Oxidation Kinetics of Fe(II) in a Eutrophic Swiss Lake

mM, [H2O2])20-50 nM, initial [Fe(II)])30 nM. Fe(II) concentrations ..... (16) Waite, T. D.; Szymaczak, R.; Espey, Q. I.; Furnas, M. J. Mar. Chem. 1995...
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Environ. Sci. Technol. 1998, 32, 2990-2996

Oxidation Kinetics of Fe(II) in a Eutrophic Swiss Lake L U K A S E M M E N E G G E R , * ,† D. WHITNEY KING,‡ LAURA SIGG,† AND BARBARA SULZBERGER† Swiss Federal Institute for Environmental Science and Technology (EAWAG), 8600 Du ¨ bendorf, Switzerland, and Department of Chemistry, Colby College, Waterville, Maine 04901

The rate of oxidation of ferrous iron was measured in samples from Lake Greifen, a eutrophic lake in Switzerland. Fe(II) concentrations were followed using an automated flow injection analysis system employing luminolbased chemiluminescence detection of Fe(II). For kinetic studies at pH > 7.8, the system was modified to allow a time resolution of less than 1 s. Oxidation rates were measured in unfiltered samples at 2-2000 nM Fe(II). The pH was varied between 6.8 and 8.3 by bubbling with CO2 and synthetic air. Above pH 7.4, rates were consistent with the rate law determined in pure carbonate systems. Between pH 6.8 and pH 7.3, however, the apparent rate was independent of pH. This surprising finding may be explained by some naturally occurring (organic, colloidal, or surface) ligand(s) that accelerate the oxidation of Fe(II). The relative importance and pH dependence of the direct reaction with O2 in comparison to that with H2O2 was determined, and the enhancement of the overall rate was attributed to the reaction of Fe(II) with oxygen.

Fe(III) and reactive oxygen species. The redox cycling of iron has been intensively studied in atmospheric waters (e.g., refs 13-15), and in marine and freshwater systems over the past 15 years (e.g., refs 9 and 16). The kinetics of Fe(III) reduction and of Fe(II) oxidation are of crucial importance in determining the speciation and thus the bioavailability of iron, since most photosynthetic aquatic organisms can take up iron only in the dissolved form (17, 18). In comparison to the oceans, the concentration of dissolved organic carbon is generally higher in lakes, especially in eutrophic waters. The kinetics of both Fe(III) reduction and Fe(II) oxidation may therefore differ in freshwater systems from that in the oceans. Regarding Fe(II) oxidation kinetics, King and co-workers (19) have determined rates of Fe(II) oxidation in bicarbonate solutions at nanomolar total iron concentrations using a flowinjection chemiluminescence system based on the Fe(II)catalyzed reaction between luminol and O2. Hitherto, the Fe(II) oxidation kinetics at natural (i.e., nanomolar) Fe(II) concentrations in water samples from lakes have not been investigated. In freshwater samples, inorganic and organic ligands and colloids may greatly affect Fe(II) oxidation rates over a broad pH range. In this study, oxidation rates of Fe(II), added to lake water samples, were followed in a concentration range of 0.5-2000 nM using an automated luminol-based chemiluminescence system (19). The system was adapted to allow measurement intervals of 9, Fe(II) is rapidly oxidized by oxygen on a millisecond time scale catalyzing the oxidation of luminol and producing blue light. The chemiluminescence system was used in two different configurations. At sampling rates of less than 1 sample/ min, a conventional injection valve was used to introduce Fe(II) samples into a carrier flow stream. The integral of the PMT signal was employed to determine Fe(II). At pH > 7.8, conventional flow injection analysis was not rapid enough to follow Fe(II). The system was therefore modified to bypass the injection valve and directly introduce the sample into VOL. 32, NO. 19, 1998 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 1. Fe(II) oxidation shows pseudo-first-order kinetics in unfiltered lake water at pH 6.80-8.26. T ) 25 °C, alkalinity ) 3.88 mM, [H2O2] ) 20-50 nM, initial [Fe(II)] ) 30 nM. Fe(II) concentrations are represented on a natural log scale. [Fe(II)] at pH 8.26 and pH 7.99 was measured with a time resolution of 1 s and plotted without linear regression. The solid lines are linear regressions with a single regression employed for the three lowest pH values. the luminescence flow cell by placing the carrier uptake tube directly in the sample. This configuration gave a continuous PMT signal proportional to the Fe(II) concentration. The time resolution in this case was less than 1 s. For rapid oxidation experiments, the baseline was sufficiently constant to allow this modification of the original setup. Calibrations were performed with the instrument set to normal FIA mode and continuous mode in lake water at pH 6.0-7.2. The calibration curves were linear in the concentration range of this study, with a detection limit below 1 nM. All concentrations given are relative to the background signal before the addition of Fe(II), which was always less than 10% of the maximum signal in oxidation experiments. Samples. All samples were collected at the point where Lake Greifen, a eutrophic lake in Switzerland, is deepest. They were taken in February/March 1997 with a Go-Flow bottle, 2 m below the surface, and stored for a maximum of 12 days at 4 °C before oxidation experiments. DOC was 3.2 mg/L, alkalinity was 3.8 mM, pH was 8.3, and H2O2 was 2040 nM. Ionic strength of water from Lake Greifen is about 7 mM. Unfiltered samples were used where not stated otherwise.

Oxidation Experiments. Fe(II) oxidation experiments were performed in a 300-mL Pyrex water-jacketed beaker by adding a standard solution of Fe(II), usually 10 mM Fe(NH4)2(SO4)2 in 1 mM HCl, to the O2-saturated unfiltered water samples. The initial Fe(II) concentration was 30 nM where not stated otherwise. pH was adjusted by bubbling a mixture of CO2, N2, and 20.9% O2 through the system. The gas stream passed through a MnO4- solution to eliminate any H2O2. A stable pH was always reached within less than 10 min. Solution temperatures were maintained at 25.0 ( 0.2 °C. UV Irradiation. Samples were irradiated with light from a mercury lamp for 6 h in a photooxidation unit (La Jolla Scientific Co., La Jolla, CA). Photoproduced H2O2 was eliminated by treatment with 1000 U/L catalase for 1-2 h at 4 °C. DOC reduction due to UV irradiation was approximately 85%.

Results and Discussion Oxidation Rate as a Function of pH. Unfiltered samples showed pseudo-first-order kinetics at a given pH (6.8-8.3) as expected from the general rate law (eq 8) (Figure 1). Figure 2 shows log (kapp) vs pH with a slope of 1.7 ( 0.1 above pH 7.3, which agrees reasonably well with literature data. However, kapp was independent of pH between pH 6.8 and pH 7.3. This is in contrast to numerous studies in natural samples confirming the rate law (eq 8). For comparison, we have plotted on the same graph the range given by a ‘universal’ rate constant, as suggested by Davison (24), and the results of the model developed by King (26), which considers not only the Fe(II) hydroxo but also the carbonato species and their reactivity. These sets of data should be compared with caution because, in contrast to the above studies, we have not reached a steady-state H2O2 concentration in our system. The data seem nevertheless consistent with the literature for the upper pH range while the trend in the lower pH range appears to be in contradiction. Ligand Effect on Oxidation Rate. A series of experiments with increasing initial Fe(II) concentrations ([Fe(II)]o ) 102000 nM) were done at pH 6.9. Since high amounts of initial Fe(II) might change H2O2 concentration, we added 500 U/L catalase 10 min before starting the experiment. Catalase (hydrogen peroxide oxidoreductase) catalyzes the transformation of H2O2 into H2O and O2. Its addition will eliminate oxidation by H2O2 (eq 3) independently of initial Fe(II)

FIGURE 2. Apparent rate of Fe(II) oxidation in unfiltered lake water as a function of pH. T ) 25 °C, pO2 ) 0.209 atm, alkalinity ) 3.88 mM, [H2O2] ) 20-50 nM, initial Fe(II) ) 30 nM. Also plotted, assuming steady-state concentrations of [H2O2], are the range for kapp given by Davison (24) and kapp calculated using the model by King (26). 2992

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FIGURE 3. Fe(II) oxidation as a function of initial [Fe(II)] at pH 6.9. 500 U/L catalase was added 10 min before addition of Fe(II), thus limiting the oxidation mechanism to eqs 1-2 (k′O2).

FIGURE 4. k′O2 as a function of initial [Fe(II)] at pH 6.9. Experimental data (b) and model fit (s). 500 U/L catalase was added 10 min before addition of Fe(II), thus limiting the oxidation mechanism to eqs 1-2 (k′O2). Inset: Fe(II) oxidation of 1 µM inital [Fe(II)] with a time resolution of 1 s. This part of the oxidation experiment does not show pseudo-first-order behavior. Experimental data (b) and model fit (s). concentration and thus limit oxidation to the reactions with O2 and O2•- (eqs 1 and 2). The minimum activity of the enzyme in our system was estimated from H2O2 determinations where catalase is in competition with horseradish peroxidase of known activity when determining blank values. Catalase is largely specific for H2O2 as the hydrogen acceptor. After the rate-limiting formation of an active catalase-H2O2 complex, other substances can compete with H2O2 to serve as hydrogen donors (42). These side reactions are, however, either slow or of no importance to Fe(II) oxidation kinetics. We never observed any effects that could not be explained by elimination of H2O2. We treated the data of the Fe(II) oxidation experiments with increasing [Fe(II)]o as pseudo-first-order (see Figure 3). The resulting t1/2 and kapp for low initial Fe(II) concentrations agree well with the first set of data we obtained at similar pH. The striking result, however, is the decrease of k′O2 with the increase of initial Fe(II).

Oxidation of Fe(II) in our natural samples is much faster at pH 6.8-7.3 (Figure 2) than expected from various other studies that were performed in synthetic samples and/or at micromolar Fe(II) concentrations. It seems reasonable to postulate that some (organic, colloidal, or surface) ligand (L) is accelerating the reaction of Fe(II) with O2. In a natural system, L might represent a variety of different ligands, which we will refer to simply as “ligand”. Independent of its type, we might however assume that the amount of L is limited and, if it is not acting purely as a catalyst, will be blocked as Fe(II) is oxidized. To further explore our findings, we developed a very simple model that describes the oxidation of Fe(II) by O2 via two parallel pathways. The first pathway involves the inorganic (hydroxo and carbonate) Fe(II) species (Fe(II)i) with a conditional rate constant ki. The second pathway considers an Fe(II)-L complex, with the corresponding conditional complex binding constant K and the rate constant kL. The VOL. 32, NO. 19, 1998 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 5. kapp for [H2O2] . [Fe(II)] at pH 7.15-8.33. The slope of each linear regression line corresponds to k′H2O2 and the y-intercept equals [O2] kO2. ligand pathway dominates oxidation at low pH. Observed total Fe(II) oxidation rates are described by eqs 9-11.

K)

-

[Fe(II) - L] [Fe(II)][L]T

(9)

d[Fe(II)T] ) ki[Fe(II)i] + kL[Fe(II)-L] dt

(10)

d[Fe(II)T] ) [Fe(II)](Riki + RLkL) dt

(11)

-

where Ri is the fraction of hydroxo and carbonate complexes, and RL is the fraction of Fe(II)-L. The overall inorganic rate constant [log (ki) ) -1.8] for pH 6.9 at 25 °C was obtained using the kinetic model published by King (26) and dividing this value by 2 to account for the elimination of H2O2. Pseudo-first-order kinetics are obtained for [Fe(II)] . [L] or [Fe(II)] , [L] and can be used to fit the model to the experimental data shown in Figure 4. Large initial Fe(II) concentrations lead to a rapid decrease of available L by formation of Fe(III)-L, and the decrease of Fe(II) as a function of time will therefore not be pseudo-

first-order until [L] becomes much smaller than [Fe(II)]. This part of the oxidation experiment has been followed with a time resolution of 1 s for 1 µM initial Fe(II) (inset Figure 4) and further constrains the fitting parameters. The simulation was done in a Microsoft Excel spreadsheet with time steps of 1 s using Excel’s Solver capability to minimize the residual sum of squares between experimental and modeled data. Using the above approach to simulate the observed oxidation kinetics, we obtained a ligand concentration of 380 ( 150 nM (compare [Fe(II)]o - [Fe(II)]6min in inset Figure 4). The conditional complex binding constant of Fe(II)-L and the corresponding oxidation rate constant are interdependent fitting parameters in the model; therefore, only the product of the two is relatively well defined with log (kL) + log K ) 9.6 ( 0.3. To gain further insight with respect to the type of ligand that is responsible for the enhanced oxidation rate at low pH, we performed a series of filtration and UV irradiation experiments. Filtration through a 0.45-µm cellulose nitrate membrane filter of one sample, taken in March 1997, reduced the oxidation rate (kapp) at pH 6.9 by a factor of 2. The same procedure on a sample taken in November reduced kapp 5-fold. From these preliminary experiments, one could conclude that surface or bacterially catalyzed reactions were more important in the fall sample. They do not, however, account for the full enhancement of the oxidation rate we would expect from the experiments done at high initial Fe(II) concentration or from extrapolation of the linear trend in log(kapp) vs pH at pH > 7.3 (Figure 2). UV irradiation experiments of the filtered fall sample showed no significant effect on the oxidation rate. We must also stress that the samples taken in November had very high concentrations of total Mn and Fe (several 100 nM filterable Fe and Mn), transported into the epilimnion by lake overturn. While catalysis by oxide surfaces seems likely, autocatalytic reaction on iron oxide, freshly precipitated during the experiment, can be neglected due to the small amounts of Fe(II) added. Further experiments need to be done to elucidate the nature of the ligand. Since the acceleration of Fe(II) oxidation is dependent on total Fe(II) added, we can nevertheless conclude that most of the ligand is inactivated and therefore

FIGURE 6. log(kH2O2) and log(kO2) as a function of pH. The data were interpolated by a linear fit for log(kH2O2) and a second-order polynom for log(kO2). In gray are the values obtained for kH2O2 by Millero (38) (I ) 7 mM, T ) 25 °C, HCO3- ) 3.8 mM) multiplied by 0.5 to account for the scavenging of HO• (see Figure 7). The range shown was determined by the standard deviation of the bicarbonate correction for salinity ) 0‰. 2994

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FIGURE 7. kapp as a function of initial [Fe(II)]. Conditions (pH, [H2O2], catalase) were chosen to strongly favor either k′O2 or k′H2O2. The observed k′O2 is independent of initial [Fe(II)], indicating that O2•is predominantly reacting with Fe(II). In contrast, k′H2O2 increases by 100% with increasing initial [Fe(II)], indicating that HO• reacts with other components rather than Fe(II) at very low Fe(II) concentrations in our natural samples. not catalytic. The calculated concentration of 380 nM is, however, rather large and may indicate unspecific organic ligands. Determination of kH2O2 and kO2. To determine kH2O2 and kO2, we performed a series of experiments with H2O2 ranging from 100 to 1000 nM for each of five different pH values between 7.1 and 8.3. The samples were brought to the desired [H2O2] by adding standard solutions of H2O2; the final H2O2 concentration being determined immediately before adding Fe(II). O2 was maintained constant in our system by continuous bubbling with the gas mixture. Since [Fe(II)] was 100 nM for all experiments, H2O2 can be considered constant during each experiment. We determined kapp as a function of pH and H2O2 concentration. The results are shown in Figure 5. Figure 6 shows that the rate kH2O2 is linearly dependent on pH, while kO2 exhibits a distinct reduction of slope around pH 7.3, which is consistent with the pattern found for the apparent overall rate (Figure 2). For comparison, we also plotted on the same graph the values obtained for kH2O2 by Millero (38). The rate was corrected for ionic strength, temperature, and bicarbonate concentration and multiplied by 0.5 to account for the competing reactions that scavenge HO• in our natural system (see following section). Extrapolation of our values to pH 6.0, where the bicarbonate correction was determined by Millero, gives excellent agreement between the two sets of data. A slight difference in the upper pH range is not surprising because the effect of

bicarbonate would not be expected to be independent of pH. As shown in Figure 6, acceleration of the oxidation rate below pH 7.3 can only be observed for kO2 and not for kH2O2. For the reaction with O2 (eq 1) at pH < 7.3 we can therefore assume that kLRL > kiRi. In contrast, kH2O2 exhibits the expected linear trend. If Fe(II)-L were a major species, this would imply that kLRL shows the same dependence on pH as does kiRi for the reaction with H2O2. Furthermore, this dependence would have to be very different for the reaction with O2. Since this is highly unlikely, it implies that Fe(II)-L is only a trace species and that exchange rates for L are fast as compared to the oxidation rates. If Fe(II)-L represents less than about 10% of total Fe(II), then it does not significantly influence kH2O2. Since [L] is well defined by the model and experiments presented above, this condition further restrains the conditional complex binding constant to log K < 5.4 and thus log (kL) > 4.4 with log(kL) + log K ) 9.8 ( 0.3. The Fate of HO• and O2•-. One issue is whether the intermediate products O2•- and HO• remain effective Fe(II) oxidants in the natural system (eqs 2 and 4). Both oxygen species are known to react with a variety of inorganic and organic substances. This is especially true for HO• which reacts rapidly and unselectively with many natural species (43). The oxidation rate of Fe(II) with O2•- and HO• should be dependent on [Fe(II)] if other substances are competing for the same oxidants. We have tested this effect by varying the initial Fe(II) concentration and choosing conditions under which either k1 and k2 (eqs 1 and 2) or k3 and k4 (eqs 3 and 4) would dominate the reaction. To limit oxidation of Fe(II) to the reaction with O2 and O2•-, we added 10 000 U/L catalase and followed [Fe(II)] at pH 8.0. H2O2 is always far below steady state under these conditions. When increasing Fe(II) concentrations from 2 to 1000 nM, we found no significant change of kapp (Figure 7). Provided that there is no sink for O2•- greater than the 1000 nM Fe(II) added to our system, then O2•- must react predominantly with Fe(II), even at nanomolar Fe(II) concentrations. To test, if HO• reacts with Fe(II) in our natural system, we chose conditions (pH 7.25 and 4 µM of added H2O2) under which oxidation is largely dominated by k′H2O2. Increasing [Fe(II)] from 2 to 1000 nM under these conditions results in doubling k′H2O2. This indicates that other substances compete successfully for HO• if Fe(II) concentration is very low (Figure 7). Our results indicate that oxidation of Fe(II) by HO• is not important at natural concentrations. The fate of HO• might also be predicted considering the reaction with DOM, HCO3-, and Fe(II) that have extensively

FIGURE 8. Main reaction pathways and lifetimes (τ) of HO• with HCO3-, DOC, and Fe(II). HCO3• is predominantly produced as a first step. The dominant reaction of HCO3• depends then on [Fe(II)]; for [Fe(II)] ) 10 nM, the radical is scavenged by DOC, while oxidation of iron dominates largely at [Fe(II)] ) 1000 nM. Literature: (1) ref 44; (2) ref 44; (3) ref 47; (4) estimated rate, see text; (5) ref 44. VOL. 32, NO. 19, 1998 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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been studied in the past, i.e., ref 44. At neutral pH and millimolar alkalinity, the formation of bicarbonate radicals and organic radicals are particularly favorable (see Figure 8). In our system, the formation of HCO3• is slightly faster than the production of DOM•+ which, additionally, will partly react to HCO3•. The issue therefore becomes whether HCO3• will react with DOM or Fe(II). Lacking literature data for the reaction of Fe(II) with HCO3•, we estimated this rate constant (eq 12) to be 5 × 108 M s-1, equal to the reaction of Fe(II) with HO•.

Fe(II) + HCO3• f Fe(III) + HCO3-

(12)

For our experimental conditions (DOC ) 3.2 mg/L and [Fe(II)]o increasing from 10 to 1000 nM), the main reaction pathway for HCO3• is expected to shift from DOM to Fe(II) based on the published rates. This is perfectly consistent with the results presented in Figure 7 and discussed above.

Significance of Results The results obtained here have important implications for the redox cycling of iron in lakes. The oxidation rate constants determined allow us to evaluate the half-life of Fe(II) under typical lake water conditions and also to estimate the relative importance of the direct reaction with O2 in comparison to that with H2O2. Typical conditions in summer in the epilimnion of Lake Greifen are pH 8.0-8.5, O2 ≈ 3 × 10-4 M, and H2O2 ) 50-200 nM. Under these conditions, the halflife of Fe(II) ranges between 7 and 60 s. The reaction with O2 is always faster than the one with H2O2; the ratio of k′O2 to k′H2O2 varies between approximately 30 (pH ) 8.5, [H2O2] ) 50) and 2 (pH ) 8.0, [H2O2] ) 200). The formation of Fe(II) by reduction processes and its reoxidation to Fe(III) are very significant reactions with regard to the availability of iron to algae, because these processes lead to the formation of easily available dissolved Fe(II) and Fe(III) species. The present results indicate that the oxidation rate of Fe(II) is reasonably predicted from experiments in synthetic solutions at pH > 7.5, taking into account the effects of hydroxo and carbonato complexes. At lower pH, however, here in the range of pH 6.8-7.3, the higher than predicted oxidation rate indicates the effect of additional ligands that are not present in simple synthetic solutions. The nature of these ligands has not yet been fully elucidated. The hypothesis that organic ligands are accelerating the oxidation of Fe(II) implies the formation of organic-Fe(II) complexes and in turn organic-Fe(III) complexes. The occurrence of strong organic-Fe(III) complexes has been recently demonstrated in the oceans (45). Fe(III) from organic complexes may become available to algae by cell-surface reduction processes, by photochemical processes, or by ligandexchange processes at the cell surface (46). Oxidation of Fe(II) on oxide surfaces would in contrast reduce the bioavailability of iron. The oxidation rate determined here is one of the key kinetic parameters to evaluate the cycling of Fe(II) and Fe(III) in the epilimnion of lakes. We plan to also determine the Fe(III) reduction rate and steady-state concentrations of Fe(II) under similar conditions to gain a better understanding of this redox cycle and its implications in natural water systems.

Acknowledgments We thank Werner Stumm for his continuous interest and Stephan Hug for critical discussions. David Kistler was always ready to help with sampling and infrastructure in the lab. This work is supported by Swiss National Science Foundation.

Literature Cited (1) Martin, J. H.; Tanner, S. J.; Barber, R. T.; Liddicoat, M. I.; Bidigare, R.; Zhang, J. Z.; Greene, R.; Yungel, J. Nature 1994, 371, 123129. 2996

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(2) Martin, J. H.; Gordon, R. M. Deep-Sea Res. 1988, 35 (2), 177196. (3) Martin, J. H.; Fitzwater, S. E. Nature 1988, 331, 341-343. (4) Davison, W. Earth Sci. Rev. 1993, 34, 119-163. (5) Sigg, L.; Johnson, C. A.; Kuhn, A. Mar. Chem. 1991, 36, 9-26. (6) Wells, M. L.; Mayer, L. M. Deep-Sea Res. 1991, 38, 1379-1395. (7) King, D. W.; Aldrich, R. A.; Charnecki, S. E. Mar. Chem. 1993, 44, 105-120. (8) Voelker, B. M.; Sedlak, D. L. Mar. Chem. 1995, 50, 93-102. (9) Voelker, B. M.; Morel, F.; Sulzberger, B. Environ. Sci. Technol. 1997, 31, 1004-1011. (10) Miller, W. L.; King, D. W.; Lin, J.; Kester, D. R. Mar. Chem. 1995, 50, 63-77. (11) Sulzberger, B.; Schnoor, J. L.; Giovanoli, R.; Hering, J. G.; Zobrist, J. Aquat. Sci. 1990, 52, 57-74. (12) McKnight, D. M.; Kimball, B. A.; Bencala, K. E. Science 1988, 240, 637-640. (13) Behra, P.; Sigg, L. Nature 1990, 344, 419-421. (14) Kotronarou, A.; Sigg, L. Environ. Sci. Technol. 1993, 27, 27252735. (15) Sedlak, D. L.; Hoigne´, J. Atmos. Environ., Part A 1993, 14, 21732185. (16) Waite, T. D.; Szymaczak, R.; Espey, Q. I.; Furnas, M. J. Mar. Chem. 1995, 50, 79-91. (17) Anderson, M. A.; Morel, F. M. M. Limnol. Oceanogr. 1982, 27, 789-813. (18) Sunda, W. G.; Huntsman, S. A. Limnol. Oceanogr. 1996, 41, 373387. (19) King, D. W.; Lounsburry, H. A.; Millero, F. J. Environ. Sci. Technol. 1995, 29, 818-824. (20) Weiss, J. Naturwissenschaften 1935, 4, 64-69. (21) Stumm, W.; Lee, G. F. Ind. Eng. Chem. 1961, 53, 143-146. (22) Singer, P. C.; Stumm, W. Science 1969, 167, 1121-1123. (23) Millero, F. J.; Sotolongo, S.; Izaguirre, M. Geochim. Cosmochim. Acta 1987, 51, 793-801. (24) Davison, W.; Seed, G. Geochim. Cosmochim. Acta 1983, 47, 6779. (25) Millero, F. J. Mar. Chem. 1988, 28, 1-18. (26) King, D. W. Environ. Sci. Technol. 1998, 32, 2997-3003. (27) Tamura, H.; Goto, K.; Nagayama, M. Corros. Sci. 1976, 16, 197207. (28) Wehrli, B. In Aquatic Chemical Kinetics; Stumm, W., Ed.; Wiley: New York, 1990; pp 311-337. (29) Sung, W.; Morgan, J. J. Environ. Sci. Technol. 1980, 14, 561-568. (30) Barry, R. C.; Schnoor, J. L.; Sulzberger, B.; Sigg, L.; Stumm, W. Water Res. 1994, 28, 323-333. (31) Harris, D. C.; Aisen, P. Biochim. Biophys. A 1973, 329, 156-158. (32) Liang, L.; McNabb, J. A.; Paulk, J. M.; Gu, B.; McCarthy, J. F. Environ. Sci. Technol. 1993, 27, 1864-1870. (33) Theis, T. L.; Singer, P. C. Environ. Sci. Technol. 1974, 8, 569573. (34) Cooper, W. J. Science 1983, 220, 711-712. (35) Walling, C. Acc. Chem. Res. 1975, 8, 125-131. (36) Goldstein, S.; Meyerstein, D.; Czapski, G. Free Radical Biol. Med. 1993, 15, 435-445. (37) Wardman, P.; Candeias, L. P. Radiat. Res. 1996, 145, 523-531. (38) Millero, F. J.; Sotolongo, S. Geochim. Cosmochim. Acta 1989, 53, 1867-1873. (39) Voelker, B. M.; Sulzberger, B. Environ. Sci. Technol. 1996, 30, 1106-1114. (40) Rush, J. D.; Bielski, B. H. J. J. Phys. Chem. 1985, 89, 5062-5066. (41) Miller, W. L.; Kester, D. R. Anal. Chem. 1988, 60, 2711-2715. (42) Bergmeyer, J. E.; Gras, M. In Catalase; Aebi, H. E., Ed.; Verlag Chemie: Weinheim, 1983-1986; Vol. III. (43) Hoigne´, J. In Process technologies for water treatment; Stucki, S., Ed.; Plenum Press: New York, 1988; pp 131-143. (44) Larson, R. A.; Zepp, R. G. Environ. Toxicol. Chem. 1988, 7, 265274. (45) Rue, E. L.; Bruland, K. W. Mar. Chem. 1995, 50, 117-138. (46) Rue, E. L.; Bruland, K. W. Limnol. Oceanogr. 1997, 42, 901-910. (47) Farahataziz; Ross, A. B. Selected specific rates of reactions of transients from water in aqueous solution, III, Hydroxyl radical and perhydroxyl radical and their radical ions; National Bureau of Standards: Washington, DC, 1977.

Received for review March 3, 1998. Revised manuscript received June 1, 1998. Accepted June 15, 1998. ES980207G