sulfite

Article pubs.acs.org/IECR

Decolorization of Orange II in Aqueous Solution by an Fe(II)/sulfite System: Replacement of Persulfate Long Chen, Xinzi Peng, Jihao Liu, Jinjun Li, and Feng Wu* Department of Environmental Science, School of Resources and Environmental Science, Wuhan University, Wuhan, 430079, P. R. China S Supporting Information *

ABSTRACT: A novel process for decolorizing dyes with sulfate radicals (SO4•−) using an Fe(II)/sulfite system is reported in this manuscript. The objective of this study was to investigate the conditions under which Fe(II) activates Na2SO3 to produce SO4•− and decolorize organic dyes. Orange II, Rhodamine B, Indigo Carmine, and Reactive Brilliant Blue X-BR could be efficiently decolorized using this novel system, which was compared with the Fe(II)/persulfate and Fenton (Fe(II)/H2O2) systems. The Fe(II)/sulfite system surpassed the other two in the decolorization of these dyes, and detailed mechanisms of the Fe(II)/sulfite system were researched. Primary radical identification through quenching experiments using tert-butyl alcohol and ethanol confirmed the existence of SO4•−, HO•, and SO5•−. A kinetic model was established for the halide ion effect, and kI−,SO4•− (3.2 × 1011 mol−1 L s−1) and RfSO4•− (10−4−10−3 mol L−1 s−1) were indirectly derived. In conclusion, the Fe(II)/sulfite system is a good candidate for use in detoxifying water contaminants. In previous studies of SO4•−, potassium persulfate (K2S2O8) was usually used because it could produce SO4•− by radiolytic,19,24,26 photolytic,23,27 or thermal activation.28 Alternatively, SO4•− can also be generated through a pathway by which persulfate or peroxymonosulfate salts are activated by transition metals.1,29,30 For example, Fe(II)/persulfate system can react according to eqs 2 and 3.1 Interestingly, both SO4•− and H2O2 could be generated in the Fe(II)/persulfate system, and hence HO• was produced from the decomposition of H2O2.

1. INTRODUCTION Currently, various oxidants, such as hydrogen peroxide, persulfate, and peroxymonosulfate salts, are used in water decontamination. By combining them with transition metals, such as Fe(II), these oxidants can propagate transient species (usually free radicals) that are more powerful in degrading organic contaminants than the parent oxidants.1 The most wellknown, Fenton reagent (Fe(II)/H2O2), generates hydroxyl radicals (HO•) through eq 1. There are two theories to explain the oxidation of the reaction, the Haber−Weiss cycle vs highvalent oxorion (ferryl) species mechanism,2−8 and Chen et al. have illustrated them in detail.9 The generation of HO• radical has been proven by electron spin resonance technique, hydroxylation of probe molecules, and kinetic models.10,11 According to the Haber− Weiss cycle, HO• is generated by one-electron reduction of H2O2 by Fe(II).12,13 Rigg et al. have determined the rate constant of the Fenton reaction, and under room temperature the theoretical value of k1 is about 70 L mol−1 s−1.14 As it is environmentally friendly, H2O2 is widely used in water treatment.

(1)

It has been discovered that the sulfate radical (SO4•−, E0 = 2.5−3.1 V, NHE)15 is as powerful as HO• (E0 = 1.8−2.7 V, NHE)16,17 and, in general, SO4•− is more selective for oxidation than HO•.1 In addition, SO4•− and HO• radicals are likely to react with substrates via different paths. HO• may react with substrates by electron transfer, hydrogen abstraction, or addition mechanisms, whereas it is generally accepted that SO4•− is more prone to electron transfer.18 However, it is rather difficult to differentiate SO4•− and HO• radicals merely from products because the majority of the products formed by sulfate radical attack on aromatics are hydroxylation products, which are also intermediates of hydroxyl radical attack.19−25 © 2012 American Chemical Society

(2)

2H 2O + S2 O82 − → 2HSO4 − + H 2O2

(3)

In the atmospheric liquid phase, Fe(II/III) catalyzes the conversion of S(IV) to S(VI) in the presence of dissolved oxygen. Many investigations of this process have extrapolated detailed mechanisms under various conditions,31−38 since it is very important for the generation of acid rain. On the basis of the mechanism of catalysis by iron, much attention has also been paid to flue gas desulfurization (FGD), which is an artificially enhanced process of transformation of S(IV) to S(VI) for the purpose of desulfurization. Kinetics and mechanisms have been proposed.39−42 Among the proposed mechanisms, the free-radical mechanism seems to be correct, since it explains the scavenging effect of alcohols or halide ions.31,39 The proposed mechanism for the Fe(II)/sulfite system is described in eqs 4−11 according to FGD.39,41 The reaction rate constants are from the literature.41−43 The stability constant of FeHSO3+ (k4) is from the estimation by Zhang et al. (104 mol−1 L),41 and the stability constant of FeSO3+

Fe2 + + H 2O2 → Fe3 + + OH− + HO• (k1 = 70 L mol−1 s−1)

Fe2 + + S2 O82 − → Fe3 + + SO4•− + SO4 2 −

Received: Revised: Accepted: Published: 13632

July 31, 2012 September 25, 2012 September 28, 2012 September 28, 2012 dx.doi.org/10.1021/ie3020389 | Ind. Eng. Chem. Res. 2012, 51, 13632−13638

Industrial & Engineering Chemistry Research

Article

selected as substrates. The information of the substrates is detailed in Table S1 (Supporting Information). KOH (AR), NaOH (AR), and H2SO4 (AR) were from Sinopharm Chemical Reagent Co., Ltd., and they were used to adjust the pH of solutions. Tert-butyl alcohol (TBA) and ethanol (EtOH, AR), used to quench the relevant radicals, were from Sinopharm Chemical Reagent Co., Ltd. The halide ions introduced into the experiments were from potassium halides (KI, KCl, and KBr, Sinopharm Chemical Reagent Co., Ltd.). All chemicals were used without further purification. The water used was Milli-Q water (18.2 MΩ·cm). 2.2. Oxidation Reaction. All experiments were performed in 500-mL open beakers (except for those conducted under N2), at ambient temperature and under normal laboratory light. Substrates at desired concentrations were spiked in a 500-mL flask. Each solution was constantly stirred with a PTFE-coated magnetic stirrer, and Fe(II) or Fe(III) was added at the desired concentration. The pH of the solution was adjusted with diluted KOH/NaOH and H2SO4 to the desired value. Each run was switched on by the careful addition of the desired concentration of Na2SO3, K2S2O8, or H2O2. The pH was then readjusted to the desired value (±0.05 unit) as quickly as possible. During the remaining reaction process, the pH was not controlled. Samples were withdrawn, and measurements were immediately made. Most of the batch experiments were conducted in triplicate, and thus some figures show error bars. The initial concentration of the substrate Orange II was 10 mg L−1 in all batch experiments. Batch experiments to investigate radical identification and the effects of halide ions were also conducted. For the latter purpose, potassium halides were introduced into the experiments. TBA and EtOH were employed to differentiate SO4•− and HO•. In this case, TBA or EtOH or halide ions at different concentrations, together with Orange II, were added to the solutions in a reaction volume of 500 mL, and the solution preparation procedure was similar to that above. The concentrations of TBA and EtOH in the solutions were 500 mmol L−1. The effect of oxygen was also investigated by purging a 1000-mL closed reaction cylinder with mixing gas (oxygen and nitrogen) which contained different ratios of oxygen. To maintain the same condition, the total flow rate of the mixing gas is constantly 1 L min−1. A glass cover was used to block air entrance but a small hole was left to allow for a slow leak. 2.3. Analysis. A UV-1601 spectrophotometer (Shimadzu, Japan) was employed to measure the concentrations of Orange II, Rhodamine B, Indigo Carmine, and RBB X-BR at wavelengths 485, 554, 610, and 596.5 nm, respectively. For the analysis of RBB X-BR, a 3-cm cuvette was used whereas, for the remaining dyes, a 1-cm cuvette was used. A 1-cm quartz cuvette was used to scan spectra of Orange II. The pH was measured using a pH meter (Mettler Toleod LE409). TOC analysis was performed on an Analytik Jena multi N/C 2100.

(k10) was estimated from the average of values cited by Lente and Fábián.42 Fe2 + + HSO3− ⇌ FeHSO3+

(log k4 = 4)

(4)

4FeHSO3+ + O2 → 4FeSO3+ + 2H 2O (rapid equilibrium)

(5)

FeSO3+ → Fe2 + + SO3•−

SO3•− + O2 → SO5•−

(k f = 0.19 s−1)

(6)

(k 7 < 109 mol−1 L s−1) (7)

SO5•− + HSO3− → SO3•− + HSO5− (k 8 = (104−107) mol−1 L s−1)

(8)

Fe2 + + HSO5− → SO4•− + Fe3 + (k 9 = (104−107) mol−1 L s−1)

Fe3 + + HSO3− ⇌ FeSO3+ + H+

(9)

(log k10 = 2.45) (10)

SO5•− + HSO3− → SO4 2 − + SO4•− + H+ (k11 = (104−107) mol−1 L s−1)

(11)

In addition, researchers have confirmed that several organic acid buffers can enhance the S(IV) removal efficiency in FGD but, unexpectedly, during this process degradation of these buffers was observed.39 Moreover, Meserole proposed a simple kinetic model for the degradation of organic acid based on the radical reaction.44 Hence, it was hypothesized that Fe(II)/sulfite system could also react with organic contaminants in water, leading to their degradation. In this work, we focused on the decolorization effect, and Orange II was selected as the model substrate to investigate the decolorization potential of Fe(II)/ sulfite system. To our knowledge, Fe(II)/sulfite system has never been used as an advanced oxidation process (AOP) in water decontamination, and there has previously been no insight into the utilization of various transient radicals produced by Fe(II)/sulfite system. The objective was to demonstrate the oxidizing strength of the Fe(II)/sulfite system, and to highlight its advantages over other AOPs, such as Fe(II)/persulfate and Fenton systems. The results showed that the Fe(II)/sulfite system is a very strong oxidant that achieves more effective decolorization than Fe(II)/ persulfate and Fenton systems under the experimental conditions. Hence, it is expected that Fe(II)/sulfite system will open up a new area of potential research in the development of AOPs, as an alternative to systems involving persulfate.

2. EXPERIMENTAL METHODS 2.1. Chemicals. Ferrous sulfate (FeSO4·7H2O), ferric sulfate (Fe2(SO4)3), and hydrogen peroxide (H2O2, 30%, AR) were supplied by Sinopharm Chemical Reagent Co., Ltd.; potassium persulfate (K2S2O8, ≥99.5%) was obtained from Shanghai Aijian Reagent Co., Ltd.; and sodium sulfite (Na2SO3, AR) was purchased from Shanghai Zhanyun Chemical Co., Ltd. Orange II (Sinopharm Chemical Reagent Co., Ltd.), Rhodamine B (Hunan Institute of Geological Experiment), Indigo Carmine (Sinopharm Chemical Reagent Co., Ltd.), and Reactive Brilliant Blue X-BR (RBB X-BR, Sinopharm Chemical Reagent Co., Ltd.) were

3. RESULTS AND DISCUSSION 3.1. Decolorization of the Dyes in Different Systems. To investigate the decolorization effect of Fe(II)/sulfite system, several dyes were selected. Table S2 (Supporting Information) listed the results of decolorization of Orange II, Rhodamine B, Indigo Carmine, and RBB X-BR. It was found that the Fe(II)/ sulfite system can achieve good decolorization of all of these organic dyes. The results also showed that the Fe(II)/sulfite system is the most powerful for decolorizing the selected dyes. 13633

dx.doi.org/10.1021/ie3020389 | Ind. Eng. Chem. Res. 2012, 51, 13632−13638


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Although the concentration of HSO3− is about 0.85 mmol L−1, this slight difference in HSO3− species concentration caused by pH variance may influence the equilibrium of eq 4, leading to different decolorization effects. Therefore, by controlling S(IV) species distribution, pH 4−6 should be the most efficient for radical generation. The results also showed that, at pH 4, both Fe(II) and Fe(III) mostly exist in the complexed form when Na2SO3 is present. This can also explain the similar decolorization effects of Fe(II)/sulfite and Fe(III)/sulfite systems. Since most of the Fe(II) is in the complexed form, the rapid transformation of Fe(II) to Fe(III) is possible through eqs 4 and 5. In this sense, the effect of Fe(II) is equivalent to that of Fe(III), and we may maintain continuous elimination of organics by proper addition of Na2SO3 in the presence of Fe(II/III). Another factor controlled by pH is the activity of iron. Although Figure S2 (Supporting Information) shows that free Fe(II) does not precipitate below pH 8.5, this does not mean that iron is still active in that pH range. As described above, Fe(II) rapidly transforms to Fe(III). However, Figure S4 (Supporting Information) shows that free Fe(III) begins to precipitate at pH 4, and almost totally precipitates when the pH is greater than 5. Hence, at pH above 4 the decolorization of Orange II decreased, largely due to the deactivation of iron. Therefore, at pH 4 the Fe(II)/sulfite system is the most effective. 3.3. Effect of Dosage Ratio in the Decolorization of Orange II by Fe(II)/Sulfite System. The dependence of the decolorization of Orange II on the dosage ratio is demonstrated in Figure 3. Generally, 10:1 is the optimum dosage ratio, and hence it was selected for the remaining batch experiments.

Hence, the potential of the Fe(II)/sulfite system was tentatively revealed. All of the decolorization processes of Orange II under various conditions are compared in Figure 1. The results show that no

Figure 1. Decolorization of Orange II by three systems and control experiments. Conditions: [Orange II]0 = 10 mg L−1, [FeSO4]0 = 0.1 mmol L−1, [Fe2(SO4)3]0 = 0.05 mmol L−1, [Na2SO3]0 = [K2S2O8]0 = [H2O2]0 = 1 mmol L−1, pHini = 4.

decolorization was observed in the mere presence of Fe(II) or Na2SO3. The Fe(II)/sulfite system decolorized Orange II more rapidly during the initial 100 min but, allowing an extended reaction time, almost the same effect was produced by the Fenton system. In addition, activation of Na2SO3 by Fe(III) was also investigated, and the results in Figure 1 show that the curves for Fe(II)/sulfite and Fe(III)/sulfite systems almost merge. This can be explained by the fact that the slow decomposition of FeSO3+ (eq 6) is the rate-determining step in the redox cycle of iron.42 Compared to the decomposition of FeSO3+, its formation from both Fe(II) (eq 4, 5) and Fe(III) (eq 10) is very fast, and hence the following reactions almost occur simultaneously. This point also provides a strategy for enhancing the reaction between Fe(II) and Na2SO3 by promoting the decomposition of FeSO3+. 3.2. Effect of pH in the Decolorization of Orange II by Fe(II)/Sulfite System. The pH mainly affects the species distribution of S(IV) and the activity of iron. Figure 2 shows that

Figure 3. Effect of dosage ratio on the decolorization of Orange II. Conditions: [Orange II]0 = 10 mg L−1, [FeSO4]0 = 0.1 mmol L−1, pHini = 4. Samples were withdrawn after reacting for 30 min.

It has also been observed that elevated concentrations of Na2SO3 reduce the decolorization effect. The reaction of SO4•− radicals or HSO5− with Na2SO3 may account for this (eq 12, 13). It has to be noted that HSO5− is a very reactive anion. Several scholars have confirmed that SO4•− radicals may be generated via the pathway of transition metal catalysis of HSO5−.1,45,46 Fe(II), for example, can catalyze the transformation of HSO5− to SO4•− radicals at a significant rate (eq 9). In addition, free Fe(II) is also a factor in depleting SO4•− radicals (eq 14). Consequently, Fe(II) is able to generate SO4•− radicals by reaction with HSO5− (eq 9) but deplete SO4•− radicals through eq 14.

Figure 2. Effect of pH on the decolorization of Orange II. Conditions: [Orange II]0 = 10 mg L−1, [FeSO4]0 = 0.1 mmol L−1, [Na2SO3]0 = 1 mmol L−1. Samples were withdrawn after reacting for 30 min.

the most effective decolorization of Orange II was at pH 4. The decolorization rate initially rose as the pH increased from 1 to 4, then decreased at pH 4 to 10. The S(IV), Fe(II), and Fe(III) species distribution curves are depicted in Figures S1−S4 (Supporting Information). MEDUSA software was used to calculate that, in Fe(II)/sulfite system at pH values below 4 and over 6, the theoretical concentrations of HSO3− are less than 0.90 mmol L−1 whereas, at pH 4−6, the concentration of HSO3− is almost constant at 0.90 mmol L−1. 13634

SO4•− + HSO3− → HSO4 − + SO3•−

(12)

HSO5− + HSO3− → 2SO4 2 − + H 2O

(13)

SO4•− + Fe 2 + → SO4 2 − + Fe3 +

(14)



Article

much as 13.9% of the decolorization of 10 mg L−1 Orange II. Many researchers have investigated the relationship between SO4•− and HO•, and it is suggested in eqs 15 and 16 that SO4•− is such a strong oxidant that it can even oxidize H2O or HO− to HO•.18 Since the reaction rate constant of SO4•− with HO− is about 10000-fold greater than with H2O, it is not surprising that under acidic conditions SO4•− is the predominant radical, whereas at alkaline pH HO• is the primary oxidative species.18

3.4. Effect of Oxygen in the Decolorization of Orange II by Fe(II)/Sulfite System. It can be seen in eq 7 that O2 reacts with SO3•− to form SO5•−, the precursor of SO4•−. Therefore, O2 is necessary for the formation of SO4•−. The effect of oxygen was depicted in Figure 4. With the ventilation of O2, the

alkaline pH: SO4•− + OH− → SO4 2 − + HO• (k15 = (4.6−9.1) mol−1 L s−1)

(15)

all pH values: SO4•− + H 2O → SO4 2 − + HO• + H+

Figure 4. Effect of oxygen on the decolorization of Orange II. Conditions: [Orange II]0 = 10 mg L−1, [FeSO4]0 = 0.1 mmol L−1, [Na2SO3]0 = 1 mmol L−1, pHini = 4. Total flow rate of gas is 1 L min−1 by mixing oxygen with nitrogen. Samples were withdrawn after reacting for 5 min.

(k16 = (103−104) mol−1 L s−1)

(16)

where the k values of eq 15 and eq 16 are taken from literature.35 However, it is very interesting to observe that the decolorization of Orange II by both Fe(II)/persulfate and Fenton systems could be totally stalled by EtOH, whereas 11.3% was still not inhibited in the Fe(II)/sulfite system. Equations 6−8 suggest that SO5•− and SO3•− were also generated. Both SO5•− and SO3•− are fairly inert toward alcohols (k ≤ 103 mol−1 L s−1).1,43 In addition, SO3•− is not stable because it is prone to oxidation by O2 into SO5•− (eq 7). Accordingly, it is hypothesized that SO5•− was responsible for the remaining decolorization (11.3%) of Orange II. In addition, quenching experiments involving the Fenton reaction proved that the decolorization was caused completely by HO•. For Fe(II)/ persulfate system the results showed that, although SO4•− accounted for approximately 66.6% of the decolorization of Orange II, HO• accounted for as much as 33.3%. This is because H2O2 was probably generated in the Fe(II)/persulfate system through eq 3, leading to the formation of HO•. Alternatively, HO• can also be produced through eqs 15 and 16. 3.6. Implication of Halide Ions in the Decolorization of Orange II by the Fe(II)/Sulfite System. The effects of halide ions on the decolorization of Orange II are shown in Figures S7− S9 (Supporting Information). These indicate that halide ions tend to quench this reaction in the increasing order: Cl− < Br− < I−.43 The residual decolorization effect of Orange II observed, despite the addition of KBr and KCl (Figures S7 and S8, Supporting Information), may be an indication of the oxidation of secondary halide radicals.53 It has to be pointed out that halide ions quenched not only SO4•− but HO•. The second-order rate constants for reactions between halide ions and SO4•− or HO• are listed in Table 2. Halide ions react with SO4•−/HO• radicals as shown in eq 17,

decolorization is rather rapid, and there are not remarkable differences among the decolorization effects under different ratios of O2 in the mixing gas after reacting for 30 min. Therefore, the samples were withdrawn after initial 5 min instead of 30 min. Under N2 ventilating condition, the decolorization rate of Orange II was almost 0. As the ratio of O2 in the mixing gas increased, the decolorization effect was enhanced and then kept almost unchanged. It was because the incoming O2 compensated for the consumed O2. 3.5. Identification of Radicals by Quenching Studies in Fe(II)/Sulfite System. To cast more light on the reaction mechanism, quenching studies were conducted by adding TBA and EtOH to identify the primary effective radicals. The concentrations of TBA and EtOH were 500 mmol L−1, corresponding to a molar ratio of 500:1 of alcohol vs oxidant. This method has been extensively used by the research group of Burrows.47−52 It is suggested that alcohols containing αhydrogen (EtOH) can react with HO• and SO4•− at significant rates. The reaction rate constant of EtOH with HO• is approximately (1.2−2.8) × 109 mol−1 L s−1, which is about 50fold greater than with SO4•−, about (1.6−7.7) × 107 mol−1 L s−1.1 In contrast, alcohols without an α-hydrogen (TBA) react with HO• and SO4•− very differently. The reaction rate constant of TBA with HO• is approximately (3.8−7.6) × 108 mol−1 L s−1, which is about 1000-fold greater than with SO4•−, about (4−9.1) × 105 mol−1 L s−1.1 On the basis of these properties, by using TBA and EtOH, different inhibitory effects were observed for the decolorization of Orange II (Table 1). The results in Table 1 confirmed the formation of SO4•− in Fe(II)/sulfite system, and it accounted for 74.8% of the decolorization of 10 mg L−1 Orange II. Surprisingly, HO• was also detected in the Fe(II)/sulfite system, and it accounted for as

SO4•− /HO• + X− → SO4 2 − /HO− + X −

−

−

(17)

−

where X is Cl , Br , or I .

Table 1. Results of Radical Identification from the Decolorization of Orange IIa system Fe(II)/sulfite Fe(II)/ persulfate Fenton a

initial 5 min decolorization of Orange II (%)

change in decolorization of Orange II due to TBA (%)

decolorization due to HO• (%)

change in decolorization of Orange II due to EtOH (%)

decolorization due to SO4•− and HO• (%)

decolorization due to SO5•− (%)

30.2 7.2

−4.2 −2.4

13.9 33.3

−26.8 −7.2

88.7 100

11.3 0

−30.2

100

−30.2

30.2 −1

100 −1

−1

0 −1

Conditions: [Orange II]0 = 10 mg L , [FeSO4]0 = 0.1 mmol L , [Na2SO3]0 = 1 mmol L , [TBA]0 = [EtOH]0 = 500 mmol L , pHini = 4. 13635



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and

Table 2. Second-Order Rate Constants for the Reactions between Halide Ions and SO4•−/HO•a kSO4•−,X−(mol

halide ion Cl− Br− I− a

−1

−1

Ls )

ref

(1.3−3.1) × 108 (1.6−3.5) × 109 NA

kHO•,X− (mol

43 43 NA

−1

−1

Ls )

4.3 × 109 1.1 × 1010 1.1 × 1010

b=

ref 57 57 57

f •− × + R SO 5

ksub,SO4•−[sub] ksub,SO4•−[sub] + k X −,SO4•−[X−] + ∑i ki[Si ] ksub,HO•[sub]

ksub,HO•[sub] + k X −,HO•[X−] + ∑j kj[Sj ] ksub,SO•− [sub] 5 ksub,SO•− [sub] + k X −,SO•− [X−] + ∑p kp[Sp] 5 5 (Formula 1)

where Rd is the initial decolorization rate of Orange II, RfSO4•− is the formation rate of SO4•−, kSO4•− is the second-order rate constant for SO4•− by Orange II, [sub] is the concentration of substrate (Orange II), and ∑iki[Si] is the pseudo-first-order rate constant for SO4•− by other scavengers. The meanings of other symbols are indicated above. Table 1 shows that SO4•− dominates the Fe(II)/sulfite system. In addition, SO•− 5 is a very weak radical so that it does not seem to oxidize halide ions.43 Hence, to estimate the value of RfSO4•−, Formula 1 was reasonably modulated as Formula 2, by omission • of SO•− 5 and HO . f •− × Rd ≈ R SO 4

ksub,SO4•−[sub] ksub,SO4•−[sub] + k X−,SO4•−[X−] + ∑i ki[Si ] (Formula 2)

By taking the reciprocal of Formula 2, Formula 3 was directly obtained. ⎛ ∑i ki[Si ] ⎞ 1 1 ⎜⎜1 + ⎟ ≈ × f ksub,SO4•−[sub] ⎟⎠ •− R SO Rd ⎝ 4 k X−,SO4•− + f × [X−] R SO4•−ksub,SO4•−[sub]

be further simplified as follows, 1 ≈ a + b[X−] d R where a=

1 f •− R SO 4

⎛ ∑i ki[Si ] ⎞ ⎟ × ⎜⎜1 + ksub,SO4•−[sub] ⎟⎠ ⎝

4. CONCLUSIONS Fe(II)/sulfite system has been shown to be an effective process to treat massive dyes. The targeted substrate Orange II was readily decolorized when O2 was present, and results showed that Fe(III) has equal effect with Fe(II) when it combines with Na2SO3, which is because the slow decomposition of FeSO3+ determines the redox cycle of iron. pH and dosage ratio effects were also researched, and pH 4 and dosage ratio 10:1 were the optimum experimental conditions for Orange II. Radicals identification batch experiments confirmed that the oxidation of Orange II was mainly due to SO4•− (74.8%), but HO• (13.9%) and SO5•− (11.3%) also played significant roles. Halide ions were also discovered to quench the Fe(II)/sulfite system, and this was because they could strongly scavenge SO4•−. Selective oxidation of SO4•− decreased TOC removal of Orange II, and this will limit its use in water remediation whose goal is to eliminate organic compounds from aqueous medium. However, since Na2SO3 (ca. $650/t) is cheaper than K2S2O8 (ca. $1300/t), able to achieve a higher decolorization efficiency, and is much more stable than K2S2O8, Na2SO3 is expected to have potential applications in wastewater treatment. In

(Formula 3)

f Since (1/RSO •− ) × (1 + (∑iki[Si])/(ksub,SO •−[sub])) 4 4 f •− (kX,SO4 /(RSO4•−ksub,SO4•−[sub])) are constants, Formula 3

(Formula 6)

−

Hence, 1/R and [X ] are theoretically linearly dependent, and the results of curve fitting are given in Table S3 (Supporting Information). They show that 1/Rd is linearly correlated to [X−] for all three halide ions (R2 ≥ 0.986). In addition, a is theoretically a constant (Formula 5), and this can be verified from Table 3, which shows that a is (0.877−1.03) × 104 mol−1 L s. Although (∑iki[Si])/(ksub,SO4•−[sub]) is unknown, it is certain that it is not very large (≤10) since most of the SO4•− radicals go to oxidize Orange II under optimum conditions. Hence, 10−4 mol L−1 s−1 ≤ RfSO4•− ≤ 10−3 mol L−1 s−1 was assumed. Formula 6 also indicates that b is proportional to kX−,SO4•−. Together with the constants b, kCl−,SO4•−, and kBr−,SO4•−, the value of kI−,SO4•− was determined to be 3.2 × 1011 mol−1 L s−1. This is significant because there is no available data for kI−,SO4•− in the literature. It shows that kI−,SO4•− ≫ kI−,HO•. This explains the phenomenon that, although kBr−,HO• is identical with kI−,HO•, Br− inhibited the decolorization of Orange II much less than I−. 3.7. TOC Removal Efficiency of Orange II in Fe(II)/ Sulfite System. The oxidation properties of SO4•− and HO• have been rationalized by Steenken.20 It shows that SO4•− and HO• undergo much different reaction mechanisms under the same condition, and this would affect the mineralization rate of the substrate. In this study, TOC removal efficiency was set as a general index to research the mineralization effect of SO4•−. It showed that 14.4% TOC removal (data not shown) of 10 mg L−1 Orange II with Fe(II)/sulfite system was achieved at pH 4 and dosage ratio 10:1 after reacting for 120 min. Anipsitakis et al. proved that the TOC removal rose with increased pH in the SO4•−-generating system from neutral to alkaline pH,54 which may result from nonselective oxidation of HO•. Therefore, the mere usage of SO4•− is better for decolorization than for mineralization. As evidenced in Figure S6 (Supporting Information), the decolorization of Orange II was due to the destruction of chromophore (e.g., azo bond), and this mechanism can be used to detoxify azo dyes in water.

In the Fe(II)/sulfite system, there are in total three major oxidizing species, SO4•−, HO•, and SO5•−. By adding halide ions, the decolorization was enervated due to competition for the oxidizing species between Orange II and the halide ion. The effect of halide ions on the decolorization of Orange II is theoretically expressed as Formula 1 based on the steady-state approximation,

f • × + RHO

f •−k R SO sub,SO4•−[sub] 4 d

NA = Not available.

f •− × Rd = R SO 4

k X−,SO4•−

and can

(Formula 4)

(Formula 5) 13636



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terms of acute toxicity, the LD50 (oral, rat) of Na2SO3 is 3560 mg/kg,55 which is more than 3 times higher than that of K2S2O8 (802 mg/kg).56 This implies that Fe(II)/sulfite system is more environmently friendly than systems containing K2S2O8.

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ASSOCIATED CONTENT

S Supporting Information *

Tables of substrates, decolorization results by three systems and linear fitting results, and figures of species distribution, pH variance, absorption spectra, and halide ion effects. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*Address correspondence to [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the Natural Science Foundation of China (No. 21077080), the Specialized Research Fund for the Doctoral Program of Higher Education of China (No. 20100141110046), and the Natural Science Foundation of Hubei Province (No. 2008CDB379). Comments from the anonymous reviewers are also appreciated.

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REFERENCES

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