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Published online 5 July 2005
Published in J Environ Qual 34:1286-1292 (2005)
DOI: 10.2134/jeq2004.0459
© 2005 American Society of Agronomy, Crop Science Society of America, and Soil Science Society of America
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TECHNICAL REPORTS

Ground Water Quality

Nitrate Reduction in the Presence of Wüstite

Sudipta Rakshit*, Christopher J. Matocha and Gerald R. Haszler

Department of Agronomy, University of Kentucky, N-122G Agricultural Science Building-North, Lexington, KY 40546-0091

* Corresponding author (sraks2{at}uky.edu)

Received for publication December 2, 2004.

   ABSTRACT
TOP
 ABSTRACT
INTRODUCTION
 MATERIALS AND METHODS
 RESULTS AND DISCUSSION
 CONCLUSIONS
 REFERENCES
 
Recent strategies to reduce elevated nitrate concentrations employ metallic Fe0 as a reductant. Secondary products of Fe0 corrosion include magnetite (Fe3O4), green rust [Fe6(OH)12SO4], and wüstite [FeO(s)]. To our knowledge, no studies have been reported on the reactivity of NO3– with FeO(s). This project was initiated to evaluate the reactivity of FeO(s) with NO3– under abiotic conditions. Stirred batch reactions were performed in an anaerobic chamber over a range of pH values (5.45, 6.45, and 7.45), initial FeO(s) concentrations (1, 5, and 10 g L–1), initial NO3– concentrations (1, 10, and 15 mM), and temperatures (3, 21, 31, and 41°C) for kinetic and thermodynamic determinations. Suspensions were periodically removed and filtered to measure dissolved nitrogen and iron species. Solid phases were characterized using X-ray diffraction (XRD) and scanning electron microscopy (SEM). Nitrate reduction by FeO was rapid and characterized by nearly stoichiometric conversion of NO3– to NH4+. Transient NO2– formation also occurred. The XRD and SEM results indicated the formation of Fe3O4 as a reaction product of the heterogeneous redox reaction. Kinetics of NO3– reduction suggested a rate equation of the type: –d/dt = k0.570.221.12 where k = 3.46 x 10–3 ± 0.38 x 10–3 M–1 s–1, at 25°C. Arrhenius and Eyring plots indicate that the reaction is surface chemical–controlled and proceeds by an associative mechanism involving a step where both NO3– and FeO(s) bind together in an intermediate complex.

Abbreviations: MES, 2-(N-morpholino)ethane sulfonic acid • SEM, scanning electron microscopy • XRD, X-ray diffraction


   INTRODUCTION
TOP
 ABSTRACT
 INTRODUCTION
MATERIALS AND METHODS
 RESULTS AND DISCUSSION
 CONCLUSIONS
 REFERENCES
 
NITRATE IS A SOIL ANION derived from both natural and anthropogenic sources. Because it is weakly sorbed, NO3– is prone to leaching through soil with infiltration water. Elevated NO3– concentrations have been linked to human health problems, prompting the USEPA to establish 44 mg NO3– L–1 as the maximum contaminant level in drinking water (Spalding and Exner, 1993). Consequently, there has been growing interest in understanding processes that lead to NO3– transformations.

Reduction of NO3– to N2 under anaerobic conditions can occur by heterotrophic and autotrophic denitrification (Korom, 1992; Tiejde, 1994). Dissimilatory NO3– reduction to NH4+ is another removal pathway that can occur under certain conditions (Korom, 1992). Historically, processes that reduce NO3– to lower oxidation states of N have been considered to be biologically mediated. However, the occurrence of N2 production and NO3– respiration in oxic surface soils and sediments raises the possibility of abiotic pathways of NO3– reduction (Lloyd et al., 1987; Seitzinger, 1988; Carter et al., 1995).

Field and laboratory studies have shown that the Fe and N cycles are closely coupled. Anschutz et al. (2000) reported that dissolved Fe(II) was oxidized where NO3– was abundant and O2 was absent in marine sediments. Ernstsen and Morup (1992) reported a decrease in NO3– concentration at the reduced zone as Fe(II) to Fe(III) ratio increased in an area dominated by clayey till. Ernstsen (1996) also observed a positive correlation between NO3– reduction and Fe(II) content in clay minerals. Chemical profiling and flux calculations led Hyacinthe et al. (2001) to suggest that NO3– is the primary oxidant for Fe(II) in sediments from the Bay of Biscay.

Coordination environment of Fe(II) has been reported to influence its reactivity with NO3–. The dissolved Fe(II) form, where Fe(II) is coordinated with six water molecules, reduces NO3– slowly even though the overall reaction is thermodynamically favorable (Ottley et al., 1997). It is well known that Fe(II) is more reactive in the complexed form, such as solid Fe(II) minerals or adsorbed Fe(II) surface species (Wehrli, 1990; Stumm and Sulzberger, 1992). In fact, green rust, a mixed Fe(II)–Fe(III) mineral, can reduce NO3– at appreciable rates with formation of NH4+ and magnetite [Fe3O4(s)] as reaction products (Hansen et al., 1996). The activation energy for the reaction was 83.9 ± 7.6 kJ mol–1, suggesting a surface chemical–controlled process (Hansen and Koch, 1998). McGuire et al. (2002) hypothesized that a green rust mineral was responsible for reducing NO3– in a contaminated aquifer. Jeon et al. (2001)(2003) noted that adsorbed Fe(II) on solid iron oxides can reduce NO3–.

Recent studies have capitalized on the effective coupling of the Fe and N cycles by employing metallic Fe0 as a reductant in permeable reactive barriers for remediation of NO3––contaminated water (Kamolpornwijit et al., 2003, 2004; Wilkin et al., 2002). Secondary products identified as a result of Fe0 corrosion include Fe3O4, green rust, and wüstite [FeO(s)] (Huang and Zhang, 2002; Huang et al., 2003; Satapanajaru et al., 2003). There are close structural relationships between Fe0, FeO(s), Fe3O4, and green rust, which suggest that FeO(s) may also exist transiently in soil environments (Bernal et al., 1959). To our knowledge, no studies have been conducted evaluating the reactivity of NO3– with FeO(s). This information is important because reactivity of secondary products of Fe0 corrosion such as FeO(s) could affect the performance of permeable reactive barriers. The objectives of this study were to measure the effects of pH, temperature, and reactant concentrations on the kinetics of NO3– removal by FeO(s).


   MATERIALS AND METHODS
TOP
 ABSTRACT
 INTRODUCTION
 MATERIALS AND METHODS
RESULTS AND DISCUSSION
 CONCLUSIONS
 REFERENCES
 
Commercially available FeO(s) (Alfa Aesar, a Johnson Matthey Company, Ward Hill, MA) was purchased and used as received. Argon-purged (3 h) nanopure water (18 {Omega}) was used to prepare all the solutions. All reactivity studies were conducted in an Ar- and H2–purged anaerobic chamber (Coy Laboratory Products, Grass Lake, MI) equipped with a palladium catalyst to remove traces of O2. The noninterfering buffers MES [2-(N-morpholino)ethane sulfonic acid] with a concentration of 0.3 M and PIEPES (1,4-piperazine diethane sulfonic acid) with a concentration of 0.05 M were added to maintain pH control (Alowitz and Scherer, 2002).

Stirred batch reactions were performed in duplicate, 30-mL glass vials at initial concentrations of 1 to 10 g FeO(s) L–1 at pH 5.45, 6.45, and 7.45. The pH values were chosen to cover the range encountered in studies involving metallic iron treating nitrate-contaminated water (pH 6.6–7.4) (Kamolpornwijit et al., 2003, 2004). Nitrate was added at initial concentrations of 1 to 15 mM (62–930 mg L–1) to initiate the reactions. Suspensions were removed at increasing time intervals and filtered using 0.2-µm membrane filter paper (Fisher Scientific, Hampton, NH). Aliquots were complexed with ferrozine [3-(2-pyridyl)-5,6 bis(4-phenylsulfonic acid)-1,2,4-triazine, monosodium salt] immediately to stabilize dissolved Fe(II). The Fe(II)–ferrozine complex was measured at 562 nm using a UV-VIS-NIR scanning spectrophotometer (UV-3101 PC; Shimadzu, Kyoto, Japan). Nitrate and NO2– concentrations were measured using a Model 792 basic ion chromatograph (Metrohm, Herisau, Switzerland) with a MetroSep A column, MetroSep RP guard disc holder, and a 3.2 mM Na2CO3 to 1 mM HCO3– eluent with conductivity detection. The retention times for NO2– and NO3– were 6 and 10 min, respectively. Ammonium was measured using the indophenol-blue method (Ngo et al., 1982) after all the uncomplexed Fe(II) had been air-oxidized to Fe(III) and filtered. Changes in NO3– concentration in control experiments with no FeO(s) added were negligible, indicating that FeO(s) was the sole contributor to NO3– reduction.

In a separate experiment to verify whether the NO3– reduction process was abiotic, all reagents were passed through sterilized 0.2-µm membrane filter paper and stored in autoclaved bottles inside the anaerobic chamber, and the reaction was performed as before using 10 mM NO3– at pH 6.45. It was found that NO3– reduction by FeO(s) using sterilized reagents and bottles resulted in similar reaction rates and NH4+ production (data not shown). In another experiment, Fe(II) was extracted from FeO(s) using MES at pH 6.45 and 10 mM NO3– was added to the extracted Fe(II) to determine whether the solid- or solution-phase Fe(II) was reactive.

Solid-phase products dried under argon on glass slides were characterized using X-ray diffraction (XRD) and scanning electron microscopy (SEM). Samples were scanned from 0 to 80 degrees 2{theta} employing CoK{alpha} radiation at 40 kV using a PW 1840 diffractometer, a PW 1729 X-ray generator (Philips, Eindhoven, the Netherlands), and a Bragg-Bretano design goniometer to detect diagnostic d-spacings. For SEM, dried samples were deposited on carbon tape attached to an aluminum holder and coated with Au/Pd. A Hitachi (Tokyo, Japan) S-3200 scanning electron microscope was used to image reacted and unreacted samples at a voltage of 20 keV and a working distance of 15 mm.

Temperature dependence of NO3– reduction was conducted using a circulating water bath at four different temperatures (3, 21, 31, and 44°C) with continuous argon purging. The initial concentrations of NO3– and FeO(s) were approximately 10 mM and 10 g L–1, respectively, and the pH was set to 6.45 using MES buffer.


   RESULTS AND DISCUSSION
TOP
 ABSTRACT
 INTRODUCTION
 MATERIALS AND METHODS
 RESULTS AND DISCUSSION
CONCLUSIONS
 REFERENCES
 
Stoichiometry
Solid FeO reduced NO3– rapidly, with nearly 50% of the NO3– lost from solution after 6 h at pH 6.45 (Fig. 1a) . During the reaction, NO3– was converted stoichiometrically to NH4+ (Fig. 1b). Similarly, Hansen et al. (1996) reported stoichiometric reduction of NO3– to NH4+ by green rust. Nitrite appeared as an intermediate in the overall reduction of NO3– (Fig. 1a). However, in the green rust study, Hansen et al. (1996) did not find any NO2– formation. When reacted alone with FeO(s), 0.6 mM NO2– was completely reduced to NH4+ within 4 h (Fig. 1a, inset). The more rapid rate of NO2– reduction by FeO(s) compared with NO3– reduction verifies that it behaves as a transient intermediate in the overall process of NO3– reduction. We also measured N2O; however, the concentration of N2O was negligible.



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  		Fig. 1. (a) Nitrate reduction to NH4+ and NO2– by FeO(s) at a constant pH 6.45. The small inset shows NO2– reduction by FeO(s) in a separate experiment over a 4-h interval under similar conditions. (b) Production of NH4+ vs. consumption of NO3– for the reaction of NO3– with FeO(s) at a constant pH 6.45. The solid line represents a linear least square regression fit of the data. Error bars indicate ±range of duplicate runs; bars not visible are smaller than the symbols.

 
The solid phase was characterized following FeO(s) oxidation by NO3– and in control experiments where NO3– was not added. X-ray diffraction indicated the presence of FeO(s) in the control experiments based on d-spacings at 2.49, 2.15, and 1.52 Å (Fig. 2a) . There are no other XRD peaks found in the control experiments, indicating that FeO(s) is the only mineral phase present under the initial reaction conditions. Magnetite [Fe3O4(s)] was found as a product of FeO(s) oxidation by NO3– as shown by diagnostic d-spacings at 2.97, 2.53, 2.10, 1.71, 1.61, and 1.48 Å (Fig. 2b). Some unreacted FeO(s) was also present in the NO3––treated experiment. Green rust oxidation by NO3– also resulted in Fe3O4 formation (Hansen et al., 1996). The broad XRD peak in the reacted sample likely indicates the presence of poorly crystalline ferrihydrite [Fe(OH3)].



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  		Fig. 2. X-ray diffraction (XRD) patterns of (a) control experiments where NO3– was not added to FeO(s) in MES [2-(N-morpholino)ethane sulfonic acid] at constant pH 6.45, and (b) reaction product at constant pH 6.45, 10 g L–1 FeO(s), and 1 mM NO3–. In (a), peaks at 2.49, 2.15, and 1.52 Å represent the diagnostic d-spacings for wüstite. In (b), peaks at 2.97, 2.53, 2.10, 1.71, 1.61, and 1.48 Å represent the d-spacings for magnetite, and peaks at 2.49, 2.15, and 1.52 Å represent the d-spacings for unreacted wüstite.

 
Scanning electron micrographs illustrate distinctly different features of the control FeO(s) experiments from the NO3––treated experiments. Unreacted FeO displayed large smooth surfaces with cubic and hexoctahedral faces (Fig. 3a) . Reacted FeO(s) showed rough surfaces and small (<0.5 µm) poorly defined Fe3O4 crystals growing out from the surface (Fig. 3b). Laboratory preparations of Fe3O4 crystals exhibit cubic forms (Schwertmann and Cornell, 1991). Koch and Hansen (1997) presented SEM micrographs of Fe3O4 particles produced during green rust oxidation by NO3– that exhibit cubic-like faces. Specific faces could not be found in our Fe3O4 particles (Fig. 3c), which could be result of a poorly crystalline nature.



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  		Fig. 3. Scanning electron micrographs of (a) unreacted FeO(s) and (b and c) reacted FeO(s) in MES [2-(N-morpholino)ethane sulfonic acid] at constant pH 6.45. Bars show the length unit.

 
A twelve-to-one (12 ± 0.56 to 1) consumption of Fe(II) for each mole of NO3– reduced was found (data not shown). The solution- and solid-phase data suggest the following reaction (Eq. [1]):

[1]

This reaction, which consumes protons, occurs without O2 and all of the NO3– appears to be reduced to NH4+ without any loss of N (Fig. 1b). For kinetic calculations a buffer was necessary to compensate for the tendency of Eq. [1] to increase pH and slow the reaction.

Reduction of NO3– by the FeO suspension could occur by either a homogeneous reaction involving dissolved Fe(II) in equilibrium with the solid FeO and dissolved NO3– or by a heterogeneous reaction with the solid FeO and dissolved NO3–. Reaction between the FeO-filtrate and NO3– was negligible (data not shown). This implies that dissolved Fe(II) was not responsible for the rapid conversion of NO3–. Previous studies have noted solid Fe(II) minerals to be more effective at reducing NO3– than dissolved Fe(II) (Ottley et al., 1997).

It seems that more than one mechanism may be operative, particularly at longer times where NO3– concentration exhibited pronounced curvature (Fig. 1a). In addition, NO2– accumulated and declined after 6 h, despite its greater reactivity with FeO(s) than NO3– (Fig. 1a, inset). One possible reason for the kinetic behavior of NO3– could be reaction with adsorbed Fe(II)–Fe(OH)3 complexes that could be present in our system, even in small amounts. Jeon et al. (2001)(2003) reported reduction of NO3– by adsorbed Fe(II) on solid iron oxides. Adsorbed Fe(II)–Fe(OH)3 complex can result in Fe3O4 production (Tronc et al., 1992) in addition to FeO(s) oxidation by NO3–. Another feasible reductant that may be present in our experiment is the FeO–Fe3O4–Fe(II) complex. Reduction of NO3– with these potential reductants may explain its complicated kinetic behavior as the reaction progressed. We assume that the initial reduction of NO3– (0–3 h) was performed by FeO(s) and kinetic calculations were limited to data from this region.

Kinetic Analysis
The method of initial rates and isolation were employed in the analysis of the kinetic data (Lasaga, 1981). The overall rate equation for NO3– reduction by FeO(s) can be expressed as:

[2]
where –d/dt is the rate of disappearance of NO3–; k is the overall rate constant; and x, y, and z are reaction orders for FeO(s), H+, and NO3–, respectively. When the experiments are performed with a constant excess of FeO(s) and varying the concentration of NO3–, Eq. [2] reduces to:

[3]
where:

Taking the log of both sides of Eq. [3] allows one to calculate z:

[4]

A plot of log vs. log [NO3–] would result in a straight line and from the slope and y intercept, the reaction order (z) and pseudo-first-order rate coefficient, kI, can be determined. Likewise, at constant solid phase and NO3– concentrations one can determine the reaction order for H+ ion from the expression below:

[5]
where:

The solid-phase dependence can also be determined using the equation below:

[6]
where:

The initial rate of NO3– disappearance was evaluated by fitting initial linear slopes over the first 20% of the reaction.

The reaction order for NO3– at pH 6.45 was 1.12 (Fig. 4a) , in contrast to zero-order dependence of NO3– above 7 mM NO3– concentration in the reaction with chloride green rust (Hansen et al., 2001). These differences can be attributed to differences in mineral structure. Chloride green rust minerals have layered structures containing external and internal sites for NO3–, with chloride functioning as a charge-balancing interlayer anion (Hansen et al., 2001). Reactive sites of FeO(s) can be classified as external (Bernal et al., 1959).



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  		Fig. 4. Initial rate plots to determine apparent reaction order for (a) NO3–, (b) [H+], and (c) [FeO]. The different initial NO3– concentrations were 0.9, 9.5, and 14.7 mM, pH values were 5.45, 6.45, and 7.45, and FeO(s) concentrations were 10, 5, and 1 g L–1. The solid line represents a linear least square regression fit of the data.

 
Evaluation of the reaction order for [H+] gave a value of 0.22 (Fig. 4b), which indicates fractional order dependence of the reaction rate on H+. The reaction order with respect to FeO(s) was 0.57 (Fig. 4c) indicating fractional order dependence on solid-phase concentration. The fractional order dependence on FeO(s) in the reduction of NO3– may result from the multistep electron transfer process as NO3– is reduced to NH4+. Therefore, the overall rate equation can be expressed by Eq. [7]:

[7]

The calculated rate constant (k = 3.46 x 10–3 ± 0.38 x 10–3 M–1 s–1, at 25°C) from Eq. [7] for 10 g L–1 (139 mM) FeO(s) concentration, pH 6.45, and 9.49 mM NO3– concentration is in good agreement with the rate constant (3.13 x 10–3 m–1 s–1) derived from the graph (Fig. 4a, from the intercept and the expression kI = k[FeO]x[H+]y). The calculated rate constant (k) is four times greater than the rate constant calculated for NO3– reduction by green rust (Hansen et al., 1996).

Temperature Study
Nitrate reduction by FeO(s) showed temperature dependence over a range of 3 to 44°C (Fig. 5a) . The rate of the reaction increased approximately 13.6 times when temperature was raised from 3 to 44°C (Fig. 5a). The activation energy (47.18 ± 6.43 kJ mol–1) determined from the slope of the graph suggests a surface chemical–controlled reaction (Sparks, 1989). Generally, surface chemical–controlled reactions have activation energies of >42 kJ mol–1 whereas diffusion-controlled reactions are characterized by activation energies of <42 kJ mol–1 (Sparks, 1999). Nitrate reduction by sulfate green rust yielded an activation energy of 83.9 ± 7.6 kJ mol–1 (Hansen and Koch, 1998), which is higher than the activation energy determined in our experiment. Entropy of activation ({Delta}S{ddagger}) calculated from the intercept of Eyring plot (Fig. 5b) showed a value of –172.8 ± 20.93 J mol–1 K–1. A value of {Delta}S{ddagger} of less than –10 J mol–1 K–1 indicates an associative reaction mechanism, whereas values greater than +10 J mol–1 K–1 indicate a dissociative reaction mechanism (Atwood, 1985). Associative reaction mechanisms predict that an intermediate of higher coordination number is formed between the entering group (in our experiment NO3–) and FeO(s). A large effect of the entering group would be expected and this is reflected in the order for NO3– in Eq. [7] (Atwood, 1985). Therefore, NO3– reduction by FeO(s) likely occurs via a surface-controlled associative mechanism. However, detailed investigation of the mechanism is needed to evaluate the exact pathways of the reaction.



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  		Fig. 5. Temperature study showing (a) Arrhenius and (b) Eyring plots describing NO3–, reduction by FeO(s) at a constant pH 6.45, solid-phase concentration 10 g L–1, and initial NO3– concentration 1 mM.

 
Environmental Implications
The long-term performance of Fe0–based permeable reactive barriers is affected by secondary corrosion products (Kamolpornwijit et al., 2003). Wüstite has been identified as a product during Fe0 corrosion (Satapanajaru et al., 2003). This finding may be due to the close structural relationship between Fe0, FeO(s), Fe3O4, and green rust (Bernal et al., 1959). If FeO(s) is present, our data indicate that it will effectively reduce NO3–.

It would be useful to evaluate NO3– reduction by FeO(s) in the presence of other competing anions typically found in normal waters such as chloride, sulfate, and phosphate. The derived rate equation (Eq. [7]) can be used to determine the probable speed of NO3– disappearance with measured environmental data. The calculated rate of reaction after assuming an initial NO3– concentration near the maximum contaminant level (44 mg NO3– L–1), a pH 6.45, at a temperature of 25°C, and an initial FeO(s) concentration of 0.59 g L–1 (8.2 mM, similar to Hansen et al., 1996) would be 4.90 x 10–3 mM NO3– h–1. This rate of NO3– disappearance or NH4+ formation (since stoichiometric NH4+ is formed) is about four times faster than the rate of NH4+ formation in presence of green rust (0.36 mg N kg–1 d–1 or 1.042 x 10–3 mM h–1 NH4+ formation) (Hansen et al., 1996), indicating that FeO(s) might be dominating the potential NO3– transformation pathways at pH 6.45. At higher pH (8.2) the calculated rate is 1.97 x 10–3, which is about two times faster than the rate of NO3– reduction by green rust under similar conditions.


   CONCLUSIONS
TOP
 ABSTRACT
 INTRODUCTION
 MATERIALS AND METHODS
 RESULTS AND DISCUSSION
 CONCLUSIONS
REFERENCES
 
In general, wüstite rapidly reduced NO3– to NH4+ with transient formation of NO2–. Characterization of the solid-phase product suggested the formation of Fe3O4 as an end product. Stoichiometric conversion of NO3– to NH4+ verified the multi-electron transfer process. Kinetic studies revealed a general rate equation of the type: –d/dt = k0.570.221.12 where k = 3.46 x 10–3 ± 0.38 x 10–3 M–1 s–1, at 25°C. The calculated rate constant (k) appeared faster than the rate constant derived from NO3– reduction by green rust. Temperature dependence of NO3– reduction resulted in an activation energy of 47.18 ± 6.43 kJ mol–1 and entropy of activation of –172.8 ± 20.93 J mol–1 K–1. The understanding of the exact reaction mechanism is complex; however, these values revealed the possibility of associative and surface chemically controlled pathways.


   ACKNOWLEDGMENTS
 
This project was supported by National Research Initiative Competitive Grant no. 2002-35107-12214 from the USDA Cooperative State Research, Education, and Extension Service. We are grateful to Dr. A.D. Karathanasis, Dr. Mark Coyne, and Dr. Alan Fryar for providing lab facilities, helpful advice, and guidance, and Yvonne Thompson for running numerous XRD samples. Comments from the anonymous reviewers greatly improved the manuscript.


   REFERENCES
TOP
 ABSTRACT
 INTRODUCTION
 MATERIALS AND METHODS
 RESULTS AND DISCUSSION
 CONCLUSIONS
 REFERENCES
 




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