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H2O2 system
Desalination 281 (2011) 306–311
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Desalination j o u r n a l h o m e p a g e : w w w. e l s ev i e r. c o m / l o c a t e / d e s a l
Decolorization of the azo dye Orange II in a montmorillonite/H2O2 system Liang Chen, Chunyan Deng, Feng Wu ⁎, Nansheng Deng Department of Environmental Science, Hubei Key Lab of Biomass Resource Chemistry and Environmental Biotechnology, School of Resources and Environmental Science, Wuhan University, 430079, China
a r t i c l e
i n f o
Article history: Received 7 September 2010 Received in revised form 25 July 2011 Accepted 8 August 2011 Available online 1 September 2011 Keywords: Orange II Decolorization Montmorillonite KSF Heterogeneous Fenton Response surface methodology (RSM)
a b s t r a c t The decolorization of azo dye Orange II using a montmorillonite/H2O2 system and a heterogeneous Fenton system, with montmorillonite KSF as a catalyst, has been studied. A series of experiments were performed to analyze the effects of several variables, including pH, KSF dosage, H2O2 dosage and the initial Orange II concentration. The results revealed that under proper conditions, relatively high decolorization efficiency (more than 90%) could be achieved in only 40 min when the initial concentration of Orange II was 20 mg/L. Response surface methodology (RSM) was applied with a Box–Behnken design (BBD) of a series of experiments, which showed that, of the range of variables studied, the order of influence is pH N KSF dosage N H2O2 dosage. A mathematical model was established and the optimal conditions were determined by RSM. Sequential experiments showed that KSF, re-used from a former reaction, performed well for 4 further runs. Comparisons between homogeneous and heterogeneous catalytic oxidation showed that unreleased free ferrous iron and structural iron in the KSF particles were responsible for the catalysis of the oxidation of Orange II in the montmorillonite/H2O2 system at pH 4. However, at pH 3, the predominant catalytic iron species was the released free ferrous iron. © 2011 Elsevier B.V. All rights reserved.
1. Introduction Azo dyes are ubiquitous commercial chemicals that are used widely in textiles, leather tanning, paper production, food industries and other areas [1,2]. However, azo dyes can penetrate organisms and react with reducible chemicals produced by metabolic processes. The \N_N\ group is reduced and broken down to form aromatic amines, some of which can penetrate through the cytomembrane and attack DNA, resulting in mutagenesis [3,4]. However, dyestuffs including azo dyes discharged into the environment are resistant to biological treatment because of their complex structure and synthetic origin [5]. All of these factors contribute to the scientific interest in them. Orange II (with structure presented in the inset of Fig. 1, formula: C16H11N2NaO4, molecular weight: 350.32 g/mol, CAS NO.: 633-96-5) is a typical azo dye and has been selected as a model dyeing pollutant in many investigations, including this study and because of its widespread application and recalcitrant nature [6,7]. The degradation of Orange II has been previously studied. Li et al. studied the decolorization of azo dye Orange II by a ferrate (VI)–hypochlorite liquid mixture, potassium ferrate (VI) and potassium permanganate, and reported that the ferrate (VI)–hypochlorite liquid mixture was more powerful than the other two reagents. They also found that the chromophore, N_N, was completely broken down, with the long conjugated π system decom-
⁎ Corresponding author. Tel./fax: +86 27 68778511. E-mail address: [email protected] (F. Wu). 0011-9164/$ – see front matter © 2011 Elsevier B.V. All rights reserved. doi:10.1016/j.desal.2011.08.006
posing at the same time. This caused decolorization in the visible region but the decay of the naphthalene and benzene rings was slow and weak [8]. Monteagudo et al. concluded that the ferrioxalate-assisted, solar photo-Fenton reaction improves the photocatalytic efficiency of Orange II because the ferrioxalate complexes absorb efficiently, and a higher portion of the solar spectrum can be used [9]. Sakakibara et al. studied the degradation of Orange II using a sonochemical reaction to determine the best degradation conditions [10]. Ray et al. researched the photocatalytic degradation of Orange II by TiO2 catalysts supported on adsorbents [11]. All of these methods gave high degradation efficiency. Ramirez et al. studied the degradation of Orange II catalyzed by saponite clay. They found that under optimal conditions, the degradation efficiency was as high as 99% in 4 h [12]. To lower the treatment cost of waste water, minimize the use of solvents and replace hazardous or expensive reagents and reactants by safer and more economical ones, catalysts, in particular clays and zeolites that are derived from soil minerals are promising choices [13–17]. Among the perspective clays, montmorillonite clay is widely utilized in various industrial processes and laboratory experiments because of its structural properties. Acid-modified montmorillonite KSF has some advantages over other catalysts: (1) it is a fine solid and possesses high specific surface area and mesoporosity, (2) montmorillonite KSF is an iron-rich catalyst with high cation exchange capacity that can result in the release of metals with strong Lewis acid properties and (3) in montmorillonite/H2O2 systems, montmorillonite KSF can be easily recovered after filtration, with sequential treatment such as purging with other chemicals or air drying [17–22]. These characteristics make it
L. Chen et al. / Desalination 281 (2011) 306–311 OH
bath to maintain the experimental temperature at 30 °C. The desired amounts of Orange II and H2O2 were added to the aqueous suspension to initiate the Fenton reaction, and then the time was noted. At regular time intervals (5, 10, 20, 30, 40 min), a 5-mL sample was withdrawn from the reactor and immediately filtered through a 0.22-μm millipore filter to remove KSF particles. The concentration of the remaining Orange II in the solution was analyzed by a spectrometric method at a wavelength of 485 nm using an UV1601 UV/VIS Spectrophotometer (Shimadzu, Japan).
2.0 NaO3S
Abs
1.5
N
N
Structure of Orange II 0min 5min 10min 20min 30min 40min
1.0
0.5
0.0 200
400
600
307
2.3. Experimental design and data analysis by RSM
800
Wavelength (nm) Fig. 1. UV–VIS spectra of Orange II as a function of reaction time with experimental conditions as follows: Orange II 20 mg/L, KSF 0.2 g/L, H2O2 4 mmol/L, pH 3. Inset: structure of Orange II.
feasible to use montmorillonite KSF as an economic and ecological catalyst for water treatment, responding to the trend of developing “green chemistry” [23]. Previous studies have showed that montmorillonite KSF is an efficient acidic catalyst for organic synthesis. Our research group has investigated the photodegradation of paracetamol, Bisphenol AF and Orange II in montmorillonite KSF suspended solutions and the results demonstrated that KSF is an alternative reagent to remove some pollutants in wastewater [15,24,25]. However, using montmorillonite KSF as the source of iron for the Fenton or Fenton-like system in natural sunlight still deserves thorough study because this method may have some advantages in terms of lower cost and greater simplicity. The novel features of our study are: (1) we use montmorillonite KSF as the source of iron for the heterogeneous Fenton-like system to remove the azo dye from water, (2) we apply response surface methodology (RSM) to optimize the parameters and analyze the degree of importance of each factor; (3) we demonstrate the stability of KSF as a catalyst for continuous treatment and (4) we investigate the roles of the iron species in the heterogeneous montmorillonite/H2O2 system. These experiments can be regarded as the fundamental study preceding further practical application. 2. Materials and methods 2.1. Chemicals The Orange II used as the substrate was purchased from Shanghai SSS reagent Co. Ltd. The KSF, purchased from Alfa Aesar (Alfa Aesar, Britain), was used without further purification. The size of the particles was 20–25 μm. The surface area of the KSF was 20–40 m 2/g (information provided by the supplier). The H2O2 was analytical reagent grade. Hydrochloric acid and sodium hydroxide were used to adjust the pH of the solutions. The orthophenanthroline, ammonium fluoride, ascorbic acid, ammonium acetate, green vitriol, and concentrated sulfuric acid were analytical grade. Purified water with 18 M Ω cm resistivity was used throughout this work, after treatment with an ultrapure system (Liyuan Electric Instrument Co., Beijing, PRC). 2.2. Oxidation reaction and analysis Ultrapure water (500 mL) was put into a glass cylinder. An appropriate amount of KSF powder was added to the water and dispersed by stirring for 5 min. The initial pH of the KSF suspension was adjusted to the desired value by adding HCl or NaOH. The reactor, with a 300-rpm stirrer located near the base, was placed in a water
Fifteen experimental runs of the Box–Behnken design (BBD), which is a three-level incomplete factorial design extensively used, were performed to optimize the initial reaction rate, the derivative of the initial portion of the kinetic curve, and to arrange in order the importance of the various influencing factors. Table 1 shows the Box– Behnken design matrix for the three variables pH, KSF dosage and H2O2 dosage. The experimental data were analyzed using Design Expert 7.1.6 software and were fitted with a second-order polynomial equation. The optimized parameters for maximizing the initial rate were also obtained by using the established equation. 2.4. Sequential decolorization experiment To test the stability of KSF as a catalyst and to explore the possibility of recycling KSF, sequential decolorization experiments were performed, consisting of four kinetic runs. Each of the first three runs lasted for 40 min, and the fourth run lasted for 100 min. To guarantee that the initial concentrations of Orange II and H2O2 were the same for all runs (Orange II 20 mg/L, H2O2 4 mmol/L), specific amounts of Orange II and 4 mmol/L of H2O2 were added to the suspension before initiating the second, third and fourth runs. The dose of Orange II added was determined by the measured concentration of the remaining Orange II after the previous run. The experimental and analytical methods for the four runs were the same as for the oxidation reaction and analysis mentioned above. 2.5. Detection of Fe (II) and Fe (III) in a suspension Ultrapure water (500 mL) was poured into a glass cylinder reactor, and a measured amount of KSF powder was added to it. The suspension was stirred for 5 min and the acidity adjusted to pH 3 by the addition of HCl. The reactor was placed in a water bath at 30 °C. The entire reaction was conducted while stirring at 300 rpm. Once the
Table 1 Box–Behnken design matrix for three variables together with the observed and predicted responses. Experimental run
Coded variables pH, X1
KSF dosage, X2 g/L
H2O2 dosage, X3 mmol/L
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15
5 3 4 3 5 4 3 4 5 4 5 3 4 4 4
0.2 0.3 0.3 0.2 0.2 0.3 0.2 0.1 0.3 0.1 0.1 0.1 0.2 0.2 0.2
3.8 2.0 0.2 3.8 0.2 3.8 0.2 0.2 2.0 3.8 2.0 2.0 2.0 2.0 2.0
L. Chen et al. / Desalination 281 (2011) 306–311
stirrer started working, the timing was started. The samples were withdrawn from the reactor at regular time intervals (5, 10, 20, 30 and 40 min) and immediately filtered through a 0.22-μm millipore filter to remove the KSF particles. Samples of the filtrate (10 mL) were taken to measure the concentration of Fe (II) and were added to a colorimeter tube containing 2 mL 1:1 (v/v) HCl. Then, 2.0 mL of 2 mol/L NH4F, 2.0 mL of 0.2% (m/V) orthophenanthroline and 5.0 mL of a NH4Ac-HAc buffer solution (pH = 4.2) were successively added to the colorimeter tube. The solution was diluted by adding water until the total volume was 25 mL. After 15 min of coloration, spectrophotometric measurements using orthophenanthroline at 510 nm were used to detect the concentration of Fe (II). Another 10 mL of filtrate was collected to measure the concentration of total iron. The procedure to detect the total iron concentration was similar to that for Fe (II), except for the following steps: (1) 1.0 mL of 0.1 mol/L ascorbic acid was added after the addition of 10 mL of filtrate to reduce Fe (III) and the solution left for 30 min and (2) NH4F, a chemical used to complex with Fe (II), was not needed to measure the total iron. After waiting 30 min, 2.0 mL of orthophenanthroline and 5.0 mL of the NH4Ac-HAc buffer solution were successively added, after which the procedure continued as above. 2.6. Decolorization of Orange II in the homogeneous Fenton system Appropriate amounts of Fe (II) and Fe (III) standard solutions were added to the reactor containing 500 mL of ultrapure water, and the mixture was stirred for 5 min. For the reaction at pH 3, the concentrations of added Fe (II) and Fe (III) were 0.4 mg/L and 0.9 mg/L, respectively. For the reaction at pH 4, the concentrations of added Fe (II) and Fe (III) were 0 mg/L and 1.0 mg/L, respectively. The amounts of Fe (II) and Fe (III) that were added were similar to the quantities released from the KSF at the corresponding pH values. After this, the procedures, including regulating the pH value, stirring the mixture in a water bath, taking samples and analyzing, were the same as the oxidation reaction and analysis method described above. 3. Results and discussion 3.1. Control experiment Setting the initial concentration of Orange II at 20 mg/L and the pH at 3, the efficiency for the decolorization of Orange II by adding 0.2 g/L of KSF and/or 4 mmol/L of H2O2 was investigated. Fig. 1 shows the results of the decolorization of Orange II in the presence of KSF and H2O2. Previous research has demonstrated that Orange II has three absorption peaks, at 230, 310 and 485 nm. The wavelength of its visible absorption peak is 485 nm. The absorption peaks in the UV region at 230 nm and 310 nm are assigned to the benzene ring and naphthalene ring [26]. The trend in variation shows that the three absorption peaks in the visible region and UV region gradually disappear. This phenomenon suggests the dramatic destruction of chromophores and the decay of the benzene ring and naphthalene ring. Fig. 2 reveals that the quantity of Orange II adsorbed on the KSF exceeded 10%. The decolorization efficiency of Orange II with KSF alone after 40 min was 15.3%, and with H2O2 alone after 30 min was 6.3%. In the presence of KSF and H2O2, the decolorization efficiency reached as much as 91.6% after reacting for 40 min. The experimental data show that a synergistically cooperative effect was observed, with KSF and H2O2 acting together. This effect was due to the Fenton reaction. Catalyzed by the Fe (II) and Fe (III), the H2O2 decomposes and hydroxyl radicals, with, high reactivity, are produced to degrade the Orange II [27]. KSF is produced from acid-modified montmorillonite [28], whose iron content is 4.8%. When the KSF dosage was 0.2 g/L, the iron concentration in the solution was about 120 μmol/L. The iron was
100
Degradation Efficiency (%)
308
KSF+H2O2 KSF H2O2
80 60 40 20 0 0
10
20
30
40
Time (min) Fig. 2. Decolorization efficiency of Orange II in control experiments with the following experimental conditions: Orange II 20 mg/L, KSF 0.2 g/L, H2O2 4 mmol/L, pH 3.
completely released from the KSF. About 8% of the iron was released when the pH value ranged from 3 to 4 [28]. In this experiment, the total iron in the suspension was about 1.3 mg/L, or 23 μmol/L, when the pH value was 3 and the KSF dosage was 0.2 g/L. Considering the production batch and the actual experimental conditions, the difference between the actual concentration and the theoretical value was not significant. The amount of iron in the suspension was far less than typical sewage discharge limits, since the total iron discharge standard requirement of the USEPA is 100 mg/L [29].
3.2. Effect of pH on decolorization of Orange II The decolorization of Orange II in aqueous suspensions of montmorillonite KSF was studied in the pH range between three and nine. Table 2 shows the decolorization efficiency of Orange II after 40 min at different pH values using initial suspensions containing 20 mg/L Orange II, 0.2 g/L KSF and 4 mmol/L H2O2. At pH 3, almost complete decolorization (91.6%) was achieved in 40 min and this was also the case at pH 4 (91.2%). However, no significant decolorization was observed at pH values higher than 5. This phenomenon was caused by the acidic pH that favored the release of Fe (II) from KSF so that the Fe (II) made more of a contribution to the production of •OH [28]. Guo et al. reported that the stability of H2O2 is irrelevant as to whether the process is a homogeneous one or heterogeneous one. However, its stability is dependent on the pH value. The degree of H2O2 decomposition at pH 3 and 4 was relatively low. The OH decrease was caused by the unsuccessful decomposition of H2O2 that produces molecular oxygen instead of •OH [30]. At a pH
Table 2 Effect of individual variables on the decolorization efficiency after 40 min. Variables pH
Montmorillonite KSF dosage g/L
H2O2 dosage mmol/L
3 4 5 6 7 8 3 3 3 3 3 3
0.2 0.2 0.2 0.2 0.2 0.2 0.10 0.30 0.2 0.2 0.2 0.2
4 4 4 4 4 4 4 4 0.2 1 2 8
Decolorization efficiency (%)
91.6 91.2 24.2 5.7 6.8 4.3 52.5 94.7 30.7 63.2 80.4 99.8
L. Chen et al. / Desalination 281 (2011) 306–311
3.3. Effect of montmorillonite KSF dosage on the decolorization of Orange II An increase in the decolorization efficiency of Orange II was observed with an increase in the KSF dosage, with experimental conditions Orange II 20 mg/L, H2O2 4 mmol/L, pH 3. As shown in Table 2, the results of series of experiments indicated that after 40 min, the decolorization efficiency of Orange II at a KSF dosage of 0.20 g/L was significantly higher than at 0.10 g/L, but it was lower than at 0.30 g/L. This phenomenon illustrates that an increase of KSF dosage favored the enhancement of Orange II decolorization. This is because the amount of Fe (II) released from the KSF was small when the dosage of KSF in the solution was low and consequently the decolorization rate was slow. At a dosage of 0.20 g/L, 25.5%, 31.3%, 58.5%, 77.0%, and 91.6% of the Orange II was removed in 5, 10, 20, 30, and 40 min, respectively. Even when the catalyst concentration was increased to 0.30 g/L, increased decolorization of Orange II did not follow as, in this case, 27.3%, 36.5%, 60.3%, 86.5%, and 94.7% of the Orange II was removed in 5, 10, 20, 30 and 40 min, respectively. Considering the small difference of the effect on the decolorization between 0.20 g/L and 0.30 g/L, the value 0.20 g/L was selected as the dosage of KSF when other individual effects were investigated. 3.4. Effect of H2O2 dosage on the decolorization of Orange II The initial concentration of H2O2 plays an important role in the decolorization of dyestuffs in the Fenton process. A preliminary examination of the influence of the H2O2 concentration in the range 0.2 to 8 mmol/L on the decolorization of Orange II after 40 min is also presented in Table 2. The results indicate that the decolorization of Orange II was remarkably dependent on the H2O2 concentration. After 40 min, an increase in the H2O2 concentration from 0.2 to 8 mmol/L increased the decolorization efficiency of the dye from 30.7% to 99.8%, which showed that H2O2 could promote the decolorization of Orange II. Given the fact that 91.6% of Orange II was removed in the presence of 4 mmol/L H2O2 and that a further increase in the dosage of H2O2 led to no obvious enhancement of the decolorization of dye, a value 4 mmol/L of H2O2 was used in the series of experiments when studying the effect of pH and KSF dosage. Generally, the degradation efficiency of organic compounds in the Fenton system increases with an increased dosage of H2O2 until a maximal concentration is reached [31]. The cause of this is the promotion of •OH production in the low concentration range and the inhibition at higher concentrations when excessive H2O2 acts as a free-radical scavenger itself and thereby decreases the production rate of •OH [32]. 3.5. Experimental design, data analysis and optimization by RSM The data of 15 experimental runs using the Box–Behnken design (BBD) were analyzed using Design Expert 7.1.6 software and were
Table 3 ANOVA results of the quadratic model of initial reaction rate. Source
Sum of squares
Degree of freedom
Mean square
F value
p-value (prob N F)
Model Residual Lack of Fit Pure error Corrected total
0.22 3.817 · 10−3 2.212 · 10−3 1.605 · 10−3 0.23
9 5 3 2 14
0.025 7.634 · 10−3 7.373 · 10−3 8.027 · 10−3
32.35
0.0007
0.92
0.5589
R2 = 0.9831; adjusted R2 = 0.9527; F value: Fisher variation ratio; p-value: significance probability value.
fitted with a second-order polynomial equation. The function presented below was employed as the prediction model: Y = 0:54742−0:31761X1 + 3:02845X2 + 0:13342X3 −0:83005X1 X2 2
2
−0:014978X1 X3 −0:12113X2 X3 + 0:047290X1 + 2:64379X2 −3
−7:98521·10
2
X3
where Y is the initial reaction rate (mg/L·min), X1 is the pH (3 to 5), X2 is the montmorillonite KSF dosage (0.1 g/L to 0.3 g/L) and X3 is the H2O2 dosage (0.2 mmol/L to 3.8 mmol/L). As shown in Table 3, the analysis of variance (ANOVA) technique was used to test the significance and adequacy of the model, giving a correlation coefficient R 2 = 0.9831, which demonstrates a satisfactory quadratic fit. In addition, the sum of squares, degrees of freedom, mean square, F value and p-value, which can be used to check the fit of the model, are also presented in Table 3. This is also reflected in Fig. 3, where the observed value (calculated using the observed data when conducting the experiment under the conditions defined by Experiment design) is plotted against the predicted value (calculated by using the formula with the same conditions). The slope of the line from the linear regression fit is very close to 1. The Model F value of 32.35 implies that the model is significant. There is only a 0.07% chance that a Model F value this large could occur due to noise. In addition, the p-value (significance probability value) less than 0.0500 indicates that the model terms are significant. We conclude from the values of the coefficient estimate and p-value that, of the 3 variables, the pH (coefficient estimate − 0.14, p-value b 0.0001) produces the largest influence on the initial rate. Next, the most influential is the montmorillonite KSF dosage (coefficient estimate 0.052, p-value 0.0030), while the H2O2 dosage (coefficient estimate is 0.031, p-value = 0.0241) shows little effect on the response. The optimized parameters for maximizing the initial rate were obtained by using the established model. The optimal conditions
0.5
Predicted vs. Observed
0.4
Predicted
N4, the precipitation of Fe (III) from the solution can also result in a decrease in reaction activity. While the influence of pH on the behavior of H2 O2 was investigated, its impact on the adsorption behavior of Orange II was also investigated because that is of greater significance to the removal of dyestuffs. Its influence was studied at pH 3 and 4 where the initial concentration of Orange II was 20 mg/L and of KSF 0.2 g/L. The results demonstrated that during the 40 min reaction time, the amount adsorbed at pH 3 was higher than at pH 4. These results can explain the difference between the decolorization efficiencies at pH 3 and pH 4. It can be concluded that, to some extent, the high adsorption at pH 3 contributed to the better decolorization effect than at pH 4. In the decolorization reaction of Orange II, the decolorization efficiency at pH 3 was a little higher than at pH 4. Consequently, pH 3 was selected as the optimal pH value for further experimentation.
309
0.3 0.2 0.1 0.0 0.0
0.1
0.2
0.3
0.4
0.5
observed Fig. 3. The observed and predicted values of the initial reaction rate (R2 = 0.9831, adjusted R2 = 0.9527).
310
L. Chen et al. / Desalination 281 (2011) 306–311
Decolorization efficiency (%)
determined from this study are as follows: pH 3; montmorillonite KSF dosage 0.30 g/L; H2O2 dosage 3.72 mmol/L. The predicted initial rate was 0.5048 mg/L·min, and the observed value was 0.5638 mg/L·min. Although the potential application of this heterogeneous Fenton system on real wastewater remains a target for the future, the present work as a basic method research was performed on synthetic solutions. With real wastewater, the negative effects of the matrix, such as adsorption and radical scavenging, could be relatively significant. Accordingly, the optimal conditions and the performance are subject to change. 3.6. Effect of initial Orange II concentration The effect of the initial Orange II concentration, which ranged from 5 mg/L to 40 mg/L, on its decolorization was also investigated. The results indicated that the decolorization efficiency at 5 mg/L was higher than at 40 mg/L of Orange II. The decolorization after 20 min was 100% at 5 mg/L and 40.5% at 40 mg/L. Even after 40 min, the final decolorization efficiency at 40 mg/L only reached 77.4%. In order to investigate the kinetics, the initial reaction rates for different concentrations were calculated. The derivative of the initial portion of the kinetic curve is the initial reaction rate r0. We used three models to describe our experimental data, firstorder, second-order and the Langmuir–Hinshelwood (L–H) equation. The results are presented in Table 4. By comparing the R 2 value for each equation, we concluded that the L–H equation fits best; the R 2 values for first-order, second-order and the L–H equations are 0.9014, 0.83201 and 0.9802, respectively. This result may be explained by the fact that the first/second-order equation suits the homogeneous system better. In the montmorillonite/H2O2 system, besides the adsorption behavior, some degradation behavior also exists. Based on this result, we chose the Langmuir–Hinshelwood (L–H) equation to describe Orange II decolorization for concentrations ranging from 5 mg/L to 40 mg/L r=−
dC kKC = ; dt 1 + KC
where C is the concentration of Orange II at time t, k is the rate constant and K is the adsorption coefficient. After calculation, k and K in this experiment were found to be 1.42 mg/L·min and 0.07014 L/mg, respectively. 3.7. Sequential decolorization experiments Sequential decolorization experiments were performed to test the stability of KSF as a catalyst. As shown in Fig. 4, after the first run, the decolorization efficiency of Orange II was 91.6% but, in subsequent runs, the decolorization efficiency decreased. At the end of the second run, the efficiency was 85.4% and it decreased to 77.5% after the third run. Nevertheless the decolorization efficiency reached as much as 94.5% after the fourth run, and the fact that the reaction time was prolonged to 100 min deserves attention. It was presumed that KSF would gradually lose its catalytic function, but it appears that a consistently high decolorization effect can be achieved by prolonging
100 80 60
1st run 2nd run 3rd run 4th run
40 20 0 0
50
100
150
200
250
Time (min) Fig. 4. Effect of recycling the catalyst on the decolorization of Orange II with the following experimental conditions: Orange II 20 mg/L, H2O2 4 mmol/L, KSF 0.2 g/L, pH 3.
the reaction time. For practical applications, further investigation on how to improve the stability of KSF could promote its wider application as a catalyst for the montmorillonite/H2O2 system and other wastewater treatment processes. 3.8. Role of iron species in the heterogeneous montmorillonite/H2O2 system To evaluate the role of both free iron (soluble iron) and structural iron (insoluble iron in crystalline lattice) in the KSF particles in the reaction, the amount of Fe (II) and the total iron released during the entire decolorization procedure was measured. A comparison of Orange II decolorization in a homogeneous system and in a montmorillonite/H2O2 system was made. The results revealed that at pH 3, about 1.3 mg/L of total iron and about 0.4 mg/L of Fe (II) were released after 5 min. However, at pH 4, there was hardly any Fe (II) released in suspension, and the total iron was about 1.0 mg/L. During the entire reaction process, slight fluctuations in Fe (II) and total iron concentrations over time occurred but, in general, these fluctuations were not significant. This variation suggested that nearly all the Fe (II) and total iron that existed in suspension was released at the beginning of the experiment. The dosage of iron in the homogeneous control experiment was the same as the iron concentration dissolving at pH 3 and 4, respectively. Fig. 5 shows that at pH 3, after 40 min, the decolorization efficiency of Orange II in a homogeneous system was 77.0% and in the montmorillonite/H2O2 system was 91.6%. However, at pH 4, after 40 min, the decolorization efficiency in the homogeneous system was only 6.7%, and in the montmorillonite/H2O2 system it was 91.2%. This phenomenon should be noted, because even though there was no Fe (II) released from the Montmorillonite KSF at pH 4, KSF could still play the role of a catalyst. This result indicates that the unreleased Fe (II) ions and structural Fe (II) in the crystalline lattice can catalyze the degradation of Orange II through the heterogeneous Fenton reaction in the absence of free Fe (II) ions. Since at pH 3, with a catalytic dose of free Fe (II) ions, the catalytic effect of structural iron was not significant, it is believed that unreleased Fe (II) ions made a greater contribution to the catalysis of
Table 4 Kinetics analysis of Orange II decolorization. C0 (mg/L)
Initial rate (mg/L·min)
Model
Kinetics equation
5 10 20 30 40
0.3584 0.6757 0.7711 0.8606 1.1255
L–H model
r=
First-order
r = 0.03134C (R2 = 0.9014)
3.13 · 10−2
Second-order
r = 0.0008361C2 (R2 = 0.83201)
8.361 · 10−4
0:0996C (R2 = 0.9608) 1 + 0:07014C
k (mg/L·min)
K (L/mg)
1.42
7.01 · 10−2
L. Chen et al. / Desalination 281 (2011) 306–311
pH 3, Montmorillonite/H2O2 system pH 3, Fe (II)/Fe (III)/H2O2 system pH 4, Montmorillonite/H2O2 system pH 4, Fe (II)/Fe (III)/H2O2 system
Decolorization efficiency (%)
100 80 60 40 20 0 0
10
20
30
40
Time (min) Fig. 5. Control experiment for the Fe (II)/Fe (III)/H2O2 system with the experimental conditions: Fe (II) 0.4 mg/L, Fe (III) 0.9 mg/L, Orange II 20 mg/L, H2O2 4 mmol/L, pH 3; Fe (II) 0 mg/L, Fe (III) 1.0 mg/L, Orange II 20 mg/L, H2O2 4 mmol/L, pH 4.
the oxidation in the montmorillonite/H2O2 system. This supports our original hypothesis that KSF has the potential to serve as a heterogeneous catalyst. 4. Conclusion In the presence of H2O2 or KSF alone, the decolorization of 20 mg/L Orange II was not clearly observed. When both H2O2 and KSF were in the solution, the decolorization efficiency was as high as 91.6% after 40 min. A second-order polynomial model based on response surface methodology using the Box–Behnken design clarified the order of importance of each individual factor on the initial reaction rate as pH N montmorillonite KSF dosage N H2O2 dosage. The model also predicted that the optimal conditions were pH 3, montmorillonite KSF 0.3 g/L and H2O2 3.72 mmol/L. At pH 3 and initial concentrations ranging from 5 to 40 mg/L, the decolorization of Orange II agreed with the L–H equation, and the efficiency decreased from almost 100% to 40.5%. In the heterogeneous reaction catalyzed by KSF at pH 3, the homogeneous reaction played the dominant role in the decolorization of Orange II, amounting to about 80% of the total decolorization. However, at pH 4, the unreleased Fe (II) ions and structural Fe (II) in the crystalline lattice can play a more important role in catalyzing the degradation of Orange II. This study indicated that KSF can be used as a new heterogeneous catalyst to induce the decolorization of some refractory organic compounds such as dyes. However, about 1.3 mg/L of total iron was released from the KSF at pH 3. Therefore, lowering the release of iron from KSF and improving the stability of KSF, which acts as a catalyst, are worth further study. Acknowledgments This work was sponsored by the Natural Science Foundation of China (NSFC No. 21077080), the Natural Science Foundation of Hubei Province (No. 2008CDB379) and the Wuhan Municipal Chenguang Project (No. 200750731254). The authors thank the anonymous reviewers for their comments on this manuscript. References [1] E. Forgacs, T. Cserhati, G. Oros, Removal of synthetic dyes from wastewaters: a review, Environ. Int. 30 (2004) 953–971.
311
[2] W. Tang, R. Jia, D.Q. Zhang, Decolorization and degradation of synthetic dyes by Schizophyllum sp. F17 in a novel system, Desalination 265 (2010) 22–27. [3] C.M. King, M.S. Lee, R.F. Jones, N. Tamura, Modification of plasmid and bacteriophage DNA by aromatic amines: effects on survival, template activity, and mutagenicity, Environ. Health Perspect. 102 (1994) 217–220. [4] P.L. Skipper, S.R. Tannenbaum, Molecular dosimetry of aromatic amines in human populations, Environ. Health Perspect. 102 (1994) 17–21. [5] S. Liakou, M. Kornaros, G. Lyberatos, Pretreatment of azo dyes using ozone, Water Sci. Technol. 36 (1997) 155–163. [6] Y. Xiong, P.J. Strunk, H. Xia, X. Zhu, H.T. Karlsson, Treatment of dye wastewater containing acid Orange II using a cell with three-phase three-dimensional electrode, Water Res. 35 (2001) 4226–4230. [7] K. Vinodgopal, J. Peller, O. Makogon, P.V. Kamat, Ultrasonic mineralization of a reactive textile azo dye, remazol black B, Water Res. 32 (1998) 3646–3650. [8] G.T. Li, N.G. Wang, B.T. Liu, X.W. Zhang, Decolorization of azo dye Orange II by ferrate(VI)–hypochlorite liquid mixture, potassium ferrate(VI) and potassium permanganate, Desalination 249 (2009) 936–941. [9] J.M. Monteagudo, A. Durán, C. López-Almodóvar, Homogeneus ferrioxalate-assisted solar photo-Fenton degradation of Orange II aqueous solutions, App. Catal. B Environ. 83 (2008) 46–55. [10] M. Inoue, F. Okada, A. Sakurai, M. Sakakibara, A new development of dyestuffs degradation system using ultrasound, Ultrason. Sonochem. 13 (2006) 313–320. [11] A. Bhattacharyya, S. Kawi, M.B. Ray, Photocatalytic degradation of Orange II by TiO2 catalysts supported on adsorbents, Catal. Today 98 (2004) 431–439. [12] J.H. Ramirez, C.A. Costa, L.M. Madeira, G. Mata, M.A. Vicente, M.L. Rojas-Cervantes, A.J. López-Peinado, R.M. Martín-Aranda, Fenton-like oxidation of Orange II solutions using heterogeneous catalysts based on saponite clay, Appl. Catal. B 71 (2007) 44–56. [13] G. Nagendrappa, Organic synthesis using clay catalysts: clays for ‘green chemistry’, Resonance 7 (2002) 64–77. [14] M. Nikpassand, M. Mamaghani, K. Tabatabaeian, M.K. Abiazi, KSF: an efficient catalyst for the regioselective synthesis of 1,5-diaryl pyrazoles using Baylis– Hillman adducts, Mol. Diversity 13 (2009) 389–393. [15] Y.X. Liu, K. Wan, N.S. Deng, F. Wu, Photodegradation of paracetamol in montmorillonite KSF suspension, React. Kinet. Mech. Catal. 99 (2010) 493–502. [16] M.M. Urrutia, E.E. Roden, J.M. Zachara, Influence of aqueous and solid-phase Fe (II) complexants on microbial reduction of crystalline iron(III) oxides, Environ. Sci. Technol. 33 (1999) 4022–4028. [17] Y.X. Liu, X. Zhang, L. Guo, F. Wu, N.S. Deng, Photodegradation of bisphenol A in the montmorillonite KSF suspended solutions, Ind. Eng. Chem. Res. 47 (2008) 7141–7146. [18] F. Bigi, L. Chesini, R. Maggi, G. Sartori, Montmorillonite KSF as an inorganic, water stable, and reusable catalyst for the Knoevenagel synthesis of coumarin-3-carboxylic acids, Org. Chem. 64 (1999) 1033–1035. [19] F. Bigi, S. Carloni, B. Frullanti, R. Maggi, G. Sartori, A revision of the Biginelli reaction under solid acid catalysis. Solvent-free synthesis of dihydropyrimidines over montmorillonite KSF, Tetrahedron Lett. 40 (1999) 3465–3468. [20] P. Kumar, R.V. Jasra, T.S.G. Bhat, Evolution of porosity and surface acidity in montmorillonite clay on acid activation, Ind. Eng. Chem. Res. 34 (1995) 1400–1448. [21] M.F. Brigatti, C. Lugli, G. Cibin, A. Marcelli, G. Giuli, E. Paris, A. Mottana, Z. Wu, Reduction and sorption of chromium by Fe (II)-bearing phyllosilicates: chemical treatment and X-ray absorption spectroscopy (XAS) studies, Clays Clay Miner. 48 (2000) 272–281. [22] S. Mendioroz, J.A. Pajares, I. Benito, C. Pesquera, F. Gonzalez, C. Blanco, Texture evolution of montmorillonite under progressive acid treatment: change from H3 to H2 type of hysteresis, Langmuir 3 (1987) 676–681. [23] J.H. Clark, D.J. Macquarrie, Catalysis of liquid phase organic reactions using chemically modified mesoporous inorganic solids, Chem. Commun. (1998) 853–860. [24] Y.X. Liu, X. Zhang, F. Wu, Photodegradation of bisphenol AF in montmorillonite dispersions: Kinetics and mechanism study, Appl. Clay Sci. 49 (2010) 182–186. [25] J. Li, F. Wu, N.S. Deng, E.M. Glebov, N.M. Bazhin, Degradation of Orange II by heterogeneous photocatalytic reaction using Montmorillonite KSF, React. Kinet. Catal. Lett. 95 (2008) 247–255. [26] F. Wu, N.S. Deng, H.L. Hua, Degradation mechanism of azo dye C. I. reactive red 2 by iron powder reduction and photooxidation in aqueous solutions, Chemosphere 41 (2000) 1233–1238. [27] R.V. Lloyd, P.M. Hanna, R.P. Mason, The origin of the hydroxyl radical oxygen in the fenton reaction, Free Radic. Biol. Med. 22 (1997) 885–888. [28] X. Zhang, F. Wu, N.S. Deng, I.P. Pozdnyakov, E.M. Glebov, V.P. Grivin, V.F. Plyusnin, N.N. Bazhin, Evidence of the hydroxyl radical formation upon the photolysis of an iron-rich clay in aqueous solutions, React. Kinet. Catal. Lett. 94 (2008) 207–218. [29] U.S. EPA regulations, 40 CFR: Protection of the Environment. Subchapter N — “Effluent Guidelines and Standards”. Parts 400–424 and 425–471, http://www.epa. gov/regulations/search/40cfr.html 2010. [30] J. Guo, M. Al-Dahhan, Catalytic wet oxidation of phenol by hydrogen peroxide over pillared clay catalyst, Ind. Eng. Chem. Res. 42 (2003) 2450–2460. [31] J. Feng, X. Hu, P.L. Yue, H.Y. Zhu, G.Q. Lu, Degradation of azo-dye Orange II by a photoassisted fenton reaction using a novel composite of iron oxide and silicate nanoparticles as a catalyst, Chem. Res. 42 (2003) 2058–2066. [32] R. Alnaizy, A. Akgerman, Advanced oxidation of phenolic compounds, Adv. Environ. Res. 4 (2000) 233–244.
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Desalination j o u r n a l h o m e p a g e : w w w. e l s ev i e r. c o m / l o c a t e / d e s a l
Decolorization of the azo dye Orange II in a montmorillonite/H2O2 system Liang Chen, Chunyan Deng, Feng Wu ⁎, Nansheng Deng Department of Environmental Science, Hubei Key Lab of Biomass Resource Chemistry and Environmental Biotechnology, School of Resources and Environmental Science, Wuhan University, 430079, China
a r t i c l e
i n f o
Article history: Received 7 September 2010 Received in revised form 25 July 2011 Accepted 8 August 2011 Available online 1 September 2011 Keywords: Orange II Decolorization Montmorillonite KSF Heterogeneous Fenton Response surface methodology (RSM)
a b s t r a c t The decolorization of azo dye Orange II using a montmorillonite/H2O2 system and a heterogeneous Fenton system, with montmorillonite KSF as a catalyst, has been studied. A series of experiments were performed to analyze the effects of several variables, including pH, KSF dosage, H2O2 dosage and the initial Orange II concentration. The results revealed that under proper conditions, relatively high decolorization efficiency (more than 90%) could be achieved in only 40 min when the initial concentration of Orange II was 20 mg/L. Response surface methodology (RSM) was applied with a Box–Behnken design (BBD) of a series of experiments, which showed that, of the range of variables studied, the order of influence is pH N KSF dosage N H2O2 dosage. A mathematical model was established and the optimal conditions were determined by RSM. Sequential experiments showed that KSF, re-used from a former reaction, performed well for 4 further runs. Comparisons between homogeneous and heterogeneous catalytic oxidation showed that unreleased free ferrous iron and structural iron in the KSF particles were responsible for the catalysis of the oxidation of Orange II in the montmorillonite/H2O2 system at pH 4. However, at pH 3, the predominant catalytic iron species was the released free ferrous iron. © 2011 Elsevier B.V. All rights reserved.
1. Introduction Azo dyes are ubiquitous commercial chemicals that are used widely in textiles, leather tanning, paper production, food industries and other areas [1,2]. However, azo dyes can penetrate organisms and react with reducible chemicals produced by metabolic processes. The \N_N\ group is reduced and broken down to form aromatic amines, some of which can penetrate through the cytomembrane and attack DNA, resulting in mutagenesis [3,4]. However, dyestuffs including azo dyes discharged into the environment are resistant to biological treatment because of their complex structure and synthetic origin [5]. All of these factors contribute to the scientific interest in them. Orange II (with structure presented in the inset of Fig. 1, formula: C16H11N2NaO4, molecular weight: 350.32 g/mol, CAS NO.: 633-96-5) is a typical azo dye and has been selected as a model dyeing pollutant in many investigations, including this study and because of its widespread application and recalcitrant nature [6,7]. The degradation of Orange II has been previously studied. Li et al. studied the decolorization of azo dye Orange II by a ferrate (VI)–hypochlorite liquid mixture, potassium ferrate (VI) and potassium permanganate, and reported that the ferrate (VI)–hypochlorite liquid mixture was more powerful than the other two reagents. They also found that the chromophore, N_N, was completely broken down, with the long conjugated π system decom-
⁎ Corresponding author. Tel./fax: +86 27 68778511. E-mail address: [email protected] (F. Wu). 0011-9164/$ – see front matter © 2011 Elsevier B.V. All rights reserved. doi:10.1016/j.desal.2011.08.006
posing at the same time. This caused decolorization in the visible region but the decay of the naphthalene and benzene rings was slow and weak [8]. Monteagudo et al. concluded that the ferrioxalate-assisted, solar photo-Fenton reaction improves the photocatalytic efficiency of Orange II because the ferrioxalate complexes absorb efficiently, and a higher portion of the solar spectrum can be used [9]. Sakakibara et al. studied the degradation of Orange II using a sonochemical reaction to determine the best degradation conditions [10]. Ray et al. researched the photocatalytic degradation of Orange II by TiO2 catalysts supported on adsorbents [11]. All of these methods gave high degradation efficiency. Ramirez et al. studied the degradation of Orange II catalyzed by saponite clay. They found that under optimal conditions, the degradation efficiency was as high as 99% in 4 h [12]. To lower the treatment cost of waste water, minimize the use of solvents and replace hazardous or expensive reagents and reactants by safer and more economical ones, catalysts, in particular clays and zeolites that are derived from soil minerals are promising choices [13–17]. Among the perspective clays, montmorillonite clay is widely utilized in various industrial processes and laboratory experiments because of its structural properties. Acid-modified montmorillonite KSF has some advantages over other catalysts: (1) it is a fine solid and possesses high specific surface area and mesoporosity, (2) montmorillonite KSF is an iron-rich catalyst with high cation exchange capacity that can result in the release of metals with strong Lewis acid properties and (3) in montmorillonite/H2O2 systems, montmorillonite KSF can be easily recovered after filtration, with sequential treatment such as purging with other chemicals or air drying [17–22]. These characteristics make it
L. Chen et al. / Desalination 281 (2011) 306–311 OH
bath to maintain the experimental temperature at 30 °C. The desired amounts of Orange II and H2O2 were added to the aqueous suspension to initiate the Fenton reaction, and then the time was noted. At regular time intervals (5, 10, 20, 30, 40 min), a 5-mL sample was withdrawn from the reactor and immediately filtered through a 0.22-μm millipore filter to remove KSF particles. The concentration of the remaining Orange II in the solution was analyzed by a spectrometric method at a wavelength of 485 nm using an UV1601 UV/VIS Spectrophotometer (Shimadzu, Japan).
2.0 NaO3S
Abs
1.5
N
N
Structure of Orange II 0min 5min 10min 20min 30min 40min
1.0
0.5
0.0 200
400
600
307
2.3. Experimental design and data analysis by RSM
800
Wavelength (nm) Fig. 1. UV–VIS spectra of Orange II as a function of reaction time with experimental conditions as follows: Orange II 20 mg/L, KSF 0.2 g/L, H2O2 4 mmol/L, pH 3. Inset: structure of Orange II.
feasible to use montmorillonite KSF as an economic and ecological catalyst for water treatment, responding to the trend of developing “green chemistry” [23]. Previous studies have showed that montmorillonite KSF is an efficient acidic catalyst for organic synthesis. Our research group has investigated the photodegradation of paracetamol, Bisphenol AF and Orange II in montmorillonite KSF suspended solutions and the results demonstrated that KSF is an alternative reagent to remove some pollutants in wastewater [15,24,25]. However, using montmorillonite KSF as the source of iron for the Fenton or Fenton-like system in natural sunlight still deserves thorough study because this method may have some advantages in terms of lower cost and greater simplicity. The novel features of our study are: (1) we use montmorillonite KSF as the source of iron for the heterogeneous Fenton-like system to remove the azo dye from water, (2) we apply response surface methodology (RSM) to optimize the parameters and analyze the degree of importance of each factor; (3) we demonstrate the stability of KSF as a catalyst for continuous treatment and (4) we investigate the roles of the iron species in the heterogeneous montmorillonite/H2O2 system. These experiments can be regarded as the fundamental study preceding further practical application. 2. Materials and methods 2.1. Chemicals The Orange II used as the substrate was purchased from Shanghai SSS reagent Co. Ltd. The KSF, purchased from Alfa Aesar (Alfa Aesar, Britain), was used without further purification. The size of the particles was 20–25 μm. The surface area of the KSF was 20–40 m 2/g (information provided by the supplier). The H2O2 was analytical reagent grade. Hydrochloric acid and sodium hydroxide were used to adjust the pH of the solutions. The orthophenanthroline, ammonium fluoride, ascorbic acid, ammonium acetate, green vitriol, and concentrated sulfuric acid were analytical grade. Purified water with 18 M Ω cm resistivity was used throughout this work, after treatment with an ultrapure system (Liyuan Electric Instrument Co., Beijing, PRC). 2.2. Oxidation reaction and analysis Ultrapure water (500 mL) was put into a glass cylinder. An appropriate amount of KSF powder was added to the water and dispersed by stirring for 5 min. The initial pH of the KSF suspension was adjusted to the desired value by adding HCl or NaOH. The reactor, with a 300-rpm stirrer located near the base, was placed in a water
Fifteen experimental runs of the Box–Behnken design (BBD), which is a three-level incomplete factorial design extensively used, were performed to optimize the initial reaction rate, the derivative of the initial portion of the kinetic curve, and to arrange in order the importance of the various influencing factors. Table 1 shows the Box– Behnken design matrix for the three variables pH, KSF dosage and H2O2 dosage. The experimental data were analyzed using Design Expert 7.1.6 software and were fitted with a second-order polynomial equation. The optimized parameters for maximizing the initial rate were also obtained by using the established equation. 2.4. Sequential decolorization experiment To test the stability of KSF as a catalyst and to explore the possibility of recycling KSF, sequential decolorization experiments were performed, consisting of four kinetic runs. Each of the first three runs lasted for 40 min, and the fourth run lasted for 100 min. To guarantee that the initial concentrations of Orange II and H2O2 were the same for all runs (Orange II 20 mg/L, H2O2 4 mmol/L), specific amounts of Orange II and 4 mmol/L of H2O2 were added to the suspension before initiating the second, third and fourth runs. The dose of Orange II added was determined by the measured concentration of the remaining Orange II after the previous run. The experimental and analytical methods for the four runs were the same as for the oxidation reaction and analysis mentioned above. 2.5. Detection of Fe (II) and Fe (III) in a suspension Ultrapure water (500 mL) was poured into a glass cylinder reactor, and a measured amount of KSF powder was added to it. The suspension was stirred for 5 min and the acidity adjusted to pH 3 by the addition of HCl. The reactor was placed in a water bath at 30 °C. The entire reaction was conducted while stirring at 300 rpm. Once the
Table 1 Box–Behnken design matrix for three variables together with the observed and predicted responses. Experimental run
Coded variables pH, X1
KSF dosage, X2 g/L
H2O2 dosage, X3 mmol/L
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15
5 3 4 3 5 4 3 4 5 4 5 3 4 4 4
0.2 0.3 0.3 0.2 0.2 0.3 0.2 0.1 0.3 0.1 0.1 0.1 0.2 0.2 0.2
3.8 2.0 0.2 3.8 0.2 3.8 0.2 0.2 2.0 3.8 2.0 2.0 2.0 2.0 2.0
L. Chen et al. / Desalination 281 (2011) 306–311
stirrer started working, the timing was started. The samples were withdrawn from the reactor at regular time intervals (5, 10, 20, 30 and 40 min) and immediately filtered through a 0.22-μm millipore filter to remove the KSF particles. Samples of the filtrate (10 mL) were taken to measure the concentration of Fe (II) and were added to a colorimeter tube containing 2 mL 1:1 (v/v) HCl. Then, 2.0 mL of 2 mol/L NH4F, 2.0 mL of 0.2% (m/V) orthophenanthroline and 5.0 mL of a NH4Ac-HAc buffer solution (pH = 4.2) were successively added to the colorimeter tube. The solution was diluted by adding water until the total volume was 25 mL. After 15 min of coloration, spectrophotometric measurements using orthophenanthroline at 510 nm were used to detect the concentration of Fe (II). Another 10 mL of filtrate was collected to measure the concentration of total iron. The procedure to detect the total iron concentration was similar to that for Fe (II), except for the following steps: (1) 1.0 mL of 0.1 mol/L ascorbic acid was added after the addition of 10 mL of filtrate to reduce Fe (III) and the solution left for 30 min and (2) NH4F, a chemical used to complex with Fe (II), was not needed to measure the total iron. After waiting 30 min, 2.0 mL of orthophenanthroline and 5.0 mL of the NH4Ac-HAc buffer solution were successively added, after which the procedure continued as above. 2.6. Decolorization of Orange II in the homogeneous Fenton system Appropriate amounts of Fe (II) and Fe (III) standard solutions were added to the reactor containing 500 mL of ultrapure water, and the mixture was stirred for 5 min. For the reaction at pH 3, the concentrations of added Fe (II) and Fe (III) were 0.4 mg/L and 0.9 mg/L, respectively. For the reaction at pH 4, the concentrations of added Fe (II) and Fe (III) were 0 mg/L and 1.0 mg/L, respectively. The amounts of Fe (II) and Fe (III) that were added were similar to the quantities released from the KSF at the corresponding pH values. After this, the procedures, including regulating the pH value, stirring the mixture in a water bath, taking samples and analyzing, were the same as the oxidation reaction and analysis method described above. 3. Results and discussion 3.1. Control experiment Setting the initial concentration of Orange II at 20 mg/L and the pH at 3, the efficiency for the decolorization of Orange II by adding 0.2 g/L of KSF and/or 4 mmol/L of H2O2 was investigated. Fig. 1 shows the results of the decolorization of Orange II in the presence of KSF and H2O2. Previous research has demonstrated that Orange II has three absorption peaks, at 230, 310 and 485 nm. The wavelength of its visible absorption peak is 485 nm. The absorption peaks in the UV region at 230 nm and 310 nm are assigned to the benzene ring and naphthalene ring [26]. The trend in variation shows that the three absorption peaks in the visible region and UV region gradually disappear. This phenomenon suggests the dramatic destruction of chromophores and the decay of the benzene ring and naphthalene ring. Fig. 2 reveals that the quantity of Orange II adsorbed on the KSF exceeded 10%. The decolorization efficiency of Orange II with KSF alone after 40 min was 15.3%, and with H2O2 alone after 30 min was 6.3%. In the presence of KSF and H2O2, the decolorization efficiency reached as much as 91.6% after reacting for 40 min. The experimental data show that a synergistically cooperative effect was observed, with KSF and H2O2 acting together. This effect was due to the Fenton reaction. Catalyzed by the Fe (II) and Fe (III), the H2O2 decomposes and hydroxyl radicals, with, high reactivity, are produced to degrade the Orange II [27]. KSF is produced from acid-modified montmorillonite [28], whose iron content is 4.8%. When the KSF dosage was 0.2 g/L, the iron concentration in the solution was about 120 μmol/L. The iron was
100
Degradation Efficiency (%)
308
KSF+H2O2 KSF H2O2
80 60 40 20 0 0
10
20
30
40
Time (min) Fig. 2. Decolorization efficiency of Orange II in control experiments with the following experimental conditions: Orange II 20 mg/L, KSF 0.2 g/L, H2O2 4 mmol/L, pH 3.
completely released from the KSF. About 8% of the iron was released when the pH value ranged from 3 to 4 [28]. In this experiment, the total iron in the suspension was about 1.3 mg/L, or 23 μmol/L, when the pH value was 3 and the KSF dosage was 0.2 g/L. Considering the production batch and the actual experimental conditions, the difference between the actual concentration and the theoretical value was not significant. The amount of iron in the suspension was far less than typical sewage discharge limits, since the total iron discharge standard requirement of the USEPA is 100 mg/L [29].
3.2. Effect of pH on decolorization of Orange II The decolorization of Orange II in aqueous suspensions of montmorillonite KSF was studied in the pH range between three and nine. Table 2 shows the decolorization efficiency of Orange II after 40 min at different pH values using initial suspensions containing 20 mg/L Orange II, 0.2 g/L KSF and 4 mmol/L H2O2. At pH 3, almost complete decolorization (91.6%) was achieved in 40 min and this was also the case at pH 4 (91.2%). However, no significant decolorization was observed at pH values higher than 5. This phenomenon was caused by the acidic pH that favored the release of Fe (II) from KSF so that the Fe (II) made more of a contribution to the production of •OH [28]. Guo et al. reported that the stability of H2O2 is irrelevant as to whether the process is a homogeneous one or heterogeneous one. However, its stability is dependent on the pH value. The degree of H2O2 decomposition at pH 3 and 4 was relatively low. The OH decrease was caused by the unsuccessful decomposition of H2O2 that produces molecular oxygen instead of •OH [30]. At a pH
Table 2 Effect of individual variables on the decolorization efficiency after 40 min. Variables pH
Montmorillonite KSF dosage g/L
H2O2 dosage mmol/L
3 4 5 6 7 8 3 3 3 3 3 3
0.2 0.2 0.2 0.2 0.2 0.2 0.10 0.30 0.2 0.2 0.2 0.2
4 4 4 4 4 4 4 4 0.2 1 2 8
Decolorization efficiency (%)
91.6 91.2 24.2 5.7 6.8 4.3 52.5 94.7 30.7 63.2 80.4 99.8
L. Chen et al. / Desalination 281 (2011) 306–311
3.3. Effect of montmorillonite KSF dosage on the decolorization of Orange II An increase in the decolorization efficiency of Orange II was observed with an increase in the KSF dosage, with experimental conditions Orange II 20 mg/L, H2O2 4 mmol/L, pH 3. As shown in Table 2, the results of series of experiments indicated that after 40 min, the decolorization efficiency of Orange II at a KSF dosage of 0.20 g/L was significantly higher than at 0.10 g/L, but it was lower than at 0.30 g/L. This phenomenon illustrates that an increase of KSF dosage favored the enhancement of Orange II decolorization. This is because the amount of Fe (II) released from the KSF was small when the dosage of KSF in the solution was low and consequently the decolorization rate was slow. At a dosage of 0.20 g/L, 25.5%, 31.3%, 58.5%, 77.0%, and 91.6% of the Orange II was removed in 5, 10, 20, 30, and 40 min, respectively. Even when the catalyst concentration was increased to 0.30 g/L, increased decolorization of Orange II did not follow as, in this case, 27.3%, 36.5%, 60.3%, 86.5%, and 94.7% of the Orange II was removed in 5, 10, 20, 30 and 40 min, respectively. Considering the small difference of the effect on the decolorization between 0.20 g/L and 0.30 g/L, the value 0.20 g/L was selected as the dosage of KSF when other individual effects were investigated. 3.4. Effect of H2O2 dosage on the decolorization of Orange II The initial concentration of H2O2 plays an important role in the decolorization of dyestuffs in the Fenton process. A preliminary examination of the influence of the H2O2 concentration in the range 0.2 to 8 mmol/L on the decolorization of Orange II after 40 min is also presented in Table 2. The results indicate that the decolorization of Orange II was remarkably dependent on the H2O2 concentration. After 40 min, an increase in the H2O2 concentration from 0.2 to 8 mmol/L increased the decolorization efficiency of the dye from 30.7% to 99.8%, which showed that H2O2 could promote the decolorization of Orange II. Given the fact that 91.6% of Orange II was removed in the presence of 4 mmol/L H2O2 and that a further increase in the dosage of H2O2 led to no obvious enhancement of the decolorization of dye, a value 4 mmol/L of H2O2 was used in the series of experiments when studying the effect of pH and KSF dosage. Generally, the degradation efficiency of organic compounds in the Fenton system increases with an increased dosage of H2O2 until a maximal concentration is reached [31]. The cause of this is the promotion of •OH production in the low concentration range and the inhibition at higher concentrations when excessive H2O2 acts as a free-radical scavenger itself and thereby decreases the production rate of •OH [32]. 3.5. Experimental design, data analysis and optimization by RSM The data of 15 experimental runs using the Box–Behnken design (BBD) were analyzed using Design Expert 7.1.6 software and were
Table 3 ANOVA results of the quadratic model of initial reaction rate. Source
Sum of squares
Degree of freedom
Mean square
F value
p-value (prob N F)
Model Residual Lack of Fit Pure error Corrected total
0.22 3.817 · 10−3 2.212 · 10−3 1.605 · 10−3 0.23
9 5 3 2 14
0.025 7.634 · 10−3 7.373 · 10−3 8.027 · 10−3
32.35
0.0007
0.92
0.5589
R2 = 0.9831; adjusted R2 = 0.9527; F value: Fisher variation ratio; p-value: significance probability value.
fitted with a second-order polynomial equation. The function presented below was employed as the prediction model: Y = 0:54742−0:31761X1 + 3:02845X2 + 0:13342X3 −0:83005X1 X2 2
2
−0:014978X1 X3 −0:12113X2 X3 + 0:047290X1 + 2:64379X2 −3
−7:98521·10
2
X3
where Y is the initial reaction rate (mg/L·min), X1 is the pH (3 to 5), X2 is the montmorillonite KSF dosage (0.1 g/L to 0.3 g/L) and X3 is the H2O2 dosage (0.2 mmol/L to 3.8 mmol/L). As shown in Table 3, the analysis of variance (ANOVA) technique was used to test the significance and adequacy of the model, giving a correlation coefficient R 2 = 0.9831, which demonstrates a satisfactory quadratic fit. In addition, the sum of squares, degrees of freedom, mean square, F value and p-value, which can be used to check the fit of the model, are also presented in Table 3. This is also reflected in Fig. 3, where the observed value (calculated using the observed data when conducting the experiment under the conditions defined by Experiment design) is plotted against the predicted value (calculated by using the formula with the same conditions). The slope of the line from the linear regression fit is very close to 1. The Model F value of 32.35 implies that the model is significant. There is only a 0.07% chance that a Model F value this large could occur due to noise. In addition, the p-value (significance probability value) less than 0.0500 indicates that the model terms are significant. We conclude from the values of the coefficient estimate and p-value that, of the 3 variables, the pH (coefficient estimate − 0.14, p-value b 0.0001) produces the largest influence on the initial rate. Next, the most influential is the montmorillonite KSF dosage (coefficient estimate 0.052, p-value 0.0030), while the H2O2 dosage (coefficient estimate is 0.031, p-value = 0.0241) shows little effect on the response. The optimized parameters for maximizing the initial rate were obtained by using the established model. The optimal conditions
0.5
Predicted vs. Observed
0.4
Predicted
N4, the precipitation of Fe (III) from the solution can also result in a decrease in reaction activity. While the influence of pH on the behavior of H2 O2 was investigated, its impact on the adsorption behavior of Orange II was also investigated because that is of greater significance to the removal of dyestuffs. Its influence was studied at pH 3 and 4 where the initial concentration of Orange II was 20 mg/L and of KSF 0.2 g/L. The results demonstrated that during the 40 min reaction time, the amount adsorbed at pH 3 was higher than at pH 4. These results can explain the difference between the decolorization efficiencies at pH 3 and pH 4. It can be concluded that, to some extent, the high adsorption at pH 3 contributed to the better decolorization effect than at pH 4. In the decolorization reaction of Orange II, the decolorization efficiency at pH 3 was a little higher than at pH 4. Consequently, pH 3 was selected as the optimal pH value for further experimentation.
309
0.3 0.2 0.1 0.0 0.0
0.1
0.2
0.3
0.4
0.5
observed Fig. 3. The observed and predicted values of the initial reaction rate (R2 = 0.9831, adjusted R2 = 0.9527).
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L. Chen et al. / Desalination 281 (2011) 306–311
Decolorization efficiency (%)
determined from this study are as follows: pH 3; montmorillonite KSF dosage 0.30 g/L; H2O2 dosage 3.72 mmol/L. The predicted initial rate was 0.5048 mg/L·min, and the observed value was 0.5638 mg/L·min. Although the potential application of this heterogeneous Fenton system on real wastewater remains a target for the future, the present work as a basic method research was performed on synthetic solutions. With real wastewater, the negative effects of the matrix, such as adsorption and radical scavenging, could be relatively significant. Accordingly, the optimal conditions and the performance are subject to change. 3.6. Effect of initial Orange II concentration The effect of the initial Orange II concentration, which ranged from 5 mg/L to 40 mg/L, on its decolorization was also investigated. The results indicated that the decolorization efficiency at 5 mg/L was higher than at 40 mg/L of Orange II. The decolorization after 20 min was 100% at 5 mg/L and 40.5% at 40 mg/L. Even after 40 min, the final decolorization efficiency at 40 mg/L only reached 77.4%. In order to investigate the kinetics, the initial reaction rates for different concentrations were calculated. The derivative of the initial portion of the kinetic curve is the initial reaction rate r0. We used three models to describe our experimental data, firstorder, second-order and the Langmuir–Hinshelwood (L–H) equation. The results are presented in Table 4. By comparing the R 2 value for each equation, we concluded that the L–H equation fits best; the R 2 values for first-order, second-order and the L–H equations are 0.9014, 0.83201 and 0.9802, respectively. This result may be explained by the fact that the first/second-order equation suits the homogeneous system better. In the montmorillonite/H2O2 system, besides the adsorption behavior, some degradation behavior also exists. Based on this result, we chose the Langmuir–Hinshelwood (L–H) equation to describe Orange II decolorization for concentrations ranging from 5 mg/L to 40 mg/L r=−
dC kKC = ; dt 1 + KC
where C is the concentration of Orange II at time t, k is the rate constant and K is the adsorption coefficient. After calculation, k and K in this experiment were found to be 1.42 mg/L·min and 0.07014 L/mg, respectively. 3.7. Sequential decolorization experiments Sequential decolorization experiments were performed to test the stability of KSF as a catalyst. As shown in Fig. 4, after the first run, the decolorization efficiency of Orange II was 91.6% but, in subsequent runs, the decolorization efficiency decreased. At the end of the second run, the efficiency was 85.4% and it decreased to 77.5% after the third run. Nevertheless the decolorization efficiency reached as much as 94.5% after the fourth run, and the fact that the reaction time was prolonged to 100 min deserves attention. It was presumed that KSF would gradually lose its catalytic function, but it appears that a consistently high decolorization effect can be achieved by prolonging
100 80 60
1st run 2nd run 3rd run 4th run
40 20 0 0
50
100
150
200
250
Time (min) Fig. 4. Effect of recycling the catalyst on the decolorization of Orange II with the following experimental conditions: Orange II 20 mg/L, H2O2 4 mmol/L, KSF 0.2 g/L, pH 3.
the reaction time. For practical applications, further investigation on how to improve the stability of KSF could promote its wider application as a catalyst for the montmorillonite/H2O2 system and other wastewater treatment processes. 3.8. Role of iron species in the heterogeneous montmorillonite/H2O2 system To evaluate the role of both free iron (soluble iron) and structural iron (insoluble iron in crystalline lattice) in the KSF particles in the reaction, the amount of Fe (II) and the total iron released during the entire decolorization procedure was measured. A comparison of Orange II decolorization in a homogeneous system and in a montmorillonite/H2O2 system was made. The results revealed that at pH 3, about 1.3 mg/L of total iron and about 0.4 mg/L of Fe (II) were released after 5 min. However, at pH 4, there was hardly any Fe (II) released in suspension, and the total iron was about 1.0 mg/L. During the entire reaction process, slight fluctuations in Fe (II) and total iron concentrations over time occurred but, in general, these fluctuations were not significant. This variation suggested that nearly all the Fe (II) and total iron that existed in suspension was released at the beginning of the experiment. The dosage of iron in the homogeneous control experiment was the same as the iron concentration dissolving at pH 3 and 4, respectively. Fig. 5 shows that at pH 3, after 40 min, the decolorization efficiency of Orange II in a homogeneous system was 77.0% and in the montmorillonite/H2O2 system was 91.6%. However, at pH 4, after 40 min, the decolorization efficiency in the homogeneous system was only 6.7%, and in the montmorillonite/H2O2 system it was 91.2%. This phenomenon should be noted, because even though there was no Fe (II) released from the Montmorillonite KSF at pH 4, KSF could still play the role of a catalyst. This result indicates that the unreleased Fe (II) ions and structural Fe (II) in the crystalline lattice can catalyze the degradation of Orange II through the heterogeneous Fenton reaction in the absence of free Fe (II) ions. Since at pH 3, with a catalytic dose of free Fe (II) ions, the catalytic effect of structural iron was not significant, it is believed that unreleased Fe (II) ions made a greater contribution to the catalysis of
Table 4 Kinetics analysis of Orange II decolorization. C0 (mg/L)
Initial rate (mg/L·min)
Model
Kinetics equation
5 10 20 30 40
0.3584 0.6757 0.7711 0.8606 1.1255
L–H model
r=
First-order
r = 0.03134C (R2 = 0.9014)
3.13 · 10−2
Second-order
r = 0.0008361C2 (R2 = 0.83201)
8.361 · 10−4
0:0996C (R2 = 0.9608) 1 + 0:07014C
k (mg/L·min)
K (L/mg)
1.42
7.01 · 10−2
L. Chen et al. / Desalination 281 (2011) 306–311
pH 3, Montmorillonite/H2O2 system pH 3, Fe (II)/Fe (III)/H2O2 system pH 4, Montmorillonite/H2O2 system pH 4, Fe (II)/Fe (III)/H2O2 system
Decolorization efficiency (%)
100 80 60 40 20 0 0
10
20
30
40
Time (min) Fig. 5. Control experiment for the Fe (II)/Fe (III)/H2O2 system with the experimental conditions: Fe (II) 0.4 mg/L, Fe (III) 0.9 mg/L, Orange II 20 mg/L, H2O2 4 mmol/L, pH 3; Fe (II) 0 mg/L, Fe (III) 1.0 mg/L, Orange II 20 mg/L, H2O2 4 mmol/L, pH 4.
the oxidation in the montmorillonite/H2O2 system. This supports our original hypothesis that KSF has the potential to serve as a heterogeneous catalyst. 4. Conclusion In the presence of H2O2 or KSF alone, the decolorization of 20 mg/L Orange II was not clearly observed. When both H2O2 and KSF were in the solution, the decolorization efficiency was as high as 91.6% after 40 min. A second-order polynomial model based on response surface methodology using the Box–Behnken design clarified the order of importance of each individual factor on the initial reaction rate as pH N montmorillonite KSF dosage N H2O2 dosage. The model also predicted that the optimal conditions were pH 3, montmorillonite KSF 0.3 g/L and H2O2 3.72 mmol/L. At pH 3 and initial concentrations ranging from 5 to 40 mg/L, the decolorization of Orange II agreed with the L–H equation, and the efficiency decreased from almost 100% to 40.5%. In the heterogeneous reaction catalyzed by KSF at pH 3, the homogeneous reaction played the dominant role in the decolorization of Orange II, amounting to about 80% of the total decolorization. However, at pH 4, the unreleased Fe (II) ions and structural Fe (II) in the crystalline lattice can play a more important role in catalyzing the degradation of Orange II. This study indicated that KSF can be used as a new heterogeneous catalyst to induce the decolorization of some refractory organic compounds such as dyes. However, about 1.3 mg/L of total iron was released from the KSF at pH 3. Therefore, lowering the release of iron from KSF and improving the stability of KSF, which acts as a catalyst, are worth further study. Acknowledgments This work was sponsored by the Natural Science Foundation of China (NSFC No. 21077080), the Natural Science Foundation of Hubei Province (No. 2008CDB379) and the Wuhan Municipal Chenguang Project (No. 200750731254). The authors thank the anonymous reviewers for their comments on this manuscript. References [1] E. Forgacs, T. Cserhati, G. Oros, Removal of synthetic dyes from wastewaters: a review, Environ. Int. 30 (2004) 953–971.
311
[2] W. Tang, R. Jia, D.Q. Zhang, Decolorization and degradation of synthetic dyes by Schizophyllum sp. F17 in a novel system, Desalination 265 (2010) 22–27. [3] C.M. King, M.S. Lee, R.F. Jones, N. Tamura, Modification of plasmid and bacteriophage DNA by aromatic amines: effects on survival, template activity, and mutagenicity, Environ. Health Perspect. 102 (1994) 217–220. [4] P.L. Skipper, S.R. Tannenbaum, Molecular dosimetry of aromatic amines in human populations, Environ. Health Perspect. 102 (1994) 17–21. [5] S. Liakou, M. Kornaros, G. Lyberatos, Pretreatment of azo dyes using ozone, Water Sci. Technol. 36 (1997) 155–163. [6] Y. Xiong, P.J. Strunk, H. Xia, X. Zhu, H.T. Karlsson, Treatment of dye wastewater containing acid Orange II using a cell with three-phase three-dimensional electrode, Water Res. 35 (2001) 4226–4230. [7] K. Vinodgopal, J. Peller, O. Makogon, P.V. Kamat, Ultrasonic mineralization of a reactive textile azo dye, remazol black B, Water Res. 32 (1998) 3646–3650. [8] G.T. Li, N.G. Wang, B.T. Liu, X.W. Zhang, Decolorization of azo dye Orange II by ferrate(VI)–hypochlorite liquid mixture, potassium ferrate(VI) and potassium permanganate, Desalination 249 (2009) 936–941. [9] J.M. Monteagudo, A. Durán, C. López-Almodóvar, Homogeneus ferrioxalate-assisted solar photo-Fenton degradation of Orange II aqueous solutions, App. Catal. B Environ. 83 (2008) 46–55. [10] M. Inoue, F. Okada, A. Sakurai, M. Sakakibara, A new development of dyestuffs degradation system using ultrasound, Ultrason. Sonochem. 13 (2006) 313–320. [11] A. Bhattacharyya, S. Kawi, M.B. Ray, Photocatalytic degradation of Orange II by TiO2 catalysts supported on adsorbents, Catal. Today 98 (2004) 431–439. [12] J.H. Ramirez, C.A. Costa, L.M. Madeira, G. Mata, M.A. Vicente, M.L. Rojas-Cervantes, A.J. López-Peinado, R.M. Martín-Aranda, Fenton-like oxidation of Orange II solutions using heterogeneous catalysts based on saponite clay, Appl. Catal. B 71 (2007) 44–56. [13] G. Nagendrappa, Organic synthesis using clay catalysts: clays for ‘green chemistry’, Resonance 7 (2002) 64–77. [14] M. Nikpassand, M. Mamaghani, K. Tabatabaeian, M.K. Abiazi, KSF: an efficient catalyst for the regioselective synthesis of 1,5-diaryl pyrazoles using Baylis– Hillman adducts, Mol. Diversity 13 (2009) 389–393. [15] Y.X. Liu, K. Wan, N.S. Deng, F. Wu, Photodegradation of paracetamol in montmorillonite KSF suspension, React. Kinet. Mech. Catal. 99 (2010) 493–502. [16] M.M. Urrutia, E.E. Roden, J.M. Zachara, Influence of aqueous and solid-phase Fe (II) complexants on microbial reduction of crystalline iron(III) oxides, Environ. Sci. Technol. 33 (1999) 4022–4028. [17] Y.X. Liu, X. Zhang, L. Guo, F. Wu, N.S. Deng, Photodegradation of bisphenol A in the montmorillonite KSF suspended solutions, Ind. Eng. Chem. Res. 47 (2008) 7141–7146. [18] F. Bigi, L. Chesini, R. Maggi, G. Sartori, Montmorillonite KSF as an inorganic, water stable, and reusable catalyst for the Knoevenagel synthesis of coumarin-3-carboxylic acids, Org. Chem. 64 (1999) 1033–1035. [19] F. Bigi, S. Carloni, B. Frullanti, R. Maggi, G. Sartori, A revision of the Biginelli reaction under solid acid catalysis. Solvent-free synthesis of dihydropyrimidines over montmorillonite KSF, Tetrahedron Lett. 40 (1999) 3465–3468. [20] P. Kumar, R.V. Jasra, T.S.G. Bhat, Evolution of porosity and surface acidity in montmorillonite clay on acid activation, Ind. Eng. Chem. Res. 34 (1995) 1400–1448. [21] M.F. Brigatti, C. Lugli, G. Cibin, A. Marcelli, G. Giuli, E. Paris, A. Mottana, Z. Wu, Reduction and sorption of chromium by Fe (II)-bearing phyllosilicates: chemical treatment and X-ray absorption spectroscopy (XAS) studies, Clays Clay Miner. 48 (2000) 272–281. [22] S. Mendioroz, J.A. Pajares, I. Benito, C. Pesquera, F. Gonzalez, C. Blanco, Texture evolution of montmorillonite under progressive acid treatment: change from H3 to H2 type of hysteresis, Langmuir 3 (1987) 676–681. [23] J.H. Clark, D.J. Macquarrie, Catalysis of liquid phase organic reactions using chemically modified mesoporous inorganic solids, Chem. Commun. (1998) 853–860. [24] Y.X. Liu, X. Zhang, F. Wu, Photodegradation of bisphenol AF in montmorillonite dispersions: Kinetics and mechanism study, Appl. Clay Sci. 49 (2010) 182–186. [25] J. Li, F. Wu, N.S. Deng, E.M. Glebov, N.M. Bazhin, Degradation of Orange II by heterogeneous photocatalytic reaction using Montmorillonite KSF, React. Kinet. Catal. Lett. 95 (2008) 247–255. [26] F. Wu, N.S. Deng, H.L. Hua, Degradation mechanism of azo dye C. I. reactive red 2 by iron powder reduction and photooxidation in aqueous solutions, Chemosphere 41 (2000) 1233–1238. [27] R.V. Lloyd, P.M. Hanna, R.P. Mason, The origin of the hydroxyl radical oxygen in the fenton reaction, Free Radic. Biol. Med. 22 (1997) 885–888. [28] X. Zhang, F. Wu, N.S. Deng, I.P. Pozdnyakov, E.M. Glebov, V.P. Grivin, V.F. Plyusnin, N.N. Bazhin, Evidence of the hydroxyl radical formation upon the photolysis of an iron-rich clay in aqueous solutions, React. Kinet. Catal. Lett. 94 (2008) 207–218. [29] U.S. EPA regulations, 40 CFR: Protection of the Environment. Subchapter N — “Effluent Guidelines and Standards”. Parts 400–424 and 425–471, http://www.epa. gov/regulations/search/40cfr.html 2010. [30] J. Guo, M. Al-Dahhan, Catalytic wet oxidation of phenol by hydrogen peroxide over pillared clay catalyst, Ind. Eng. Chem. Res. 42 (2003) 2450–2460. [31] J. Feng, X. Hu, P.L. Yue, H.Y. Zhu, G.Q. Lu, Degradation of azo-dye Orange II by a photoassisted fenton reaction using a novel composite of iron oxide and silicate nanoparticles as a catalyst, Chem. Res. 42 (2003) 2058–2066. [32] R. Alnaizy, A. Akgerman, Advanced oxidation of phenolic compounds, Adv. Environ. Res. 4 (2000) 233–244.