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Journal of Environmental Quality 32:1306-1315 (2003)
© 2003 American Society of Agronomy, Crop Science Society of America, and Soil Science Society of America

TECHNICAL REPORTS
Ground Water Quality

Effects of Oxide Coating and Selected Cations on Nitrate Reduction by Iron Metal

Yong H. Huanga, Tian C. Zhang*,a, Patrick J. Sheab and Steve D. Comfortb

a Civil Engineering Dep., Univ. of Nebraska-Lincoln, Omaha Campus, Omaha, NE 68182-0178
b School of Natural Resource Sciences, 309 Biochemistry, University of Nebraska-Lincoln, Lincoln, NE 68583-0728

* Corresponding author (tzhang{at}unomaha.edu)

Received for publication June 11, 2002.

   ABSTRACT
TOP
 ABSTRACT
INTRODUCTION
 MATERIALS AND METHODS
 RESULTS
 DISCUSSION
 REFERENCES
 
Under anoxic conditions, zerovalent iron (Fe0) reduces nitrate to ammonium and magnetite (Fe3O4) is produced at near-neutral pH. Nitrate removal was most rapid at low pH (2–4); however, the formation of a black oxide film at pH 5 to 8 temporarily halted or slowed the reaction unless the system was augmented with Fe2+, Cu2+, or Al3+. Bathing the corroding Fe0 in a Fe2+ solution greatly enhanced nitrate reduction at near-neutral pH and coincided with the formation of a black precipitate. X-ray diffractometry and scanning electron microscopy confirmed that both the black precipitate and black oxide coating on the iron surface were magnetite. In this system, ferrous iron was determined to be a partial contributor to nitrate removal, but nitrate reduction was not observed in the absence of Fe0. Nitrate removal was also enhanced by augmenting the Fe0–H2O system with Fe3+, Cu2+, or Al3+ but not Ca2+, Mg2+, or Zn2+. Our research indicates that a magnetite coating is not a hindrance to nitrate reduction by Fe0, provided sufficient aqueous Fe2+ is present in the system.


   INTRODUCTION
TOP
 ABSTRACT
 INTRODUCTION
MATERIALS AND METHODS
 RESULTS
 DISCUSSION
 REFERENCES
 
THE WIDESPREAD occurrence of nitrate-contaminated ground water and the expense of mitigating this problem are major concerns for many communities throughout the USA. Although several treatment processes have been developed for nitrate removal, in situ remediation of nitrate-contaminated ground water remains formidable. Recent research indicates many potential uses of zerovalent iron (Fe0) in environmental remediation (Stucki, 1988; Matheson and Tratnyek, 1994; Weber, 1996; Scherer et al., 1997, 1998, 2000; Till et al., 1998; Singh et al., 1998). Using Fe0 to remove nitrate from contaminated waters would be an attractive approach provided it is effective at environmentally relevant pH values.

Past research has demonstrated that the efficiency of nitrate removal by iron metal is largely dependent on pH (Siantar et al., 1995; Singh et al., 1996; Cheng et al., 1997; Hansen and Koch, 1998; Huang et al., 1998; Zawaideh and Zhang, 1998; Alowitz and Scherer, 2002). In acidic solutions (pH < 2–3), nitrate removal in Fe0–H2O systems is fast and efficient (>=95%) (Singh et al., 1996; Huang et al., 1998; Zawaideh and Zhang, 1998). However, when solution pH increases above 5, nitrate removal efficiencies decline and are usually <50% (Singh et al., 1996; Cheng et al., 1997; Huang et al., 1998; Zawaideh and Zhang, 1998) unless organic buffers, such as acetic acid, sodium acetate, HEPES [N-(2-hydroxyethyl)piperazine-N'-(2-ethanesulfonic acid)], or MES [2-(N-morpholino)ethanesulfonic acid], are used (Cheng et al., 1997; Zawaideh and Zhang, 1998; Alowitz and Scherer, 2002). Because near-neutral pH will be encountered in many soil–water environments, enhancing nitrate reduction by Fe0 without an organic buffer is desirable.

Nitrate reduction by Fe2+, Fe(OH)2, and hydrous ferrous oxides has been investigated (Sidgwick, 1950; Szabo and Bartha, 1952; Bremner and Shaw, 1955; Brown and Drury, 1967; Bremner and Bundy, 1973; Buresh and Moraghan, 1976; Van Hecke et al., 1990). These studies indicate that at alkaline pH, Fe(OH)2 rather than free Fe2+ is the species responsible for nitrate reduction to ammonia in the presence of Cu2+, Ag+, or MgO (Sidgwick, 1950; Szabo and Bartha, 1952; Bremner and Bundy, 1973; Buresh and Moraghan, 1976) and that magnetite is the product of the reaction (Buresh and Moraghan, 1976; Van Hecke et al., 1990). Nitrate reduction has also been associated with the oxidation of green rust [mixed Fe(II)–Fe(III) double hydroxides], which can form when Cl-, SO2-4, or CO2-3 is present during Fe0 corrosion (Hansen and Koch, 1998).

Favorable empirical observations have shifted more attention to mechanisms of contaminant destruction by Fe0. Klausen et al. (1995) proposed that Fe2+ adsorbed on iron oxide surfaces (including magnetite, goethite, and lepidocrocite) or mixed-oxide films plays a key role in the reductive transformation of organic pollutants. Most contaminant destruction is believed to result from interaction with Fe0 or surface Fe2+. Once exposed to an oxidizing environment, however, all Fe0 surfaces become coated with (hydr)oxides that will continue to form and oxidize further with time. To account for this, Scherer et al. (1998) proposed models depicting the iron oxides as passive films, semiconductors, and coordinating surfaces. Surface-coordinated Fe2+ species are essential to the coordinating-surface model and can originate from a variety of sources, such as reductive dissolution of Fe3+ phases (Klausen et al., 1995), "structural" Fe2+ ions within the oxide lattice (Stucki, 1988), or freshly precipitated mixed Fe(II)–Fe(III) oxide or hydroxide coatings (magnetite or green rust). According to Stumm (1992), at pH of >=7.0, inner-sphere complexation of Fe2+ to metal oxides can create a stronger reductant. Hydrolysis of Fe2+ to FeOH+ can also increase reducing power (Schultz and Grundl, 2000). Thus, the nature of the iron (hydr)oxides and their relationships to Fe0 and Fe2+ are critical to understanding mechanisms of nitrate reduction by Fe0.

Our objective was to determine nitrate reduction by Fe0 under anoxic conditions leading to magnetite formation. This study focuses on the oxide coating forming on granular iron and enhanced nitrate removal with iron by adding selected cations in bulk solution at varying pH. We also discuss possible mechanism(s) and the environmental significance of the treatment process.


   MATERIALS AND METHODS
TOP
 ABSTRACT
 INTRODUCTION
 MATERIALS AND METHODS
RESULTS
 DISCUSSION
 REFERENCES
 
Chemicals
Unless otherwise indicated, all aqueous solutions were prepared with ultrapure water (Nanopure Series 550; Barnstead/Thermolyne Co., Dubuque, IA). The water is deionized water of ultrahigh quality with resistivity up to 18.3 megohm-cm. All commercially available chemicals and minerals were used as received.

The industrial iron powder consisted of filings and shavings that were largely free from visible rust and retained a metallic glaze (US Metals Co., Chicago, IL). The iron particles were approximately 0.5 mm in diameter, irregular in shape, with a slightly rough surface and a Brunauer–Emmett–Teller (BET) surface area of 0.04 m2 g-1.

Ferrous iron (Fe2+) was prepared from FeCl2·4H2O (J.T. Baker Co., Phillipsburg, NJ) and NaNO3 (Baker) was used for preparing the NO-3 stock. The selected cations (Fe3+, Cu2+, Al3+, Ca2+, Mg2+, Zn2+) were used in the following salt forms: FeCl3·6H2O (Fisher Scientific, Pittsburgh, PA), CuCl2·2H2O (Fisher), Al2(SO4)3·15H2O (Fisher), CaCl2·2H2O (Mallinckrodt Chemical, Paris, KY), MgSO4·7H2O (Baker), and ZnSO4·7H2O (EM Science, Darmstadt, Germany). Ferric–ferrous oxide black (Fisher) was the source of magnetite (Fe3O4) powder used.

In some batch tests, the iron powder was precoated with magnetite. To accomplish this, industrial iron powder (5% w/v) was added to a 150-mL serum bottle filled with 100 mL acidified (pH of approximately 2.3) and deoxygenated NaNO3 solution (30 mg NO-3–N L-1). The serum bottle was purged with argon gas for 10 min and shaken for 10 h in an opaque box rotating at 30 rpm to form a magnetite coating on the Fe0. We then removed the bulk solution from the bottle, rinsed the iron powder with deoxygenated water, and dried the powder by purging with argon gas.

Experiment Methods
All batch tests employed 10-mL serum bottles (VWR Int., West Chester, PA) with rubber stoppers as batch reactors. Initial test conditions are listed in Table 1 (or otherwise specified). In each test, multiple reactors were prepared using the following procedures:

  1. Chemicals (NaNO3 and FeCl2·4H2O) were dissolved in ultrapure water and adjusted to the desired pH using 2 M HCl or 2 M NaOH.
  2. Solutions were purged with argon gas for 15 min to eliminate dissolved oxygen and 10 mL was transferred to reactors containing the solid reactants (iron powder and/or magnetite powder).
  3. The reactors were immediately capped with stoppers and the headspace was flushed with argon for 30 s by inserting (two) needles through the stopper.
  4. The reactors were placed in a 30- x 45-cm box rotating at 30 rpm to provide complete mixing in the dark.
  5. At selected times, one reactor was sacrificed for analyses.


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  		Table 1. Experimental conditions and results for Batch Tests 1 through 16.{dagger}

 
In these experiments, only the initial conditions in the reactors were controlled. Ionic strengths ranged from 2.14 mM (Test 1) to 12.8 mM (Test 8). All tests and analyses were conducted at room temperature (24 ± 1°C).

Analytical Methods
Nitrate N, nitrite N, pH, oxidation–reduction potential (ORP), Fe2+, and Fe3+ were measured at regular intervals. A DX 500 high performance liquid chromatography (HPLC)–ion chromatography (IC) system (Dionex Corp., Sunnyvale, CA) was used to analyze NO-3, NO-2, NH+4, Fe2+, Fe3+, and Cu2+. When measuring NO-3 and NO-2, a conductivity detector (CD20) and self-regenerating suppressor (SRS) (ASRS-ULTRA 4-mm) with a 100-mA current were used with an IonPac AG14 4-mm precolumn and separation column. The flow rate of the eluent (2.7 mM Na2CO3 + 1.0 mM NaHCO3) was 1.2 mL min-1. When measuring Fe2+, Fe3+, and Cu2+, a UV-VIS detector (AD20) was used with an IonPac CG5A 4-mm precolumn and separation column. The flow rate of the eluent (MetPac PDCA; Dionex) was 1.2 mL min-1, while the flow rate of the MetPac PAC postcolumn reagent was 0.6 mL min-1. When measuring NH+4, a conductivity detector (CD20) and SRS (CSRS-II 4-mm) with a 100-mA current were used. An IonPac CG12A 4- x 50-mm precolumn and 4- x 250-mm separation column were used with 22 mM H2SO4 eluent at a flow rate of 1.0 mL min-1. The ORP was measured by a microredox electrode (MI-800-407) together with a reference electrode (MI-401; Microelectrodes, Bedford, NH). A semimicro pH probe was used for pH measurements (Thermo Orion, Beverly, MA). Scanning electron microscopy (SEM) (S-3000 N; Hitachi, Tokyo, Japan) was used to discern the features of the source iron and the oxide coating(s) formed under the test conditions.

X-ray diffraction (XRD) was used to determine the iron oxide composition (PAD V model equipped with a peltier-cooled, solid-state detector; Scintag, Cupertino, CA). The operating wavelength was 1.5405 µm (Cu source). The library used for peak identification was from the JCPDS International Center for Diffraction Data (PCPDFWIN v. 2.01). To form a thin film, the commercial magnetite powder was sprayed onto a piece of glass. To analyze the black precipitate formed in our reaction flasks, the suspension was filtered through a 0.45-µm filter paper.

To ascertain the nature of the oxide film forming on the granular iron, a piece of pure iron foil (25 x 25 x 2 mm, 99.9985% certified; Alfa Aesar, Ward Hill, MA) was used because the granular Fe0 powder was not fine enough to provide a smooth surface for X-ray diffraction analysis. The iron foil was placed in a jar containing 100 mL of solution with 100 mg Fe2+ L-1 and 30 mg NO-3–N L-1 and held under anoxic conditions. After 24 h of exposure, a black coating covered the iron foil and nitrate loss was similar to that observed when granular Fe0 was used. The coated foil was dried by piercing a needle through the cap and flushing the reactor with ultrapure argon gas (Linweld, Lincoln, NE). Once a positive pressure was created inside the reactor, a second needle was pierced through the cap of the reactor. The second needle was used to decant the liquid from the reactor. The oxide-coated iron foil was then rinsed by introducing deoxygenated water into the reactor via the second needle while argon gas was still flushing via the first needle. After the rinse, the second needle served as a vent for flushing with argon gas to dry the sample (about 1 h). The dry sample was preserved against oxidation under argon until analysis.


   RESULTS
TOP
 ABSTRACT
 INTRODUCTION
 MATERIALS AND METHODS
 RESULTS
DISCUSSION
 REFERENCES
 
Effects of Initial pH
When solution pH was unbuffered (pH 7.1; Test 1 in Table 1), use of Fe0 alone resulted in <10% NO-3–N removal within 24 h. The bulk solution pH increased from 7.1 to 9.4 during the reaction and no Fe2+ or Fe3+ was detected in solution (Fig. 1a) . After 24 h, the iron surface was coated with a black iron oxide and no additional nitrate removal was observed.



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  		Fig. 1. Results of batch tests under differing experimental conditions (defined in Table 1). (a) Test 1; (b,c) Test 2 (note that nitrate reduction occurred in three stages, TN = NH+4+ NO-3); (d) Test 4; (e) Test 12; and (f) Test 13. Values and error bars in (a), (b), and (c) represent the average and plus or minus one standard deviation, respectively, from three replicate runs.

 
When the initial pH was lowered to 2.3 (Test 2 in Table 1), nitrate reduction occurred in three stages (Fig. 1b,c). In Stage 1 (the first 30 min), nitrate concentration rapidly decreased (>5 mg NO-3–N L-1 removed within 0.5 h) while the pH sharply increased from 2.3 to 4.8 as a result of the reaction between Fe0 and H+ (Fe0 + 2H+ -> Fe2+ + H2{uparrow}). Small bubbles, presumably H2, were observed on the iron surface and about 105 mg Fe2+ L-1 was detected in solution. Stage 2 was characterized by a slower rate of nitrate removal. During this stage (0.5–10 h), the pH increased from 4.8 to 6.2 and corresponded to a decrease in Fe2+ (105 to about 75 mg L-1). A black coating on the Fe0 surface was visible at about 5 h. A black precipitate, arising from the solution phase or detachment of the black coating from the Fe0 surface, became visible at about 10 h. The X-ray diffraction analysis (see below) identified both the black precipitate and the black coating as magnetite (Fe3O4). The appearance of the black precipitate in solution denoted the beginning of Stage 3, which was characterized by rapid loss of nitrate from solution and a concurrent disappearance of aqueous Fe2+. The remaining nitrate (approximately 67%) was removed within 20 h at near-neutral pH during Stage 3. During the entire test, NH+4 was the only nitrogen product detected and it resulted in a near complete (100%) nitrogen mass balance (Fig. 1b).

Screening Tests
Stage 3 indicated that nitrate could be rapidly reduced at neutral pH but not until the iron became coated with magnetite and Fe2+ was present in the bulk solution. In addition to pH, three test parameters were varied during a series of 16 batch experiments (hereinafter referred to as Tests 1–16). In these tests, no additional chemicals were added to control the ionic strength in each batch reactor. The initial ionic strengths, calculated based on the test conditions shown in Table 1, were 2.14 mM for Tests 1, 3, 6, 7, and 11; 7.14 mM for Tests 2, 4, and 5; 5.4 mM for Tests 15 and 16; 7.54 mM for Tests 9, 10, and 12 through 14; and 12.54 mM for Test 8. These tests were designed to determine the effects of the magnetite coating on the Fe0 surface, the addition of magnetite as a separate phase in the batch reactor, and the presence of aqueous Fe2+ during corrosion of the Fe0. To determine whether Fe2+ in solution or Fe2+ sorbed onto magnetite would reduce nitrate in a magnetite–nitrate–Fe2+–water system, tests with batch reactors containing 1% (w/v) commercial magnetite, approximately 90 mg Fe2+ L-1, and 30 mg NO-3–N L-1 were conducted under anoxic conditions and a range of pH values (2–10) that crossed the zero point of charge (ZPC) of the iron oxides (see Table 2) . Although green rust can facilitate rapid nitrate removal at neutral pH (Hansen and Koch, 1998), green rust was not visually observed before or after the formation of the black precipitate in these experiments.


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  		Table 2. Effect of Fe2+ on nitrate removal in magnetite–nitrate–water batch system.{dagger}

 
When the initial pH was 8.6 (Test 3 in Table 1), adding magnetite did not enhance nitrate reduction. When the pH was lowered to 2.3 and magnetite was present (Test 4), nitrate was completely removed from the reactor within 12 h (Fig. 1d). The removal rate was similar to that observed in Stage 3 of Test 2 (Fig. 1b). Moreover, addition of the magnetite eliminated the stage with moderate nitrate removal (Stage 2, Fig. 1b).

No nitrate removal was detected in the absence of Fe0 (Tests 5–10 in Tables 1 and 2), indicating that neither magnetite (Tests 5–7) nor magnetite combined with Fe2+ (Tests 8–10 in Tables 1 and 2) was responsible for the rapid loss of nitrate from solution during Stage 3 (of Test 2). When Fe0 was precoated with magnetite without an initial pH adjustment, no nitrate removal was observed (Test 11), indicating that the black coating on the surface of granular Fe0 was not directly responsible for rapid nitrate reduction. When magnetite-coated Fe0 (Test 12) or noncoated Fe0 (Test 13) was used with Fe2+ at an initial pH of 5.4, almost all of the initial 30 mg L-1 nitrate was removed within 24 h and accompanied by the formation of a black precipitate. Moreover, the noncoated Fe0 (Test 13) was covered with a black oxide within the first 2 h. Results from the magnetite-coated and noncoated Fe0 were nearly identical (Fig. 1e,f), indicating that the magnetite coating was not a hindrance to nitrate removal as long as Fe2+ was available. This was further supported by Test 14, in which 100 mg Fe2+ L-1 was added to magnetite-coated Fe0 and then the pH was raised to >=8.5 to precipitate the aqueous Fe2+ as Fe(OH)2. Results indicated no dissolved Fe2+ in solution and only 10% nitrate removal in 32 h.

Results from the control tests (no NO-3 added) demonstrated the important role nitrate plays as an electron acceptor in the anoxic Fe0–H2O system. In these tests, the added Fe2+ remained dissolved and stable for a comparable test period with no pH change (Test 15) and no black oxide coating was observed on the bare Fe0 surface (Test 16). This is in sharp contrast to the results of Tests 12 and 13, where adding nitrate resulted in aqueous Fe2+ disappearance from solution and formation of a black oxide coating.

X-Ray Diffraction and Scanning Electron Microscopy Analysis
The X-ray diffraction analysis of the commercial magnetite powder (Fig. 2a) matched the magnetite spectrum from the X-ray diffraction library (peaks at 18, 30, 35.5, 37, 43, 53, 57, 62.5, and 74 degree in 2{theta}, Cu source). In addition to the magnetite peaks, the diffractograms from the black oxide coating (Fig. 2b) and the black precipitate (Fig. 2c) included peaks at 44.6 and 65 degrees, which were identified by the library as metallic iron. The black precipitate also contains some background peaks (e.g., peak at 52 degrees) belonging to the filter paper (Fig. 2c). A comparison of all diffractograms (Fig. 2) provides evidence that both the black precipitate and the black coating were magnetite. While the commercial magnetite powder, the black coating, and magnetite precipitate appeared equivalent in composition and morphology, some differences are likely (Sidhu et al., 1978). The black coating that initially formed on Fe0 (Test 13) may be an amorphous (hydr)oxide, which transforms to magnetite in the presence of Fe(II) at alkaline pH (Tamaura et al., 1983).



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  		Fig. 2. X-ray diffraction spectra analyses indicate that the black oxide coating and black precipitate generated from nitrate reduction by Fe0 under the anoxic condition are magnetite. (a) Commercial magnetite powder; (b) black coating on Fe0; and (c) black precipitate on filter paper.

 
The morphological features of the black coating and black precipitate were evaluated by scanning electron microscopy (Fig. 3) . The similarity between the black coating and black precipitate suggests that the black precipitate originated from the black coating or precipitated directly from solution.



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  		Fig. 3. Scanning electron microscope (SEM) image of (a) granular Fe0 before use, bare surface (x60); (b) granular iron coated with magnetite after 20 h in the reactor with the initial conditions of Test 12 (x18k); and (c) black precipitate on a filter paper, obtained after 10 h from the reactor with the initial conditions of Test 13 (x20k).

 
Ferrous Iron Consumption and Nitrate Reduction to Ammonium
Because bathing the corroding Fe0 in a Fe2+ solution facilitated nitrate reduction and its loss paralleled loss of nitrate (Fig. 1b,c), we determined the relationship between Fe2+ loss and nitrate reduction. Batch reactors in these experiments contained 100 mg NO-3–N L-1 (as opposed to 30 mg L-1), 100 mg Fe2+ L-1, and 5% (w/v) magnetite-coated Fe0 and were agitated without additional pH control (initial pH = 5.4). Within 19 h, 35 mg NO-3–N L-1 was transformed to NH+4–N (approximately 35 mg L-1) and all of the initial 100 mg Fe2+ L-1 was consumed. Once all Fe2+ was lost from solution, nitrate removal stopped (Curves 1 and 1' in Fig. 4) . However, when a second dose of 100 mg Fe2+ L-1 was added at 19 h, an additional 35 mg NO-3–N L-1 was transformed to NH+4–N by 42 h (Curves 2 and 2' in Fig. 4). Based on the molar ratio (35/14:100/56 = 1:0.72), transforming 1 mol NO-3–N to NH+4–N under the test conditions would consume 0.72 mol Fe2+.



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  		Fig. 4. Results of the stoichiometric test. Transforming 1 mol of NO-3–N to NH+4–N consumed 0.75 mol of aqueous Fe2+. Results are averages of duplicates.

 
To further explore the relationship between Fe2+ and NO-3 removal, a series of batch tests were conducted. The initial conditions were 5% magnetite-coated Fe0 + 7.14 mM NO-3–N (100 mg NO-3–N L-1) + 0, 0.9, 1.8, 2.7, or 3.6 mM Fe2+ (0, 50, 100, 150, or 200 mg L-1) without pH control (initial pH = 5.0–5.5). As shown in Fig. 5 , results indicated that nitrate removals were proportional to the initial concentration of aqueous Fe2+. Based on these results, 0.75 mM Fe2+ corresponded to 1 mM NO-3–N removed, indicated by the slope (1/0.75) of the line [Fe2+ added = 0.75(NO-3–N removed) - 0.22; R2 = 0.99] in Fig. 5b. The new value (0.75 mM Fe2+ for 1 mM NO-3–N transformation) can be considered refined from the preliminary result (approximately 0.72 mM Fe2+ for 1 mM NO-3–N) based on Fig. 4.



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  		Fig. 5. Series of tests with various initial Fe2+ concentrations (0–3.6 mM or 0–200 mg L-1) show the stoichiometry between nitrate reduction and Fe2+ depletion in the anoxic magnetite-coated Fe0–nitrate–H2O system.

 
Cations
Based on results obtained with Fe2+, we speculated that other cations may also accelerate nitrate reduction and selected Fe3+, Cu2+, Al3+, Ca2+, Mg2+, and Zn2+ for further testing. Experiments were conducted using magnetite-coated Fe0 with no pH adjustment. To be consistent, 1.8 mM selected cation was used in each of these tests. The initial pH differed slightly from that in the tests with Fe2+ due to the different hydrolysis reactions associated with the selected cations. The initial ionic strengths were 17.94 mM in the tests with Fe3+, 12.54 mM with Cu2+ and Ca2+, 20.64 mM with Al3+, and 14.34 mM with Mg2+ and Zn2+. Results indicated that Fe3+, Cu2+, and Al3+ accelerated nitrate reduction by precoated iron at near-neutral pH with a similar pH change (e.g., for Cu2+, pH changed from 4.9 to 9.1) as observed in tests with Fe2+ (below), while Ca2+, Mg2+, and Zn2+ had no significant effect (<5% nitrate removed, pH increased from <6 to 9–10).

When 100 mg Fe3+ L-1 (1.8 mM) and 100 mg NO-3–N L-1 were added to the Fe0–H2O, almost all of the Fe3+ disappeared within the first 30 min while an equivalent charge balance of Fe2+ (150 mg L-1) was detected in solution (Fig. 6a) . On release of Fe2+ (following loss of Fe3+), observations were similar to previous tests (rapid nitrate reduction, Fe2+ depletion, gradual pH increase, and formation of a black precipitate). When Fe2+ became depleted, nitrate reduction had plateaued, with about 55 mg NO-3–N L-1 removed. When nitrate was omitted in a control test, adding 100 mg Fe3+ L-1 to the Fe0–H2O system still resulted in release of 150 mg Fe2+ L-1 into solution, but the Fe2+ concentration remained stable throughout the same test period. By using uncoated Fe0 instead of magnetite-coated Fe0, almost identical results were obtained (data not shown). When 114 mg Cu2+ L-1 (1.8 mM) was used in lieu of Fe3+, instant release of 100 mg Fe2+ L-1 (1.8 mM) was observed (Fig. 6b) and metallic Cu (Cu0) was observed in both the magnetite film and the black precipitate. Approximately 35 mg NO-3–N L-1 was rapidly reduced and paralleled Fe2+ depletion and black precipitate formation. Without nitrate addition, the released Fe2+ remained constant (100 mg L-1) throughout the test (Fig. 6b). When 48.6 mg Al3+ L-1 (1.8 mM) was used in a similar experiment, Al3+ addition resulted in the release of 150 mg Fe2+ L-1 (2.7 mM) (similar to the test with Fe3+), and approximately 50 mg NO-3–N L-1 was removed (data not shown).



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  		Fig. 6. Effect of Fe3+ or Cu2+ addition on nitrate removal. Initial conditions: 5% precoated iron + 30 mg NO-3–N L-1 + (a) 100 mg Fe3+ L-1 (FeCl3) or (b) 115 mg Cu2+ L-1 (CuCl2), without initial pH adjustment.

 
Chloride
Chloride enhances iron pitting and corrosion and has been suggested as a possible mechanism for contaminant reduction (Scherer et al., 1998, 2000). In tests with Fe2+, Fe3+, Cu2+, and Ca2+, chloride salts were used, and therefore Cl- was the major anion in the system. No change in Cl- concentration was detected in any of the tests. These tests indicate that Cl- is not a major contributing factor under the conditions of our experiments. Results using 5% magnetite-coated Fe0 + 1000 mg Cl- L-1 (as NaCl) + 30 mg NO-3–N L-1 without pH control also did not enhance nitrate reduction (<10% nitrate removed in 72 h).


   DISCUSSION
TOP
 ABSTRACT
 INTRODUCTION
 MATERIALS AND METHODS
 RESULTS
 DISCUSSION
REFERENCES
 
The results of this study indicate that in anoxic Fe0–H2O systems, nitrate is removed very rapidly at low pH (2–4), while oxide formation at higher pH will potentially halt the transformation of nitrate. Under our experimental conditions, magnetite (as a surface oxide or precipitate) was a product of nitrate reduction because it was not formed in the absence of nitrate.

At neutral pH, we demonstrated that neither magnetite-coated Fe0 nor magnetite used alone or in combination with Fe2+ was directly responsible for enhanced nitrate removal. However, if Fe2+ or other selected cations (Fe3+, Cu2+, or Al3+) coexisted in the bulk solution with Fe0 or magnetite-coated Fe0, efficient reduction of nitrate to ammonium occurred.

We also demonstrated that reducing 1 mol of NO-3–N to NH+4–N in the Fe0–H2O system at near-neutral pH requires 0.75 mol of Fe2+. Based on this stoichiometric coefficient (0.75) and some additional assumptions, we describe the redox reaction occurring in our systems as follows. In our calculations, we have assumed that (i) NH+4 is produced from nitrate, (ii) both Fe0 and Fe2+ can serve as electron donors, and (iii) the primary oxidation product of Fe0 and Fe2+ is Fe3O4 at a Fe(III) to Fe(II) ratio of 2:1. We also recognized that the conversion of one mole of NO-3–N to NH+4–N will produce X moles of Fe3O4. Each X moles of Fe3O4 contains 3X moles of iron, of which 2X moles will be Fe(III) and 1X moles will be Fe(II). For the Fe(II), our experiments support that 0.75 (stoichiometric coefficient) will come from the aqueous Fe2+ bathing the corroding iron and the rest from the Fe0. Therefore, 3X - 0.75 mol of Fe0 will be oxidized. The oxidation of Fe0 will yield both Fe(II) and Fe(III) and eight electrons are required to convert nitrate to ammonium. Consequently, 3(2X) + 2(X - 0.75) = 8 mol electrons and X = 1.19. Correspondingly, 2.38 mol Fe0 will become Fe(III), while 0.44 (i.e., 1.19 - 0.75) moles Fe0 become Fe (II). The corresponding half-reactions are as follows:

[1]

[2]

[3]

Equations [2] and [3] take control at Stage 2, while under acidic conditions, Fe0 -> Fe2+ + 2e- may be the main reaction. In addition, iron corrosion at near-neutral pH under anoxic conditions can be represented by the equation Fe0 + 2H2O -> Fe2+ + 2OH- + H2. Little change in pH for the control batch Tests 15 and 16, however, suggests that this reaction is slow at pH 5.3 (unadjusted) compared with nitrate-induced iron corrosion. Following mass and charge balance, the overall reaction would be:

[4]

Equations [1] through [4] describe the stoichiometric relationship and do not indicate the specific reaction mechanism(s). Adsorbed Fe2+ (or hydrolyzed adsorbed Fe2+) may react with iron (hydr)oxides (e.g., ferrihydrite, lepidocrocite, maghemite) to form intermediates that are transformed to magnetite (Tamaura et al., 1983, 1984; Charlet et al., 1998; Odziemkowski et al., 2000b; Schultz and Grundl, 2000; Vikesland and Valentine, 2000). The overall reaction (Eq. [4]) will be the same because magnetite is the sole end product of iron oxidation in our tests. Deviation from a Fe(III) to Fe(II) ratio of 2:1 and/or production of {gamma}-Fe2O3, Fe5HO8·4H2O, or {gamma}-FeOOH would alter the stoichiometry. Equation [4] indicates that Fe2+ is a contributor but not the primary reductant of nitrate. From a stoichiometric standpoint, 8 mol of Fe2+ are needed to reduce 1 mol of nitrate (62 g mol-1) or nitrate N (14 g mol-1). This was not observed in our experiments, suggesting that the added Fe2+ is not the major electron donor.

The release of Fe2+ after adding Fe3+, Cu2+, and Al3+ may result from redox reactions or substitution into the corroding iron matrix (Sidhu et al., 1978; Schwertmann and Cornell, 1991). The added Fe3+ may have simply been reduced to Fe2+ by the Fe0 core but a similar release of Fe2+ was noted when Al3+ was used, suggesting substitution. Immediate release of Fe2+ following addition of Cu2+ may involve a thermodynamically favorable redox reaction between Fe0 and Cu2+: Cu2+ + Fe0 -> Fe2+ + Cu0. In fact, metallic Cu was visible at the end of our experiments in which Cu2+ was added. An alternative or additional mechanism in the presence of Cu2+ involves nitrate reduction by Fe2+ or Fe(OH)2, as demonstrated by Buresh and Moraghan (1976). Unlike our results in which treatment of nitrate with Fe0 and Fe2+ produced NH+4 as the primary product, NO-2, N2O, N2, NH2OH, as well as NH+4 were found after treating nitrate with Fe2+ and small quantities of Cu2+, with product distribution dependent on pH (Buresh and Moraghan, 1976; Van Hecke et al., 1990).

While the formation of green rust is a possibility after adding Al3+ because Al2(SO4)3 was used (Hansen and Koch, 1998), green rust was not observed in our systems. An abundance of aluminum during Fe0 oxidation promotes its incorporation into the oxidized iron structure (Schwertmann and Cornell, 1991). The smaller ionic radius of Al3+ disrupts crystallization, yielding amorphous iron (hydr)oxide (perhaps as aluminous ferrihydrite; Taylor and Schwertmann, 1978) and facilitating release of Fe2+.

In this study, we observed three stages of nitrate removal if the initial pH was lowered to 2.3. The acidic pH in Stage 1 favored removal of passivating (nonconducting) layers from the iron grains, allowed rapid nitrate removal (Stage 1), and enhanced iron dissolution and release of Fe2+ into the bulk solution. As the iron continued to corrode and pH increased, we observed a surface layer of corrosion products forming on the iron grains (Stage 2). The slow nitrate removal rate observed in Stage 2 could be due to the formation of amorphous iron oxide (Tamaura et al., 1983; Scherer et al., 1998), which passivates the iron surface until Fe2+ is adsorbed and transformed to magnetite. Based on the literature (Scherer et al., 1998), the surface layer of corrosion products will evolve into a mixture of amorphous iron oxides. This thin, unstable layer readily transforms to magnetite and is difficult to detect by X-ray diffraction (Tamaura et al., 1983). As the pH continued to increase, sustained adsorption and hydrolysis of Fe2+ produced a coating of magnetite on the Fe that resulted in rapid nitrate removal (Stage 3).

Our experimental results are consistent with a Fe2+ adsorption edge at about pH 5.5 (Stage 2) and hydrolysis of adsorbed Fe2+ at pH 6.5 (Stage 3), as reported by Charlet et al. (1998) for Fe2+ adsorption on hematite and magnetite (Vikesland and Valentine, 2000). In our study magnetite was identified as the end product of Fe0 oxidation, consistent with previous research demonstrating magnetite formation following Fe(II) adsorption on {gamma}-FeOOH (Tamaura et al., 1983) and iron oxide colloids in solution (Tronc and Jolivet, 1984). Because Fe2+ was prepared from FeCl2·4H2O, a chloride green rust (GR) may be transiently forming in some of our tests. This GR can reduce nitrate to ammonium with magnetite as the sole Fe oxidation product (Hansen et al., 2000). The highly unstable GR intermediate may not be readily detectable in our experiments.

Given that oxide films will be initially present and evolve with time as the iron corrodes, Scherer et al. (1998) presented three conceptual models for electron transfer at the oxide–iron–water interface: (i) bare iron exposure by pitting, (ii) electron transfer from adsorbed or lattice Fe(II), and (iii) electron transfer through the conductance bands of semiconducting oxide layers. Our tests with added Cl- (a corroding anion) dismiss pitting as a significant mechanism under the conditions of our experiments, because nitrate removal was not accelerated by increasing the Cl- concentration. Electron transfer from adsorbed–lattice Fe(II) or hydrolyzed Fe2+ (FeOH+) is a possibility (Klausen et al., 1995; Schultz and Grundl, 2000), but results from our tests showed no nitrate removal when magnetite and Fe2+ were used alone or in combination, within a wide pH range (Tests 8–10 in Tables 1 and 2). In addition, the stoichiometric relationship between Fe2+ and nitrate does not support Fe2+ as the sole reductant. This leaves electron transfer (semiconductor model) and catalytic hydrogenation via adsorbed atomic hydrogen (Moshtev and Hristova, 1967; Odziemkowski et al., 2000a) as possibilities.

Information generated from this study is important for the application of Fe0–promoted remediation processes. While researchers have been concerned about the inefficiency of Fe0 once iron oxides are formed, our study indicates that under anoxic conditions magnetite will be the product of iron oxidation and the magnetite coating will not hinder nitrate reduction provided sufficient aqueous Fe2+ is present in the system. In situ treatment with permeable iron barriers may be improved by adding Fe2+ to contaminated ground water before it passes through the barrier.


   ACKNOWLEDGMENTS
 
The authors gratefully acknowledge Dr. J. Betanadhatla, Physics Dep., University of Nebraska-Omaha, for X-ray diffraction analysis, Dr. K.J. Klabunde, Dep. of Chemistry, Kansas State University, for Brunauer–Emmett–Teller surface area analysis, and Dr. K. Lee, School of Biological Sciences, Univ. of Nebraska-Lincoln (UNL) for scanning electron microscopy analysis. This research was supported in part by the USEPA Region 7 and 8 Great Plains–Rocky Mountain Hazardous Substance Research Center (Project 95-32) and EPA/EPSCoR Project R-829422-010. The UNL Center for Infrastructure Research and the Water Center provided matching funds for the project. This paper (no. 13812) is a contribution of Agric. Res. Div. Projects NEB-40-002 and -019.


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