Environ. Sci. Technol. 2005, 39, 7158-7163
EDTA Degradation Induced by Oxygen Activation in a Zerovalent Iron/Air/Water System CHRISTINA E. NORADOUN AND I. FRANCIS CHENG* Department of Chemistry, University of Idaho, Moscow, Idaho 83844-2343
A method for the removal of ethylenediaminetetraacetic acid (EDTA) at room temperature and 1 atm is demonstrated. EDTA (1 mM, 50 mL) containing 2.5 g of granular zerovalent iron (ZVI) (20-40 mesh) was degraded in 2.5 h. Using a recently developed form of O2 activation, reactive oxygen species are generated in situ, resulting in the degradation of EDTA when complexed with FeII. ESI-MS measurements indicate that degradation of EDTA yields low-molecular carboxylic acids. The presence of oxygen is crucial: the observed pseudo-first-order rate constants for EDTA removal are kobs ) 1.02 h-1 (kSA ) 1.85 ( 0.046 L h-1 m-2) and kobs ) 0.04 h-1 (kSA ) 0.00724 ( 0.002 L h-1 m-2) under air and under N2 purge, respectively. kSA represents surface area normalized rate constants. Large excesses of EDTA in the reaction mixture slow the rate of degradation. Increasing the concentration of EDTA from 1.0 to 10.0 mM while holding all other parameters constant gave observed rates of kobs ) 1.02 ( 0.26 h-1 (kSA ) 1.85 ( 0.046 L h-1 m-2) and kobs ) 0.044 ( 0.01 h-1 (kSA ) 0.00796 ( 0.002 L h-1 m-2), respectively. The rate-limiting step is determined to be homogeneous oxygen activation.
Introduction Wastewaters containing complexing agents such as diethylenetriaminepentaacetic acid (DTPA) and ethylenediaminetetraacetic acid (EDTA) have become an environmental concern (1). One of the reasons for this alarm is the possibility of toxic heavy metal mobilization, extended biological availability to aquatic life, and the subsequent risks these metals pose to groundwater and drinking water. The use of chelating agents has become widespread over the past decade. They are used to moderate the adverse effects of transition metal ions on the performance of laundry detergents, cosmetics, and photochemicals. Chelating agents have also proven useful in the manufacture of textiles and in paperpulp bleaching (2). Transition metal ions such as iron, zinc, manganese, and copper can trigger chemical processes such as corrosion, catalytic degradation, polymerization inhibition, and redox reactivity, necessitating the control of their activity in manufacturing (2). The growing popularity of green bleaching technologies, which include hydrogen peroxide and ozone centered processes, has also added to the concern regarding increased EDTA release (3). Increased consumption of EDTA is attributable to its ability to control reaction rates in green oxidation schemes. The use of EDTA in high [EDTA]: * Corresponding author phone: (208)885-6387; fax: (208)885-6173; e-mail:
[email protected]. 7158
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 39, NO. 18, 2005
[metal] ratios is an inexpensive way to effectively bind transition metals and suppress their catalytic activity (4). Molar ratios of 2.5:1 ([EDTA]:[FeII/III]) and above have been shown to inhibit the electrocatalytic reduction of H2O2 in aqueous solutions (4). An example of the increasing use of anthropogenic chelating agents comes from the Swedish pulping industry where the use of EDTA has increased from 2000 metric tons to about 8000 metric tons per year from 1990 to 2000 (2). Many industrial chelating agents are not degradable by methods currently found in wastewater treatment facilities and consequently are not being removed (2, 5-8). Kari and Giger have demonstrated that considerable amounts of EDTA pass through wastewater treatment facilities in the form of FeIIIEDTA. These investigators found EDTA concentrations as high as 18 µM in effluent from sewage treatment centers (9). Numerous studies have investigated a wide range of systems for the degradation/removal of EDTA, e.g., enriched bacterial cultures (10), H2O2 (11, 12), hydrolysis, and photolysis (7). The present investigation uses a recently developed oxygen activation scheme for the removal of EDTA. It is unique in that it is a nonbiological system employing lowcost materials under mild reaction conditions at room temperature and 1 atm pressure (RTOP) (13). The system only requires zerovalent iron, EDTA, and air (ZEA) under aqueous RTOP conditions. A previous investigation reported by this research group demonstrated that the ZEA reaction was able to destroy chlorinated phenols, producing low molecular weight carboxylates. In that previous investigation, it was discovered that EDTA was concomitantly degraded (13). Such a technique appears promising for the treatment of EDTA contaminated wastewaters. This study, therefore, examines EDTA degradation in greater detail, providing optimal parameters for degradation and offering insight into the reaction kinetics and mechanism.
Experimental Section A stock EDTA (J.T. Baker, 99.9%) solution was prepared by dissolving Na2H2EDTA in deionized water (Fisher, HPLC grade, New Jersey). The reaction vessel was a 150 mL roundbottom flask, and the total suspension volumes were 50 mL. BET surface area analysis of the granular zerovalent iron (ZVI) (Aldrich, ACS grade, 20-40 mesh) was performed by Porous Materials, Inc. (Ithaca, NY) and found to be 0.1105 m2/g. The initial concentration of EDTA was set at 1.00 mM (unless noted). A pH of 5.5-6.5 was maintained by the reaction mixture without the use of a buffer. Due to the selfbuffering nature of the reaction products, no attempts were made to regulate the pH of these reaction mixtures. Experiments were conducted at room temperature (20 ( 2 °C) and open to the atmosphere. Reproducible stirring was accomplished using a Bioanalytical Systems Controlled Growth Mercury Electrode Apparatus (West Lafayette, IN). Temperature control was maintained by a custom-made 100 mL water-jacketed vessel with a circulating water bath (Haake GH-D1). The uncertainty in the activation energy measurements was calculated using a propagation of errors method (14). Concentration of FeIIIEDTA was measured based on an HPLC method developed by Nowack et al. (1). The mobile phase consisted of 92% 0.02 M formate buffer (formic acid sodium salt 99%, Acros, New Jersey, and formic acid 89.5%, Fisher, Fairlawn, NJ), and 8% acetonitrile (Fisher, HPLC grade) with 0.001 M tetrabutylammonium bromide (TBA-Br) (Acros, 99+%, New Jersey). The pH was adjusted to 3.5. The HPLC 10.1021/es050137v CCC: $30.25
2005 American Chemical Society Published on Web 08/17/2005
included a 20 µL sample loop, an Alltech (Deerfield, IL) Econosphere C18 column 150 mm in length, 4.6 mm in diameter, 5 µm packing diameter with corresponding guard column. Detection was by means of a UV absorbance detector (Hewlett-Packard 1050 series) at a wavelength of 258 nm. For the control experiments run in the absence of ZVI, the FeIIIEDTA complex was created by adding equal millimolar amounts of Fe(NO3)3 (Fisher, 99+%, New Jersey) to the EDTA sample prior to HPLC analysis. Controls showed that neither EDTA nor Fe3+ salts alone gave UV absorbance at 258 nm. All samples were filtered using 0.22 µm nylon syringe filters (Fisher, New Jersey) prior to analysis. Gaseous headspace analyses were conducted in 15 mL sealed glass vials with Teflon septa. To examine EDTA mineralization to carbon dioxide and other volatile organic compounds, Tenax (Fisher, Fairlawn, NJ) traps were attached to the effluent gas stream from headspace vials. The Tenax traps were then analyzed using a cryogenic trapping system with collection on a HP 5890 GC-MS equipped with a 30 m 5% phenyl column with an inner diameter of 0.25 mm (J&W Scientific). The mass spectra were collected in selective ion mode (SIM). Total organic carbon (TOC) analysis was preformed by Analytical Science Laboratory (University of Idaho, Moscow, Idaho) using EPA method 415.1 on a Tekmar Apollo 9000. Tafel analysis was performed using a Gamry Instruments PC4 potentiostat (Warminster, PA) equipped with a Gamry corrosion cell which included a saturated calomel reference electrode (Gamry) and a ZVI working electrode (99+% pure, Metal Samples, Munford, AL). Polishing was performed on the working electrode between each measurement using 600 grit sandpaper. Direct measurement of the low molecular weight acids (propionic, oxalic, and iminodiacetic acid) were preformed on an Agilent 1100 LC-ESI-MS using an XTerra MS C18 column (part # 18600045) 2.1 × 100 mm with 5 µm packing. The mobile phase used in the low molecular weight analysis was 50% acetonitrile and 50% 0.1% NH4OH and was pH adjusted to pH 9 using HCl. The column flow rate was 0.300 mL/min and the injection volume was 5 µL. ESI-MS detection was conducted in negative ion mode using a cone voltage of 3500 V and a desolvation temperature of 350 °C. All data were collected in both scan and selected ion mode (SIM) using masses 73, 89, and 132. Oxygen concentrations were measured in aqueous solutions using a Vernier dissolved oxygen probe (part #DOBTA) and Vernier Lab Pro Software (Vernier, Beaverton, OR).
Results and Discussion The mechanism by which the ZEA reaction proceeds is hypothesized to involve reactive oxygen intermediates. Intermediates such as hydrogen peroxide are postulated to be continuously produced by the reduction of aqueous oxygen, which may take place either on the iron surface or in solution (13, 15, 16). Many FeII complexes can react with O2 to form the superoxide radical (eq 2) which leads to the production of H2O2 (eq 3) and eventually the Fenton reaction (eq 4) (17-19).
Fe0 f Fe2+ + 2e2+
Fe II
+ L f Fe III
(1)
II
(1b) •-
Fe + O2 f Fe + O2
(2)
FeII + O2•- + 2H+ f FeIII + H2O2
(3)
FeII + H2O2 f FeIII + OH• + OH-
(4)
FIGURE 1. (a) Three identical EDTA degradation runs shown using 9, b, 2 symbols. ZVI ) 2.5 g, [EDTA] ) 1 mM, 450 rpm, air, 50 mL total volume (2). (b) pseudo-first-order plot showing linearity for EDTA degradation from 10 min-2.5 h. (c) 24 h study showing final [EDTA] concentration in 1 µM range.
Waite et al. have shown that humic acids in seawater accelerate the rate of FeIII reduction. They attribute this to the high density of iron-binding carboxylate sites (20). Free iron ions are precluded from taking part in the O2 activation sequence shown in eqs 2-4. Conditions Required for EDTA Destruction. The degradation of EDTA was attempted both in the presence of air and under N2 purge. EDTA degradation is negligible under a N2 purge. In the reaction containing oxygen from air, greater than 95% degradation of 1 mM EDTA was achieved in 2.5 h, Figure 1a. As shown in Figure 1b, pseudo-first-order loss of EDTA was seen after an induction period of 10 min until 2.5 h reaction time. The slope of the line for the first-order plots of natural logarithm versus time was used to obtain the pseudo-first-order rate constants kobs. Pseudo-first-order rate constants of the system purged with N2 and open to air were kobs ) 0.044 ((0.01) h-1 and kobs ) 1.02 ((0.26) h-1, respectively, or kSA ) 0.00796 ( 0.002 L h-1 m-2 and kSA ) VOL. 39, NO. 18, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
7159
FIGURE 2. ZVI maintains EDTA degradation over 12 h after the addition of three aliquots of 1 mM EDTA. All systems mixed at 450 rpm, open to the atmosphere, unbuffered using 2.5 g of 20-mesh ZVI (induction period not shown for clarity). The overall changes in concentration versus time ∆C/∆t are 2.4 × 10-6, 2.3 × 10-6, and 4.5 × 10-6 M/min for the 1st, 2nd, and 3rd runs, respectively. 1.85 ( 0.046 L h-1 m-2 when normalized with respect to the initial Fe(0) surface area. The induction period is attributed to the time needed for the in situ formation of FeIIEDTA complex through the oxidative dissolution of ZVI and the ferrous iron complexation by EDTA as represented in eqs 1 and 1b. The control consisting of a N2 purge in place of air shows no removal of EDTA by volatilization over a 6 h period. Therefore, it is understood that no loss of EDTA can be attributed to volatilization during the time scales of the experiments conducted in this investigation. Control studies done in the absence of oxygen illustrate the major role of dissolved oxygen in the reaction mechanism and further indicate that the mechanism for EDTA removal is not simply adsorption onto the iron surface. Studies similar to those shown in Figure 1a were carried out for 24 h and gave final EDTA concentrations of 1 µM, Figure 1c, demonstrating the ability of the ZEA system to degrade EDTA to low residual concentrations. The robustness of the ZEA system to degrade EDTA over several decay profiles can be seen in Figure 2. Three successive aliquots of 1.00 mM EDTA were added to the same reaction mixture upon significant degradation of the previous EDTA spike and all three were degraded, demonstrating the ability of the system to maintain EDTA degradation over several hours. Compared to Figure 1a, significant deviations from pseudo-first-order kinetics are apparent, particularly for the second and third EDTA additions. However, ratios of ∆C/∆t, the overall change in concentration versus time for the first, second, and third decay profiles are 2.4 × 10-6, 2.3 × 10-6, and 4.5 × 10-6 M/min. These values indicate continuation and perhaps an increase in reactivity over this time period. Deviations from first-order kinetic behavior may be attributed to accumulation of corrosion products on the surface of Fe(0) and the sloughing breakup of Fe(0) and corrosion products from surfaces. Also, the effective surface area of Fe(0) may increase because of the finer particles produced by the breakup of the Fe(0) particles from the corrosion processes, thus increasing overall reaction rates. Production of Radical Intermediates. Evidence for radical activity is supported by the suppression of the reaction by a known hydroxyl radical scavenger, 1-butanol (21, 22). The ZEA reaction is effectively suppressed by the addition of 5 mM 1-butanol (see Supporting Information). This is not exclusive proof that hydroxyl radical species are present in the reaction but corroborates earlier work using a TBARS assay (23, 24) showing evidence of reactive oxygen species present in the reaction mixture (13). 7160
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 39, NO. 18, 2005
TABLE 1. Carbon Balance 1 mM EDTA, 2.5 g of ZVI, Air, Reaction Volume 50 mL, 6 h %C + CO + CO2 iminodiacetic acid oxalic acid propionic acid EDTA total eC3a
35% ((5) 28% ((3) 17% ((2) 14% ((2) 2% ((2) 96% ((5)
a eC3 represents unidentified compounds of C3 or less excluding oxalic and propionic acid.
Products. Analysis of the degradation products by ESIMS show low molecular weight carboxylic acids as the only degradation products in solution with iminodiacetic, propionic, and oxalic acid being the three primary products produced (see Supporting Information). Trapping of the volatile gases released from the system over the course of the reaction using Tenax revealed no volatile organic compounds containing four or more carbons. Total Organic Carbon (TOC) analysis of the postreaction aqueous phase yielded 35% removal of organic material as CO2, CO, and other unidentified C3 and smaller compounds. As seen in Table 1, the observed system (1 mM EDTA, 2.5 g of ZVI, air) has the ability to degrade relatively large amounts of EDTA to benign organic acids and CO2 in 6 h. Reaction Kinetics. All three components of this system, air, ZVI, and aqueous EDTA, are critical for the reaction to proceed as demonstrated by the control studies presented in Figure 1. To enhance the efficiency of the reaction, all parameters and their ratios in the reaction were examined. The most significant experimental variables affecting the kobs are the amount of oxygen in the reaction vessel controlled by the rate of mixing and the ratio of [EDTA]:[Fe2+]. At 450 rpm a steady state dissolved oxygen concentration was achieved (6.7 ppm). Oxygen solubility under these conditions is 8.0 ppm. This stir rate was used in all subsequent studies yielding pseudo-first-order conditions. At the slower mixing speeds the measured oxygen concentrations achieved steadystate conditions; however, they were significantly lower in concentration. The role of ZVI in this system is 2-fold, a source for the reduction of FeIII to FeII, thus redox cycling eqs 2-4, and as a source of Fe2+ for eq 1. It can be expected that the quantity and condition of the metal surface should influence the
FIGURE 3. Pseudo-first-order rate constants for EDTA degradation affected by zerovalent iron (ZVI) mass. All systems mixed at 450 rpm, open to the atmosphere, unbuffered. For a mass of 0.25 g (surface area 0.028 m2) the observed rate is 0.014 h-1 and for a mass of 2.5 g (surface area 0.29 m2) the rate is 1.1 h-1. Error bars are calculated from the standard deviation of the slopes of each corresponding pseudo-first-order decay curve.
FIGURE 4. Pseudo-first-order rate constants for EDTA degradation affected by EDTA concentration. All systems mixed at 450 rpm, open to the atmosphere, unbuffered using 2.5 g of ZVI, 1 mM EDTA, 50 mL total volume. observed rate of EDTA degradation. Figure 3 shows a distinct increase in the kobs with an increase in ZVI mass (surface area) within the system. For a mass of 0.25 g (surface area 0.028 m2) the observed rate is 0.014 h-1 and for a mass of 2.5 g (surface area 0.29 m2) the rate is 1.1 h-1. An analysis of the slope between each data point beyond 0.25 g in Figure 3 yields an average order of 1.75 or approximately 2. This demonstrates that the surface area has an approximate second-order effect with respect to the overall rate. Additional kinetic studies show a marked decrease in kobs with an increase in concentration of EDTA to a point at which the reaction essentially ceases above 10.00 mM, Figure 4. Three mechanisms may explain this observation: I. EDTA may hinder the rate of iron dissolution at high concentrations (eq 1) (25-27). II. The reduction of FeIIIEDTA at the iron surface may be inhibited by the excess EDTA ligand by creating a passive iron oxide EDTA surface. III. ZEA is inhibited by high [EDTA]:[FeII/III] ratios as the result of a kinetic barrier to either O2 activation or H2O2 reduction (eqs 2 and 4).
Hypothetical mechanism II may be of negligible importance since cyclic voltammetric studies of FeII/IIIEDTA at glassy carbon electrodes indicate that the heterogeneous electrontransfer rate of this redox species is not influenced by excess EDTA (4). Also, Figure 5 demonstrates that higher EDTA concentrations increases or does not hinder the rate of iron corrosion. On the other hand, previous investigations add support to mechanism III. It has been previously observed that the rate of reduction of H2O2, i.e., the Fenton reaction (eq 4), is strongly influenced by the ratio of [EDTA]:[FeII]. Higher relative concentrations of EDTA strongly inhibited the Fenton reaction (4, 28). It also is significant to note that EDTA can behave as either an anti-oxidant or a pro-oxidant depending upon the ratios of [EDTA]:[FeII] (4, 29). EDTA Influences the Rate of ZVI Dissolution. To distinguish between postulated mechanisms I and III, the effect of EDTA concentration on the rate of ZVI dissolution was measured. An important factor influencing the rate of dissolution or corrosion of ZVI is the oxide film that forms on the metal surface. Ligands that form bidentate-mononuclear complexes such as oxalate, citrate, and others with adjacent hydroxyl groups often enhance dissolution of metal surface oxides (30). While ligands that form bidentatebinuclear surface complexes (borate and phosphate) inhibit dissolution. As the coordinative environment of the metal oxide changes with pH and with ligand interaction, the dissolution rate is affected. It has been theorized that polarized ligands can weaken and break the metal oxygen bonds in the lattice of the surface oxide layer, allowing for enhanced dissolution. It also has been shown that the addition of EDTA to metal oxide (β-FeOOH) systems can increase dissolution within certain pH regions (31). Electrochemical corrosion measurements were conducted to examine the role of EDTA concentration in ZVI dissolution. Tafel measurements conducted using ZVI in the presence of N2 and air are shown in Figure 5. In both the air- and N2sparged systems, ZVI dissolution increases with increasing EDTA concentrations from 0.1 to 5 mM. Above an EDTA concentration of 5 mM, steady-state rates of 8.9 mm/yr (N2sparged) and 3.9 mm/yr (air-sparged) were observed. The observation that the rate of corrosion is lower in air than N2 is most likely attributable to the passivation layer that forms on the iron surface in the presence of oxygen. The Negative Influence of High [EDTA]:[Fe] Ratios on the ZEA Reaction Is Due to Interferences with O2 Activation and/or Fenton Reaction. The corrosion measurements summarized in Figure 5 indicate that increasing EDTA concentrations increase iron dissolution rates. This runs counter to the observation that higher EDTA concentrations inhibit the ZEA reaction. It is on this basis that mechanism I is rejected. Therefore, the phenomenon of higher EDTA concentrations suppressing the observed degradation rates (Figure 4) is best attributed to mechanism III. Previous investigators have found that the oxidation of FeIIEDTA requires the formation of an O2-FeIIEDTA complex with a coordination number of 7 on the metal center (15-17, 32). This is consistent with the above analysis that the presence of excess EDTA, i.e., a molar ratio greater than 1:1 [EDTA]: [FeII/III], may prevent the formation of H2O2-FeII/IIIEDTA and/ or O2-FeII/ IIIEDTA adducts necessary for the activation of dioxygen or H2O2 (eqs 1-4). Activation Energy. An examination of the rate-limiting step of the ZEA reaction was conducted by measuring the activation energy (Ea). An investigation of the degradation of EDTA at 3.2, 14.2, 22.2, 32.7, and 42.2 °C with all other parameters held constant reveals an increase in the kobs for EDTA degradation with an increase in temperature. The data when plotted using the Arrhenius equation k ) A exp(-Ea/RT), and normalized with respect to the solubility of oxygen in water at the noted temperatures, established VOL. 39, NO. 18, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
7161
FIGURE 5. Tafel corrosion analysis showing corrosion rate of Fe0 in the presence of N2 (b) and air (O), pH 5.5-6.
TABLE 2. Activation Energy Comparison of ZEA System to Literature Values
ref 32 ref 32 ref 32 this study a
Ea (kJ/mol)
[FeEDTA] (mM)
47.6 ((0.6) 33.9 ((1.4) 27.2 ((2.3) 39.3 ((0.6)14
2.5 20 100.0 1.0
[O2] (mM) 0.125 0.125 0.20-0.40a
pH
temp. range
5 5 7.5 5.5-6
13.5-49 °C 16-46 °C 3.2-42 °C
Reaction vessel open to the atmosphere; therefore, oxygen solubility varied with temperature. This was accounted for in Ea calculations.
Literature values for reactions 10 and 11 indicate that they are not rate-limiting relative to reaction 6 (34-36).
FeIIEDTA + O2 f O2-FeIIEDTA
(5)
k1 ) 103 M-1 s-1 k-1) 106 (K1 ) k1/k-1 ) 10-3) (32) O2-FeIIEDTA f FeIIIEDTA + O2-
(6)
k2 ) 102 M-1 s-1 (32) FeIIEDTA + O2- f O22--FeIIIEDTA FIGURE 6. Arrhenius plot demonstrating temperature dependence of observed rate constants (kobs). (2.5 g of ZVI, 1 mM EDTA, 50 mL total volume).
(7)
k3 ) 106 M-1 s-1 (32) O22--FeIIIEDTA + 2H+ f FeIIIEDTA + H2O2
(8)
k4 ) fast (32) the linear behavior shown in Figure 6. An activation energy of Ea ) 39.3 ((0.6) kJ/mol was obtained in this temperature range (14). The activation energy measured can be compared with results obtained by previous investigators measuring the rate of oxidation of FeIIEDTA to FeIIIEDTA by O2 (32, 33). Both studies indicate the rate-limiting step to be the O2 reduction step represented in eq 6 (32). The activation energy measured for this study is in good agreement (Table 2); the discrepancies may be attributed to the different methods and conditions used in measuring the reaction kinetics. The following reaction scheme shows the rate-limiting step as being oxygen activation (eq 6). Equations 5-9 have been theorized as the sequence of reactions important to the reduction of O2 by FeIIEDTA (32). In our studies we hypothesize, in addition to eqs 5-9, eqs 10 and 11, which explain the production of the hydroxyl radical through the Fenton reaction and the subsequent degradation of EDTA. 7162
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 39, NO. 18, 2005
II
2Fe EDTA + H2O2 f 2FeIII + 2H2O
(9)
k5 ) 104 M-1 s-1 (32) FeIIEDTA + H2O2 f FeIIIEDTA + OH- + •OH
(10)
k6 ) 103 M-1 s-1 (36) FeII/IIIEDTA + •OH f LMW acids + CO2
(11)
k7 ) fast low molecular weight (LMW) acids ) propionic, oxalic and iminodiacetic acids Due to EDTA’s low cost and relative effectiveness in emerging green technologies, its industrial usage has increased. The trade-off is the increasing accumulation of EDTA
in the environment which is entirely anthropogenic. Elimination of EDTA from wastewater treatment facilities by our proposed treatment may be an appropriate, cost-effective approach. The system in this study exhibits the desirable characteristics for the treatment of EDTA in wastewaters, i.e., the use of inexpensive reagents and room temperature and atmospheric pressure reaction conditions. Moreover, the EDTA degradation products have far less metal chelation abilities and are subject to faster biodegradation kinetics. The ZEA reaction chemical scheme (eqs 5-11) is a rare example of a nonbiological room temperature and 1 atm pressure O2 activation reaction.
Acknowledgments This investigation was supported by NSF award number BES0328827.
Supporting Information Available Figure S1 demonstrates the suppression of EDTA degradation by the ZEA system with addition of the radical scavenger, 1-butanol. Figure S2 is an electrospray mass spectrum of products of EDTA degradation by the ZEA system. This material is available free of charge via the Internet at http:// pubs.acs.org.
Literature Cited (1) Nowack, B. Environmental Chemistry of Aminopolycarboxylate Chelating Agents. Environ. Sci. Technol. 2002, 36, 4009. (2) Ekland, B.; Bruno, E.; Lithner, G.; Hans, B. Use of Etheylenediaminetetraacetic acid in Pulp Mills and Effects on Metal Mobility and Primary Products. Environ. Toxicol. Chem. 2002, 21, 1040. (3) Sillanpaa, M.; Pirkanniemi, K.; Sorokin, A. Degradative hydrogen peroxide oxidation of chelates catalyzed by metellophthalocyanines. Sci. Total Environ. 2003, 307, 11. (4) Engelmann, M.; Bobier, R.; Hiatt, T.; Cheng, F. Variability of the Fenton reaction Characteristics of EDTA, DTPA, and Citrate complexes of Iron. Biometals 2003, 16, 519. (5) Sillanpaa, M.; Orma, M.; Ramo, J.; Oikair. The importance of ligand speciation in environmental research: a case study. Sci. Total Environ. 2001, 267, 23. (6) Sillanpaa, M.; Pirkanniemi, K. Recent Developments in Chelate Degradation. Environ. Technol. 2001, 22, 791. (7) Kari, F. G.; Giger, W. Modeling the Photochemical Degradation of Ethylenediaminetetraacetate in the River Glatt. Environ. Sci. Technol. 1995, 29, 2814. (8) Nirel, P. et al. Method for EDTA Speciation Deteremination: Application to Sewage Treatment Plant Effluents. Water Res. 1998, 32, 3615. (9) Kari, F. G.; Giger, W. Speciation and fate of EDTA in municipal wastewater treatment. Water Res. 1996, 30, 122. (10) Nortemann, B. Biodegradation of EDTA. Appl. Microbiol. Biotechnol. 1999, 51, 751. (11) Ramo, J.; Sillanpaa, M. Degradation of EDTA by hydrogen peroxide in alkaline condtions. J. Cleaner Prod. 2001, 9, 191. (12) Oviedo, C.; Contreras, D.; Juanita, F.; Rodriquez, J. Degradation of Fe(III)EDTA Complex by Catechol-driven Fenton Reaction. Fresenius Environ. Bull. 2003, 12, 1323. (13) Noradoun, C.; Engelmann, M.; McLaughlin, M.; Hutcheson, R.; Breen, K.; Paszczynski, A.; Cheng, F. Destruction of Chlorinated Phenols by Dioxygen Activation under Aqueous Room Temperature and Pressure Conditions. Ind. Eng. Chem. Res. 2003, 42, 5024. (14) Bevington, P.; Robinson, D. K. Chapter 3: Error Analysis. In Data Reduction and Error Analysis for the Physical Sciences, 3rd ed.; McGraw-Hill: New York, 2003; pp 36-50. (15) Seibig, S.; van Eldik, R. [Multistep Oxidation Kinetics of [FeII(cdta)][cdta ) N,N′,N′′,N′′′-(1,2-Cyclohexanediamine)tetraacetate] with Molecular Oxygen. Eur. J. Inorg. Chem. 1999, 3, 447.
(16) Seibig, S.; van Eldik, R. Kinetics of [FeII(edta)] Oxidation by Molecular Oxygen Revisited. New Evidence for a Multistep Mechanism. Inorg. Chem. 1997, 36, 4115. (17) Bull, C.; McClune, G. J.; Fee, J. A. The Mechanism of Fe-EDTA Catalyzed Superoxide Dismutation. J. Am. Chem. Soc. 1983, 105, 5290. (18) Buettner, G. R.; Doherty, T. P.; Patterson, L. K. The Kinetics of the Reaction of Superoxide Radical with Fe(III) complexes of EDTA, DETAPAC, and HEDTA. FEBS Lett. 1983, 158, 143. (19) Rose, A.; Waite, D. Kinetic Model for Fe(II) Oxidation in Seawater in Absence and Presence in Natural Organic Matter. Environ. Sci. Technol. 2002, 36, 433. (20) Waite, D.; Rose, A. Effect of Dissolved Natural Organic Matter on The Kinetics of Ferrous Iron Oxygenation in Seawater. Environ. Sci. Technol. 2003, 37, 4877. (21) Vassilakis, C.; Pantidou, A.; Psillakis, E.; Kalogerakis, N. Mantzavinos, D; Sonolysis of natural phenolic compounds in aqueous solutions: degradation pathways and biodegradability. Water Res. 2004, 38, 3110. (22) Psillakis, E.; Goula, G.; Kalogerakis, N.; Mantzavinos, D. Degradation of polycyclic aromatic hydrocarbons in aqueous solutions by ultrasonic irradiation. J. Hazard. Mater. 2004, 108, 95. (23) Cheng, Z.; Li, Y.; Chang, W. Kinetic deoxyribose degradation assay and its application in assessing the antioxidant activities of phenolic compounds in a Fenton-type reaction system. Anal. Chim. Acta. 2003, 478, 129. (24) Halliwell, B.; Gutteridge, J.; Aruoma, O. The Deoxyribose Method: A Simple “Test-Tube” Assay for Determination of Rate Constants for Reactions of Hydroxyl Radicals. Anal. Chem. 1987, 165, 215-219. (25) Rubio, J.; Matijevic, E. Interactions of Metal Hydrous Oxides with Chelating Agents: I β-FeOOH-EDTA. J. Colloid Interface Sci. 1979, 68, 408. (26) Chang, H.-C.; Healy, T. W.; Matijevic, E. Interactions of Metal Hydrous Oxides with Chelating Agents: III Adsorption on Spherical Colloidal Hematite Particles. J. Colloid Interface Sci. 1983, 92, 469. (27) Blesa, M. A.; Borghi, E. B.; Maroto, J. G.; Regazzoni, A. E. Adsorption of EDTA and Iron-EDTA Complexes on Magnetite and the Mechanism of Dissolution by EDTA. J. Colloid Interface Sci. 1984, 98, 295. (28) Stadtman, E. R. Oxidation of Free Amino Acids and Amino Acid Residues in Proteins by Radiolysis and Metal-Catalyzed Reactions. Annu. Rev. Biochem. 1993, 62, 797. (29) Feig, A.; Becke, M.; Siegfried, S.; van Eldik, R.; Lippard, S. Mechanistic Studies of the Formation and Decay of Diiron(III) Peroxo Complexes in the Reaction of Diiron(II) Precursors with Dioxygen. Inorg. Chem. 1996, 35, 2590. (30) Stumm, W. Chemistry of the Solid-Water Interface; John Wiley & Sons: New York, 1992; pp 160-167. (31) Rubio, J.; Matijevic, E. Interaction of Metal Hydrous Oxides with Chelating Agents. J. Colloid Interface Sci. 1979, 68, 408. (32) Zang, V.; van Eldik, R. Kinetics and Mechanism of the Autoxidation of Iron(II) Induced through Chelation by EDTA and Related Ligands. Inorg. Chem. 1990, 29, 1705. (33) Wubs, H.; Beenackers, A. A. C. M. Kinetics of the Oxidation of Ferrous Chelate of EDTA and HEDTA in Aqueous Solution. Ing. Eng. Chem. Res. 1993, 32, 2580. (34) Arouma, O. I.; Halliwell, B. The Iron-Binding and Hydroxyl Radical Scavenging Action of Anti-inflammatory Drugs. Xenobiotica 1988, 18, 459. (35) Shi, X.; Dalal, N. S.; Jain, A. C. Antioxidant Behavior of Caffeine: Efficient Scaveging of Hydroxyl Radicals. Food Chem. Toxicol. 1991, 29, 1. (36) Croft, S.; Gilbert, B. C.; Smith, L.; Whitwood, A. C. An ESR Investigation of the Reactive Intermediate Generated in the Reaction Between FeII and H2O2 in Aqueous Solution. Free Radical Res. Commun. 1992, 17, 22.
Received for review January 20, 2005. Revised manuscript received July 18, 2005. Accepted July 18, 2005. ES050137V
VOL. 39, NO. 18, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
7163
