 (a), 2.30 (b) and 2.96 (c). Comparison between measured (symbols) and calculated values (lines). The inset shows the pseudo-first-order rate constants (kobs) as a function of [SO42-]T. The fact that the overall rate of decomposition of H2O2 follows a pseudo-first-order kinetics suggests that the overall rate of decomposition of H2O2 by Fe(III) is governed by a kinetically limiting reaction step and that many reactions listed in Tables 1 and 2 play a minor role in the Fe(III)/ H2O2/SO42- system. Under the conditions used ([H2O2]0: 1050 mM, [SO42-]0: 0-200 mM), a comparison of rate constants for reactions R21 (Table 1) and R32 (Table 2) indicates that most of the •OH radicals formed by the reaction of H2O2 with Fe(II) are consumed by H2O2 to form HO2•/O2•- radicals, and, consequently, the production of SO4•- is not significant (k21 [H2O2] . k32 [HSO4-]). Computer calculations confirmed that i) steady-state concentrations of SO4•- are much lower than those of •OH ([•OH]ss/[SO4•-]ss > 50; Table 4) and ii) all the reactions involving sulfate radicals in Table 2 (reactions R32R38) as well as reactions R22-R27 in Table 1 are of minor importance and therefore could be ignored under our 1816

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TABLE 5. Simplified Model for the Fenton’s Reaction in the Presence of Sulfate in Acidic Aqueous Solution R50 R51 R52 R53 R54 R55 R56 R57 R58

Fe(III)+ H2O2 f Fe2+ + (HO2•/O2•-) Fe2+ + H2O2 f Fe(III) + •OH + OHFeSO4 + H2O2 f Fe(III) + •OH + OH- + SO42H2O2 + •OH f (HO2•/O2•-)+ H2O (HO2•/O2•-)+ Fe2+ f Fe(III) + (HO2-/O22-) (HO2•/O2•-) + FeSO4 f Fe(III) + (HO2-/O22-)+ SO42(HO2•/O2•-) + Fe(III) f Fe2+ + O2 (+ H+) (HO2•/O2•-) + FeSO4+ f Fe2+ + O2 (+ H+) + SO42(HO2•/O2•-) + Fe(SO4)2- f Fe2+ + O2 (+ H+) + SO42-

k50 k51 k52 k53 k54 k55 k56 k57 k58

experimental conditions. By ignoring these reactions, a simplified reaction scheme can be postulated for the Fe(III)/ H2O2/SO42- system (Table 5) from which the following expression for the overall rate of decomposition of H2O2 can be derived

d[H2O2] ) k’obsRFe(III)[Fe(III)]T[H2O2] dt

-

with

(6)

RFe(III) )

[Fe(III)] ) [Fe(III)]T

[Fe3+] + [Fe(OH)2+] + [Fe(OH)2+] + 2[Fe2(OH)24+] [Fe(III)]T

(7)

[Fe(III)]T ) [Fe3+] + [Fe(OH)2+] + [Fe(OH)2+] +

2[Fe2(OH)24+] + [FeSO4+] + [Fe(SO4)2-] + [Ia]+ [Ib] (8)

k’obs ) 2(k50(k51 + k52K31RSO42-[SO42-]0)‚ (k56 + k57K29RFe3+RSO42-[SO42-]0 + k58K30RFe3+ (RSO42-[SO42-]0)2)/(k54 + k55K31RSO42-[SO42-]0))0.5 (9) In eq 9, RSO42- and RFe3+ represent the molar fractions for SO42- and Fe3+, respectively:

RSO42- )

[SO42-] [SO42-]0

)

FIGURE 9. Measured and modeled atrazine (At) and H2O2 concentrations as a function of time in the absence and in the presence of sulfate ([Fe(III)]0 ) 1.0 mM; [SO42-] ) 0 or 33.33 mM; pH ) 2.95; I ) 0.1 M). Symbols are for experimental data. Solid lines are for model fits.

[SO42-] [HSO4-] + [SO42-] + [FeSO4+] + 2[Fe(SO4)2-] + [FeSO4] (10) RFe3+ )

[Fe3+] [Fe(III)]

(11)

RSO42- and RFe3+ depend on [Fe(III)]0, [SO42-]0, [H2O2]0, pH and ionic strength. Additional information on the kinetic equations can be found in the Supporting Information. Under the conditions used in the present study, the concentrations of Fe(II) (Fe2+ and FeSO4) and of the iron(III)-peroxo complexes represent less than 1% of [Fe(III)]0 ([Fe(III)]T ≈ [Fe(III)]0). Therefore, the rate of decomposition of H2O2 by Fe(III) can be described by the following pseudofirst-order kinetic law with respect to H2O2

d[H2O2] ) k’obsRFe(III)[Fe(III)]0[H2O2] ) k’’obs[H2O2] dt (12) Eq 12 shows that the pseudo-first-order rate constant should be proportional to the concentration of uncomplexed Fe(III) (RFe(III) [Fe(III)]0) if iron(III)-sulfato complexes do not react with H2O2. This is in excellent agreement with the experimental results presented in Figure 8. Furthermore, Figure 8 shows that the rate constants determined from our kinetic model by GEPASI calculations (Tables 1 and 2) or by our simplified reaction scheme (Table 5, Eq 12) were in excellent agreement with the values calculated from experimental data. Effects of Sulfate Concentration on the Rate of Oxidation of Atrazine. To confirm the validity of our kinetic model, the kinetic model has been tested to predict the effect of sulfate on the rate of oxidation of atrazine by Fe(III)/H2O2. Experiments were conducted by using very low initial concentrations of atrazine ([Atrazine]0 < 1 µM) in order to ignore possible secondary reactions involving oxidation byproducts of atrazine (Experiments E1 and E2, Table 4). Computer calculations were done by assuming that the rate constants for the reaction of •OH and of SO4•- with the protonated form of atrazine (k ) 1.2 × 109 M-1 s-1, pKa of atrazine ) 1.65) or with the molecular form of atrazine (k ) 3 × 109 M-1 s-1) are identical (23). Figure 9 shows that our kinetic model correctly predicts the rates of decomposition of atrazine and of H2O2 in the absence and in the presence of sulfate.

FIGURE 10. Influence of the concentration of sulfate on [SO4•-]/ [•OH] at the beginning of the reaction for I ) 1 M and [Fe(III)]0 ) 0.1 mM. Symbols represent computed values for the following conditions: a: pH ) 1, [H2O2]0 ) 0.1 mM; b: pH ) 2, [H2O2]0 ) 0.1 mM; c: pH ) 3, [H2O2]0 ) 0.1 mM; d: pH ) 3, [H2O2]0 ) 1 mM; e: pH ) 3, [H2O2]0 ) 10 mM. Furthermore, computer calculations indicate that the decrease of rate of the oxidation of atrazine can be mostly attributed to a decrease of the rate of generation of hydroxyl radicals because of the formation of unreactive Fe(III)-sulfato complexes. The contribution of sulfate radicals to the oxidation of atrazine is also insignificant because steadystate concentrations of SO4•- were always lower than those of •OH (Table 5, Experiment E2). As previously suggested by Pignatello (24), the sulfate radical may contribute to the degradation of organic solutes only if solutions contain very high concentrations of sulfate (. 1 M). By using our kinetic model, the computed values in Figure 10 indicate that the sulfate radical should be the predominant oxidant species in the solution ([SO4•-]ss/[•OH]ss >10) under the following conditions: acidic pH (pH < 1), low concentrations of Fe(III) (, 1 mM) and H2O2 (, 1 mM) and high concentrations of sulfate (> 0.1 M at pH 1, . 0.5 M at pH 3). Environmental and Technical Significance. For the first time, a generalized kinetic model successfully describes the overall rate of decomposition of H2O2 by Fe(III) at pH e 3 and in the presence of various concentrations of sulfate. The kinetic model and the experimental data demonstrate that the formation of iron(III)-sulfato complexes decreases the overall rates of decomposition of H2O2 by Fe(III) and of oxidation of a probe organic solute by •OH radicals. These results suggest that the efficiency of the Fenton and FentonVOL. 39, NO. 6, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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like oxidation processes for the degradation of organic pollutants in industrial wastewaters might be markedly decreased in the presence of high concentrations of sulfate (> 5 mM). Finally, this work may also have broader implications. Dissolved iron species and many reactive species (H2O2, •OH, HO2•/O2•-,...) which are present in natural waters and in atmospheric water droplets play an important role in the redox reactions and in the fate of organic pollutants in the environment. Since iron(III)-sulfato complexes can also be formed at pH > 3, sulfate ion might also have an effect on the overall rates for the chemical and photochemical reactions occurring in the environment. The kinetic model presented in this paper cannot be used at pH > 3 since precipitation reactions of iron(III) have not been considered in our model.

Acknowledgments The research was supported by the French Ministe`re des Affaires Etrange`res (Program: “Fond de Solidarite´ Prioritaire, FSP ESPOIR”) and the Centre National de la Recherche Scientifique (CNRS).

Supporting Information Available A detailed description of the kinetic model and a table giving pseudo-first-order rate constants calculated with eq 12 under various experimental conditions. This material is available free of charge via the Internet at http://pubs.acs.org.

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Received for review April 26, 2004. Revised manuscript received November 24, 2004. Accepted November 30, 2004. ES0493648