Theoretical Stability Calculation and Experimental Investigation

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Toward a Better Understanding of Fe(III)−EDDS Photochemistry: Theoretical Stability Calculation and Experimental Investigation of 4-tert-Butylphenol Degradation Yanlin Wu,†,‡,§ Marcello Brigante,‡,§ Wenbo Dong,*,† Pascal de Sainte-Claire,‡,§ and Gilles Mailhot*,‡,§ †

Shanghai Key Laboratory of Atmospheric Particle Pollution and Prevention, Department of Environmental Science & Engineering, Fudan University, Shanghai 200433, China ‡ Clermont Université, Université Blaise Pascal, Institut de Chimie de Clermont-Ferrand (ICCF)-ENSCCF, BP 10448, F-63000 Clermont-Ferrand, France § CNRS, UMR 6296, Institut de Chimie de Clermont-Ferrand, F-63171 Aubière, France S Supporting Information *

ABSTRACT: The present work describes in detail the chemical structure of the complex Fe(III)−EDDS and the predominance of different species with respect to pH. These results were obtained with ab initio calculations. From the photoredox process, the formation of hydroxyl radical was confirmed, and HO• is the main species responsible for the degradation of the organic compound present in aqueous solution. The degradation of 4-tert-butylphenol (4-t-BP), used as a model pollutant, was investigated in different conditions. For the first time, the second-order rate constant of the reaction between HO• and 4-tBP and the formation rate of HO• (RfHO•) from the photochemical process were evaluated. Through the degradation of 4-t-BP, the effect of Fe(III)−EDDS concentration, oxygen, and pH was also investigated. The pH, which plays a role in the iron cycle and in the Fe(III)−EDDS speciation, was noticed as an important parameter for the efficiency of 4-t-BP degradation. Such a result could be explained by taking into account the complex speciation and presence of a predominant form (FeL−) up to pH 8. These results are very useful for the use and optimization of such iron complexes in water treatment processes.

1. INTRODUCTION In recent years, the degradation of organic pollutants in water by green photochemical processes has become a very active research topic. Numerous studies have reported the application of Fe(III) complexes in the photodegradation of organic pollutants. They are successfully used to reduce the toxicity of toxic chemicals, convert toxic and biorecalcitrant contaminants into biodegradable byproducts, remove color, or obtain a complete mineralization of organic pollutants.1−4 Such positive results in terms of water decontamination are due to the irradiation of Fe(III) complexes which could produce oxidative radicals (like HO•, HO2•, O2•−) by the ligand-to-metal charge transfer (LMCT) reactions.5 The most obvious advantages of the application of Fe(III) complexes are the low process costs, wide irradiation wavelengths (λ < 580 nm), and a pH range that is relatively wider than that of photo-Fenton processes (optimal pH 2.8).6 Even though the application of Fe(III) complexes is considered to be an improved photochemical process, there are still some drawbacks with certain kinds of Fe(III) complexes. For example, the optimal pH in Fe(III)-citrate,1 Fe(III)-oxalate,7 and [Fe(III)salen]Cl3 photochemical processes is 5.0, 4.3, and 4.1, respectively. They are still in the acid pH range, and the treated water must be neutralized before its reintroduction to natural aquatic environments. Ethylenediaminetetraacetic acid (EDTA) can form stable water-soluble complexes with Fe(III) in a wide pH © 2013 American Chemical Society

range, but this substance is rather recalcitrant to biodegradation and is considered as a persistent organic pollutant.8 Ethylenediamine-N,N′-disuccinic acid (EDDS) is a structural isomer of EDTA and is also a strong complexing agent.9−11 However, it is biodegradable and has been reported as a safe and environmentally benign replacement for EDTA and can be used for environmental applications. Our group has studied the physicochemical properties of the Fe(III)−EDDS complex. Fe(III) is complexed by EDDS with a ratio of 1:1. Under irradiation, Fe(III)−EDDS was easily photolyzed and produced HO• in the wide pH range between 3 and 9.5,9 However, the chemical structure of the Fe(III)−EDDS complex and its speciation as a function of pH is not very well-known and seems to have an important impact on the efficiency of the process. For this reason, a detailed study by theoretical calculation to determine the different chemical structures was performed for the first time. Concomitantly, its photochemical efficiency under irradiation has been attentively investigated. 4-tert-butylphenol (4-t-BP) is an alkylphenols (APs) and is one of the endocrine-disrupting chemicals (EDCs) with highly estrogenic effects12−14 and poor biological degradability. 4-t-BP has been largely Received: September 10, 2013 Revised: December 17, 2013 Published: December 17, 2013 396

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solution to show the spectral overlap with the tubes emission are presented in Figure SM1 of the Supporting Information). The solutions were magnetically stirred with a magnetic bar during irradiation, and the total volume of the solution was 50 mL. All the experiments were carried out at room temperature (293 ± 2 K). Samples were taken from the reaction tube at fixed interval times. For the determination of the second-order reaction rate constant of the reaction between 4-t-BP and HO•, the solutions were irradiated in a monochromatic parallel beam using a 1 cm path length quartz cell. The light source was a Hg Oriel lamp (200 W). The monochromatic irradiation was carried out at a wavelength of 313 nm to selectively photolyze H2O2, generating hydroxyl radicals. The irradiations in oxygen- or nitrogensaturated solutions were performed bubbling the appropriate gas for at least 20 min before and during all irradiations. 2.4. Analytical Methods. The concentration of the 4-t-BP remaining in the aqueous solution was determined by highperformance liquid chromatography (HPLC) (Alliance, Waters, USA) equipped with a photodiode array detector (Waters 2998). The flow rate was 1 mL min−1, and the mobile phase was a mixture of water and methanol (20/80, v/v). The column was a Nucleodur 100-5 C18 of 250 × 4.6 mm. In this condition, the retention time of 4-t-BP (detected at 221 nm) was 6.5 min. The initial rate of 4-t-BP degradation was determined with the slope formed by the first five points obtained during the first 10 min of irradiation. The concentration of Fe(III)−EDDS complex was determined using the same HPLC apparatus, and the mobile phase was a mixture of an aqueous solution with tetrabutylamonium hydrogen sulfate (2 mM), and sodium formiate (15 mM) at pH 4.0, and methanol (95/5, v/v). The column was an Agilent Eclipse XDB-C18 of 150 × 4.6 mm. The retention time of Fe(III)−EDDS was 6.8 min in these conditions; the wavelength detection was fixed at 240 nm. UV−vis spectra were recorded with a Cary 300 UV−visible spectrophotometer. pH values of the solutions were measured using a Cyberscan 510 pH meter.

used as a raw material for polymerization inhibitors and stabilizing agents in the chemical industry, and so it is widely detected in food,15 aquatic animals,16 human urine,17 and rivers.18−20 The serious threat to human health has already been made clear, and it would be of great urgency to degrade 4-t-BP artificially in our living surroundings. In this work, we used 4-t-BP as a target pollutant to investigate the photocatalytic activity of Fe(III)−EDDS. The effect of irradiation time, pH, Fe(III)−EDDS concentration, and O2/N2 on the photodegradation performance of 4-t-BP under UV light irradiation (300 nm < λ < 500 nm) was investigated. The secondorder reaction rate constant between 4-t-BP and HO• was also evaluated for the first time.

2. EXPERIMENTAL SECTION 2.1. Chemicals. 4-tert-butylphenol (4-t-BP), S,S′-ethylenediamine-N,N′-disuccinic acid trisodium salt (EDDS-Na) solution (35% in water), and 2-propanol were obtained from Sigma, France. Ferric perchlorate (Fe(ClO4)3) was from Fluka, France. Perchloric acid (HClO4) and sodium hydroxide (NaOH) were used to adjust the pH of the solutions. All chemicals were used without further purification. Fe(III)−EDDS solutions were prepared by mixing appropriate volumes of freshly prepared aqueous solutions of Fe(ClO4)3 and EDDS. 2.2. Computational Method. The calculations were performed at the B3LYP/TZVP hybrid density functional theory level with the Gaussian09 series of programs.21 The SMD framework, based on the polarizable continuum model (PCM) and updated parameter sets, was used to simulate implicitly the solvent medium.22,23 Both gas-phase and solvated structures were fully optimized. Frequency analyses were systematically performed to ensure that energy minima were obtained and to compute free energies. In addition, iron complexes were taken in the high-spin sextet configuration. The complexation energies were computed through the use of thermodynamic cycles because continuum solvation models are not designed to reproduce accurate total free energies in solution.24 Moreover, such cycles benefit from the accuracy that is usually reached from high-level gas-phase calculations. In a recent review,25 it was shown that the choice of a cycle is just a matter of convenience as long as the most accurate experimental values for the free energy of hydrogen ions (H+, HO−) and water are used and no explicit water molecules are added. Recommended values for a standard state of 1 M and 298 K are ΔGsolv(H+) = −265.9 kcal/mol, ΔGsolv(HO−) = −104.7 kcal/mol,26,27 and ΔGsolv(H2O) = −6.32 kcal/mol.28 This procedure minimizes the inherent theoretical uncertainties and is used in our work. Thus, the mixed implicit−explicit or cluster continuum model was not used here. Last, the issue of correct accounting of standard state corrections for the water molecules or water clusters was recently raised and clarified.24,25,29 These corrections are included in our calculations and detailed in this paper. 2.3. Photochemical Experiments. For the degradation of 4-t-BP, the irradiation experiments were performed in a homemade photoreactor placed in a cylindrical stainless steel container. Four fluorescent lightbulb lamps (Philips TL D 15W/05) whose emission spectrum ranges from 300 to 500 nm were separately placed in the four different axes while the photoreactor, a water-jacketed Pyrex tube of 2.8 cm internal diameter, was placed in the center of the setup. The measured spectral irradiance of the four tubes used during these experiments, as well as a UV−vis spectrum of 4-t-BP and Fe(III)−EDDS in water

3. RESULTS AND DISCUSSION 3.1. Characteristics of Fe(III)−EDDS Complex: Theoretical Calculations. Ab initio calculations were performed to characterize the minimum energy structures of FeL− and Fe(OH)nL−(1+n) species at different pH values for n = 1−3 and L4− = [S,S]-EDDS. Recent literature data30 suggests that the species identified as FeL−, FeOHL2−, Fe(OH)2L3−, and Fe(OH)4− are predominant for pH values of [2−7.5], [7.5−9.75], [9.75−10.75], and pH > 10.75, respectively. The experimental complexation energy of FeL− is ΔGc = −27.7 kcal/mol at 298 K33, i.e., log(Kc) = 20.7 (Kc = [FeL−]/ [Fe3+][L4−]). The goals of the theoretical calculations presented in this section are to first characterize the minimum energy structures of FeL− and Fe(OH)nL−(1+n) in solution and then discuss the predominance of these species with respect to pH. 3.1.1. Calculation of Fe3+ Hydration Free Energy. In this section, the above theoretical method is used to calculate the hydration free energy of Fe3+. Experimental values for the hydration free energy of Fe3+ were reported between −103231 and −1043 kcal/mol,32 relative to the older reference hydration energy of −259.2 kcal/mol for H+. When the most accurate value26,27 for ΔGsolv(H+) is used (−265.9 kcal/mol), the experimental hydration free energy of Fe3+ is expected to be between −1052 and −1062 kcal/mol.29 397

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Scheme 1. Thermodynamic Cycle Used for the Calculation of Fe3+ Hydration Free Energy

ΔGhydration(Fe3+) was calculated using the thermodynamic cycle shown in Scheme 1.29 In this scheme, ΔG0‑* is the free energy change of 1 mol of an ideal gas from 1 atm to 1 M and ΔG*solv(H2O) = −6.32 kcal/mol; [H2O] = 55.34 M is the standard state concentration of pure water at T = 298 K. Such a procedure yields ΔGhydration(Fe3+) = −1075.3 kcal/mol, i.e. a discrepancy of 23.2 kcal/mol with the experimental value. This difference arises mainly from the basis set superposition error (BSSE). The BSSE was calculated with the counterpoise correction at the B3LYP/TZVP level. It is +10.1 kcal/mol, and compares well with the BSSE of 15.5 kcal/mol that was computed at the BP86/TZVP level for the solvation energy of Fe3+ ions.33 Including this value for BSSE in our calculation yields ΔGhydration(Fe3+) = −1065.2 kcal/mol, a value which is in good agreement with the experimental value. 3.1.2. The Stability and Structure of FeL−. The complexation of EDDS like ligands proceeds through a multistep mechanism that includes outer-sphere and inner-sphere complexes.34 First, EDDS reacts with the solvated Fe3+ ion to yield rapidly the outersphere complex, where the ligand (L4−) and metal ion are separated by the first solvation shell (Fe(H2O)6·L−). Such an association reaction is expected to proceed without an electronic energy barrier. The inner-sphere complex FeL− is obtained by the successive removal of water molecules. However, L4− is not predominant for pH < 9−10 because the first protonation constant of EDDS is KHL3−/L4− = 10.35 Thus, other outer-sphere complexes like Fe(H2O)6·HL and inner-sphere species like FeHL are necessarily involved in the complexation mechanism (see Scheme 2).

Figure 1. Optimized structure of solvated FeL− at the B3LYP/TZVP level and SMD continuum model (L = [S,S]-EDDS). Hydrogens were removed for clarity. Only those born by nitrogen atoms are shown in this Figure. R(Fe−Oeq) = 2.040 Å, R(Fe−Oax) = 2.077 Å, and R(Fe−N) = 2.240 Å.

3.1.3. Hydroxylated Complexes Fe(OH)nL−(1+n). FeOHL2− and Fe(OH)2L3−. The structures of several hydroxylated complexes were investigated. It was found that the most stable structure for FeOHL2− was obtained with the hydroxyl ligand in the equatorial position (see Figure 2a). At the same time, one of

Scheme 2. FeL− Complexation Mechanism

Figure 2. Optimized structure of solvated FeOHL2− (a) and Fe(OH)2L3− at the B3LYP/TZVP level and SMD continuum model (L = [S,S]-EDDS). Hydrogens were removed for clarity. Only those born by nitrogen atoms and that of the hydroxy ligands are shown in this Figure. (a) Structure for FeOHL2−: R(Fe−OH) = 1.902 Å, R(Fe−Oeq) = 2.040 Å, R(Fe−Oax,top) = 2.088 Å R(Fe−Oax,bottom) = 2.263 Å, R(Fe−Nleft) = 2.313 Å, and R(Fe−Nright) = 2.238 Å. (b) Structure for Fe(OH)2L3−: R(Fe− OHfront) = 1.945 Å, R(Fe−OHleft) = 1.895 Å, R(Fe−Oright) = 2.074 Å, R(Fe−Otop) = 2.058 Å, R(Fe−N) = 2.326 Å and 3.314 Å.

This scheme shows that the formation of FeL− might proceed through different routes depending on pH. In acidic conditions for the complexation reaction of Cd2+ and EDTA (named Y hereafter), it was shown34 that the distribution of outer-sphere complexes was shifted toward lower degrees of protonation, with Cd(H2O)6·HY− being the predominant species for pH > 5 during the course of the reaction. However, at equilibrium, this mechanism is significantly driven toward FeL− for moderate and large pH values. This is because FeL− + H+ and FeHL have similar stabilities. At T = 298.15 K, G(aq)(FeL−) − G(aq)(FeHL) + G(aq)(HL3−) − G(aq)(L4−) − RT log(KHL3−/L4−) = −0.47 kcal/mol at the B3LYP/TZVP level and the SMD continuum model. Thus [FeL−][H+]/[FeHL] ∼ 1, and for pH > 2, [FeL−] ≫ [FeHL]. Thus, FeL− is the predominant species for pH values greater than 2. The optimized structure of FeL− is shown in Figure 1.

the COO− ligands is displaced. Further hydroxylation yielded a structure (see Figure 2b) where one nitrogen atom is displaced and stabilized by HO−. The respective N···HOFe hydrogen bond is 1.919 Å. In this structure, two COO− groups are displaced. The second hydroxy lies in the axial position. The metal−ligand bond lengths increase consistently with the degree of hydroxylation. The stability constants pK1 and pK2 and respective free energies ΔG1 and ΔG2 were computed with the thermodynamic cycle shown in Scheme 3. 398

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ΔG1 was set to 7.63 kcal/mol, i.e., a difference of 2 kcal/mol with our results. This small discrepancy lies within the intrinsic error that is expected from the quantum calculations. The theoretical results are shown in Figure 3. Our results show that the

Scheme 3. Thermodynamic Cycle Used for the Calculation of Fe(OH)nL−(1+n) Free Energy of Complexation

In this scheme, the notations are similar to those used by Bryantsev and co-workers.36 The recommended experimental solvation energy was used for the hydroxyl ion,25 i.e., ΔG*solv(HO−) = −104.7 kcal/mol. ΔG1= −5.63 kcal/mol (log K1 = 4.13) and ΔG2 = −5.72 kcal/mol (log K2 = 4.19) were found from our calculations. Fe(OH)3L4−. Optimization of the trihydroxylated complex yielded a species where EDDS was unfolded and bonded to the metal ion through one carboxy group only. This result shows that for large pH values, Fe(OH)3L4− might not be a stable species. To further address this point, the solvent was modeled explicitly in Fe(OH)3L4− to keep the hexadentate character of the metal species. This enabled us to compare the stability of Fe(OH)2L3− and that of Fe(OH)3L4−. The respective free-energy change of the reaction shown in Scheme 4 was ΔG3 = +8.02 kcal/mol.

Figure 3. Proposed theoretical model distribution of the predominant species for the Fe(III)-[S,S]-EDDS complexation reaction as a function of pH.

Scheme 4. Formation of the Tri-hydroxylated Fe(OH)3L4− Complex

concentration of Fe(OH)3 remains small at all pH values in agreement with the proposed experimental assignment of the dominant species for large pH values. In addition, the predominant species are identical to those proposed by Orama et al.30 This good agreement between theoretical and experimental results shows that the complexation mechanism of [S,S]EDDS and Fe(III) may be described by the detailed suite of reactions discussed in our work. 3.2. Photochemical Reactivity of Fe(III)−EDDS Complex. The photoactivity of Fe(III)−EDDS was investigated following the degradation of 4-t-BP (0.05 mM) in the presence of Fe(III)−EDDS (0.1 mM) under polychromatic radiation. The results are shown in Figure 4. Controlled experiments were also performed and the results showed its stability under direct photolysis and dark reaction. In the presence of UV light and Fe(III)−EDDS, approximately 17% of the 4-t-BP was transformed after 10 min of irradiation but only 5% of it was removed in the following 50 min. The degradation of 4-t-BP is due to the reaction with oxidative radicals (HO•, HO2•, O2•−) which were produced during the rapid photochemical reactions involving the Fe(III)−EDDS complex under UV light.9 In fact, the concentration of Fe(III)−EDDS was also detected and its degradation is very fast and important, around 90% degradation after 10 min of irradiation. The reaction was almost suspended after 10 min of irradiation, and it was obviously caused by the fast consumption of Fe(III)−EDDS complex at the beginning of the reaction. 3.3. Determination of the Second-Order Rate Constant between HO• and 4-t-BP. The rate constant of the reaction between HO• and 4-t-BP was measured by competition kinetics with 2-propanol, a compound of known reaction rate constant with HO•. H2O2 (1.0 mM) photolysis was used as source of HO•, which would induce the following main reactions (R1−R4) in the system containing both 2-propanol and 4-t-BP.

Thus, further hydroxylation might rather yield the Fe(OH)3 and Fe(OH)4− species. It was proposed in earlier work30 that only Fe(OH)4− was predominant for large pH values. This point is now investigated. 3.1.4. Equilibrium Speciation of Fe(III)-[S,S]-EDDS Complexes. The equilibrium speciation of Fe(III)-[S,S]-EDDS complexes (see Scheme5) was investigated by solving numeriScheme 5. Equilibrium Speciation of Fe(III)-[S,S]-EDDS

cally the respective set of equilibrium relationships. In this scheme, the experimental value Kexp (log K0 = 20.3),35 the 0 theoretical constants calculated in our work, Kth 1 = exp(−ΔG1/ RT) and Kth 2 = exp(−ΔG2/RT), and the experimental stability constants of Fe(OH)3 and Fe(OH)4−, i.e., log(K)exp Fe(OH)3 = 14.3 37 and log(K)exp = 34.30, respectively, were used. K3 = Fe(OH)4− th th

exp K1 K2 exp (K)exp and K4 = Kexp Fe(OH)3/K0 Fe(OH)4−/KFe(OH)3 were determined from the above values. Our calculations showed that excellent agreement between experimental data and the theoretical results could be obtained if

H 2O2 + hν → 2HO• 399

R •fOH(M s−1)

(R1)

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Figure 4. Photodegradation of 4-t-BP (5 × 10−5 M) in the presence of Fe(III)−EDDS (1 × 10−4 M) at pH 4.0 under polychromatic wavelengths. Filled triangles represent the 4-t-BP degradation and empty triangles the Fe(III)−EDDS degradation.

HO• + 4‐t ‐BP → products H 2O2 + HO• → H 2O + HO•2

k4 − t − BP, •OH

Figure 5. Degradation rate of 4-t-BP (5 × 10−5 M) under monochromatic irradiation (λ = 313 nm) in the presence of different 2-propanol concentration. The line represents the fit using the Eq 1.

at pH 4.0 (Figure SM2 of the Supporting Information). The rate of 4-t-BP degradation increased with increased concentrations of Fe(III)−EDDS in the range of 5 × 10−5 to 4 × 10−4 M, but much higher concentration of Fe(III)−EDDS (such as 5 × 10−4 and 6 × 10−4 M) inhibited 4-t-BP degradation. Li et al.5 reported similar degradation rate variation under various Fe(III)−EDDS complex concentrations. These results gave clear evidence that photolysis of Fe(III)− EDDS could induce the degradation of 4-t-BP due to the reaction between 4-t-BP and HO• which was formed during the photolysis of Fe(III)−EDDS. But HO• could also react with Fe(III)−EDDS and the oxidation products of EDDS. When the concentration of Fe(III)−EDDS was high, Fe(III)−EDDS played the role as a competitor for HO•. In fact, considering that the second-order rate constant between Fe(III)−EDDS and hydroxyl radical has been estimated around 4 × 108 M−1 s−1 (ref 40), we can argue that at lower concentration of Fe(III)− EDDS (5 × 10−5 M) about 98% of photogenerated hydroxyl radicals react with 4-t-BP. On the other hand, at 6 × 10−4 M Fe(III)−EDDS, 23% of HO• reacts with the iron complex. Therefore, the appropriate concentration of Fe(III)−EDDS should be chosen for 4-t-BP degradation. 3.5. Effect of pH and Oxygen. To better understand the effect of pH value during the photodegradation of 4-t-BP in the presence of Fe(III)−EDDS, experiments at different pHs between 2.6 and 9.3 were conducted. The results reported in Figure 6 show the rapidly increasing degradation rate of 4-t-BP (R4‑t‑BP) between pH 2.6 and 4.5 and the slow increase of R4‑t‑BP between pH 4.5 and 8.0, whereas R4‑t‑BP started to decrease at pH higher than 8.0. R4‑t‑BP was calculated from the data obtained at the beginning, after 10 min of irradiation to ensure the pH was unchanged. Li et al.5 reported similar results at pH between 3.1 and 8.0. The effect of oxygen on 4-t-BP photodegradation in the presence of Fe(III)−EDDS is shown in Figure 7. The faster decay of 4-t-BP was noticed at higher oxygen concentration (bubbling oxygen), while at the lower oxygen concentration (bubbling nitrogen) the 4-t-BP removal was retarded. Another control experiment was performed to confirm the mechanism of the photochemical process of the Fe(III)−EDDS complex. 2Propanol (5 mM) was added to the solution which was bubbled with nitrogen. The result showed that 4-t-BP was not degraded because of the quenching of HO• by 2-propanol, confirming the

(R2)

k H2O2, •OH = 2.7 × 107 M−1s−1

(R3)

2‐propanol + HO• → products k 2‐propanol, •OH = 1.9 × 109 M−1s−1

(R4)

Values for the rate constants are taken from ref 40. The application of the steady-state approximation to the concentration of HO• yielded the following equation (eq 1) to describe the degradation rate of 4-t-BP (y = Rd4‑t‑BP) as a function of 2-propanol concentration (x = [2pr]). R 4d‐t ‐BP =

R •fOHk4‐t ‐BP,•OH[4‐t ‐BP] k4‐t ‐BP,•OH[4‐t ‐BP] + k H2O2,•OH[H 2O2 ] + k 2pr,•OH[2pr]

(1)

The data reported in Figure 5 were fitted with a rational equation type y = A/(1 + Bx) with A=

R •fOHk4‐t ‐BP, •OH[4‐t ‐BP] k4‐t ‐BP, •OH[4‐t ‐BP] + k H2O2,OH[H 2O2 ]

and B=

k 2pr, •OH k4‐t ‐BP, OH[4‐t ‐BP] + k H2O2, •OH[H 2O2 ] •

The formation rate of HO• (R•f OH) and the second-order rate constant of the reaction between HO• and 4-t-BP (k4‑t‑BP,•OH) were estimated to be (1.24 ± 0.07) × 10−9 M s−1 and (1.61 ± 0.26) × 1010 M−1 s−1, respectively. According to the literature, the second-order rate constants of the reaction between HO• and o-cresol and p-cresol were 1.1 × 1010 M−1 s−1 and 1.2 × 1010 M−1 s−1, respectively.38 Both of them had a chemical structure similar to that of 4-t-BP; therefore, the result obtained for k4‑t‑BP,•OH seems reasonable. Contrary to previously reported works,39 in which radicals generated by the reaction between 2-propanol and hydroxyl radical could react with selected molecules, no degradation of 4-t-BP has been detected at high 2pr concentrations under our adopted conditions. 3.4. Effect of Fe(III)−EDDS Complex Concentrations. 4t-BP disappearance was followed at different initial concentrations of Fe(III)−EDDS ranging from 5 × 10−5 to 6 × 10−4 M 400

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HO2• ↔ O2•− + H+

pK a = 4.88

(R7)

• − O•− 2 + HO2 + H 2O → H 2O2 + O2 + OH

k = 9.7 × 107 M−1s−1

(R8)

Fe2 + + H 2O2 → Fe3 + + HO• + OH−

k = 76 M−1s−1 (R9)

As we mention in the paper of Huang et al.,40 the observed effect of pH could be due to the formation of HO2•/O2•− radicals and/or to the presence of different forms of the complex Fe(III)−EDDS as a function of pH. However, for the first part of the effect of pH, until pH 6.0, we can exclude the effect of Fe(III)−EDDS speciation. Indeed, as we evaluate by theoretical calculation the second form of the complex appears from pH 6.0 (Figure 3). On the contrary, the decrease of the degradation rate of 4-t-BP (R4‑t‑BP) from pH 8.0, corresponds to the presence of the second form FeOHL2− (with L4− = [S,S]-EDDS) at 50% and 50% of the starting form in acidic pH FeL−. For the first time, we prove that the hydroxylated form FeOHL2− of such a complex is less efficient photochemicaly in terms of a photoredox process (R5). The observed increase of the degradation rate of 4-t-BP (R4‑t‑BP) until pH 8.0 is certainly due to the iron cycle and the relative concentration between Fe(III) and Fe(II) species. These relative concentrations are strongly impacted (reactions R10−R13) by the presence of HO2•/O2•− radicals photogenerated from the complex Fe(III)−EDDS (reactions R5 and R6).

Figure 6. Effect of pH value on the degradation rate of 4-t-BP (5 × 10−5 M) in the presence of Fe(III)−EDDS (1 × 10−4 M) under polychromatic irradiation.

3+ Fe 2 + + O•− + H 2O2 + 2OH− 2 + 2H 2O → Fe

k = 1.0 × 107 M−1s−1

(R10)

Fe2 + + HO•2 + H 2O → Fe3 + + H 2O2 + OH− k = 1.2 × 106 M−1s−1 Fe3 + + O2•− → Fe2 + + O2

Figure 7. Effect of oxygen on 4-t-BP degradation using Fe(III)−EDDS (1 × 10−4 M) at pH 4.0 under polychromatic irradiation.

(R11)

k = 5.0 × 107 M−1s−1 (R12)

Fe3 + + HO•2 → Fe 2 + + O2 + H+ k < 1.0 × 103 M−1s−1

formation of this radical in the photochemical process from the Fe(III)−EDDS complex. Moreover, we can conclude that the effect of oxygen could be mainly attributed to the reactivity of oxygen on the 4-t-BP radical formed subsequently by the hydroxyl radical attack. On the other hand, photogenerated EDDS• (R5) can play a crucial role during the 4-t-BP degradation in the absence of oxygen. Under such conditions, the 4-t-BP degradation is still reported and such a trend can be attributed to an unknown reactivity. To explain experimental data reported in Figure 7, we consider that the reaction (i.e., oxidation) between EDDS• and 4-t-BP can be relevant in deaerated solutions. Moreover EDDS• should enhance the hydroxyl radical formation via reoxidation of Fe2+ into Fe3+ followed by its direct photolysis in water. Both reactivities could be expected and are supported by experimental work in the presence of 2-propanol. In fact, under such conditions the EDDS• can be quantitatively quenched by 2-propanol at high concentrations. The formation of HO• can be explained by reactions R5−R9 (as discussed in refs 38 and 42): Fe(III)−EDDS + hν → Fe2 + + EDDS•

(R5)

EDDS• + O2 → product + O2•−

(R6)

(R13)

Values for the rate constants are taken from ref 43 (R10, R11, and R13) and ref 44 (R12). To generate the maximum concentration of HO• we need a photochemically efficient Fe(III) species, and so the formation of Fe(II) species is essential for the Fenton process and the production of OH radical (R9). Moreover, in the iron cycle the oxidation of Fe(II) into Fe(III) is strongly pH-dependent in terms of efficiency and also in terms of Fe(III) species formed. The slowing increment of the 4-t-BP degradation rate observed after pH 4.5 can be due to the higher oxidation rate of Fe(II) observed at higher pH leading to the regeneration of Fe(III). Moreover, for pH > 4.0, Fe(III) is unstable and formation of insoluble iron oxides, which present the lowest photoactivity, is expected.

4. CONCLUSION Theoretical and experimental approaches were used to investigate the stability and photoreactivity of Fe(III)−EDDS complexes between 300 and 500 nm. This study was performed in a large range of pH, which is the main parameter influencing the efficiency of the photochemical process. 401

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From this work we can conclude that before pH 8 the main process responsible for the observed trend is attributed to the iron cycle reactivity between Fe(III)/Fe(II). The relative concentration between Fe(III) and Fe(II) is mainly due to their reactivity with the radicals HO2•/O2•−, photogenerated from the Fe(III)−EDDS complex, and to their oxido-reduction and solubility in aqueous solution. Moreover, the main new result is coming from theoretical calculation for the distribution of the predominant species of the complex Fe(III)−EDDS as a function of pH. We demonstrated that the most photoactive form is the nonhydroxylated one present as the majority at pH lower than 8.0.

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ASSOCIATED CONTENT

S Supporting Information *

Measured spectral irradiance of the four tubes used during these experiments, UV−vis spectrum of 4-t-BP and Fe(III)−EDDS in water solution, plot of 4-t-BP degradation versus irradiation time for different Fe(III)−EDDS concentrations, and plot of effect of Fe(III)−EDDS concentration on the degradation of 4-t-BP. This material is available free of charge via the Internet at http://pubs. acs.org.

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AUTHOR INFORMATION

Corresponding Authors

*Phone: +33 (0)4 73 40 71 73. E-mail: gilles.mailhot@ univ-bpclermont.fr. *Phone: +86-21-6564-2030. E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS

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REFERENCES

The authors gratefully acknowledge financial support from China Scholarship Council for Y.W. to study at the Blaise Pascal University in Clermont-Ferrand, France. This work was supported by the National Natural Science Foundation of China (NSFC 21077027), Shanghai Natural Science Fund (12ZR1402000), Shanghai Tongji Gao Tingyao Environmental Science & Technology Development Foundation (STGEF), and the Graduate Innovative Fund of Fudan University (13). This work was also supported by one project CMIRA founded by the Rhône-Alpes Région, the “Federation des Recherches en Environnement” through the CPER “Environnement” founded by the “Région Auvergne,” the French government, FEDER from the European community, and the CNRS program EC2CO.

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