Environ. Sci. Technol. 1997, 31, 160-170
Iron(III) Catalyzed Photochemical Reduction of Chromium(VI) by Oxalate and Citrate in Aqueous Solutions STEPHAN J. HUG* AND HANS-ULRICH LAUBSCHER Swiss Federal Institute for Environmental Science and Technology, EAWAG, CH-8600 Du ¨ bendorf, Switzerland BRUCE R. JAMES University of Maryland, College Park, Maryland 20742
Cr(VI) reduction in soil solution, wastewater, and natural waters is poorly understood in complex systems containing Fe(II,III) and dissolved organic C, especially when influenced by sunlight. The Fe(II,III)-mediated photochemical reduction of Cr(VI) was investigated in a laboratory study with 5-200 µM Cr(VI), 25-1300 µM oxalate or citrate, 0.13-6.7 µM Fe(III), 10 mM KCl, pH 3-7, and a xenon light source (720 W/m2 between 300 and 800 nm) at 25 °C. In situ UV-VIS multicomponent analysis avoided addition of interfering reagents. At higher [Cr(VI)], photochemical Cr(VI) reduction was zero order in [Cr], with quantum yields [relative to light absorption by Fe(III)-oxalate] of up to 0.53. Over 95% Cr(VI) reduction was observed within 2040 min in 5-cm cells. At lower [Cr(VI)], the reaction order became complex due to slow reaction of Fe(II) with Cr(VI) compared to photochemical Fe(II) production. The thermal reaction of Cr(VI) with Fe(II) at pH 5 was measured and described by -d[Cr(VI)]/dt ) 1.2 ((0.3) × 107 M-2 s-1 × [Cr(VI)][Fe(II)][Ox]. A tentative mechanistic kinetic model is presented that fits the results of the dark and photochemical experiments with oxalate. Photochemically formed superoxide (O2•-), and hydroperoxyl radical (HO2), also appeared to be important reductants, reducing Cr(VI) with bimolecular rates of 5-8 × 104 M-1 s-1. Fe(II), HO2/O2•-, and H2O2 were likely reductants of Cr(V) and Cr(IV) intermediates. The reaction product with oxalate was mainly soluble Cr(III)oxalate. The results are applicable to understanding how rapidly Cr(VI) may be reduced in natural water and soil environments.
Introduction Reactions that reduce Cr(VI) to Cr(III) in contaminated waters and soils have received considerable attention because they diminish or eliminate the threat to aquatic life and to human health posed by Cr(VI) contaminations (1). Chromium(VI) is acutely toxic and a known human carcinogen and mutagen (2), and due to its solubility, it is highly mobile in aquatic systems. Chromium(III) is considered nontoxic in most forms and is an essential trace metal in human nutrition (3). At pH values above 5-7, aqueous uncomplexed Cr(III) hydrolyzes to sparsely soluble chromium(III) hydroxides, adsorbs strongly to mineral and organic surfaces, and co-precipitates with other minerals. Precipitation and sedimentation of Cr(III) * E-mail address:
[email protected].
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constitute a known pathway for removal of chromium from the water column in wetlands, lakes, and estuaries (4). The redox chemistry and surface reactions of Cr have recently been reviewed (5). Reduction of Cr(VI) by dissolved organic compounds at pH values from 4-8 is a slow process, occurring on a time scale of days at micromolar (6) and up to months at nanomolar Cr(VI) concentrations (7). Cr(VI) contamination in wetlands, lakes, and subsurface waters can spread over large areas before substantial reduction occurs. For these reasons, it is important to understand Cr(VI) reduction pathways and rates. Deng and Stone (8) studied the catalytic effect of surfaces on Cr(VI) reduction by a number of low molecular weight organic compounds in the pH range from 3 to 7. TiO2, Al2O3, and FeOOH surfaces reduced Cr(VI) half-life times from typically hundreds of days down to tens to hundreds of hours, and TiO2 was found most effective. For example, while at pH 3.5-6.5, no reduction of 20 µM Cr(VI) in homogeneous 200 µM oxalate was observed even after 400 days, 30% reduction occured in 200 h in the presence of TiO2 at pH 4.7. In remediation procedures for Cr(VI)-contaminated natural waters and soils, Fe in its zerovalent or divalent state is often suggested as an efficient reductant for Cr(VI) (9). Fe(II) is generally considered to reduce Cr(VI) quickly, but rate expressions and rate constants at near-neutral pH’s have not been reported. Recently, Kieber and Helz (10) observed a diurnal cycle of Cr(VI)/Cr(III) in estuarine waters. In laboratory experiments, they have shown that Fe(II), which is photochemically produced from organics and solid Fe(III) phases, was the most likely reductant of Cr(VI) at nanomolar concentrations at pH 8. Kaczynski and Kieber studied photochemical production of Cr(III) in freshwater lakes (11). Photochemical reduction of Cr(VI) might be useful in remediation procedures and is likely to be important in natural aquatic systems where Fe(II,III) and organic compounds are present, particularly for slightly acidic to neutral surface waters (e.g., in contaminated wetlands) and in atmospheric cloud and fog droplets. In such systems, photochemical production of Fe(II) from both solid Fe(III) phases and dissolved Fe(III) in the presence of iron(III)-complexing organic compounds is known to be efficient (12, 13). However, photochemical production of Fe(II) from Fe(III) organic complexes is accompanied by the production of HO2/O2•-, H2O2, and other reactive oxygen species that, together with O2, reoxidize Fe(II) (14). The net results are photochemical Fe(II,III) cycles in which steadystate concentrations of Fe(II), Fe(III), HO2/O2•-, and H2O2 are established, and the complexing organic ligands are ultimately oxidized (14, 15). Strong Fe(III)-complexing ligands shift the steady-state equilibria to Fe(III), thereby lowering the p for Fe(II) oxidation at a given pH [Fe(II) becomes a stronger reductant]. Fe(II,III) cycling in the presence of oxalate was described by Sedlak and Hoigne (16). Reduction of Cr(VI) by Fe(II) is thus in competition with oxidation of Fe(II) by O2 and other oxygen species. While rate laws for reduction of Cr(VI) by Fe(II) between pH 1 and 2 have been determined (17), little is known about the rates of Cr(VI) reduction by Fe(II) at higher pH values. With respect to the mobility in the aqueous environment, the speciation of Cr(III) formed from the reduction of Cr(VI) is of interest. During the reduction of Cr(VI) to Cr(III), the coordination changes from tetragonal to octahedral. If the two additional ligands are H2O, formation of chromium(III) hydroxides and precipitation would be expected at pH values >5. In contrast, complexation with organic ligands might form soluble or insoluble Cr(III) complexes that do not readily hydrolyze. The speciation of Cr(III) is relevant in connection
S0013-936X(96)00253-2 CCC: $14.00
1996 American Chemical Society
to its reoxidation, e.g., by manganese oxides (18, 19, 20). In order to address these questions in more detail, we have conducted laboratory studies at conditions that are similar to those found in contaminated wetlands and soils or in slightly acidic atmospheric water: 5-20 µM Cr(VI), pH 5-7, 25-100 µM oxalate or 100 µM citrate, with ionic strength maintained constant with 0.01 M KCl. We used oxalate and citrate as model organic ligands. These and other low molecular weight acids are present in carbon-rich surface waters, atmospheric water, and drainage water from forested and agricultural watersheds. They also serve as models for polycarboxylic groups in fulvic and humic acids. Leenheer et. al showed that di- and polycarboxylic acid functional groups are abundant in fulvic acids (21, 22). We also conducted experiments at higher Cr(VI) and oxalate concentrations at pH 3-4 to simulate conditions in more heavily polluted natural waters, in acid mine drainage, in acidic organic horizons of forest soils, and in acid sulfate soils. The issues addressed in this study are as follows: (1) quantum yields for Fe(II,III)-mediated photochemical Cr(VI) reduction, (2) reaction stoichiometry and possible mechanisms, (3) rates of Cr(VI) reduction by Fe(II), and (4) role of other possible reductants for Cr(VI).
Experimental Section All chemicals used were at least reagent grade and were used as supplied. High purity 18MΩ water (Q-H2O grade Barnstead Nanopure) was used for all experiments. Solutions or suspensions were prepared by dilution of stock solutions [1 × 106 >1 × 106 1 × 109
M12 M13 M14
Fe(II) + O2•- + 2H+ Fe(II) + H2O2 Fe(II)Ox + H2O2
Reactions of Fe(II) in Presence of Oxalate f Fe(III) + H2O2 f Fe(III) + OH• + OHf Fe(III)Ox + OH• + OH-
1 × 107 63 3.1 × 104
(36) (37) (16)
M15 M16 M17
CH3COO- + OH• Cr(VI) + Red Ox + OH•
Reactions with Hydroxyl Radicals f Red f Cr(V) f OH- + CO2 + CO2•-
1 × 108 >1 × 104 7.7 × 106
a
M18 M19
HO2 + O2•- + H+ H2O2 + OH•
Reactions with Reactive Oxygen Species f H2O2 + O2 f H2O + HO2
9.7 × 107 2.7 × 107
(39) (40)
M20 M21 M22
Fe(III) + Ox Fe(III)Ox + Ox
Formation of Fe(III)Ox f Fe(III)Ox f Fe(III)Ox2 f O2
1 × 1010 1 × 1010 3 × 10-10 M s-1
b c
Reactions of Cr(VI) and Cr Intermediates with Fe(II) a Fe(II)Ox f Cr(V) + Fe(III)Ox a Cr(V)-Fe(III) f Cr(V) + Fe(III)Ox
0.037 ( 0.007 s-1
1 × 109 2.4 × 109 1 × 1010/1 × 107 (1.2 ( 0.3) × 104 1 × 1010/1 × 1010 (1.2 ( 0.3) × 107
(35) (26)
Reactions of Cr(VI) with Oxygen and Formate Radicals f Cr(V) + O2 (5-8) × 104 f Cr(V) + CO2 (6-12) × 107
(38)
a Possible pathway of reductant (Red) formation: CH COO- + OH• f •CH COO- + H O(1 × 108 M-1 s-1) (41) and •CH COO- + O f •O CH COO 3 2 2 2 2 2 2 2 (1.7 × 109 M-1 s-1) (42), where •CH2COO- and/or •O2CH2COO2- ) Red. b The back reaction is omitted since virtually all Fe(III) is complexed. c O2 input by purging with 0.5 mL/s N2, which contains 0.2 ppm O2. [O2]0 ) 250 µM for aerated and 0.25µM for deareated solutions. See text for typeface explanation.
our surprise, Cr(VI) reacts very slowly with Fe(II) at pH 5 in the absence of oxalate at micromolar concentrations of Cr(VI) and Fe(II). In solutions with 20 µM Cr(VI) and 60 µM Fe(II) at pH 5, less than 50% of Cr(VI) was reduced in 30 min. Our preliminary experiments have shown a strong pH dependence of this reaction rate, and we are currently studying the role of pH and of organic ligands in more detail. We point out that we have not added any reagents to the reaction solutions. In many previous studies, DPC (1,5diphenylcarbazide) was used to measure Cr(VI), which might lead to errors due to the normally very low pH of the DPC reagent solutions. [The reaction of Cr(VI) with Fe(II) is fast at low pH and might compete with the formation of the colored complex from Cr(VI) and DPC]. Postulated Model with Photochemical and Thermal Reactions. Our model (see Table 2) produced the fits (solid and dashed lines) shown in Figures 4a-c and 6. Rate constants printed in bold were obtained by simultaneous least square fitting, rate constants printed in normal type are from the literature references indicated, and rate constants printed in italics could not be found in the literature and were estimated. Equilibria are represented by fast rate constants for forward and back reactions. For easier readability, we have abbreviated C2O42- with Ox and have omitted the charges of the complexes. In ACUCHEM, reactions M8 and M9, where several reactants and products are separated by commas, had to be entered as separate lines for each reactant and product. CO2, H2O, and H+ were omitted because CO2 is unreactive and [H2O] and [H+] were constant. The influence of pH was accounted for by considering the speciations of HO2/O2•and Cr(VI). At pH 5, we have 62% O2•- and 38% HO2 (pKa )
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4.7). Since O2•- is the major and more reactive component for oxidations of Fe(II) species, we have listed only the reactions with O2•- and have multiplied the rate constant in reaction M12 by 0.62. The rate constant for M18 is multiplied by 0.62 × 0.38 to account for the speciation at pH 5. Discussion of the Model. M1-M3: Photolysis of FeOx and Formation of Products. The dominant species in our concentration ranges were FeOx2 and FeOx3, which were assumed to be photolyzed with the same quantum yields. The quantum yield of Fe(II) production from FeOx under actinometric conditions ([FeOx] > 6 mM, pH 0-1) has been characterized extensively (27). Under these conditions, [FeOx] . [O2] and CO2•- or C2O4•- reacts quantitatively with FeOx and forms a second Fe(II), resulting in quantum yields for Fe(II) of 1.11-1.25 between 430 and 300 nm. The question whether C2O4•- or CO2•- is the primary photoproduct of FeOx photolysis is inconsequential in our systems: C2O4•- is reported to decay with a unimolecular rate of 2 × 106 s-1 to CO2 and CO2•- (43), which is faster than bimolecular reactions of C2O4•- with reactants at micromolar concentrations. In our systems, [FeOx] is between 125 nM and 6.7 µM and CO2•reacts almost quantitatively with O2 to O2•- in aerated systems. Thus, Fe(II) and O2•- are formed at equal rates with expected quantum yields of about 0.56-0.63. In deaerated systems, the reaction of CO2•- with FeOx, Cr(VI), and residual O2 are assumed to be competitive. In our model, we have left the specific photolysis rate of FeOx (reaction M1) as an adjustable parameter. It was well determined from fitting of the experiments, with deviations from the optimal values by (10% resulting in higher Χ2 and in visibly worse fits. The specific photolysis rate of FeOx divided by the specific rate of photon
absorption by FeOx, ka (see Table 1), gives the quantum yield of FeOx photolysis. We obtained a value of 0.76 with an estimated error range of (20%, in the upper range expected from the comparison with actinometer values (0.56-0.63). M4-M5 and M4b-M5b: Reaction of Fe(II) with Cr(VI) in the Presence of Oxalate. The two reaction mechanisms (M4M5 and M4b-M5b) result in the same overall rate law for the reaction of Cr(VI) with Fe(II) in the presence of oxalate and explain the kinetic dark experiments shown in Figure 6 equally well. The apparent rate constant in eq 14 is the product of KM4kM5 or KM4bkM5b, respectively. There are other models involving Cr(V) and Cr(IV) intermediate equilibria that result in the same overall rate law, but they are chemically less sensible. We cannot currently differentiate between the various possibilities. Model M4M5 is attractive because it states that the reaction of Fe(II)Ox with Cr(VI) is much faster than the reaction with Fe(II). This is expected, because Fe(II)Ox is a much better reductant than Fe(II). Model M4b-M5b is also plausible. Given the similar one electron potentials for Cr(VI)/Cr(V) and for Fe(III)/Fe(II), an equilibrium between a Cr(V)-Fe(III) intermediate and educts is a reasonable proposition. Oxalate would mediate the dissociation of such an intermediate by stabilizing Fe(III). Espenson (17) suggests a similar mechanism with a Cr(V)-Fe(III) intermediate, which dissociates involving one or two protons. In any case, an overall mechanisms where Cr(VI), Fe(II), and oxalate are part of the expression is essential. Models where only Cr(VI) and Fe(II) react in a bimolecular reaction cannot explain the experiments in Figures 4 and 6. M6: Reaction of Cr(VI) with HO2/O2•-. The reaction rate constant was estimated from fitting of all experiments in Figure 4a-c simultaneously. This reaction was needed for good fits, but the rate constant obtained depends on reaction rate constants chosen for Cr(V) and Cr(IV) intermediates. Good fits could be obtained with values between 5 × 104 and 8 × 104 M-1 s-1, slightly less than the reported rate (1 × 105 M-1 s-1) for the reaction of Cr(VI) with HO2 at pH 1 (33). At pH 5, we have 38% HO2 and at pH 6 only 4%. The model also predicts the experimental findings well at pH 6, with the rate constants for Cr(VI) reduction by Fe(II) and HO2/O2•- adjusted by multiplication with 0.7 (fraction of HCrO4- at pH 6). If only HO2 reacted with HCrO4-, the rate for the reaction with HO2/O2•- at pH 6 would be 9-10 times slower than at pH 5. Such low values, however, lead to bad fits of the experiments. These findings suggest that both HO2 and O2•- are important reductants for Cr(VI) in our aerated systems between pH 5 and 6. M7: Reaction of Cr(VI) with CO2•-. This rate constant is very tentative because of residual O2 in the experiments. At high [Cr(VI)], where residual O2 is less important, the model is insensitive to the rate constant for Cr(VI) reduction by CO2•because the photolysis of FeOx is rate determining. For very low Cr(VI) concentrations, it was not possible to achieve sufficient exclusion of O2. The effect of deaeration was most noticeable at 20 µM Cr(VI) and 0.67 µM Fe(III). To explain the faster Fe(II) reduction in deaerated solutions, inclusion of reaction M7 was needed. M8-M11: Reactions of Cr(V) and Cr(IV) with Fe(II), CO2•-, HO2/O2•-, and H2O2. These reactions were assumed to be fast, and we thus required that the model was insensitive to the absolute value of their rate constants. Variation from 1 × 106 M-1 s-1 up to diffusion-controlled (1010 M-1 s-1) did not affect the model output as long as their relative values were not changed. A change of relative rate constants for the reactions involving Fe(II) and the ones involving HO2/O2•and H2O2 affected the rates of Cr(VI) reduction by changing the steady-state concentration of Fe(II) and the rate for reformation of FeOx. The choice of the relative rate constants influenced the rate constants obtained for reactions M6 and
M7, but not more than by a factor of 2, as Fe(II) and HO2/O2•must be produced at equal rates in aerated solutions and H2O2 is not allowed to accumulate. X represents Ox or an unstable coordination state that ultimately leads to a coordination with Ox. It is not known at which state coordination with Ox occurs. M12-M14: Reactions of Fe(II) in the Presence of Oxalate. We used rate constants reported by Sedlak and Hoigne (16). In accordance with their observations, we neglected reduction of Fe(III)Ox by HO2/O2•-. M15-M17: Reactions with OH•. These reactions were needed only for a good fit of the one experiment with 20 µM Cr(VI) and 2µM Fe(III) in aerated solution. In only this experiment, H2O2 might be produced faster than it is oxidized by the Cr(V) intermediate, and thus the Fenton reaction with Fe(II) produces significant amounts of OH• radicals. We assume that OH• radicals react indiscriminantly with the most abundant species (the acetate buffer) and that a product is formed which is able to reduce Cr(VI). Omission of these reactions did not significantly change the fit for all other experiments. M18-M22: Reaction of Reactive Oxygen Species. These reactions are described in the cited literature and in the footnotes under the kinetic model. Despite of the numerous assumptions and estimated rate constants, the model provides a framework for the rationalization of the measured reactions under diverse conditions: (a) With [Cr(VI)]0 ) 200 µM and [Ox] ) 200-1300 µM, the Fe(III)-photocatalyzed Cr(VI) reduction is zero-order in [Cr] and does not depend on [O2] because all photoproducts [Fe(II), CO2•-, and HO2/O2•-] react with Cr(VI) faster than they are produced photochemically. (b) With [Cr(VI)]0 ) 20 µM and [Ox]0 ) 100 µM, the reaction order changes from zero order to complex order because, at [Cr] < 10 µM, the thermal reactions of Cr(VI) with photoproducts [Fe(II), CO2•-, and HO2/O2•-] become rate determining. Cr(VI) reduction in deaerated systems is faster because CO2•- reduces Cr(VI) faster than HO2/O2•-. (c) At [Cr(VI)]0 ) 5 µM and [Ox]0 ) 25 µM, the thermal reactions of Cr(VI) with photoproducts are rate determining and the reaction order is complex. The effect of deaeration was small because the interference of residual O2 was more severe in these experiments where the concentrations of the other reactants were low. Residual O2 concentrations of 0.25 µM were sufficient to scavenge most CO2•-. The reduction speeds up considerably by increasing the oxalate concentration to 1 mM. While the reduction of Cr(VI) at [Ox]0 ) 25 µM must occur predominantly by HO2/O2•-, reduction of Cr(VI) by Fe(II) becomes important at [Ox]0 ) 1000 µM. pH Dependence. Cr(VI) reduction with Fe and oxalate decreases almost linearly with pH above pH 5.5 (Figure 8). An explanation for this pH dependence is difficult, because the speciation of Cr(VI), Fe(III), and the intermediates all change with pH. For example, while HCrO4- is the dominant species for Cr(VI) at pH 5, CrO42- is the major species at pH 7. It is possible that CrO42- is less reactive than HCrO4-. According to the speciation calculations shown in Figure 2, FeOx2 and FeOx3 are still the main species of Fe(III) at pH 7, so that photochemical Fe(II) production should occur at similar rates up to pH 7, if the quantum yields for Fe(II) production from FeOx are pH independent in this range. However, there are too many possible reasons for the observed pH dependence, and more studies will be needed to find a mechanistic explanation. Cr(VI) Reduction with Citrate. Citrate, like oxalate, forms complexes with Fe(III) that can be photolyzed to yield Fe(II) and organic radicals. We did not study the Fe(III)-catalyzed photochemical Cr(VI) reduction with citrate as detailed as with oxalate, but a comparison of the data with citrate (Figure 9) with the data with oxalate (Figures 3, 4, and 7) shows that the photochemical Cr(VI) reduction occurs at similar rates
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between pH 5 and pH 7. Analogous to the experiments with oxalate, no precipitation of chromium(III) hydroxides was observed. Instead, an organic Cr(III) complex with the oxidation product of citrate must have formed (with a corresponding absorbance increase in the UV spectrum at 280 nm, see Figure 9) and remained soluble at least for hours at pH values up to 7. Possible Consequences for Environmental Systems and Remediation. We would expect that many organic ligands that form photolyzable complexes with Fe(III) play a similar role as oxalate or citrate. Most aquatic systems contain a variety of Fe(III)-complexing, low molecular weight organic acids and fulvic and humic acids that can complex Fe(III) and engage in photochemistry similar to oxalate and citrate (15, 44, 45). Based on the present study, it seems likely that photochemical reduction of Cr(VI) in contaminated wetlands and in atmospheric water is important and that a variety of Cr(III) complexes are formed in this process. Addition of small amounts of Fe(II) or Fe(III) to contaminated DOC-rich surface waters might accelerate photochemical Cr(VI) reduction. Also, one could envision the addition of Fe(II,III) and oxalate (or similar ligands) to highly contaminated wastewaters for photochemical treatment. However, such approaches should not be undertaken before the fate of Cr(III)oxalate and other Cr(III)-organic complexes in the environment is understood. Unlike in the direct reaction of Cr(VI) with Fe(II), where insoluble chromium(III) or mixed chromium(III)/iron(III) hydroxides are formed, the Cr(III)-organic complexes that form during the photochemical reduction are soluble and thus mobile, and they might be more easily re-oxidized to Cr(VI) than solid chromium hydroxides or chromium/iron hydroxides. Similar concerns might be important for non-photochemical reductions, when Fe(II) is added as a reductant for Cr(VI) in soils and sediments with high DOC contents.
Acknowledgments We thank Barbara Sulzberger, Ju ¨ rg Hoigne´, Werner Stumm (EAWAG), Carrick Eggleston (University of Wyoming), and David Sedlak (UC Berkeley) for useful discussions.
Literature Cited (1) Calder, L. M. Chromium in the Natural and Human Environments; Nriagu, J. O., Nieboer, E., Eds.; Advances in Environmental Science and Technology 20; Wiley Interscience: New York, 1988; pp 215-229. (2) National Research Council, Committee on Biologic Effects of Atmospheric Pollutants. Chromium; National Academy of Science: Washington, DC, 1974. (3) Anderson, R. A. Sci. Total Environ. 1989, 86, 75-81. (4) Bartlett, R. J., James, B. R. In Mobility and bioavailabilty of chromium in soils; Nriagu, J. O., Nieboer, E., Eds.; Wiley: New York, 1988; pp 267-301. (5) Fendorf, S. E. Geoderma 1995, 67, 55-71. (6) Wittbrodt,P. R.; Palmer,C. D. Environ. Sci. Technol. 1995, 29, 255-263. (7) Johnson, C. A.; Sigg, L.; Lindauer, U. Limnol. Oceanogr. 1992, 37, 315-321. (8) Deng, B.; Stone, A. T. Environ. Sci. Technol. 1996, 30, 463-472. (9) Powell, R. M.; Puls, R. W.; Hightower, R. K.; Sabatini, D. A. Environ. Sci. Technol. 1995, 29, 1913-1922.
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(10) Kieber, R. J.; Helz, G. R. Environ. Sci. Technol. 1992, 26, 307312. (11) Kaczynski, S. E.; Kieber, R. J. Environ. Sci. Technol. 1993, 27, 1572-1576. (12) Miles, C. J.; Brezonik, P. L. Environ. Sci. Technol. 1981, 15, 10891095. (13) Siffert, C; Sulzberger, B. Langmuir 1991, 7, 1627-1634. (14) Zuo, Y; Hoigne, J. Environ. Sci. Technol. 1992, 26, 1014-1022. (15) Voelker, B. M., Sulzberger, B. Environ. Sci. Technol. 1996, 30, 1106-1114. (16) Sedlak, D. L.; Hoigne´, J. Atmos. Environ. 1993, 27A, 2173-2185. (17) Espenson, J. H. J. Am. Chem. Soc. 1970, 92, 1880-1883. (18) James B. R.; Bartlett, R. J. J. Environ. Quality 1983, 12, 173-176. (19) Johnson, C. A.; Xyla, A. G. Geochim. Cosmochim. Acta 1991, 55, 2861-2866. (20) Silvester, E.; Charlet L.; Manceau, A. J. Phys. Chem. 1995, 99, 16662-16669. (21) Leenheer, J. A.; Wershaw, R. L; Reddy, M. M. Environ. Sci. Technol. 1995, 29, 393-398. (22) Leenheer, J. A.; Wershaw, R. L.; Reddy, M. M. Environ. Sci. Technol. 1995, 29, 399-405. (23) LI-1800UW Underwater Spectroradiometer LI-COR, Inc., 4421 Superior Street, P.O. Box 4425, Lincoln, NE 68504. (24) Press, W. H.; Flannery, B. P.; Teukolsky, S. A.; Vetterling, W. T. Numerical recipies in C; Cambridge University Press: Cambridge, 1988. (25) Vince, L.; Papp, S. J. Photochem. 1987, 36, 289-296. (26) Martell, A. E.; Smith, R. M. Critical Stability Constants; Plenum Press: New York, 1977. (27) Hatchard, C. G.; Parker,C. A. Roy. Soc. London Proc. 1956, Ser A, 235, 518-536. (28) Vincze, L.; Papp, S. J. Photochem. 1987, 36, 289-296. (29) Schlafer, H. L.; Gausmann, H.; Tausch, W. Z. Phys. Chem. 1962, 34, 8. (30) Braun, W.; Herron J. T; Kahaner, D. ACUCHEM, Computer Program for Modeling Complex Reaction Systems; National Bureau of Standards: Gaithersburg, MD 20899. (31) Matlab; The Math Works, Inc.: Natick, MA. (32) Bard, A. J.; Parsons. R.; Jordan J. Standard Potentials in aqeuous solutions; Marcel Dekker Inc.: New York, 1985. (33) Sharpe, P. H. G.; Sehested, K. Radiat. Phys. Chem. 1989, 34, 763-768. (34) Seigneur, C.; Constantinou, E. Environ. Sci. Technol. 1995, 29, 222-231. (35) Adams, G. E.; Wilson, R. L. Trans. Faraday Soc. 1969, 65, 2981. (36) Rush, J. D.; Bielski, B. H. J. J. Phys. Chem. 1985, 89, 5062-5066. (37) Hartwick, T. J. Can. J. Chem. 1957, 35, 428-436. (38) Getoff, N.; Schwoerer, F.; Markovic, V. M.; Sehested, K.; Nielsen, S. O. J. Phys. Chem. 1971, 75, 749-755. (39) Bielski, B. H. J.; Cabelli, D. E.; Arudi, R. L.; Ross, A. B. J. Phys. Chem. Ref. Data 1985, 14, 1041-1100. (40) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. J. Phys. Chem. Ref. Data 1988, 17, 513-886. (41) Fisher, M. M.; Hamill, W. H. J. Phys. Chem. 1973, 77, 171-177. (42) Schuchmann, M. N.; Zegota, H.; von Sonntag, C. Z. Naturforsch. B, Anorg. Chem., Org. Chem. 1985, 40B, 215-21. (43) Mulazzani, Q. G.; D’Angelantonio, M.; Venturi, M.; Hoffmann, M. C.; Rodgers, A. J. J. Phys. Chem. 1986, 90, 5347-5352. (44) Zuo, Y.; Hoigne´, J. Atmos. Environ. 1994, 28, 1231-1239. (45) Faust, B. C.; Zepp, R. G. Environ. Sci. Technol. 1993, 27, 25172522.
Received for review March 18, 1996. Revised manuscript received August 7, 1996. Accepted August 30, 1996.X ES960253L X
Abstract published in Advance ACS Abstracts, November 1, 1996.
