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Modeling the Performance of an Innovative Membrane-Based Reactor. Abatement of Azo Dye (Orange II) up to Biocompatibility A. Lopez† and J. Kiwi*,‡ CNR-IRSA, Water Research Institute, Department of Water Chemistry and Technology, Via F. de Blasio 5, 70123 Bari, Italy, and Swiss Federal Institute of Technology (EPFL), Institute of Physical Chemistry II, Lausanne 1015, Switzerland
This study reports the kinetically fast and efficient decoloration/degradation of solutions of the azo dye Orange II photocatalyzed by a Nafion-Fe(1.78%) membrane in a photoreactor. Orange II solutions (30 ppm C), decolored under 36-W light in the presence of H2O2 (5.0 mM) in less than 0.5 h. The decoloration process was optimized via statistical modeling to attain the most economic use of energy, chemicals and time. Experimental evidence is presented for the dye sensitizing the degradation process. Orange II degradation is controlled by mass transfer and not by the chemical reaction of the oxidative radicals in solution. It was possible to attain a BOD5/COD ratio of ∼0.4 during the reactor pretreatment. A steep increase in the value of BOD5 was observed for pretreated solutions of Orange II because of the more readily biodegradable intermediates produced in solution. Introduction Finding better ways of treating and recycling effluents from the textile industry is a topic of timely interest because of the increased water shortages and escalating costs in many European countries. Nonbiodegradable azo dyes represent 15% of the world dye production. The application of innovative advanced oxidation technologies (AOTs) is of interest in reducing the treatment costs through destructive techniques. Reactors for pollutant abatement via Fe3+/H2O2 photocatalysis have recently been constructed.1 The dye abatement process is based on the generation of OH radicals (and other oxidative radicals) from H2O2 in the presence of added Fe3+ ions.2 This process is enhanced by light through the decomposition of photoactive Fe(OH)2+, leading to the generation of additional OH radicals in solution with concomitant partial recycling of the Fe3+ to Fe2+. 3,4 The light-induced reaction is shown below.
Fe(OH)2+ + hν f Fe2+ + ‚OH
(1)
However, Fe3+/H2O2 homogeneously catalyzed reactions need up to 80 ppm of Fe ions in solution, which is well above the European Economic Community directives that allow for a content of 2-4 ppm Fe3+ in treated wastewaters. Only in this case can the treated waters be discharged directly into the environment.5 To remove sizable amounts of Fe ions from solution at the end of the treatment, the precipitation and redissolution of the Fe ions have proven to be necessary. This involves added costs in chemicals and labor. To avoid these two drawbacks, Nafion perfluorinated membranes have recently been developed in which the Fe ion clusters have been fixed and shown to be active in H2O2 decomposition.6 These Nafion-Fe membranes resist OH radical attack (OH/OH- E° ) 1.90 V NHE) and do not leach out of the Fe exchanged on the sulfonic groups of
the Nafion within the 3000-h testing period. In the present study, the Nafion-Fe membrane is not used to perform a physical separation but to immobilize Feoxy-hydroxy cluster species that decompose H2O2 catalytically during the abatement of the azo dye Orange II. The present study intends to (a) investigate the solution parameters intervening in the decoloration/ degradation of Orange II in kinetically acceptable processes mediated by Nafion-Fe(1.78%) membranes in an immersion-type concentric reactor and (b) carry out statistical modeling for the optimization of the solution parameters during reaction operation. Modeling is especially meaningful in the case in which the effect of one reaction parameter depends on the setting of another one and vice versa. This is the case of immobilized Fenton systems. The optimization method used is appropriate no matter what the reactivities of the reagents or the radical intermediates (HO‚, HO2‚) are during the oxidation process and is independent of the nature of the different reaction steps. Experimental Section Materials and Techniques Employed. Mohr’s salt was used as the source of iron ions (ammonium ferrous sulfate hexahydrate, FeH8N2O8S2‚6H2O) to load the Nafion membrane with Fe ions. The details of the procedure for loading the Nafion with Fe ions has been reported recently in the literature.6 Orange II and the oxidant H2O2 (30% w/w) were from Fluka p.a. and were used as received. The formula of Orange II is shown below.
* Author to whom correspondence should be addressed. † CNR-IRSA. ‡ Swiss Federal Institute of Technology (EPFL). 10.1021/ie000420x CCC: $20.00 © 2001 American Chemical Society Published on Web 03/16/2001
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Figure 1. Schematic of the reactor used to decolor/degrade Orange II solutions by Nafion-Fe(1.78%) membranes in the presence of H2O2.
Total organic carbon (TOC) was monitored with a Shimadzu 500 provided with an automatic autosampler. The data processing of the TOC analyzer automatically rejects values with deviations >3% during the determination of the TOC value. Spectrophotometric measurements were carried out with a Hewlett-Packard 386/20 N diode array. Biological oxygen demand (BOD) measurements were carried out by means of a mercuryfree WTW 2000 Oxytop unit thermostated at 20 °C. Chemical oxygen demand (COD) was determined by a HACH reactor model CE provided with a DR/890 colorimeter. Reactor. A Philips 36-W (1.20 m long and 26 mm in diameter, TLD 36 W/08) black actinic light source (referred to as BL from this point on) was employed in such a way that its center passed through the focal axis of the reactor. The lamp radiation was centered at λ ) 366 nm and had a λ distribution between 330 and 390 nm. A second BL light with an output of 18 W was also tested to assess the effect of light intensity on the decoloration/degradation process. Figure 1 shows the schematic of the photoreactor used. During batch-mode operation, H2O2 was added to the mixing flask. The solution was recycled through the reactor loop by the pump, and samples were taken periodically from a sampling port available in the mixing flask. The NafionFe membranes (two pieces of dimensions 30 × 10 cm) floated freely in the space between the axial tube light source and the outer wall of the reactor jacket. Chemical Actinometry. The radiant flux was measured with a power meter from Yellow Springs, CO. Chemical actinometry experiments were carried out by means of an Aberchrome instrument in the λ range 310-545 nm. After photocoloration, the solution was bleached with white light and reused. The actinometer was irradiated under conditions similar to those found in the photoreactor. This latter step eliminated the need to make corrections for the reflectance and nonuniformity of the incident light beam. Determination of the quantum yields at various wavelengths was perfromed with a Baush and Lomb monochromator blazed at 300 nm. The incident beam on the solution had an area of 1 cm2. Results and Discussion Effect of H2O2 Concentration on the Degradation/Decoloration of Orange II. Figure 2 shows the results of some reactor runs in the presence of NafionFe membranes with solutions containing three different Orange II concentrations keeping a molar ratio of 31.2:1
Figure 2. TOC decrease for Orange II solutions in the presence of H2O2 under 36-W lamp irradiation (3 h) followed by dark reaction at pH 3. Batch-mode reactor operation with a recirculation rate of 280 mL/min on Nafion-Fe(1.78%). The solid lines from experimental results. The dotted lines are the theoretical curves from modeling. Inset: Decoloration of the solutions determined in a 1-mm cell.
between Orange II and H2O2. All of the FeOx/Fe3+ was fixed at the Nafion surface. Control experiments with homogeneous solutions of Fe3+/H2O2 showed a rate about twice as fast as that found when the Nafion-Fe membrane was used to catalyze H2O2 decomposition and consequently Orange II degradation. For the experiments of Figure 2, the Orange II solutions in the presence of the membrane were irradiated for 3 h at pH 3. After the initial stage of the lightinduced degradation a dark process of up to 20 h was allowed to further decrease the concentrations of Orange II and the intermediates generated in the initial stage. The slower degradation in the dark can also be ascribed to the reoxidation of Fe2+ ion to Fe3+ ion. The steeper decline observed in TOC values at higher Orange II concentrations is ascribed to the mass transfer limitation of the degradation. The three concentrations of Orange II (Figure 2) appear to give a transition from pseudo-first-order to zero-order kinetics within the reaction time. Laminar flow takes place in the concentric immersion reactor where the Reynolds number was found to be 160 for a recirculation rate of 280 mL/min or 11.2 × 10-3 m s-1. In this case, the mass transfer between the solution and the Nafion-Fe(1.78%) membrane will be directly proportional to the difference in Orange II concentration existing in the diffusion layer between the bulk of the solution and the membrane
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Figure 3. Reduction in optical absorbance (A) of Orange II under the same conditions as used in Figure 2 in the presence of the Nafion-Fe(1.78%) membrane.
interface. The diffusion distance (x) of the OH- (or any other oxidative radical formed in solution) away from the Nafion-Fe(1.78%) membrane can be estimated from the Smoluschowski diffusion relation
x2 ≈ Dτ
(2)
Taking the approximate values in eq 2 of D ≈ 10-5 cm2/s and τ ≈ 10-9 s for the oxidative radical, a diffusion length of ∼1 nm can be estimated for the diffusion layer of these radicals. Because the flow is laminar, the mass transfer between neighboring fluid layers (within the diffusion layer) proceeds entirely through molecular diffusion. Supposing that the concentration the radicals (e.g., OH radical) away from the Nafion-Fe follows a smooth function, the decrease in the OH radical concentration can be estimated from
d[OH] ) k[c][OH] dt
Figure 4. Reduction in TOC of a solution of Orange II (0.16 mM) as a function of H2O2 concentration under light irradiation of a 36-W lamp. Batch-mode reactor operation with a recirculation rate of 280 mL/min on Nafion-Fe(1.78%).
Figure 3 shows the reduction in optical density for the three different concentrations of Orange II in solution suggesting that (a) Orange II photosensitizes the degradation/decoloration process and (b) no new colored species appear during Orange II abatement. As in Figure 2, the black fluorescent 36-W light irradiation was applied for 3 h to induce the precursors active during the subsequent dark degradation of up to 20 h. Figure 4 shows that the degradation of Orange II solutions is seen to be more efficient as the H2O2 concentration increases up to 5 mM. At higher H2O2 concentrations, the abatement process becomes less efficient because of the scavenging of the H2O2 by the ‚OH radicals that are present at higher concentration in solution.
‚OH + H2O2 w H2O + HO2‚ (3)
Substituting numerical values in eq 3 for the concentration of OH radicals ≈ 10-12 M, the concentration of Orange II ) 0.25 mM, and k(Orange II + OH) ) 6 × 109, a value close to the diffusion-controlled rate,7 then d[OH]/ dt can be estimated by eq 3 as 1.5 × 10-9 M/s. As the reaction progresses with time, less Orange II is available in solution and the rate of decrease decays. This is shown in Figures 2 and 3. The theoretical curves shown by the traces in Figure 2 will be discussed in the modeling section below.
(4)
The overall stoichiometry for the mineralization of Orange II is in agreement with the above results and can be written as
C16H11N2NaO4S + 37/2O2 + 9/2H2O2 w 16CO2 + 8H2O + 2NO3- + NaHSO4- + 3H+ (5) Figure 5 suggests a reaction scheme for Orange II degradation under light irradiation catalyzed by the Nafion-Fe membrane. The proposed mechanism in Figure 5 is substantiated by the following experimental
Ind. Eng. Chem. Res., Vol. 40, No. 8, 2001 1855
Figure 5. Scheme for the reaction of Orange II at the NafionFe(1.78%) surface under light.
observations: (a) methanol added in solution (1%) precluded the degradation of the azo dye, (b) the pH of the solutions increased by 0.4 units (∼4 times) through the generation of OH- in solution, and (c) the initial redbrown Fe3+ of the Nafion-Fe membrane decolored to a beige color because of the increase in the colorless Fe2+ ion during the time of reaction. The membrane became catalytically faster with time because of the increase in the Fe2+ concentration. At times >1000 h, the initial red-brown color of the membrane due to Fe2O3 or Fe(III) oxyhydroxide was regenerated by immersion in NaOH (1 M) solution. Also in this way, the initial Orange II decoloration kinetics were reinstated. This finding has been described in work from our laboratory involving the photochemistry of Orange II with Nafion-Fe membranes in static reactors.6 To further substantiate the beneficial effect of light activation, the reduction of TOC was studied under two different light intensities. The results are shown in Figure 6. The two Nafion-Fe membranes used in the reactor were 30 cm long and 10 cm wide, amounting to a total surface of 600 cm2. When all of the Fe ions of the Nafion-Fe membrane were dissolved, a loading of 2.5 × 1018 molecules Fe3+ ions/cm2 was found.6 For a 600-cm2 membrane, this corresponds to about 1.5 × 1021 Fe3+ ions occupying about half of the total reactor surface (1636 cm2). The actinic light source (36 W) had a photon flux of 0.8 × 1015 photons s-1 cm-2 as determined by actinometry (see Experimental Section). The total number of photons reaching the membrane per second was then (0.8 × 1016) × 818 ) 0.65 × 1018. This number of photons is far below the number of Fe3+ ions available on the membrane surface, as mentioned before. Therefore, the work is carried out far below the saturation limit for the incoming light reaching the reactor. Recently, it has been shown that UV-visible light strongly accelerates Fenton reactions as stated in eq 1. The light enhancement is explained by Fe3+-sensitized reactions, mainly the photolysis of hydroxyde complexes of Fe3+ (see eq 1) yielding hydroxyl radicals and regenerating Fe2+. However, photochemical reactions of complexes formed between Fe3+ and Orange II8 or its degradation intermediates, especially organic acids,9 can also take place and cannot be ignored during Orange II degradation.10 Figure 6 show that a decrease in the degradation efficiency of Orange II occurs when the intensity of the lamp is decreased from 36 to 18 W. This is readily understood in terms of the available excited states created by light irradiation reaching the membrane surface. Light excitation induces Fe (d-d) transitions
Figure 6. TOC reduction of the solution of Orange II as a function of the intensity of the applied black fluorescent light. Batch mode reactor operation using a Nafion-Fe(1.78%) membrane with a recirculation rate of 280 mL/min and H2O2 (5 mM).
that are short-lived and occur on the nanosecond scale, leading to lower degradation rates when a lower photon density is applied.3 The overall degradation kinetics seem to be a function of the composition of the reactor solution, the loading of the membrane, the recirculation rate, and, as shown in Figure 6, the intensity of the applied light. Figure 7 shows the effect of the recirculation rate on the decoloration when the recirculation was varied between 140 and 420 mL/min in the reactor shown in Figure 1. When the recirculation rate is decreased, the slower recirculation increases the contact time of Orange II and the Nafion-Fe(1.78%) membrane. This, in turn, leads to the formation of an increased amount of oxidative radicals, leading to a higher decoloration rate. Pretreatment Leading to Biocompatibility by Means of the Nafion-Fe(1.78%) Membrane. Figure 8 shows the increase in biodegradability due to the Nafion-Fe(1.78%) membrane as a function of the Orange II abated in solution. The BOD5 values are seen to steeply increase up to 30 mg (O2/L) within 120 min, reaching a plateau. Because of the reactor pretreatment, Orange II is abated immediately from the beginning of the reaction. More biodegradable intermediates are generated in solution. The color of the solution fades away during the treatment. At times between 120 and 1200 min, no further increase in biodegradability was observed. The increase in BOD stops when Orange II has been completely removed from solution (see Figure
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Figure 9. Increase in the ratio BOD5/COD as a function of pretreatment time of an Orange II (0.40 mM) solution under 36-W lamp irradiation in the presence of H2O2 (10 mM) and NafionFe(1.78%) at pH 3. The insert shows the ratio of COD/TOC as a function of time. For other experimental details, see text.
Figure 7. Decoloration of a solution of Orange II (0.16 mM) in the presence of a Nafion (1.78%) membrane and H2O2 (10 mM) under 36-W lamp irradiation. Batch-mode reactor operation. Recirculation rates: (a) 140, (b) 280, and (c) 420 mL/min.
Figure 8. Increase in BOD5 as a function of time. Batch-mode reactor operation for an Orange II (0.40 mM) solution in the presence of H2O2 (10 mM) and Nafion-Fe(1.78%) under black light irradiation (36 W) at pH 3.
3), indicating that the dye degradation gives rise to two types of intermediates: (a) intermediates that can be easily mineralized, as shown by the continuous decrease of the TOC (Figure 2, initial stages), and (b) intermediates that cannot be mineralized, since, at times g1200 min, there are still about 20 ppm of C in solution (Figure 2, final stages). Figure 9 shows the increase in the BOD5/COD ratio for a solution of Orange II (0.25 mM) as a function of pretreatment time. The increase in the BOD5/COD ratio indicates a higher biodegradability of the pretreated
solutions until a BOD5/COD ratio of 0.4 is reached, which is the reference value for truly biodegradable effluents.5 Taking the chemical formula of Orange II [C16H11N2NaO4S], the average oxidation number of the organic carbon is -0.5 according to the relation oxidation number of C ) [4(TOC - COD)/TOC].11 Using the same formula, the average oxidation number of the organic carbon with TOC ) 76/12 ) 6.3 mmol and COD ) 202/ 32 ) 6.3 mmol is close to 0 after 3 h of pretreatment. This indicates that the remaining organic carbon after pretreatment is more oxidized than that at the beginning. Mathematical Modeling of the Degradation Parameters. Exponential Function. Experimental design methodology for the simulation of reaction variables is a modern approach in industrial reactor design.12 By following such an approach, a statistical model can be built that requires only a minimum number of experiments. Our group has used this approach for the modeling of the degradation of pollutants aiming at the most economical use of the oxidant, reaction time, light intensity, and electrical energy.1,13,14 More recently, the use of reduced centered dimensionless variables has gained wider attention. It was experimentally observed that the TOC decrease with time (Figure 2) followed an exponential decay. Therefore, a mathematical single-exponential expression was taken to fit the experimental results observed. This methodology is applied to the main variables contributing to the degradation process of Orange II, namely, H2O2 concentration, Orange II concentration, recirculation rate, and time. Experimental points are taken in pairs, x1 and x2, and labeled for the present treatment as the reaction parameters. Each graph represents a set of 48 pairs of values for TOC as a function of the five reaction parameters. Reduced centered dimensionless variables of these parameters are used in order to avoid having different units for different variables. Each reduced centered variable xi was specif-
Ind. Eng. Chem. Res., Vol. 40, No. 8, 2001 1857 Table 1. Reduced Centered Variables Used in the Modeling of the Degradation of Orange II with Different Concentrations of H2O2 as a Function of Time
Figure 10. Minimum regions for TOC attained during Orange II degradation found by statistical modeling of the degradation process. These TOC minimum regions correspond to 2-4 mg of C/L attained within 3 h of irradiation using the minimum time and oxidant at the most suitable recirculation rate in the reactor.
ically associated with a specific solution parameter ui in such a way that xi ) (ui - uio)/∆ui, where uio is the value of ui at the center of the experimental region and ∆ui )(uimax-uimin)/2. The treatment of the data takes the reaction parameters in pairs, x1 and x2. Each graph represents a set of 48 pairs of values for the TOC as a function of the reaction parameters. Then, a simple exponential expression can be constructed
∑(bixi + ∑biixi2 + ∑∑bijxixj)]
Z ) b0 exp[s
(6)
where
∑i Zi/N, the average of the values of the TOC
b0 )
over N experimental points
bi )
∑i Zixi/N, the coefficients for the main effect
of the variable xi
bii )
∑i Zixi /N, the coefficients for the quadratic 2
effect of the variable xi
bij)
∑i Zixixj/N, the coefficient for the first-order
interaction effect of xi and xj
and s is a scaling factor that is used to adjust the fitting of the curves for the initial and final concentrations of the reagents. The method for calculating the coefficients bi and bij is based on the coefficients bi representing first-order effects. A set number of six experimental points are taken with eight TOC values corresponding to each of the values found for xi, that is, 48 values in total. In each case, the TOC values are multiplied by the values of the corresponding xi. The 48 products found were added and divided by 48 to find the coefficient for bi. Coefficients denoted by bii represent the quadratic effect. Contour plots were obtained using the Igor 3.0 program on a Power Macintosh 8200/120 computer. The minimum regions for the Z values were located as a function of the combinations of variables taken in pairs in eq 6. The modeling allows for the drawing of contour plots or levels of constant response values through eq 7. The plots obtained allow the prediction at any point
time (min)
x2 (5 mM H2O2)
x2 (10 mM H2O2)
x2 (20 mM H2O2)
0 20 40 60 120 210 300 1200
1 13/14 6/7 5/7 2/7 0 -4/7 -1
1 13/14 6/7 5/7 2/7 0 -4/7 -1
1 14/15 13/15 11/15 1/3 1/15 -7/15 -1
(e.g., concentration of Orange II) in the experimental region of interest. The plots were obtained by calculating the coefficients of the exponential function in eq 7 and subsequently drawing the contour plot for the respective pair of variables. The central regions for the three contour plots indicate the minimum value for the function Z(TOC). The time variable has been fixed at 3 h (Figure 10). The central regions at the four contour plots indicate two regions in which the TOC attained a minimum. The central regions in Figure 10 represent degradation close to 95%. Therefore, the exponential function allows for the prediction of the optimal conditions for the TOC decrease up to almost complete TOC abatement in solution. The dotted lines in Figure 2 show the calculated values from the exponential model via eq 7. The experimental results are shown by the solid lines in Figure 2. The values of the adjustable parameters in eq 7 are obtained by fitting the initial and final TOC values. Each value of TOC can be calculated by way of the coefficients b0, b1, b2, b11, b12, and b22 to fit the experimental data.
Z(TOC) ) 29.7 exp[s(-0.249x1 + 16.4x2 + 24.2x12 + 17.69x22 + 1.26x1x2)] (7) where s is the scaling factor adjusted to fit the initial and final experimental values of the TOC in Figure 2. The values of xi, xi2, and xixj correspond to the experimental values at any time between 0 and 1200 min when Z(TOC) in eq 7 is optimized. The value of xi refers to the amount of H2O2 (in mL/min) present in the time interval chosen. That is, the real variable u is given by the expression xi ) (ui-uio)/∆ui, where uio is the value of ui at the center of the experimental region and ∆ui ) (uimax - uimin)/2. The value of x2 relates to u2 (the Orange II concentration), and this value is a constant as Orange II is added only at the beginning of the reaction. The product of x1x2 will vary according to x1 with time, since x2 stays constant during the treatment. In terms of reduced centered variables, the values of x1 for the three Orange II concentrations were selected as -1, -1/3, and +1, respectively. The values of x2 for H2O2 are given in Table 1. By way of the reduced centered values, the three theoretical curves (dotted lines) in Figure 2 were plotted. The experimental result (solid line) relates x1 (H2O2 concentration) and the values of x2 (Orange II concentration). The modeling through eq 7 has a predictive value for the lowest TOC value related to the three solution variables and allows the optimal combination of initial H2O2 concentration, recirculation rate, and initial
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Orange II concentration (∼60 mg/L) to be determined. This indicates how a given amount of Orange II can be degraded to 2-4 mg of C/L within 3 h of irradiation with the minimum of time and oxidant at the most suitable reactor recircultaion rate. Each of the two ellipsoids represents a solution of the system Z(TOC) in eq 7 leading to the lowest TOC values possible during the degradation of Orange II. These results are shown in Figure 10. Conclusions The Nafion-Fe reactor catalyzed treatment with H2O2 under light is shown to be effective in decoloring/ degrading the model azo dye Orange II. The kinetics of the Orange II abatement were initially first-order and varied with initial dye concentration in a way that suggests a reaction controlled by mass transfer. The intermediates formed under light during an initial irradiation period allow for Orange II abatement in the dark. For the treatment of wastewaters, a membranemediated process avoids the Fe-sludge posttreatment inherent in homogeneous systems. Modeling was carried out to achieve the most effective process in terms of the solution variables. The model is seen to fit the experimental data in a range better than 5%. The model used allows for optimization of the reaction parameters for the degradation process without a detailed knowledge of the reaction intermediates. Acknowledgment This work was supported by the E.U. ENV-CT95-0064 Program (OFES N° 96.350, Bern). We thank M. Manzano and J.-M. Quiroga of the Department of Chemical Engineering of the University of Cadiz, Spain, for their help in the initial stages of this work. Literature Cited (1) Balanosky, E.; Kiwi, J. Mathematical Modeling of the Photocatalytic Reactor Degradation of p-Nitrotoluene Sulfonate. Ind. Eng. Chem. Res. 1998, 37, 347-356.
(2) Walling, Ch. The Fenton Reaction Revisited. Acc. Chem. Res. 1975, 8, 125-131. (3) Ruppert, G.; Bauer, R.; Heisler, G. 1993 The photo-Fenton, an effective photochemical wastewater treatment process. J. Photochem. Photobiol. A 1993, 73, 75-78. (4) Nadtochenko, V.; Kiwi, J. Photoinduced adduct formation between Orange II and Fe3+(aq). Photocatalytic degradation and laser spectroscopy. Faraday Trans. 1997, 93, 2373-2378. (5) EEC, European Economic Community. List 1 of Council Directives 76/4647, Brussels, Belgium, 1982. (6) Fernandez, J.; Bandara, J.; Lopez, A.; Buffat, Ph.; Kiwi, J. Photoassisted Fenton Degradation of Nonbiodegradable Azo Dye (Orange II) in Fe-Free Solutions Mediated by Cation Transfer Membranes. Langmuir 1999, 15, 185-192. (7) Halmann, M. Photodegradation of Water Pollutants; CRC Press: Boca Raton, FL, 1996. (8) Bandara, J.; Morrison, C.; Kiwi, J. Degradation/decoloration of concentrated solutions of Orange II. Kinetics and quantum yields. J. Photochem. Photobiol. A 1996, 99, 57-66. (9) Zuo, Y; Hoigne´, J. Formation of H2O2 and Depletion of Oxalic Acid in Atmospheric Water by Photolysis of Iron(III)Oxalate Complexes. Environ. Sci. Technol. 1992, 26, 1014-1022. (10) Bandara, J.; Kiwi, J. Fast kinetic spectroscopy, decoloration and production of H2O2 induced by visible light in oxygenated solutions of the azo dye Orange II. New J. Chem. 1999, 23, 717724. (11) Scott, P. J.; Ollis, F. D. Integration of Chemical and Biological Oxidation Processes for Water Treatment: Review and Recommendations. Environ. Prog. 1995, 14, 88-103. (12) Box, G.; Hunter, W.; Hunter, J. Statistics for Experimenters. An Introduction to Design, Data Analysis and Model Building; John Wiley & Sons: New York, 1987. (13) Balanosky, E.; Fernandez, J.; Kiwi, J.; Lopez, A. Degradation of membrane concentrates of the textile industry by Fentonlike reactions in Fe-free solutions at biocompatible pH (7-8). Water Sci Technol. 1999, 40, 417-424. (14) Balanosky, E.; Herrera, A.; Lopez, A.; Kiwi, J. Degradation of Textile Wastewaters. Modeling the Reaction Performance. Water Res. 2000, 34, 582-596.
Received for review April 19, 2000 Revised manuscript received October 10, 2000 Accepted October 24, 2000 IE000420X
