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Enhanced Biodegradation of Azo Dyes Using an Integrated Elemental Iron-Activated Sludge System: II. Effects of Physical-Chemical Parameters

Posted on: Wednesday, 15 February 2006, 06:00 CST

By Saxe, Jennie Perey; Lubenow, Brian L; Chiu, Pei C; Huang, Chin-Pao; Cha, Daniel K

ABSTRACT:

As part of a study to evaluate an integrated zero-valent iron (Fe^sup 0^)-biological oxidation process for treating azo dye wastewaters, we conducted batch and column experiments with the azo dye orange G to assess the effects of solution conditions on the performance of iron pretreatment. The influence of iron type and surface area, solution pH, dissolved inorganic salts, and phosphate ion on the reduction (decolorization) of orange G solution were examined. In batch experiments, increased iron surface area, decreased pH, and chloride and sulfate salts enhanced dye decolorization, whereas high pH (9.9) and phosphate concentrations (>3 mg/L PO^sub 4^-P) inhibited dye reduction. Results from batch experiments were confirmed in column experiments. An increase in temperature from 22 to 35C resulted in a near doubling of the reduction rate constant in a column study. The abiotic reduction results illustrate the feasibility and potential limitations of an integrated iron column, activated sludge treatment process for wastewaters containing azo dyes. Water Environ. Res., 78, 26 (2006).

KEYWORDS: azo dyes, elemental iron, chemical reduction, wastewater treatment.

doi: 10.2175/106143005X84486

Introduction

Azo dye production worldwide has been estimated at 6.35 10^sup 8^ kg (7 10^sup 5^ tons) per year (Zollinger, 1987). These dyes, containing the azo function (R-N=N-R'), are designed to be stable in the media to which they are applied, resulting in molecules that are resistant to light and degradation under aerobic conditions. This stability is essential for dyeing applications, but treatment of azo dyecontaining wastes has been problematic because of the refractory nature of these compounds in traditional aerobic biological processes (Dohnyos et al., 1978; Hitz et al., 1978; Shaul et al., 1991). Other treatment methods have been investigated, including chemical reduction (Cao et al., 1999; Nam and Tratnyek, 2000), advanced oxidation (Arslan and Balcioglu, 1999; Goncalves et al., 1999), and anaerobic biological treatment (Chung and Stevens, 1993; Jiang and Bishop, 1994; Keck et al., 1997; Kudlich et al., 1997).

The chemical reduction of azo dyes with zero-valent iron (Fe^sup 0^) has been studied (Cao et al., 1999; Nam and Tratnyek, 2000; Percy et al., 2002), and the results were encouraging. Azo dyes have been shown, in our previous study (Perey et al., 2002), to decolorize in a relatively short time through iron treatment, yielding aromatic amines as products that are more amenable to aerobic biodegradation. This result suggests that sequential iron treatment and aerobic biodegradation may be a feasible method to treat wastewaters containing azo dyes.

In this study, batch and column experiments were performed to investigate several physical and chemical factors that may influence the reduction of orange G under anaerobic conditions. These factors include iron surface area, solution pH, temperature, and the anions chloride, sulfate, and phosphate. The objective of these experiments was to determine whether these parameters are important in potential application of iron for the pretreatment of azo dye wastewater before aerobic biological oxidation processes, such as activated sludge.

Materials and Methods

Chemicals. Orange G dye (C.I. 16230, 86% purity) was purchased from Sigma (St. Louis, Missouri). Cast iron filings were obtained from Master Builders, Inc. (Aurora, Ohio). High-purity iron granules (1 to 2 mm, 99.98%, metal basis) were purchased from Alfa Aesar (Ward Hill, Massachusetts). Master Builders iron is known to contain 3.2% carbon, 0.65% manganese, and 0.6% sulfur (Reardon, 1995). The surface area was 1.28 m^sup 2^/g for Master Builders iron and 0.0417 m^sup 2^/g for the high-purity iron granules, as reported previously (Perey et al, 2002).

Batch Reduction Experiments. Batch reduction experiments were conducted in an anaerobic glove box (Coy Laboratory, Grass Lake, Michigan), with an atmosphere of 95% nitrogen (N^sub 2^) and 5% hydrogen (H^sub 2^). Borosilicate vials (8 mL) were used, which contained 5 mL of aqueous orange G solution (50 mg/L) and varying amounts (0.5 to 2.5 g) of iron. Test solutions were purged with N^sub 2^ and then stirred in the anaerobic glove box for complete deoxygenation before experiments. Reaction vials were placed horizontally in a rack and rotated on an orbital mixer at 100 rpm. All batch experiments were conducted at room temperature (21 1C). Replicate vials were set up for each experiment, and, at each sampling time, one of the vials was sacrificed, supernatant was vacuum filtered through a glass-fiber filter, and the filtrate was analyzed by spectrophotometry or high-pressure liquid chromatography (HPLC). The changes in orange G concentration, under different experimental conditions, were monitored as a function of time.

To investigate the effect of pH on the reduction rate of orange G, the following pH buffers were used at 0:2 M to give an initial pH of 7.7, 4.3, and 9.9, respectively: HEPES [N- (2-hydroxyethyl)- piperazine-N'-(2-ethanesulfonic acid)], acetate, and CAPSO [3- (cyclohexylamino)-2-hydroxy-1 -propanesulfonic acid].

Because inorganic salts are commonly present in dye wastewaters (Carliell et al., 1998), experiments with chloride and sulfate salts were conducted to assess the effect of these solutes on orange G reduction rate. Also, because phosphate is known to be inhibitory to iron corrosion processes (McNeill and Edwards, 2000; Mutombo and Hackerman, 1998), experiments were also performed in phosphate solutions at different concentrations. Orange G was prepared in a phosphate buffer and in monobasic (KH^sub 2^PO^sub 4^) and dibasic (K^sub 2^HPO^sub 4^) potassium phosphate solutions. To separate the pH effect from the effect of phosphate, orange G and potassium phosphate were also prepared in O. IM HEPES buffer (initial pH approximately 7.7). Batch reduction experiments with chloride, sulfate, and phosphate were performed, as described above, with 0.5 g cast iron in 5 mL test solution.

Figure 1-Effects of (a) iron type and (b) pH on ironmediated reduction of orange G in batch experiments. Test conditions: 2.5 g iron/5 mL solution and 100 rpm mixing. Data points are average values from triplicate vials and error bars represent standard deviations.

Column Experiments. The effects of chloride and phosphate ions and temperature on orange G reduction efficiency were further tested in a continuous-flow column system. Glass columns (8 cm L 1.3 cm i.d.) were packed with 20 to 30 mesh (0.5 to 0.7 mm) Ottawa Sand (Fisher Scientific, Pittsburgh, Pennsylvania) and Master Builders cast iron and solution was pumped at a flowrate between 0.35 and 0.55 mL/min, resulting in a residence time of 22 to 36 minutes. For the phosphate experiments, the column was packed with a 2:1 (wt:wt) mix of sand and cast iron. For the chloride experiment, the front three-fourths of the column was packed with sand only, and the last one-fourth was the 2:1 sand/iron mixture. (A similar setup was used to study the effect of temperature, with the front two-thirds of the column being sand and the remaining onethird being 2:1 sand/iron mixture). In the chloride and temperature experiments, the residence time of the solution in the iron-containing portion of the column ranged from 5 to 12 min. These alterations were made to prevent complete reduction of the dye in the column to more clearly demonstrate the effects of increased chloride concentration and temperature. To ensure anaerobicity, the test solutions were continuously purged with N^sub 2^ during the experiment.

The orange G solution was prepared at 100 mg/L for all column experiments. Flowrates were adjusted during the initial period of the experiment to establish steady-state conditions, which were taken to be when effluent pH and dye concentration became invariant. The initial pH of the unbuffered orange G and sodium chloride solution was approximately 5.7. The effect of chloride was examined at CF concentrations of 0.01, 0.1, and 1.0 M. Phosphate experiments were performed in dibasic sodium phosphate (Na^sub 2^HPO^sub 4^) solution at concentrations of 0.1, 1.0, and 10 mM. The initial solution pH was adjusted to 7.2 after Na^sub 2^HPO^sub 4^ addition. The effect of temperature was examined at room temperature (22C) and in a 35C temperature-controlled room. Samples were filtered through a Whatman GF/A glass-fiber filter (Kent, United Kingdom; nominal pore size 1.6 m) before analysis.

Analytical Protocol. Routine quantification of orange G was performed using a Hach DR/2010 spectrophotometer (Hach Company, Loveland, Colorado) or a Hewlett Packard (Boston, Massachusetts) 8452A diode array spectrophotometer at a detection wavelength of 475 nm. Additional analysis was also performed using a Varian HPLC equipped with a guard column and reversephase Supelcosil C-18 (Supelco, Bellefonte, Pennsylvania) column. At a flowrate of 1 mL/ min and using an eluent of 50/50 (v/v) methanol/water, the retention time of orange G was 2 min. Calibration curves were constructed using samples of known concentrations and a detection wavelength of 254 nm.

Results and \Discussion

Batch Reduction Experiments. Effect of Iron Type. Two types of iron were compared in batch reduction experiments with orange G: Alfa Aesar high-purity iron granules (Ward Hill, Massachusetts) and Master Builders cast iron. Figure Ia illustrates the increased reduction efficacy of cast iron (on a mass basis) over pure iron for the azo dye orange G. The faster reduction of orange G by the cast iron was most likely because of the difference in surface area between cast iron (surface area = 1.28 m^sup 2^/g) and high-purity iron (surface area = 0.042 m^sup 2^/g). Because the cast iron has a higher surface for reaction and is available at a fraction of the cost of the high-purity iron granules, it is a more cost-effective and practical choice for treatment applications. Consequently, cast iron was used for the majority of the batch-reduction experiments and all of the column experiments.

Effect ofpH. Solution pH is another important parameter, which was expected to affect the reduction rate. Reduction of orange G was rapid (within 5 min) in batch reactors with cast iron filings at pH 7.7 and 4.3 (Figure Ib). Increasing the pH of the test solution to 9.9 markedly decreased the rate of organic G reduction. Similar pH effect has been observed by other researchers (Agrawal and Tratnyek, 1996; Cao et al., 1999). The result suggests that acidic and circum- neutral pH favors azo dye reduction, but the reaction may be retarded at alkaline pH.

Effect of Sulfate and Chloride Salts. Because chloride and sulfate salts are commonly used in dye-manufacturing processes (Carliell et al., 1998), several of these salts were selected to investigate the effect of salts on orange G reduction with cast iron. All six chloride and sulfate salts tested (at concentrations up to 0.1 M) slightly enhanced the dye reduction rate (Figure 2). According to MacDougall and Graham (1995), the presence of "aggressive anions" can disrupt the passivation (iron oxide) layer at the iron surface, thus enhancing iron corrosion. As shown in Figure 2a, higher concentrations of magnesium chloride (MgCl^sub 2^) and magnesium sulfate (MgSO^sub 4^) resulted in more rapid reaction. Figure 2b shows the effect of five chloride salts on orange G reduction rate, relative to the deionized water control. All of these salts, at 0.1 M, increased the reaction rate, regardless of the cation type.

Figure 2-Fraction of orange G remaining in solution with (a) increased concentrations of magnesium chloride (MgCI^sub 2^) and magnesium sulfate (MgSO^sub 4^) and (b) 0.1 M chloride salts in batch experiments. Test conditions: initial dye concentration 50 mg/ L; 0.5 g cast iron/5 mL test solution; 100 rpm mixing; initial pH = 7.

Effect of Phosphate Ion. A wide range of K^sub 2^HPO^sub 4^ concentrations (in unbuffered solution) was used to examine the effect of phosphate on the dye-reduction reaction. Using 11 concentrations of K^sub 2^HPO^sub 4^, ranging from 0.003 mM to 5 mM, the data clearly show that higher phosphate concentration resulted in decreased dyereduction efficiency (Figure 3). There appears to be a (empirical) linear correlation between percent orange G reduction and the logarithm of phosphate concentration.

This result is not surprising, because phosphate has been used extensively as a corrosion inhibitor for ferrous and nonferrous metals. The phosphating of an iron surface involves multiple steps (cleaning, sealing, and rinsing) and produces a surface that is not only resistant to corrosion, but improved with respect to adhesion of paints to the metal surface (Grass, 2000). Persson et al. (1996) analyzed IR spectra of phosphate on goethite and hematite surfaces and showed that the bonding of phosphate at the iron-water interface occurs because of either phosphate adsorption to Fe(III) oxides, precipitation of iron phosphates on the iron surface, or a combination of both.

Figure 3-Effect of phosphate concentration on orange G reduction by cast iron in batch experiments. Test conditions: 0.5 g cast iron/ 5 mL test solution, initial orange G concentration = 50 mg/L, initial pH = 8.62 0.12, and reaction time = 30 min. Error bars represent standard deviation of triplicate samples at each concentration.

Additional batch experiments (Saxe, 2003), in phosphate-buffer solutions at lower pH, yielded some interesting results. When the pH of a 10-mM phosphate solution was adjusted to 6.8 (near the second phosphate pKa of 7.2) before the experiment, 92.7% orange G reduction was achieved under the same conditions (Saxe, 2003). Because of the relatively high buffering capacity, the pH remained at approximately neutral throughout the test, and the inhibitory effect on dye reduction was minimal, despite the high phosphate concentration. This result, which is in sharp contrast to that in Figure 3, suggests that the extent of phosphate inhibition on iron reactivity depends strongly on pH and, possibly, phosphate speciation.

Column Experiments. Effect of Chloride Ion. Nonbuffered orange G solutions containing 0.01, 0.1, and 1.0 M sodium chloride (NaCl) were continuously fed to the iron columns for 18 days. Over 3500 pore volumes, no decrease in dye reduction efficiency was observed in any of these columns (in fact, a slight increase in removal was observed in all columns). Relatively small amounts of iron were used to pack the columns to prevent complete reduction of orange G and to better discern the chloride effect. Figure 4a illustrates that increasing chloride concentration increased the extent of orange G reduction, although the effect was small, which is consistent with the batch experiment result. The effluent pH for the control, 0.01,0.1, and 1.0 M chloride experiments were 6.1,7.1, 8.7, and 9.4, respectively, because of iron corrosion in the columns. Thus, the effluent pH reflects the rate of iron corrosion as influenced by the different chloride concentrations. Had the solutions been buffered to maintain a constant pH, the effect of chloride would have been more pronounced, because the pH increase might have offset some of the enhancement because of chloride. Figure 4b shows the corresponding pseudo-first-order rate constants calculated from initial and effluent dye concentrations. For most data points, including the deionized water blank, the rate constant falls between 0.09 and 0.12 min^sup -1^. With increasing chloride concentration, and most notably at 1.0 M, the percent reduction and the reaction rate constant both increased. Under the conditions used, NaCl at concentrations up to 1 M did not inhibit, and may enhance slightly, the transformation of azo dyes with iron filings.

Figure 4-Effect of chloride on orange G reduction in column experiments: (a) effluent concentration of orange G in O, 0.01, 0.1, and 1.0 M sodium chloride solutions and (b) average pseudo-first- order rate constants for each chloride concentration. Initial pH of test solution = 5.7 and effluent pH for the control, 0.01, 0.1, and 1 M Cl- experiments were 6.1, 7.1, 8.7, and 9.4, respectively. Error bars represent standard deviations.

Figure 5-Effect of phosphate on orange G reduction with iron in column experiments, (a) Percent reduction of orange G in O, 0.1, 1.0, and 10 m M sodium phosphate solutions and (b) average pseudo- first-order rate constants for each concentration. Initial pH of test solutions = 7.2. Effluent pH values were 9.7,9.5, and 7.3 for 0.1,1.0, and 10 mM phosphate concentration, respectively. Error bars represents standard deviations.

Effect of Phosphate Ion. Results of the column experiments, with orange G in phosphate buffer solution (initial pH 7.2), are shown in Figure 5. Steady-state conditions were established in approximately 5 days (680 pore volumes), at which point the effluent pH was 9.7, 9.5, and 7.3 for 0.1, 1.0, and 10 mM phosphate. Apparently, more than 1 mM of phosphate buffering capacity was consumed in the iron column under the experimental conditions. The reduction efficiency decreased with increasing phosphate concentration (Figure 5a), consistent with the result from batch experiments using potassium phosphate. In contrast to the batch result (Saxe, 2003), however, low pH failed to negate the adverse effect of high phosphate concentration over extended operation. The column receiving 10 mM phosphate buffer exhibited the lowest steady-state orange G reduction efficiency (<20%), despite its neutral effluent pH. This suggests that the low pH effect was relatively short-term and could not prevent inactivation of iron surface over time because of phosphate adsorption and/or precipitation. Note that, in the early stage of the column experiment (e.g., the first 100 pore volumes), the 10mM phosphate column exhibited an efficiency (>95%) as high as that of the control (deionized water) column, which is consistent with the batch result.

The influent and effluent orange G concentrations were used to calculate the pseudo-first-order rate constants, which decreased with increasing phosphate concentration (Figure 5b). The inhibitory effect of phosphate on iron reactivity is abundantly clear; even at 0.1 mM, phosphate decreased the steady-state orange G reduction rate constant by over 70%, from 0.329 0.038 min^sup -1^ to below 0.1 min^sup -1^.

Effect of Temperature. Increased temperature has been implicated in passive film failure in corrosion-prevention applications (Jones, 1996). To examine the effect of temperature on orange G reduction, column experiments were conducted at ambient temperature (22 C) and at 35 C in a temperature-controlled room. Because the reaction was rapid, only one-third of the column was packed with iron to decrease percent orange G reduction and to better discern the temperature effect. At 22C, the percent orange G reduction was 74.6 4.7%, and the calculated pseudo-first-order reduction rate constant was 0.336 0.043 min^sup -1^. In comparison, the percent orange G reduction was 91.2 0.5% at 35\C, while the calculated k value was 0.657 0.006 min^sup -1^, which nearly doubled as a result of a 13C increase in temperature. An apparent activation energy of 38.9 kj/mol was calculated based on these data. This activation energy corresponds to an 80% increase in rate constant for a 10C increase in temperature (Schwarzenbach et al., 2002). While the temperature effect was pronounced, what process this activation energy corresponds to is unclear, because temperature can have effects on multiple processes, including iron corrosion, solute transport, and surface reactions.

Some dyeing waste streams may have elevated temperatures because of the high-temperature processes involved in dye manufacturing. For example, the acid coupling step for the production of basic orange 2 (C.I. 11270) requires heating the solution to 60C (Azo Dyes, 1978). Our result suggests that effluent temperature is a process-specific parameter that should be taken into account when designing an iron treatment process for dye effluents.

Summary

This study examined a number of chemical and physical parameters that may affect the cast iron-mediated reduction of the azo dye orange G. Results of the batch and column experiments indicate that high iron surface area, low and neutral solution pH, addition of chloride and sulfate salts, and high temperature would increase orange G reduction efficiency. In contrast, alkaline pH and the presence of phosphate can inhibit azo dye reduction.

These solution parameters are important in feasibility evaluation and design of an iron pretreatment process for enhancing aerobic biodegradation of azo dyes in wastewaters. The applicability and long-term performance of the iron process may be further affected by permeability change over time because of, for example, mineral formation and microbial activities, which were not investigated in this study. Pilot-scale and longer-term studies are necessary to better evaluate the feasibility of the sequential iron treatment- biological oxidation process and the service life of iron.

References

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Acknowledgments

Credits. The authors gratefully acknowledge the Water Environment Research Foundation (Project 99-CTS-3-UR) (Alexandria, Virginia) for funding this research.

Authors. Jennie Perey Saxe received her Ph.D. from the Department of Civil and Environmental Engineering at the University of Delaware (Newark). Brian L. Lubenow received his M.C.E. from the University of Delaware in 2002 and is currently with Camp Dresser & McKee, Inc. Pei C. Chiu is an associate professor, C. P. Huang is the distinguished professor, and Daniel K. Cha is an associate professor in the Department of Civil and Environmental Engineering at University of Delaware. Correspondence should be addressed to Daniel K. Cha, Department of Civil and Environmental Engineering, 301 DuPont Hall, University of Delaware, Newark, DE 19716; e-mail: cha@ce.udel.edu.

Submitted for publication April 1, 2003; revised manuscript submitted August 10, 2004; accepted for publication November 15, 2004.

The deadline to submit Discussions of this paper is April 15, 2006.

Copyright Water Environment Federation Jan 2006

Source: Water Environment Research

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