Environ. Sci. Technol. 2000, 34, 3495-3500
Reduction of N-Nitrosodimethylamine with Granular Iron and Nickel-Enhanced Iron. 2. Mechanistic Studies MAREK S. ODZIEMKOWSKI, LAI GUI, AND ROBERT W. GILLHAM* Department of Earth Sciences, University of Waterloo, Waterloo, Ontario, Canada N2L 3G1
To elucidate the reduction mechanism of N-nitrosodimethylamine (NDMA) by granular iron, various electrochemical experiments using a mercury electrode were conducted. The studies included differential pulse voltammetry and exhaustive potentiostatic electrolysis. The results of the NDMA electroreduction experiments were compared with the results obtained in the column and batch experiments of Part 1 of this study. The results show that (1) electroreduction of NDMA occurs at potentials more negative than -1.3 V and this potential cannot be achieved under the conditions of the column and batch experiments and (2) different reduction products of NDMA were observed in the electrochemical tests relative to the column and batch tests. That is, dimethylamine (DMA) and nitrous oxide were formed in the electrochemical reduction experiments, whereas ammonia and DMA were produced in the column and batch experiments. The difference in product formation and more importantly the fact that the iron cannot reach the potentials required for electroreduction indicate that the reduction of NDMA on iron cannot take place by direct electron transfer. The process of catalytic hydrogenation was found to be consistent with all experimental observations and is proposed as the alternative mechanism.
Introduction Recognition of the effectiveness of granular iron for destruction of chlorinated solvents and many other reducible contaminants in groundwater has led to numerous investigations of the kinetics and mechanisms of these reactions. Most researchers observe pseudo-first-order kinetics with respect to the contaminant in batch and column experiments (e.g., refs 1-3) with reduction rates for a particular compound dependent on the type and surface area of the granular iron (4). Based on observations from batch experiments, Matheson and Tratnyek (2) proposed three possible dechlorination mechanisms: (1) direct electron transfer from the iron surface to the chlorinated compound, (2) reduction by Fe2+ produced from iron corrosion, and (3) catalyzed hydrogenolysis by H2 produced from the reduction of water. To explain the observed experimental results with chlorinated solvents, most researchers suggest direct electron transfer to be the most likely process (2, 5, 6). To support this mechanism, thermodynamic calculations using standard redox potentials of * Corresponding author phone: (519)888-4658; fax: (519)746-1829; e-mail:
[email protected]. 10.1021/es9909780 CCC: $19.00 Published on Web 07/13/2000
2000 American Chemical Society
chlorinated compounds are often performed (7), and the results are compared to the standard reduction potential of Fe/Fe2+ (2) or, more recently, standard reduction potentials of various Fe2+ complexes (8). There is, however, growing evidence that catalyzed hydrogenolysis, or more generally hydrogenation, may also be an important reduction process in some situations. For example, based on experimental and theoretical work done by Marcus and Protopopoff (9), Bonin et al. (10) suggested that atomic hydrogen adsorbed on Fe, or less likely on magnetite surfaces, may take part in the catalytic dechlorination reaction of CCl4. Furthermore, though not directly related to chlorinated solvents, it has been known for many years that hydrogen adsorbed on Fe surfaces can participate in the reduction of nitrate (11, 12). More recently through electrochemical investigations, it was proposed that the reduction of trichloroethylene by iron and palladized iron is indirect and involves atomic hydrogen (13). The purpose of the present study was to elucidate the reduction mechanism of N-nitrosodimethylamine (NDMA) by granular iron. This was pursued through electrochemical experiments, with the results subsequently being compared with those of the column and batch experiments described in Part 1 of this series of papers. NDMA proved to be advantageous for the study because it is very soluble and thus suitable for spectroelectrochemical studies in aqueous solution. To date, the electrochemical reduction (i.e., direct electron transfer mechanism in the absence of charge transfer mediators) of nitrosamines has only been studied by polarographic methods using alcohols as cosolvents (14-16). Depending on the solution pH, two overall electroreduction processes have been suggested (14)
for pH ) 1-5 RR′N - N ) OH+ + 4e- + 4H+ f RR′N - NH3+ + H2O (1) for pH > 5
2RR′N - N ) O + 4e- + 3H2O f
2RR′ - NH + N2O + 4OH- (2)
To determine the experimental conditions for carrying out electroreduction of NDMA, the reduction potential of NDMA in aqueous solution was measured using differential pulse voltammetry (DPV). Due to low detection limits of DPV (10-9 mol/L), this technique is well suited for studies which concern low level organic contamination (17). We were prompted to use this method by the work of Merica et al. (18), who measured the reduction potentials of hexachlorobenzene very precisely. The measured reduction potential of NDMA was compared to the potential measured on a grain of iron. In addition, the products of electrochemical reduction of NDMA were compared with the products formed through contact with iron in the column and batch experiments. The methodology used in this work to distinguish the electrochemical reduction (i.e., charge transfer model in the absence of charge transfer mediator) of NDMA from catalytic hydrogenation is not entirely new. Recently, Cote` et al. (19) used a similar experimental approach to distinguish electrochemical reduction from catalytic hydrogenation of 2-iodonitrobenzene.
Experimental Methods Chemicals. All solutions were prepared using Milli-Q water. Mercury was obtained with a purity of 99.9998% (Johnson & Matthey), and potassium perchlorate (Aldrich) was used VOL. 34, NO. 16, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 1. Spectroelectrochemical cell. TP, Teflon plug; Pt, platinum ring sealed in glass, counter electrode; Ag/AgCl, Ag wire coated with AgCl, reference electrode; GC, 1.89 mm glass capillary containing mercury, working electrode; SSR, stainless steel rod, electrical contact for the working electrode (Hg); ITGC, internally threaded glass connector; ETR, externally threaded end of stainless steel rod; OFW, optically flat window. as a supporting electrolyte at a concentration of 0.1 mole/L. The KClO4 was purified by calcinating at 300 °C, recrystallizing twice from Milli-Q water and then drying. (CH3)2NNO (NDMA), (CH3)2NH‚HCl (dimethylamine hydrochloride, DMA), were used as received. The granular iron used in the test was obtained from Connelly-GPM Inc. (Chicago, IL) and was used as received. Cells and Electrodes. A three-electrode spectroelectrochemical cell, described in detail elsewhere (20), was modified for this study (Figure 1). Ag/AgCl/Clsat, electrode was used as an external reference electrode. The electrode was prepared by anodic polarization of a Ag wire in 0.1 M NaCl for 24 h with a current density of 0.1 mA/cm2. The AgCl-coated Ag wire was then placed in a solid electrolyte (gel) saturated with NaCl in a glass tube which terminated with a porous Vycor frit. The reference electrode was sealed into a Luggin capillary by means of a Teflon plug (TP) and Kalrez O-ring. To avoid contamination of the working electrode by Cl- ions, the Luggin capillary ended with a second glass-sealed asbestos frit. The working electrode was assembled by filling a 1.89 mm glass capillary (GC) with Hg. The area of the capillary exposed to the electrolyte was flat polished, while the second end terminated in an internally threaded glass connector (ITGC). Electrical contact to the mercury working electrode was made by a stainless steel rod (SSR). As shown in Figure 1, the working electrode was sealed into the main cell body by a Kalrez O-ring. One filling of the glass capillary allowed 4-6 independent measurements on fresh Hg half drops to be performed. The assembled cell was purged with ultrapure argon. The test solution was then injected into the cell through the injection port (Figure 1) and further deaerated for 15 min. Mercury was chosen as the working electrode for several reasons. Mercury, having a very high overpotential for hydrogen evolution, is more useful than other electrode materials to study reduction reactions, both in neutral and alkaline aqueous solutions. In nonaqueous systems, its useful potential reaches as low as -2 V (SHE) (21). In contrast to Pd, Cu, Ni, Fe, etc., Hg is not a hydrogenation metal catalyst (19, 22); therefore the use of this material for our study is essential because it allows us to study exclusively the electrochemical mechanism (i.e., direct charge transfer in the absence of charge-transfer mediators) of the reduction process as demonstrated earlier by other research groups (18, 19, 23). 3496
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The open-circuit potential of a single grain of Connelly iron was measured in a simple two electrode cell with a Ag/ AgCl/Clsat electrode used as an external reference electrode. The cell was purged with ultrapure argon and filled with distilled water containing 25 mg/L of NDMA. The solution was introduced into the cell and deaerated for 30 min. To simulate the geochemical conditions of the batch/column experiments 150 g of sieved granular Connelly iron was added at the bottom of the cell. As in the batch/column investigations, no electrolytic salt was used in this experiment. All potential data were converted to the standard hydrogen electrode value (SHE) as a common point of reference. In this work, we also attempted to use in-situ Raman spectroscopy to study mechanistic aspects of NDMA electroreduction. This method was previously used successfully to study the electrochemical reduction of quinoline (24). Electrochemical Instrumentation. Differential pulse voltammetry (DPV) and cyclic voltammetry (CV) were carried out with an EG&G, PAR, model 273, Potentiostat/Galvanostat interfaced to a PC computer and operated by M270 version 4.1 electrochemical software. The Hg sessile half-drop was subject to continuous cleaning cycles between the potential limits of 0.1 and 0.9 V at 20 mV s-1. This cycling persisted until the cyclic voltammograms were reproducible, which was usually achieved after the second cycle. No peaks characteristic of oxygen reduction were observed, indicating proper deaeration of the solution. The cathodic electrolysis and open-circuit potential measurements were carried out using an AutoLab PGSTAT 30 Potentiostat/Galvanostat interfaced to a personal computer. Raman Spectroscopy Instrumentation. Raman spectra were obtained using a Renishaw 1000 microscope system. The instrument consists of an Olympus microscope, a single spectrograph fitted with holographic notch filters for spectroscopic mode, and a Peltier-cooled charged-couple device detector (CCD) detector. The optical throughput of the instrument is high, as is its sensitivity. The detection of very weak signals (1 photon/s) is therefore possible. Excitation was achieved using the 632.8 nm line of a Melles Griot 35 mW HeNe laser. Raman measurements were carried out using backscattering geometry. Raman spectra of surface films were obtained in-situ using the long working-length objective lens, having a magnification of 50. Using the nonconfocal setting of the instrument, the focal depth of field for this objective is 25.7 µm. Thus, the
FIGURE 3. The dependence of the cathodic peak currents on the concentration of NDMA. Limiting current marked as 2 symbol.
FIGURE 2. Example of differential pulse voltammograms of NDMA at various concentrations. Supporting electrolyte 0.1 M KClO4 in H2O, scan rate 1 mV/s, pulse height 100 mV, pulse width 50 ms, current range 0.1 mA. majority of the vibrational information detected by the Raman microprobe arises from a region of about 25.7 µm around the focal point.
Results and Discussion
FIGURE 4. The dependence of the cathodic reduction peak potentials on the concentration of NDMA. The vertical dotted lines indicate NDMA concentrations used in the iron column experiment.
Determination of Reduction Potential of NDMA at the Mercury Electrode. Differential pulse voltammetry on a sessile Hg drop electrode of 1.76 µL was used to determine the precise reduction potential over a broad range of NDMA concentrations. This information was required to construct a calibration curve (E vs concentration) used later to calculate the electrolysis potentials for given NDMA concentrations. Typical differential pulse voltammograms for the supporting electrolyte (KClO4, 0.1 mol/L) and various concentrations of NDMA are presented in Figure 2. The reduction current for the supporting electrolyte (background current marked as a solid line in Figure 2) started to increase at about -1.2 V and increased sharply at -1.5 V. Merica et al. (18) observed a similar background current increase during electroreduction of hexachlorobenzene in methanol and in methanol-water mixtures (23). A single current maximum, strongly dependent upon NDMA concentration, is assigned to the electroreduction of NDMA. As expected, and as shown in Figure 3, the peak current was proportional to the concentration of NDMA. For concentrations lower than 1.05 mM of analyte, the slope of the peak current vs concentration curve was 2.74 µA/mM. The observed slope is 1 order of magnitude lower than observed for electroreduction of hexachlorobenzene in methanol (18). When methanol is used as the solvent, peak current-concentration curves (Figure 3) intersect the ordinate very close to the origin (18). In our work, the background currents for water reduction are significantly greater; therefore, the peak current vs concentration curve does not intersect at the origin. However, due to the high solubility of NDMA in comparison to hexachlorobenzene (18, 23) it was easier to study a broader concentration range of the analyte. For concentrations in the range of 1.05 to 4.62 mM, the slope decreased to 1.43 µA/mM.
When the concentration of NDMA is further increased, diffusion becomes the limiting step in the course of electroreduction. The diffusion-limited current for 0.1 M NDMA in aqueous solution was measured to be 21.3 µA. An increase in NDMA concentration resulted in a shift of reduction current maxima toward a slightly less negative potential, as shown in Figures 2 and 4. The results of the electroreduction potentials of NDMA obtained by differential pulse voltammetry are compared to available polarographic data in Table 1. As shown, the measured reduction potentials correlate well with those reported in the literature. The reduction potentials of NDMA in water are only slightly more negative than those observed in the water-alcohol solutions using the polarographic method. The influence of water on the reduction potential of hexachlorobenzene was studied by Merica et al. (23). Those authors also observed a shift in reduction potential for the second chlorine removal toward more negative potentials upon addition of water into a methanol organic contaminant mixture (see Figure 1 in ref 23). The reduction potentials at various NDMA concentrations were calculated using the calibration curve of Figure 4. These values were used in the subsequent electrolysis experiments. Electrolysis Experiment. The electrolysis experiments were repeated five times, varying the initial concentrations of NDMA and the duration time of electrolysis. The analyses of reduction products for controlled-potential electrolysis indicated the formation of dimethylamine (DMA) and nitrous oxide (N2O) as the only products. This observation is consistent with the mechanism proposed by Lund for solutions of pH > 5 (reaction 2). Based on the appearance of products, using the Faraday electrolysis equation, current VOL. 34, NO. 16, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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this work not buffered, no additives 0.1 M KClO4 (salt twice recrystallized) [NDMA] ) 10-1 - 10-4
(15)
(14)
(13)
Poorly defined wave, gelatine added as maximum suppressor. b
No polarographic wave. a
differential pulse voltammetry
polarography
polarography
polarography
pH -E1/2 (V) -id (µA) pH -E1/2 (V) -id (µA) pH -E1/2 (V) -id (µA) pH -Epeak (V) -i (µA)
0.0 0.5
not given pH concn dependent ∼ 5.5 concn dependent from 1.31 to 1.38 diffusion limiting current ) 21.3, observed for 0.1 M NDMA only
9.4 1.34
7.15 1.29b 8 5.05 a a 4.34 0.99 17.5 3.6 0.97 13.3 3.62 0.87 20.5 2.75 0.88 13.4 2.8 0.8 17.1 0.95 0.7 13.5 1.41 0.66 23.4 0.2 0.63 13.5
0.6 0.67 13.5
electrochemical reduction of DMNA method
TABLE 1. Reduction Potentials and Diffusion Currents for the Electroreduction of N-Nitrosodimethylamine
10.35 1.3 6.5
12.5 1.31 6.4
H2O + 20% EtOH + 0.01% gelatine 1 M KCl + buffer [NDMA] ) 10-3 H2O + 50% EtOH + 0.1 M KCl, buffered with HCl & NaHPO4, CH3COOH, & CH3COONa [NDMA] ) 3 * 10-4 - 10-3 H2O + 20% EtOH Britton-Robinson buffer [NDMA] ) 2 * 10-3
ref solution - electrolyte
efficiency Ceff ) Qr/Qtotal of the reaction was calculated. Qr is the charge consumed by reaction 2, while Qtotal is the total charge passed during electrolysis. For initial NDMA concentrations of 9.8 and 11.7 mM the average current efficiency was 74% at the potentials of the reduction peaks. The potentials of the reduction peaks were calculated using the calibration curve of Figure 4. For lower NDMA concentrations (0.3 and 0.7 mM) the average current efficiency dropped to 64%. The drop in current efficiency results from the fact that lower concentrations of NDMA require application of more negative potentials (see Figure 4). Therefore, for the total reduction reaction (Qtotal ) Qr + charge consumed for water reduction) there is stronger competition from the reduction of water. For 3 h of electrolysis, the percent of NDMA conversion was ∼4.5%, while for 9 h the conversion reached 11.7%. Yields of DMA were 3.8% and 9.85%, and yields for N2O were 1.1% and 1.94% for 3 and 9 h of potentiostatic electrolysis, respectively. Considering the high current efficiencies, the product yields and percent of NDMA conversion are surprisingly low. Very high current efficiencies are due to (a) a very high overpotential for hydrogen evolution on the Hg electrode and (b) very small current losses for iR drops, noting that the surface area of the Hg cathode is very small (Figure 1). The small surface area not only will contribute to high current efficiency but also will result in low rates of NDMA reduction. These results demonstrate that the overall electrochemical reduction reaction (i.e., direct electron-transfer mechanism) of NDMA in water, free of other additives or charge-transfer mediators, can be written as
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where the NdO double bond has a delocalized character. Mechanistic Considerations. As shown in Part 1 of this study, NDMA reduction as a consequence of contact with Connelly Fe leads to the formation of DMA and ammonia. While DMA is a common product for reduction reactions on Hg and Fe, ammonia was not detected during electroreduction of NDMA. Similarly, nitrous oxide, characteristic of the electroreduction mechanism (14), was not detected in the column and batch experiments. This observation suggests that the reaction occurring in the column and batch experiments was not a consequence of direct electron transfer to the organic molecule. Second, an even more convincing argument can be drawn from the analysis of the results from the electrochemical experiment, as summarized in Figure 5. In this figure, the solid line represents the change of open-circuit potential (OCP) of a single, arbitrarily chosen, grain of Connelly iron immersed in the solution used for the column and batch experiments. Note that column experiments are conducted at OCP (i.e. no current or potential applied). The potential of the grain decayed from -0.102 V (not shown on graph) to -0.52 V, the potential region of magnetite (Fe3O4) thermodynamic stability (25). This behavior, particularly a potential arrest (see inset in Figure 5) is the electrochemical evidence of autoreduction of the Fe2O3 surface film (26-28). The important feature in Figure 5 is that the potential decay passes the equilibrium potentials of hydrogen evolution for pH values observed in the batch tests (pH ) 6.8). The pH range observed in the column experiment increased from 6.3 at the influent to less than 9.2 at the effluent. As demonstrated by Marcus and Protopoppoff (9) for Fe and Ni underpotential deposition (UPD), hydrogen evolution can take place at potentials more positive than the equilibrium potential. Consequently, in the column there is no thermodynamic restraint on hydrogen evolution. While the potential of the iron grain reaches the potential necessary for hydrogen
FIGURE 5. Open-circuit potential time dependence of the single grain of Connelly iron immersed in distilled water containing 25 mg/L of NDMA. The potential spike results from the addition of 10 mL of nondeaerated NDMA solution. evolution (UPD for the entire column and OPD for the majority of the column), it does not approach the potentials (-1.31 to -1.38 V) required for electroreduction of NDMA molecules, marked as the broken lines in Figure 5. Therefore, in our particular case, the charge-transfer reaction does not take place between the iron material and NDMA but rather between iron and water molecules. In the corrosion and electrochemical literature (29, 30) there is agreement that hydrogen evolution takes place through several steps:
Fe + H2O + e- h FeHads + OHFeHads + FeHads F 2Fe + H2
(4)
FeHads + H2O + e- F Fe + H2 + OHThe adsorbed atomic hydrogen is then available to take part in catalytic hydrogenation reactions. It appears that these reactions are responsible for reduction of NDMA according to
FIGURE 6. In-situ Raman spectra of the iron grain after open-circuit measurement. Spectrum A and B represent two randomly chosen grain areas. Raman band assignments are after Hart et al. (32). film. Therefore, reaction 5 might be modified to
The above reaction scheme refers to the situation when bare Fe metal is exposed to the aqueous solution containing NDMA. In normal batch tests such an ideal case can exist only when the grain is mechanically scratched, for example by shaking the vials during batch experiment. As shown in earlier work, in most situations we deal with iron covered by a magnetite (Fe3O4) film (28). These earlier results were confirmed in this work. As demonstrated in Figure 6 by Raman spectral measurements, after ca. 6500 min of immersion in the solution used for the column and batch experiments, the grain of Connelly iron appeared to be covered by magnetite
While there exists a considerable amount of spectroscopic evidence concerning adsorption of atomic hydrogen on transition metals such as Fe, Co, and Ni, the adsorption of H on magnetite-covered Fe surfaces is more speculative and requires confirmation by independent spectroscopic experiments. Thermodynamic calculations by Marcus and Protopopoff (9), however, indicate such a possibility. The thermodynamic stability region of Hads (i.e., UPD H) on Fe clearly overlaps the region of Fe3O4 stability. Calculations in ref 9 assume a simplified model of Langmuir competitive adsorpVOL. 34, NO. 16, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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tion of H and O atoms. In the case when adsorbed O has a strong blocking effect, the domain of FeHads would be shifted to more cathodic potentials, outside the stability region of Fe3O4 (see Figure 4 in ref 9). While the latter may be true in normal aqueous solutions, O2 is rapidly consumed by granular iron and thus is not available as a competitor for hydrogen adsorption. Therefore, in eq 6 the symbol (Fe/Fe3O4)Hads indicates the coexistence of FeHads and magnetite. Let us consider possible reduction mechanisms such as (1) direct electron transfer from Fe to the organic molecule, (2) electron transfer by oxide tunneling, (3) mediated electron transfer, and (4) catalytic hydrogenation. Connelly iron is a product of high temperature thermal treatment of scrap iron (i.e. low alloyed steel and/or cast iron) in rotary kiln ovens and as a result is covered with an inner layer of Fe3O4 and an outer layer of Fe2O3. Direct electron transfer from the metallic surface to the organic molecule is therefore not probable. Electron transfer by oxide tunneling must also be consider. Reductive dissolution, or at low pH chemical dissolution, of Fe2O3 leads to formation or exposure of electronically conducting Fe3O4. Electrochemical and spectroscopic evidences are presented in Figures 5 and 6. Because magnetite is conducting with respect to electrons, the charge transfer through oxide film is a possibility. However the mismatch between the electroreduction potential of NDMA (-1.31 to -1.38 V) and the potential naturally adopted by the Connelly iron (OCP potential -0.51 V) is too large to allow direct charge transfer or transfer by oxide tunneling. The possibility that the observed potential mismatch was overcome by the presence of an electron-transfer mediator must also be considered. The principle underlying mediated charge transfer is that the mediator is a substance having one electron reversible redox properties (24, 31) and the redox potential is close to the peak potential for reduction of the substrate (31). In our experiments only the redox couple Fe2+/ Fe3+ (E° ) +0.77 V vs SHE) meets the first criterion. The complexity of the heat-treated Connelly iron requires consideration of other Fe2+ complexes (8) such as Fe2+/R - Fe2O3 (E° ) -0.28 V at pH 7) or Fe2SiO4/Fe3O4 (E° ) -0.43 V at pH 7) (8). It is evident that the second criterion is not met by these redox couples. Furthermore, it was recently demonstrated (23) that in the case of organic molecules that are difficult to reduce, such as hexachlorbenzene, metal complexes are not effective charge transfer mediators. In addition, the most effective organic mediators (absent in our experiments) can lift the potential mismatch only by a maximum of 0.4 V (31) which is much less than required for NDMA reduction. In addition to the electrochemical evidence, the difference in reaction products between electroreduction and reduction in the presence of iron provides further evidence that the reduction of NDMA by metallic iron occurs by a process other than direct or mediated electron transfer. This study clearly shows that direct charge transfer from iron to the organic molecule does not apply in the case of NDMA reduction. Alternatively, and consistent with the results of a previous study (13), we propose catalytic hydrogenation as the mechanism for NDMA reduction.
Acknowledgments The authors wish to thank Drs. E. Protopopoff and P. Marcus, from Laboratoire de Physico-Chimie des Surfaces CNRS, Universite´ Pierre et Marie Curie, for helpful discussion. This research was funded by the NSERC, Motorola, and ETI Industrial Research Chair held by R. W. Gillham.
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Literature Cited (1) Gillham, R. W.; O’Hannesin, S. F. Ground Water 1994, 32, 958967. (2) Matheson, L. J.; Tratnyek, P. G. Environ. Sci. Technol. 1994, 28, 2045-2053. (3) Orth, W. S.; Gillham, R. W. Environ. Sci. Technol. 1996, 30, 6671. (4) Johnson, T. L.; Scherer, M. M.; Tratnyek, P. G. Environ. Sci. Technol. 1996, 30, 2634-2640. (5) Scherer, M. M.; Westall, J. C.; Ziomek-Moroz, M.; Tratnyek, P. G. Environ. Sci. Technol. 1997, 31, 2385-2391. (6) Deng, B.; Hu, S.; Burris, D. R. In Physical, chemical, and thermal technologies: remediation of chlorinated and recalcitrant compounds; Wickramanayake, G. B., Hinchee, R. E., Eds.; Battelle Press: Columbus, OH, 1998; pp 341-346. (7) Vogel, T. M.; Criddle C. S.; McCarty P. L. Environ. Sci. Technol. 1987, 21, 722-736. (8) Balko, B. A.; Tratnyek, P. G. J. Phys. Chem. B 1998, 102, 14591465. (9) Marcus, P.; Protopopoff, E. In Electrochemical Surface Science of Hydrogen Adsorption and Absorption; Jerkiewicz, G., Marcus, P., Eds.; The Electrochemical Society Inc.: Pennington, NJ, 1997; Vol. 97-16, pp 211-224. (10) Bonin, P. M. L.; Odziemkowski, M.; Gillham, R. W. Corrosion Sci. 1998, 40(8), 1391. (11) Smialowski, M.; Szklarska-Smialowska, Z. Corrosion 1962, 18(1), 1t. (12) Moshtev, R. V.; Hristova, N. I. Corrosion Sci. 1967, 7, 255. (13) Li, T.; Farrell, J. Environ. Sci. Technol. 2000, 34, 173-179. (14) Lund, H. Acta Chem. Scand. 1957, 11, 990. (15) Martin, R. B.; Tashdijian, M. O. J. Phys. Chem. 1956, 60, 1028. (16) Borghesani, G.; Pulidori, R.; Pedriali, R.; Bighi, C. J. Electroanal. Chem. 1971, 32, 303. (17) Greef, R.; Peat, R.; Peter, L. M.; Pletcher, D.; Robinson, J. In Instrumental Methods in Electrochemistry; Kemp, T. J., Ed.; John Wiley & Sons: 1985; Chapter 2, pp 68-73. (18) Merica, S. G.; Bunce, N. J.; Je¸ ndral W.; Lipkowski J. J. Appl. Electrochem. 1998, 28, 645. (19) Cote´, B. J.; Despres, D.; Labrecque, R.; Lamothe, J.; Chapuzet, J.-M.; Lessard, J. J. Electroanal. Chem. 1993, 355, 219. (20) Odziemkowski, M.; Krell, M.; Irish, D. E. J. Electrochem. Soc. 1992, 139, 3052. (21) Galus, Z. In Fundamentals of Electrochemical Analysis; Ellis Horwood: 1994; p 59. (22) Hydrogen Effects in Catalysis. Fundamentals and Practical Applications; Paa`l, Z., Menon, P. G., Eds.; Marcer Dekker: NY, 1988. (23) Merica, S. G.; Bunce, N. J.; Je¸ ndral W.; Lipkowski J. Environ. Sci. Technol. 1998, 32, 1509. (24) Ozeki, T.; Odziemkowski, M.; Irish, D. E. submitted to J. Solution Chem. (25) Pourbaix, M. In Lectrures on Electrochemical Corrosion; Plenum Press: 1973; p 143. (26) Oswin, H. G.; Cohen, M. J. Electrochem. Soc. 1957, 104, 9. (27) Haruyama, S.; Ogura, K.; Nagasaki, K. J. Electrochem. Soc. Jpn. 1969, 37(1), 38. (28) Odziemkowski, M.; Gillham, R. W. In Abstracts, 213th National Meeting of the American Chemical Society; April 13-17, Division of Environmental Chemistry, San Francisco, CA; American Chemical Society: Washington, DC, 1997; pp 177-180. (29) Zakroczymski, T. In Hydrogen Degradation of Ferrous Alloys; Oriani, R. A., Hirth, J. P., Smialowski, M., Eds.; Noyes Publications: Park Ridge, NJ, 1985; p 215. (30) Mizuno, T.; Enyo, M. In Modern Aspects of Electrochemistry; White, R. E., Conway, B. E., Bockris, J. O’M., Eds.; Plenum Press: New York: 1985; No. 30, p 415. (31) Kargina, O.; MacDougall, B.; Kargin, M. Yu.; Wang, L. J. Electrochem. Soc. 1997, 144, 3715. (32) Hart, R. T.; Temkin, H.; Adams, S. B. In Light Scattering in Solids; Balkanski, M., Leite, R., Porto, S., Eds.; Flammarion Sciences: Paris, 1976; p 254.
Received for review August 18, 1999. Revised manuscript received May 10, 2000. Accepted May 17, 2000. ES9909780
