Environ. Sci. Technol. 2008, 42, 2420–2425
Anaerobic Corrosion Reaction Kinetics of Nanosized Iron E R I C J . R E A R D O N , * ,† R A N D A L F A G A N , † JOHN L. VOGAN,‡ AND ANDRZEJ PRZEPIORA‡ Department of Earth Sciences, University of Waterloo, Waterloo, Ontario, N2L 3G1, Canada and EnviroMetal Technologies Inc., 745 Bridge St. W., Suite 7, Waterloo, Ontario, N2V 2G6, Canada
Received May 23, 2007. Revised manuscript received December 15, 2007. Accepted December 21, 2007.
Nanosized Fe0 exhibits markedly different anaerobic corrosion rates in water compared to that disseminated in moist quartz sand. In water, hydrogen production from corrosion exhibits an autocatalytic style, attaining a maximum rate of 1.9 mol kg-1 d-1 within 2 d of reaction. The rate then drops sharply over the next 20 d and enters a period of uniformly decreasingrate,representedequallywellbyfirst-orderordiffusioncontrolled kinetic expressions. In quartz sand, hydrogen production exhibits a double maximum over the first 20 d, similar to the hydration reaction of Portland cement, and the highest rate attained is less than 0.5 mol kg-1 d-1. We ascribe this difference in early time corrosion behavior to the ability of the released hydrogen gas to convect both water and iron particles in an iron/water system and to its inability to do so when the iron particles are disseminated in sand. By 30 d, the hydrogen production rate of iron in quartz sand exhibits a uniform decrease as in the iron/water system, which also can be described by first-order or diffusion-controlled kinetic expressions. However, the corrosion resistance of the iron in moist sand is 4 times greater than in pure water (viz. t1/2 of 365 d vs 78 d, respectively). The lower rate for iron in sand is likely due to the effect of dissolved silica sorbing onto iron reaction sites and acting as an anodic inhibitor, which reduces the iron’s susceptibility to oxidation by water. This study indicates that short-term laboratory corrosion tests of nanosized Fe0/ water slurries will substantially underestimate both the material’s longevity as an electron source and its potential as a longterm source of hydrogen gas in groundwater remediation applications.
Introduction The potential role of nanosized zerovalent iron (NZVI) in remediation technology has spawned considerable research. Zhang (1) gives an overview of the prospects and limitations of this new technology. Many studies report greater reactivity of NZVI compared to most commercial irons, which are typically mixtures of macro and microsized particles. Nurmi et al. (2) noted that these reports should be considered preliminary because a host of potentially significant questions concerning dehalogenation performance and longevity re* Corresponding author e-mail:
[email protected]; phone: 519-888-4567ext 33234. † University of Waterloo. ‡ EnviroMetal Technologies Inc. 2420
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main unresolved. The susceptibility of NZVI particles to oxidation in air is one factor that limits longevity. Huber (3) concludes the prevention of passivation associated with this oxidation is a key technology missing today. One question that has not received a lot of attention is the rate at which NZVI particles lose their ability to dechlorinate hydrocarbons due to progressive electron loss from anaerobic corrosion. Fe(s) + 2H2O(l) S Fe2+ + 2OH- + H2(g)
(1)
Liu et al. (4) reported that an analysis of the Fe0 content of a Koda Togyo Corp. NZVI, which had been stored underwater for 7 mo, was only 27%, compared to the company’s initial analysis of 60%. These authors also synthesized NZVI using the borohydride reduction method (5) and measured hydrogen production rates in deionized water. They found that all the Fe0 was consumed by corrosion in 11 d. Although the results indicate markedly different reactivities of the two materials to water, neither result presents a promising outlook for NZVI as a long-term electron source. This is reflected in current ideas that the likely role of NZVI will be as a highly reactive, short-term treatment of source contamination (1, 6). In this study, we examine the reactivity of a NZVI material under several conditions: (1) to oxygen when the material is in its original dry state, (2) to water under water-saturated, anaerobic conditions, and (3) to water when the NZVI is disseminated in quartz sand under both water unsaturated and water-saturated, anaerobic conditions.
Experimental Section Materials. The variety of methods to produce NZVI have increased over the past decade. The techniques that show promise to produce large batches of material fall within three general categories: (1) low-temperature chemical methods, such as reduction of a ferrous or ferric salt solutions using a strong reducing agent like sodium borohydride (7); (2) hightemperature reduction of an oxidized iron source, such as hematite or goethite in a stream of hydrogen gas (8); and (3) ball-milling of iron or iron oxides in the presence of reducing organic solvents or hydrogen gas (9), or annealing the material in hot hydrogen gas after the ball-milling step (10). A promising new method involves heating a metastable iron compound, such as iron pentacarbonyl (Fe(CO)5), in the presence of a surfactant, which catalyzes the formation of iron nanoparticles (11). Iron pentacarbonyl has also been used to produce iron nanoparticles by vapor condensation (12). A review of synthesis methods for NZVI and other nanosized materials is given by Baer et al. (13). Any NZVI synthesis method that involves ball-milling creates a material with a large number of lattice defects (14) and thus a potentially large reservoir for entrapment of the H2(g) produced when it corrodes (15). This is a drawback if H2(g) generated by corrosion at the iron surfaces is desirable for effecting reductive reactions in the particular remediation application. However, if the iron is ball milled in a stream of H2(g) or is annealed in H2(g) afterward (10), the produced NZVI could be a valuable source of additional hydrogen — released from the lattice as the NZVI corrodes by the contaminant groundwater. In this study, two NZVI materials were selected for study. One was produced by Dr. Wei-xian Zhang’s research group at Lehigh University using the borohydride reduction method and shipped under ethanol. Following the nomenclature proposed by Nurmi et al. (2), this iron will be referred to as FeBH. The other was produced by the Toda Kogyo Corporation 10.1021/es0712120 CCC: $40.75
2008 American Chemical Society
Published on Web 02/23/2008
TABLE 1. Information on NZVI Samples: Surface Area, Initial Conditions of Corrosion Tests (Masses of NZVI, Quartz Sand, Water and Cell Void Space), Apparent Corrosion Rate (% of the Initial Fe0 Content Per Day) and % Fe0 Content at the End of the Corrosion Tests run No.
NZVI source
surface area (m2/g)
NZVI (g)
quartz (g)
H2O (g)
cell void vol. (cm3)
Rcorr (test end) (% d-1)
1 2 3 4
Toda Kogyo Corp. (FeH2)
24.0
2.17 0.501 0.77 0.512
0 0 150.7 100.1
15.0 15.0 20.5 40.0
112 112 50.3 50.0
0.48 (80 d) 1.27 (8 d) 0.037 (325 d) 0.57 (57 d)
5
Lehigh Univ. (FeBH)
1.64
0
15.0
113
0.0 (80 d)
(Japan) using hydrogen gas to reduce ferric oxide at high temperature. This material will be referred to as FeH2. It was shipped dry in a sealed metal drum. Upon opening, the bagged material was stored in an anaerobic chamber. A record of the manufacturer; BET surface area; and mass of NZVI, water, and void space used for the corrosion tests are listed in Table 1. Also included are the corrosion rate and % Fe0 remaining at the end of the test periods. The initial materials were examined with scanning electron microscopy (SEM) (see Figure SI-1 in the Supporting Information) and were analyzed by low angle X-ray diffraction (XRD), both before and after the corrosion tests using a Bruker AXS D8 Advance X-ray diffractometer. Corrosion Test Apparatus. All corrosion tests were conducted with the NZVI sample placed in a 125 cm3 stainless-steel cell outfitted with an upper Swagelok valve to allow evacuation before adding water through the lower valve. Because the FeBH material was shipped in ethanol, the actual mass of dry sample added to the test cell was determined by subtracting the mass lost during an evacuation step performed prior to the addition of water. During the test period, the cells were immersed in a water bath maintained at 25 ( 0.1 °C. Cell pressures were measured with Omega PX 302, 0–15 psi, absolute vacuum pressure transducers, and readings were collected at 1 min intervals using a PC outfitted with a 16-bit data acquisition board. Full details on the corrosion test procedure and calculations of corrosion rates are given by Reardon (16, 17).
Results FeBH. Oxygen uptake and corrosion rate measurements for the FeBH sample showed it to be unreactive. XRD analyses indicated no compositional change over an 80 d corrosion test (Figure SI-2), and except for weak evidence for the major line of iron, only lines for magnetite were observed. SEM analysis of the material (Figure SI-1) showed only octahedral crystals, characteristic of magnetite. An analysis was not provided with the received sample but we performed triplicate analyses by measuring hydrogen evolution upon addition of 1:1 HCl. These results indicated less than 0.1% Fe0 was present. Somehow, the sample had oxidized during or after synthesis. We tested the remote possibility that the iron had reacted with the ethanol storage solution, since methanol solutions are known to corrode iron (18). We placed two samples of the dry FeH2 in test cells with pure ethanol and ethanol mixed with 5% water. No measurable hydrogen gas was produced over 7 d of monitoring. We are uncertain what led to the oxidation of the FeBH sample, but further testing of this material was terminated. Oxygen Reactivity of FeH2. Establishing the reactivity of NZVI with oxygen is important to determine if special handling is necessary prior to its use. The reactivity of dry FeH2 was determined by placing a 4 g sample in a sealed, 125 mL stainless-steel cell in contact with air. Continuous monitoring of the cell pressure enabled appraisal of the reactivity of the material to oxygen by calculating the % mass
Fe metal remaining (%) 34 77 40 66 0.0
uptake of oxygen from the decrease in cell pressure over time. These results are presented as the lower curve in Figure SI-3. The pressure drop in the cell over 20 h was only 1.0 kPa, which corresponds to the reaction of 1 mL of O2 with the 4 g of iron. This is a mass uptake of only 0.033% and indicates oxygen uptake by this NZVI is negligible if handling times in air are short. Anaerobic Corrosion of FeH2 in Water. Several corrosion tests of FeH2 were performed (see Table 1). Runs 1 and 2 were performed in water at two water/solid ratios, and Runs 3 and 4 were performed with the FeH2 disseminated in Ottawa quartz sand, one water-unsaturated and the other water-saturated. The cell pressure for the 80 d Run 1 experiment is plotted in Figure 1. Hydrogen production was evident within 3 h following the addition of water. By 92 h, the cell pressure had climbed to over 150 kPa, and the cell was evacuated as a precaution to prevent excessive pressure buildup, which could damage the transducer (burst pressure of 400 kPa). Two additional pressure reduction steps were carried out on days 14 and 33. The apparent rates of H2(g) production, that is, uncorrected for hydrogen entering the iron lattice, are shown in Figure 2. The maximum rate was nearly 2000 mmol kg-1 d-1 and occurred at 40 h. By contrast, Fisher Electrolytic iron, which is a microsized material that shows a similar style of corrosion, attains a maximum rate of only 70 mmol kg-1 d-1 after 75 h of reaction (17). Even after 80 d, the rate for FeH2 was 80 mmol kg-1 d-1, still higher than Fisher’s maximum rate. Because of the OH- production associated with corrosion (1), NZVI material emplaced into an aquifer or permeable reactive barrier (PRB)could have a dominant and prolonged effect on the pH of inflowing groundwater. At the end of Run 1, the cell was placed under vacuum to remove residual water. After attaining constant mass, the cell was filled with air, sealed, and pressure monitoring was begun to determine the reactivity of the corroded FeH2 to oxygen. The results are presented in Figure SI-1 and show that oxygen uptake by the reacted FeH2 was nearly 20 times greater over the 20 h of reaction than the unreacted material. Much of this uptake is likely a result of oxygen reaction with
FIGURE 1. Cell pressure change versus time for FeH2 in water (Run 1). A gap in the data around 10 d was due to a power outage. VOL. 42, NO. 7, 2008 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 2. Apparent rate of H2(g) production vs time for FeH2 in water (Run 1). Rate expressed per kg of the initial Fe0 metal content of the FeH2 sample.
FIGURE 3. A comparison of the first 40 d of H2(g) production for Runs 1-4. (Rate expressed per kg of the initial Fe0 metal content of the FeH2 samples. ferrous-containing alteration products formed during the 80 d of anaerobic corrosion. X-ray diffraction patterns were obtained for both the original, “as received” FeH2 and the material recovered from Run 1 after 80 d of reaction. These data are presented in Figure SI-2a and indicate that unreacted FeH2 is principally iron metal, showing only a weak peak for the major magnetite line. In the reacted sample, the intensity of the major iron metal line dropped to about 25% of the initial intensity. Magnetite is the dominant phase present. All identified peaks are accounted for by these two phases. Anaerobic Corrosion of FeH2 in Quartz Sand. In Run 3, 0.5 mass % FeH2 was disseminated within quartz sand (see Table 1), and the mixture was packed into a test cell. The cell was sealed, evacuated, and filled with 70 mL of DI water. After allowing several minutes for the water to soak into the material, the water was drained by pressurizing with nitrogen gas. This produced a wet, but water-unsaturated iron/sand mixture. The cell was then re-evacuated, and pressure monitoring was initiated to record H2(g) generation. The initial water content and void space are provided in Table 1. The cell pressure was monitored for 325 days. The H2(g) production rate results for the first 40 d are compared to the Run 1 results in Figure 3. The maximum rate (475 mmol kg-1 d-1) occurred at 15 h. Calculation of the Amount of FeH2 Corroded. The most direct method for calculating the quantity of iron corroded in Run 1 is to total the amount of H2(g) released to the atmosphere as a result of the three degassing events shown in Figure 1 and to add it to that present in the cell at the end of the corrosion test. These four values are: 6.75, 6.15, 2.61, and 6.77 mmol, respectively, for a total of 0.0223 mol. Because the XRD data indicates the major product of corrosion is magnetite (Fe3O4), the overall corrosion reaction can be represented by eq 2, 3Fe(s) + 4H2O(l) S Fe3O4 + 4H2(g) 2422
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which indicates that 1.333 moles of H2(g) were released for every mole of iron corroded. So the total iron corroded at the end of the corrosion test is 55.85 × 0.0223/1.333 ) 0.934 g. The initial quantity of FeH2 used was 2.17 g (Table 1), of which 65–75 mass % was Fe metal as reported on the container by Toda Corp. The 25–35% balance was listed as Fe3O4. We performed triplicate analyses of the Fe0 metal content by measuring the volume of H2(g) evolved by acid digestion. These results averaged 65 ( 3% Fe0. Using this value, the Fe0 content of the FeH2 material used in Run 1 is 0.65 × 2.17 g ) 1.41 g Fe, and the percent corroded over the 80 d experiment is 0.934/1.41 × 100 ) 66%. This is in qualitative agreement with the XRD results, which show an approximate 75% reduction in the peak intensity of the major Fe0 line before and after the corrosion test (see Figure SI-2a). For Run 3 (FeH2 mixed with moist quartz sand), the initial Fe0 content was 0.65 × 0.77 g ) 0.50 g. The total H2(g) gas produced during the 325 d experiment (0.00714 mol) corresponds to 0.299 g Fe0 corroded or 60 mass %, which is less than the 66% corroded in pure water in only 80 d. Toda FeH2 Slurry pH Values. To get some idea of the pH values inside the corrosion cell, pH measurements were recorded for a FeH2/water and a FeH2/sand/water slurry over the early time period of high corrosion rates (see Table SI-1). These ranged from 11.2 to 11.7 for the FeH2/water slurry and from 10.6 to 11.4 for the FeH2/sand/water slurry. The pH values are slightly lower with sand present, possibly due to some pH buffering due to dissociation of silicic acid (H4SiO4°) formed from the dissolution of quartz: SiO2 + 2H2O ⇒ H4SiO4°. These measured pHs are high compared to predicted values for a pure Fe0/water system (∼9.2) but agree with the pH reported by Liu and Lowry (21) for their 300 g/L slurry of Toda FeH2 (10.6). Apparently, some alkali component from the synthesis procedure is present in Toda FeH2.
Discussion Lattice Hydrogen and the Sievert Rate Constant. Reductions in the cell pressure carried out during a corrosion test produce discontinuities in the hydrogen production rate (17). These discontinuities reflect a decrease in the rate of hydrogen entering the iron lattice (Rentry) due to the decrease in PH2 (eq 3), Rentry ) kSPH2
(3)
where kS is the Sievert rate constant, which has units of mmol kg-l d-1 kPa-0.5. This constant allows one to estimate the proportion of hydrogen entering the iron lattice at a given PH2. Three pressure reductions were conducted during Run 1 but only the first produced a discontinuity. This indicates that hydrogen was entering the iron lattice only at early time. The Sievert rate constant (derived by removing this discontinuity) was 11.0 mmol kg-l d-1 kPa-0.5 and was used to calculate corrected hydrogen production rates over the entire test period. These data are shown in Figure 4. A comparison of the apparent and corrected curves indicates that about 10% of the produced hydrogen was entering the iron at the time of the first pressure reduction. However, the derived kS that allows this conclusion generates a discontinuity in the corrected corrosion rate curve at 330 h, where none existed. This indicates that little or no hydrogen was entering the iron lattice at 330 h. In fact, it is possible that hydrogen that previously entered the Fe0 lattice is released with progressive corrosion. Consequently, we consider the apparent rate of hydrogen production a better estimator of the overall corrosion rate than the corrected rate curve shown in Figure 4. For the FeH2 in moist sand experiment (Run 3), two pressure reductions were carried out (one is shown in Figure 3 at 29 d), and neither produced discontinuities in the rate curve. We conclude that lattice entry of hydrogen is not a significant process in the long term anaerobic corrosion of FeH2.
FIGURE 4. Apparent and “corrected” H2(g) production rates per kg of initial Fe0 metal content in Run 1. Correction was applied using the Sievert rate constant derived by eliminating the discontinuity in the apparent rate curve at 100 h. The correction, however, is inappropriate at later time because it creates a discontinuity at 330 h where none existed. Early Time Corrosion Kinetics. H2(g) production over the first 10 d of FeH2 in water (Run 1 in Figure 3) is substantially greater but similar in style to Fisher electrolytic iron (17). There is a rapid rise to a maximum rate over several days and then a marked drop to a long period of slowly decreasing rate. The reaction kinetics during this early period is best described as autocatalytic. Initially, the corrosion rate is low as the added water penetrates the protective oxide coatings and begins to corrode the iron metal beneath. The generated H2(g) escapes, causes convection of the water and particles, which helps remove alteration products from the particles or prevent them from accumulating, thus helping keep fresh surfaces exposed to water. After a period of rapid acceleration in rate, a maximum is attained and the agitation (and transport) of the particles due to the escaping hydrogen gas begins to subside. With progressively less bubbling, convection of water and iron particles is markedly reduced, and particles corrode in situ. To verify the effect of gas bubbling on enhancing corrosion rates, a short-term test (Run 2 in Table 1) was conducted with the same amount of water but only ¼ the amount of FeH2 as Run 1. The results in Figure 3 demonstrate that the corrosion rates are indeed lower at early times because the intensity of gas release is lower, thus reducing the contribution of convection of FeH2 particles to the corrosion process. A visualization of this effect is provided in Figure SI-4. It shows a photograph taken of two glass vials containing a high and a low water/FeH2 ratio three days into the corrosion process. The corrosion of FeH2 disseminated in moist sand differs markedly from FeH2 in pure water. Run 3 in Figure 3 shows that, after a period of increasing rate of H2(g) production, a decline to a minimum occurs, followed by a rise to a second maximum. The rate then falls again and enters a period of slowly decreasing rate. This behavior is akin to the hydration kinetics of Portland cement (19). Upon addition of water to cement, a period of increasing hydration occurs as water wets the particles. A maximum rate is attained as hydration products build up on the cement particle surfaces. This causes a sharp drop in the hydration rate, called the “dormant period”. After several hours, the protective coatings break down or recrystallize, exposing fresh unhydrated cement to the porewater and a second maximum in hydration rate is attained. Several models are presented by Taylor (20). For the corrosion of FeH2 disseminated in sand, rather than the onset of the second maximum being due to surface alteration products recrystallizing, it is more likely a result of their disaggregation by outward escaping hydrogen gas, allowing water in to contact fresh Fe0 surfaces. The ability of the particles to convect and keep alteration products from accumulating when FeH2 corrodes in water only is likely why this double maximum and dormant period does not occur.
FIGURE 5. FeH2 in water (Run 1): Comparison of predicted firstorder and diffusion-controlled corrosion rates to experimental determinations from H2(g) evolution. Rate expressions in the figure yield values in units of % of initial Fe0 content per day.
FIGURE 6. FeH2 in moist quartz sand (Run 3): Comparison of predicted first-order and diffusion-controlled corrosion rates to experimental determinations from H2(g) evolution. Rate expressions in the figure yield values in units of % of initial Fe0 content per day. Later Time Corrosion Kinetics. Figure 3 (and 2) shows that after about 15 d for FeH2 in water and 30 d for FeH2 in sand, H2(g) production has slowed to 20 d for Run 1 and >30 d for Run 3). The rate equations shown in the figures were derived by integrating the following linear equations, which were fit to plots of either ln Fe (g) versus t (d) for the first-order model or Fe (g) versus t1/2 (d) for the diffusion model. Regression equations for Run 1 data (FeH2 in water; t > 20 d). First Order:
ln Fe (g) ) -0.0089t 0.0120; R2 ) 0.9982 (5)
Diffusion:
Fe (g) ) -0.077t1/2 + 1.1765; R2 ) 0.9962 (6)
Regression equations for Run 3 data (FeH2 in moist sand; t > 30 d). First Order:
ln Fe (g) ) -0.0019t 0.9988; R2 ) 0.9984 (7)
Diffusion:
Fe (g) ) -0.0126t1/2 + 0.4307; R2 ) 0.9985 (8)
Plots of the linear fits to the data are provided in the Supporting Information (Figure SI-5 and SI-6). The closeness of the R2 values and the visual fits of the rate equation predictions to the experimental data in Figures 5 and 6 indicate that the data is well represented by both models and does not provide a basis to distinguish which model more likely reflects the actual kinetics. Liu and Lowry (21) derived a first-order reaction constant (kobs ) k[SA][H+]) of 0.0060 d-1 for a 300 g/L slurry. This was based on analyses of residual Fe0 contents versus time. In an experiment using a far less concentrated slurry (0.5 g/L), they observed a much lower rate constant of 0.0031 d-1 based on hydrogen gas evolution measurements. The kobs values in our study are comparable: 0.0089 d-1 for a FeH2/water slurry (Run 1) and 0.0019 d-1 for FeH2 dispersed in moist sand (Run 3). Corrosion in Water versus Moist Sand. The first-order model predicts a half-life of 365 d for FeH2 disseminated in moist sand compared to 78 d in water only. Clearly, the controls on corrosion rate are very different. Three factors may be involved: (1) convection of water by escaping H2(g) in the FeH2/water sample continues to enhance the corrosion rate beyond 30 d; (2) water transport to corrosion sites in the FeH2/sand sample becomes rate limiting; and (3) silica dissolved from the quartz sand sorbs to and blocks corrosion sites on the FeH2 particles as the overall corrosion rate subsides. The first explanation appears unlikely. After 30 d, H2(g) production in the 15 mL FeH2/water sample (Run 1) only 3.5 mL/d, hardly enough to produce sufficient convection of water and particles to account for a 4 times shorter corrosion half-life as compared to the sample of FeH2 in wet sand (Run 3). A water transport problem also seems unlikely. We calculate that 1 cm3 of the moist sand contains 100 times more water than required to completely corrode the FeH2 particles it contains. We also repeated the FeH2/sand run but under water-saturated conditions to provide further evidence. Unlike the former water-unsaturated run, produced hydrogen 2424
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must diffuse or bubble up through 0-10 cm of the cell’s saturated water/sand mixture before the pressure transducer responds to its production. So the transducer response will be delayed and less sensitive to rapid changes in corrosion rate. The results (Run 4) are shown in Figure 3 and, although lacking the detail obtained in Run 3, corroborate those results rather well. Explanation 3 seems most probable. PHREEQCI (22) using the MINTEQ.V4 database (23) predicts a pure water/Fe0 system at saturation with respect to Fe(OH)2(s) and quartz has a pH of 9.20 and a Si content of 0.12 mmol/l. Kohn et al. (24) showed that, at a similar concentration (0.17 mmol/ L), sufficient Si sorbed onto microscale ZVI to reduce TCE degradation rates by as much as 30% compared to when no Si is present. Kohn et al. (24) ascribe the effect of adsorbed silica to its known role as an anodic inhibitor, preventing the metal from releasing electrons and thus reducing its ability to undergo anaerobic corrosion (25). As Si dissolves from the quartz, it forms a hydrated gel on the iron surfaces, which inhibits the anodic dissolution reaction (Fe0 ⇒ Fe2+ + 2e-) (25, 26). The measured pHs of Toda FeH2/water/sand mixtures (10.6–11.4; see Table SI-1) are much higher than the predicted pure system pH of 9.2, and because quartz solubility increases markedly above a pH of 9, dissolved Si could be substantially higher than 0.12 mmol/L. Thus, the iron surface poisoning effect noted by Kohn et al. (24) appears to be a probable explanation for the substantially lower anaerobic corrosion rate we observe for FeH2 when quartz sand is present. This study has shown that both the early time and longterm anaerobic corrosion rates of FeH2 as measured by H2(g) production are markedly lower when disseminated in a quartz sand as compared to in water alone. The principle reasons are (1) an early time enhancement of corrosion due to gasinduced convection of NZVI particles in water and (2) the probable role of silica as an anodic inhibitor of corrosion when quartz sand is present. Thus, short-term, laboratorybased corrosion tests of NZVI/water slurries may substantially underestimate both the material’s longevity and its potential as a long-term source of hydrogen gas in groundwater remediation applications.
Supporting Information Available Some discussion, tables, and figures were referred to in this paper but not presented. This information is available free of charge via the Internet at http://pubs.acs.org.
Literature Cited (1) Zhang, W-X. Nanoscale iron particles for environmental remediation: An overview. J. of Nanopart. Res. 2003, 5, 232–332. (2) Nurmi, J. T.; Tratnyek, P. G.; Visarathy, V.; Baer, D. R.; Amonetter, J. E.; Pecher, K.; Wang, C.; Linehan, J. C.; Matson, D. W.; Penn, R. L.; Driessen, M. D. Characterization and properties of metallic iron nanoparticles: Spectroscopy, electrochemistry, and kinetics. Environ. Sci. Technol. 2005, 39, 1221–1230. (3) Huber, D. L. Synthesis, properties, and applications of iron nanoparticles. Small 2005, 1, 482–501. (4) Liu, Y.; Majetich, S. A.; Tilton, R. D.; Sholl, D. S.; Lowry, G. V. TCE Dechlorination rates, pathways, and efficiency of nanoscale iron particles with different properties. Environ. Sci. Technol. 2005, 39, 1338–1345. (5) Ponder, S. M.; Darab, J. G.; Mallouk, T. E. Remediation of Cr(VI) and Pb(II) aqueous solutions using supported, nanoscale zero valent iron. Environ. Sci. Technol. 2000, 34, 2564–2569. (6) Tratnyek, P. G.; Johnson, R. L. Nanotechnologies for environmental cleanup. Nanotoday 2006, 1, 44–48. (7) Wang, B. W.; Zhang, W-X. Synthesizing nanoscale iron particles for rapid and complete dechlorination of TCE and PCB’s. Environ. Sci. Technol. 1997, 31, 2154–2156. (8) Hisano, S.; Saito, K. Research and development of metal powder for magnetic recording. J. of Magn. Magn. Mater. 1998, 190, 371–381.
(9) Nasu, T.; Tokumitsu, K.; Konno, T.; Suzuki, K. Reduction of iron-oxide by ball-milling with hydrogen gas flow. Mater. Sci. Forum. 2000, 343 (I), 435–440. (10) Pardavi-Horvath, M.; Takacs, L. Iron-Alumina nanocomposites prepared by ball milling. IEEE Trans. Magn. 1992, 28, 3186– 3188. (11) Huber, D. L.; Venturini, E. L.; Martin, J. E.; Provencio, P. P.; Patel, R. J. Synthesis of highly magnetic iron nanoparticles suitable for field structuring using a beta-diketone surfactant. J. Magn. Magn. Mater. 2004, 278, 311–316. (12) Choi, C. J.; Tolochko, O.; Kim, B. K. Preparation of iron nanoparticles by chemical vapor condensation. Mater. Lett. 2002, 56, 289–294. (13) Baer, D. R.; Tratnyek, P. G.; Qiang, Y.; Amonette, J. E.; Linehan, J.; Sarathy, V.; Nurmi, J. T.; Wang, C.; Anthony, J. Synthesis, characterization, and properties of zero-valent iron nanoparticles. In: Environmental Applications of Nanomaterials: Synthesis, Sorbents, and Sensors; Fryxell, G. E. , Ed.; Imperial College Press: London, (in press). (14) Di Cicco, A.; Berrettoni, M.; Stizza, S.; Bonetti, E.; Cocco, G. Microstructural defects in nanocrystalline iron probed by x-rayabsorption spectroscopy. Phys. Rev. B 1994, 50, 12386–12397. (15) Novakova, A. A.; Kiseleva, T. Yu.; Agladze, O. V.; Perov, N. S.; Tarasov, B. P. The effect of hydrogen incorporation in the nanocrystalline iron particles on their magnetic properties. Intl. J. Hydrogen Energy. 2001, 26, 503–505. (16) Reardon, E. J. Anaerobic corrosion of granular iron: Measurement and interpretation of hydrogen evolution rates. Environ. Sci. Technol. 1995, 29, 2936–2945. (17) Reardon, E. J. Anaerobic corrosion of zero-valent irons: Inorganic controls on hydrogen pressures. Environ. Sci. Technol. 2005, 39, 7311–7317.
(18) Kawai, T.; Nishihara, H.; Aramaki, K. Inhibition effects of amines and thiols on iron corrosion in anhydrous methanol solution containing FeCl3. J. Electrochem. Soc. 1996, 143, 3866–3873. (19) Sujata, K.; Jennings, H. M. Formation of a protective layer during the hydration of cement. J. Am. Ceram. Soc. 1992, 75, 1669– 1673. (20) Taylor, H. F. W. Cement Chemistry, 2nd ed.; Thomas Telford: London, United Kingdom, 1997. (21) Liu, Y.; Lowry, G. V. Effect of particle age (Fe0 content) and solution pH on NZVI reactivity: H2 evolution and TCE dechlorination. Environ. Sci. Technol. 2006, 40, 6085–6090. (22) Charlton, S. R.; Parkhurst, D. L. PHREEQCl—A graphical user interface to the geochemical model PHREEQC: U.S. Geological Survey Fact Sheet FS-031–02, 2002. (23) Allison, J. D.; Brown, D. S.; Novo-Gradac. K. J. MINTEQA2/ PRODEF2, A Geochemical Assessment Model for Environmental Systems: Version 3.0 User’s Manual. US EPA, Athens, Georgia, EPA/600/3–91/021,1991. (24) Kohn, T.; Kane, S. R.; Fairbrother, D. H.; Roberts, A. L. Investigation of the inhibitory effect of silica on the degradation of 1,1,1-trichloroethane by granular iron. Environ. Sci. Technol. 2003, 37, 5806–5812. (25) Katsanis, E. P.; Esmonde, W. B.; Spencer, R. W. Soluble silicate corrosion inhibitors in water systems. Mater. Perform. 1986, 25, 19–25. (26) Fujita, N.; Matsuura, C.; Ishigure, K. The effect of silica on hydrogen evolution and corrosion of iron in high-temperature water. Corrosion. 1990, 46, 804–812.
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