Methods for Accelerating Nitrate Reduction Using Zerovalent Iron at

Westerhoff, P. Reduction of nitrate, bromate, and chlorate by zero valent iron (Fe0). ...... Nitrate in Potable Water Supplies: Alternative Management...
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Environ. Sci. Technol. 2005, 39, 9643-9648

Methods for Accelerating Nitrate Reduction Using Zerovalent Iron at Near-Neutral pH: Effects of H2-Reducing Pretreatment and Copper Deposition Y . H . L I O U , † S . L . L O , † C . J . L I N , * ,† C. Y. HU,† W. H. KUAN,‡ AND S. C. WENG† Research Center for Environmental Pollution Prevention and Control Technology, Graduate Institute of Environmental Engineering, National Taiwan University, Taipei 106, Taiwan, and Department of Environmental and Safety Engineering, Ming-Chi Institute of Technology, Taishan, Taipei hsien 243, Taiwan

Both surface treatments, H2-reducing pretreatment at 400 °C and the deposition of copper as a catalyst, were attempted to enhance the removal of nitrate (40 (mg N) L-1) using zerovalent iron in a HEPES buffered solution at a pH of between 6.5 and 7.5. After the iron surface was pretreated with hydrogen gas, the removal of the passive oxide layers that covered the iron was indicated by the decline in the oxygen fraction (energy dispersive X-ray analysis) and the overlap of the cyclic polarization curves. The reaction rate was doubled, and the lag of the early period disappeared. Then, the deposition of copper onto freshly pretreated iron promoted nitrate degradation more effectively than that onto a nonpretreated iron surface, because of the high dispersion and small size of the copper particles. An optimum of 0.25-0.5% (w/w) Cu/Fe accelerated the rate by more than six times that of the nonpretreated iron. The aged 0.5% (w/w) Cu/Fe with continual dipping in nitrate solution for 20 days completely restored its reactivity by a regeneration process with H2 reduction. Hence, these two iron surface treatments considerably promoted the removal of nitrate from near-neutral water; the reactivity of Cu/Fe was effectively recovered.

Introduction The use of zerovalent iron as a reductant of organohalides (1-5), nitrate (6-10), heavy metals (9, 11, 12), and radioactive elements (13) has been extensively investigated over the last 15 years. The disappearance of contaminants is caused by a corrosion-like process, in which the iron donates electrons to reduce the target pollutants, a process that is accompanied by the dissociation of water. The formation of an adherent oxide layer as a result of the anaerobic corrosion of iron in water negatively affects the ability of the iron to reduce contaminants. Different types of passive films are present on iron, each with a different structure and composition, which depend on the mechanism of formation (14). Previous investigations (3, 15) have demonstrated that the performance * Corresponding author phone: +886-2-23625373; fax:+886-223928821; e-mail: [email protected]. † National Taiwan University. ‡ Ming-Chi Institute of Technology. 10.1021/es048038p CCC: $30.25 Published on Web 11/16/2005

 2005 American Chemical Society

of the zerovalent iron in remediating the chlorinated solvent did not decrease with elapsed time over the several months of the experiments. However, nitrate has been demonstrated to increase iron surface passivation and reduce rates of iron corrosion (16). The pH value controls the reduction of nitrate by iron and affects the formation of the passive oxide layer. Huang et al. (6) reported that iron powder effectively reduced the nitrate only at pH < 4. They also found that the nitrate removal was insignificant except when the solution pH was low enough to dissolve the passive oxide layers. Hence, the iron corrosion rate and the solubility of ferrous hydroxide on the iron surface strongly depended on the pH of the solution (17). The pH control by adding acid solution is not convenient and practical for use in applications. Thus, an enhanced method is required to treat nitrate in near-neutral solutions. The deposition of a small amount of a second metal, such as Pd, Pt, Ag, Ni, and Cu, onto iron has been demonstrated to accelerate the reaction rate (18-26). If iron acts alone as the anode, then the contaminants are reductively degraded by the electrons transferred from the iron. While the iron is deposited with the second metal, a relative potential difference drives the electron from iron to the second metal, before an adsorbed hydrogen ion accepts an electron and is thus converted into adsorbed atomic hydrogen (Hads). The Hads recombination reaction results in the release of hydrogen gas from the second metal surface. Hads, which acts as a reducing agent, has been found to reduce simultaneously various organohalides more quickly than direct electron transfer in bimetallic systems (20, 24, 26-29). Therefore, the nitrate is reduced on the reactive sites of the second metal surface, suggesting that iron is not directly oxidized by pollutants but by water. Although the results obtained using bimetallic systems are encouraging, declining reaction rates over time caused by the loss of loosely bound second metal particles and the formation of a thick oxide layer demonstrate that the benefit of bimetals is short-lived (30). Therefore, the frequency of replacement and the high cost of noble metals result in little practical application without effective methods of regeneration. A few investigators attempted to recover the reactivity of the aged iron using acid washing or ultrasound (24, 31). The aged iron is washed using acidic solutions to clean the surface by dissolving metal and breaking down the passivating oxide layer. The promotion of the reaction in the earlier times is strongest, but the improvement at later times is slight. Moore et al. (31) indicated no significant difference between the reaction rate of the sonicated aged iron and that of the unsonicated aged iron. The sonication procedure did not successfully remove superficial oxides to recover the reactivity of the iron. In this work, both iron surface treatmentssperforming hydrogen treatment at 400 °C for 3 h and depositing copper as a catalystswere attempted to increase the rate of reduction of nitrate at neutral pH (pH ) 6.5-7.5). The aims of this work are to develop an effective reductive material for removing nitrate at neutral pH and an effective surface treatment to activate and restore the reactivity of this material.

Experimental Section Chemicals. Potassium nitrate and sodium nitrite were purchased from Aldrich (99+%, Milwaukee, WI). The chemicals used were Nessler’s reagent (Fluka) and N-[2-hydroxyethyl]piperazine-N′-[2-ethanesulfonic acid] acid (HEPES, Sigma) for ammonia measurement and pH control, respectively. The zerovalent iron used was iron powder (99.6%, VOL. 39, NO. 24, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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electrolytic and finer than 100 mesh) obtained from J. T. Baker. Additionally, the copper precursors, copper(II) chloride, were supplied by Alfa. All aqueous solutions were made in water purified with a Milli-Q system (18.2 MΩ cm-1). The desired concentrations of nitrate, 40 (mg N) L-1, in Ar-purged water were prepared by dilution of a 1000 (mg N) L-1 stock solution, adding the buffer, 40 mM HEPES, to control the pH of the solution in the range 6.5-7.5. Surface Treatment. The iron was heated in a flow of H2/ N2 (20 vol %, 50 mL min-1) from ambient to 400 °C. The temperature was maintained at 400 °C for 3 h to completely reduce the aged oxide layer on the iron surface into zerovalent iron. After the iron was cooled to room temperature, the flow of H2/N2 was then replaced by helium gas (50 mL min-1) to purge the reduced sample for 10 min. The H2-reduced iron must be stored in a dry box. When the loss of reactivity of bimetal particles occurs due to a buildup of iron oxide layers or chemisorption of compounds to reactive sites, the same process was used to recover its activity after drying the bimetallic particles. Bimetals Preparation. Bimetallic particles were prepared by mixing the solution of copper precursor with iron particles, similar to the previous study (18, 19). The deposition of copper onto the surface of iron occurred through the following redox reaction.

Fe0 + Cu2+ f Fe2+ + Cu0

(1)

Initially, the copper precursors were dissolved in Arpurged Milli-Q water at a concentration of 1000 mg L-1. The desired amount of copper precursor (2.5, 5.0, 12.5, 25.0, and 50.0 mL of 1000 mg L-1 copper), associated with the mass ratio of copper to iron (0.05, 0.10, 0.25, 0.50, and 1.00% (w/w)), was added with stirring to 5 g of iron particles. Then, the mixture was washed twice with Milli-Q water. The mixture was filtered and then dried via a vacuum freeze-drying technique (0.2 Torr and -56 °C for 24 h). X-Ray Photoelectron Spectroscopy Analysis. The oxidation states of the surface of iron were identified by electron spectroscopy for chemical analysis. The X-ray photoelectron spectroscopy (XPS) measurements were performed by using a Vacuum Generators ECSALAB MKΠ photoelectron spectrometer (East Grinsted, U. K.) with an ALKR1,2 (1486.6 eV) X-ray source and a hemispherical 150 mm mean radius electron analyzer with a takeoff angle of 90°. The binding energies of the photoelectrons were determined by assuming the energies of the carbon 1s electrons to be 284.5 eV. During the data acquisition, the pressure in the sample chamber did not exceed 5 × 10-10 Torr. For the analysis of multiple peaks in the XPS spectra, the VGX 900 software permitted simultaneous fitting of components with adjustable Gaussian line shape contributions and asymmetries. Scanning Electron Microscopy/Energy Dispersive X-Ray Spectroscopy. The morphology of the resulting bimetallic particles was viewed by scanning electron microscopy (SEM), and localized elemental information from the chosem region with energy dispersive X-ray (EDX) spectroscopy in conjunction with SEM. Electrochemical Analysis. Experiments were performed using a standard three-electrode cell and a microprocessorcontrolled electronic potentiostat (EG&G, Princeton Applied Research, model 273A) with the software Zplot2 and Zview2. The working electrode was a cavity microelectrode (CME), similar to that described by Vivier et al. (32). The cavity was then filled up with pretreated or nonpretreated iron powder using the electrode as a pestle. A Pt wire was used as the counter electrode, and an Ag/AgCl electrode was used as the reference electrode. The cyclic polarization curves were recorded in the potential range from -1.0 to 1.0 V versus -1.0 V at a scan rate of 0.1 mV/s. An Ar-purged buffered 40 9644

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FIGURE 1. Nitrate removal by nonpretreated iron, pretreated iron, and bimetallic particles with various copper loadings in the pH range of 6.5-7.5. The initial nitrate concentration was 40 (mg N) L-1. Error represents a 95% confidence interval. The iron was pretreated at 400 °C in a flow of H2/N2 (20 vol %, 50 mL min-1). (mg N) L-1 nitrate solution was employed to examine the electrochemical corrosion properties of the pretreated and nonpretreated iron powder. All experiments were performed stirring at 400 rpm at 25 ( 0.1 °C. Reactor System. All experiments as a function of time were performed with 65 mL serum bottles. In each bottle, 0.5 g of bimetallic particles and 65 ((0.1) mL of Ar-purged buffered 40 (mg N) L-1 solution were added, leaving no headspace. Immediately, the vials were capped with Teflon silicone septa and aluminum seals and then mixed at 200 rpm using a reciprocal shaker water bath (Yihder, BT-350R) at ambient temperature (25 °C). Sample Analysis. Approximately 1 mL of aqueous solution was collected from the serum bottle by a syringe through the septa, and simultaneously another needle was used to inject argon gas to replace the liquid removal. The target pollutant, nitrate, and the intermediate, nitrite, were measured using an ion chromatograph (Dionex DX-100 TM). Ammonium was analyzed by indophenol method using a spectrometer (UV Spectronic 20 Genesys). Copper ions were measured using an atomic absorption spectrophotometer (PerkinElmer, AAS 800 TM). The pH was measured with a Beckman model 71 pH meter. Because of the possibility of change in the Fe(II)/Fe(III) ratio with time, ferrous ion determination must be at the sampling site. Samples were filtered through a 0.22 µm MillexGS syringe filter before analysis. The procedure consisted of adding 20 mL of phenanthroline solution and 10 mL of NH4C2H3O2 solution to 50 mL of an acidified sample with vigorous stirring. The absorbance at 510 nm was measured using a 5 cm absorption cell.

Results and Discussion H2-Reducing Pretreatment. The reaction rate was evaluated using a nitrate solution of 0.5 g of pretreated iron for comparison with nonpretreated iron (Figure 1). The reduction of nitrate exhibited pseudo-first-order kinetics with respect to the concentration of nitrate

r)

-d[NO3-] ) kobs[NO3-] dt

(2)

where kobs is the observed pseudo-first-order reaction rate constant (min-1). The reduction of nitrate using the nonpretreated iron exhibited stagnation during the first 20 min of the reaction and attenuation at a rate of 0.0093 ( 0.0014

TABLE 1. EDX-Determined Surface Compositions of Various Reductants elemental fraction (wt %) reductants

Fe

Cu

O

nonpretreated Fe pretreated Fe 0.25% Cu/Fe 0.25% Cu/Fe with dipping for 20 days regenerated 0.25% Cu/Feb

95.4 98.1 87.5 86.2 93.4

N.D.a N.D. 10.4 3.3 3.7

4.6 1.9 2.2 10.5 2.9

a Not determined. b After the Cu/Fe was dipped for 20 days, 20% H2/N2 at 400 °C was used to activate the surface of 0.25% Cu/Fe.

FIGURE 2. Sample Tafel profiles for nonpretreated and pretreated iron foils in Ar-purged buffered 40 (mg N) L-1 of nitrate solutions (pH ) 6.5-7.5). min-1. Nonpretreated commercial iron is covered with a discontinuously passive layer, formed during the hightemperature manufacturing process (14). Additionally, a mixture of nonstoichiometric iron oxide and oxyhydroxide species may form in storage (33, 23). Hence, early in the reaction, iron oxides with various valences dominated the reaction rate by restricting the diffusion of nitrate to the reactive sites on the iron surface. Pretreatment methods, including washing in acid (2, 4, 34), treatment with chloride ions (35), and sonication (33), were applied to remove the passive oxide layer before the reaction and thus increase the number of available reactive sites for accelerating the degradation early in the reaction. Subsequently, acid washing and treatment with chloride ions only slightly improved the reaction rate. Evidence demonstrates that numerous fine iron particles were lost (23), and the increases in concentrations of adsorbed H+ and Cl- greatly increased the oxidation rate of iron (4). A large amount of energy was input for sonication pretreatment (33). Accordingly, these pretreatment methods may not be effective for conveniently overcoming the disadvantages of iron. An inexpensive and practical method was employed to pretreat the surface of iron. Heating the iron in flowing H2/N2 gas at 400 °C reduced the passive iron oxides to zerovalent iron. The results of this study revealed that not only was kobs approximately doubled to 0.0214 ( 0.0011 min-1 but also initial lag disappeared. Figure 2 demonstrates that the corrosion potentials (Ecorr) in the forward scan curves are -0.3775 and -0.267 V for the CME with pretreated and nonpretreated iron, respectively. Meanwhile, the Ecorr values in the reverse scan curves are -0.3726 and -0.1987 V, respectively. The shift in Ecorr appeared to be significant for the CME with nonpretreated iron, whereas the CME with pretreated iron yielded highly reproducible Ecorr values. The cathodic region of Figure 2 exhibits no significant difference between the currents associated with the forward and reverse scan curves of the nonpretreated iron, indicating that the predominant halfreaction under this condition is kinetically limited. In contrast, the anodic region of Figure 2 shows a drop in the current from the reverse scan curve of the nonpretreated iron. The cathodic current is constant while the anodic current decreases, yielding the observed shift in Ecorr between the forward and the reverse scan curves of the nonpretreated iron. The shift demonstrates that the oxide film stabilizes as a passivating oxide film in the near-neutral nitrate solution. When the H2-pretreated iron surface is exposed to nitrate solution, Fe2+ tended to pass rapidly from the iron to the solution. Overlaying the forward and reverse scan curves of

the pretreated iron indicates that the oxide film, being in the initial stage of formation, was probably too thin to prevent the corrosion of the iron. Additionally, the EDX analysis (Table 1) shows that H2 pretreatment abruptly reduced the oxygen fraction of nonpretreated iron. This finding indicates that most of the air-formed passive film on the commercial iron was converted to zerovalent iron. The corrosion current densities (icorr) measured from Figure 2 are 5.18 × 10-7 and 10.37 × 10-7 A cm-2 for nonpretreated and pretreated iron, respectively. The larger icorr demonstrates that the pretreated iron is more reactive in the nitrate solution, which is consistent with the increase of the denitrification rate exhibited by the pretreated iron. Furthermore, Figures 3a and 3b present images of the bimetallic 0.25% (w/w) Cu/Fe particles obtained by SEM. The knobs in the figure are individual copper particles, identified by EDX (data not shown). Most copper particles on the surface of the nonpretreated iron had diameters in the range of 750-850 nm (Figure 3a). EDX revealed no significant amount of copper on the iron particles, labeled A and C in Figure 3a. In contrast, the smaller diameter of approximately 150 nm and the pretreated iron surface exhibited more uniformly distributed copper particles, as shown in Figure 3b. This result reveals that the pretreated iron presented a more uniform and reactive surface to the reducible species. Hence both the destruction of the passive film and the increase in the reactive site concentration on the iron surface increased the rate of denitrification and the disappearance of the initial lag. Deposition of the Second Metal. Three noble metals, Pd, Pt, and Cu, were deposited onto an H2-pretreated iron surface to test their reactivity in nitrate reduction. The atomic ratio was considered for comparing the catalytic reduction rate of these three noble metals. Approximately 0.5 g of 0.44% (atomic ratio) bimetallic particles (Pd/Fe, Pt/Fe, and Cu/Fe) was added into a 65 mL serum bottle containing Ar-purged buffered 40 (mg N) L-1 nitrate solution. In the Pd/Fe and Pt/Fe systems, the kobs values were 0.0242 ( 0.0021 min-1 for Pd/Fe and 0.0208 ( 0.0011 min-1 for Pt/Fe, resembling that of the system with only pretreated iron (0.0214 ( 0.0011 min-1). When copper was deposited onto the pretreated iron surface, the nitrate reduction rate was tripled (0.0664 ( 0.0036 min-1). Therefore, copper was chosen as a second metal for further experiments. Various bulk loadings (0.05, 0.10, 0.25, 0.50, and 1.00% (w/w)) of copper were used to evaluate the rate of nitrate reduction in this experiment. Figure 1 reveals that the reaction rates were increased with the copper loading up to 0.25% (w/w). The kobs values were from 0.0367 ( 0.0037 to 0.0656 ( 0.0043 min-1, as presented in Figure 4. The maximum kobs was in the copper loading range from 0.25 to 0.50% (w/w), and the concentration of nitrate was reduced after 45 min below the detection limit of the detector, which was 0.2 (mg N) L-1. When the copper loadings exceeded 0.50% (w/w), the rate dropped. Therefore, further increasing the copper VOL. 39, NO. 24, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 4. Reaction rate constants for both the fresh and the regenerated reductive metals. Error represents a 95% confidence interval

TABLE 2. Average Generation Rate of Ferrous Ion in Water and Nitrate Solutions average ferrous ion generation rate (mM h-1)a

metallic iron 0.25% Cu/Fe 0.50% Cu/Fe 1.00% Cu/Fe a

Ar-purged Milli-Q water

Ar-purged Milli-Q water with 40 (mg N) L-1 nitrate

0.28 0.59 1.69 1.87

12.31 22.75 23.29 15.41

Average results from duplicate tests were reported.

When nitrate was added to the solution, nitrate may be adsorbed onto the Cu surface and rapidly reduced to ammonia by neighboring Hads. The proposed mechanisms are as follows

NO3-aq T Cu-NO3-ads (in equilibrium)

(6)

Cu-NO3-ads + 6Cu-Hads + 2e- f NH3 + 3OH- (7)

FIGURE 3. SEM images of bimetallic particles with various copper loadings: (a) 0.25% copper loading on nonpretreated iron particles, (b) 0.25% copper loading on pretreated iron particles, and (c) 1.00% copper loading on pretreated iron particles. loadings above 0.50% (w/w) did not further increase the removal rate. The data in Figure 4 almost yield a volcano curve. When contaminant-free water is anaerobic and neutral, the reactions between water and Cu/Fe can be divided into anodic and cathodic reactions: anodic reaction

Fe f Fe2+ + 2e-

(3)

H2O + e- f Cu-Hads + OH-

(4)

Cu-Hads + Cu-Hads f H2(g)

(5)

cathodic reaction

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Therefore, both water and nitrate oxidize Fe0 into Fe2+, which enters solution. The rate of production of ferrous iron, the initial iron corrosion rate, exhibits a significant difference between contaminant-free water and 40 (mg N) L-1 nitrate solution. Table 2 shows that, in a contaminant-free water system, the average rate of generation of ferrous ions using iron alone is lower than those of all Cu/Fe systems. Moreover, the rate of generation increased with copper loading up to 1.0% (w/w), indicating that the iron corrosion rate was cathodically controlled under these conditions. After nitrate was added, the iron corrosion reaction was accelerated by more than 1 order of magnitude in all of the tested systems. In 40 (mg N) L-1 nitrate solution, 0.25% (w/w) Cu/Fe almost doubled the average generation rate of ferrous ions above that of iron alone. However, further increasing the copper loading to 0.50% (w/w) did not substantially further increase the average generation rate of ferrous ions. In contrast, the average ferrous ion generation rate declined when the bulk loading of copper was 1.0% (w/w). This result demonstrated that the iron corrosion rate in nitrate solution was anodically controlled when the copper loading exceeded 0.5% (w/w). Figures 3b and 3c present the SEM images of 0.25% (w/w) Cu/Fe and 1.0% (w/w) Cu/Fe, respectively. Copper particles were densely deposited on the surface of 1.0% (w/w) Cu/Fe particles. Thus one process that may account for these results

TABLE 3. Characteristics of the Iron Surfaces Using XPS Analysis sample

state

binding energy (eV)

concentration (%)

0.5% Cu/Fe in NO3- solution

Fe(II) Fe(III) Fe(II) Fe(III) Fe(II) Fe(III)

709.5 711.5 709.7 711.4 710.5 711.5

44.8 42.9 32.5 37.8 18.3 54.8

0.5% Cu/Fe in water Fe in NO3- solution

is transport of Fe2+ out of the Cu/Fe cavity. The iron dissolution may reach a maximum and then decline to a small value, in a manner similar to the variation in kobs with the bulk loading of copper. Consequently, the left and right “legs” of the volcano curve do not suffice, and the copper that covers the iron surface is in excess (Figure 4). Regeneration. The loss of reactivity of a bimetal over time, caused by a buildup of corrosion products or other precipitates on its surface, is of great concern. The removal of a passive oxide layer using acid washing and sonication to restore the reactivity of metallic iron has been reported. Gui et al. (24) noted that the acid solution (H2SO4) used in Ni/Fe regeneration removed some of the covered corrosion products, making iron and nickel more accessible to the N-nitrosodimethylamine (NDMA) molecules in the solution. However, the loss of recoverable reactive sites on Ni was caused by the regeneration in acid solution. Sonication to regenerate surface reactivity and the removal of superficial deposits did not successfully remove oxide layers, and so the reaction rate did not differ from that of the nonsonicated samples (31). This study employed a mixture of hydrogen and nitrogen gas in a ratio of 1:4 to flush the aged bimetal surface at 400 °C in a closed oven. This process was similar to the pretreatment method elucidated above. The regeneration experiment was conducted at 0.25, 0.50, and 1.00% (w/w) Cu/Fe with continual dipping in nitrate solution over 20 days. Before the H2 regeneration, the reactivity of the aged Cu/Fe decreased by 70-80% (0.017 ( 0.002 min-1 for 0.25% (w/w) Cu/Fe, 0.021 ( 0.002 min-1 for 0.5% (w/w) Cu/Fe, and 0.010 ( 0.001 min-1 for 1.0% (w/w) Cu/Fe) below that of fresh Cu/Fe. After regeneration, the reaction rate constants increased to 0.041 ( 0.004, 0.066 ( 0.008, and 0.038 ( 0.004 min-1 for 0.25, 0.50, and 1.00% (w/w), respectively. The reactivity was completely restored for both 0.50 and 1.00% (w/w) Cu/Fe, whereas approximately 60% of the reactivity was restored for 0.25% (w/w) Cu/Fe (Figure 4). The elemental fractions (Fe, Cu, and O) of the surface compositions of the fresh, aged, and regenerated 0.25% (w/w) Cu/Fe were determined by EDX (Table 1). After the Cu/Fe was dipped for 20 days, the O fraction significantly increased relative to that of the fresh 0.25% (w/w) Cu/Fe, indicating the formation of oxide iron on the surface of iron. When the aged 0.25% (w/w) Cu/Fe was treated by H2 regeneration, the oxide iron was completely converted into zerovalent iron, as indicated by the low O fraction. However, part of the copper particles dropped from the surface of the iron during the reaction in which the Cu fraction after the reaction was less than that of the fresh 0.25% (w/w) Cu/Fe. In the cathodically controlled reaction, a lower exposure area of copper to the solution corresponded to a lower reactivity, as compared to that of the nitrate reduction. Thus a drop in reactivity of about 40% for 0.25% (w/w) Cu/Fe was observed after H2 regeneration. But at the higher loadings, 0.5 and 1.0% (w/w) Cu/Fe, the necessarily reactive copper sites were still enough, even though the exposed area of copper to the solution was decreased. Reaction Intermediates and Products. In reactions with metallic iron, trace amounts of nitrite indeed accumulated as an intermediate in the system at a final pH of 7.5-9.0 (36,

Fe(II)/Fe(III) ratio 1 1 1/3

37). In this case of denitrification by Cu/Fe at near-neutral pH, no detectable amount of nitrite (the detection limit was 0.1 (mg N) L-1) was observed, and ammonia was the only final product, indicating that copper as a catalyst not only rapidly degrades nitrate but also completely transforms nitrate into ammonia. However, the experiments of Cu/Fe with continual dipping for 20 days revealed not only that the reaction rate had fell but also that a small amount of nitrite was detected (0.5 (mg N) L-1 at maximum). Less active and inactive catalysts, resulting from a buildup of adherent oxide layers, cause the formation of toxic intermediates. After the reduction of nitrate, the oxidation states of iron on the surface were examined by XPS, which is an ex situ method. Inevitably, the oxidation states of the surface of iron change when the sample is removed from the reaction vessel to the ultrahigh vacuum chamber of the XPS instrument. In this study, the iron sample was dried by a vacuum freeze-drying technique (0.2 Torr and -55 °C for 24 h) and then was closed off by an air-proof membrane before transfer to the ultrahigh vacuum chamber of the XPS instrument. The results (Table 3) show that the Fe(II)/Fe(III) ratio was about 1 on the surface of bimetallic Cu/Fe but only approximately 1/3 for zerovalent iron. Consequently, the fraction of Fe(II) on the Cu/Fe surface exceeded that on the iron surface. The significant difference between the Fe(II)/ Fe(III) ratios may be caused by the fact that Cu-Hads rather than direct electron transfer on the iron surface dominated the reduction of nitrate, so the anodic reaction was similar to that in water. Bimetallic Cu/Fe is a potentially reductive material for nitrate reduction at near-neutral pH, at which nitrate was converted completely into ammonia and no nitrite was detected. Heating the aged Cu/Fe in flowing H2/N2 at 400 °C for 3 h almost completely restored its reactivity. Both surface treatments facilitated the development of a process that could effectively remove nitrate without acidification and conveniently regenerate the reductive medium.

Acknowledgments The authors thank the National Science Council of the Republic of China for financially supporting this research under Contract No. NSC 93-2211-E-002-035. We thank Miss Chao-Ling Lai from the Department of Chemical Engineering (NTU) for her help with the XPS analysis.

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Received for review December 12, 2004. Revised manuscript received July 20, 2005. Accepted September 2, 2005. ES048038P