Superoxide Mediated Reduction of Organically Complexed Iron(III

Mar 30, 2007 - from the complex (dissociative reduction) or by direct reaction of superoxide with the ... superoxide can reduce iron(III) bound to a w...
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Environ. Sci. Technol. 2007, 41, 3205-3212

Superoxide Mediated Reduction of Organically Complexed Iron(III): Comparison of Non-Dissociative and Dissociative Reduction Pathways SHIKHA GARG,† ANDREW L. ROSE,‡ AND T . D A V I D W A I T E * ,†,‡ School of Civil and Environmental Engineering, and Centre for Water and Waste Technology, The University of New South Wales, Sydney 2052, Australia

We have investigated the mechanism of reduction of organically complexed iron(III) in the presence of superoxide, the one-electron reduced form of dioxygen that is produced in natural waters by thermal, photochemical, and biological pathways. Experimental results show that reduction of organically complexed iron(III) by superoxide may occur by either (or, in some instances, both) reaction of superoxide with inorganic iron(III) after its dissociation from the complex (dissociative reduction) or by direct reaction of superoxide with the complex (non-dissociative reduction). In the presence of low concentrations of ligands such as citrate and sulfosalicylate that bind iron(III) relatively weakly and result in complexes with high dissociation rate constants (kd > 1 × 10-4 s-1), a dissociative reduction pathway dominates. However, in the presence of strong ligands or high concentrations of weak ligands, only non-dissociative reduction of complexed iron(III) occurs. The relative contribution of each pathway has major implications for the lability and hence potential bioavailablity of iron in natural waters. The simple kinetic model developed here can be used to correctly predict the superoxide-mediated formation rates of iron(II) in natural systems.

Introduction Superoxide, a known reductant of iron(III) (1-4), may be produced in natural waters by abiotic thermal processes (e.g., iron(II) oxygenation or reduction of Cu(II) by hydrogen peroxide), photochemically, and biologically in natural waters. In an earlier study (1, 4), we have shown that superoxide can reduce iron(III) bound to a wide range of organic ligands and, in so doing, form substantial concentrations of iron(II) which, because it forms weaker complexes with organic ligands than iron(III), is more bioavailable. As biological production of superoxide appears relatively widespread in the phytoplankton community (5-8), superoxidemediated reduction of organically complexed iron(III) (Fe(III)L) could increase the bioavailability of iron to these and other organisms in marine environments where bioavailable iron is scarce. Consistent with this idea, it was shown (8) that superoxide production is the precursor step for the majority * Corresponding author phone: +61-2-9385 5060; fax: +61-29385 6139; e-mail: [email protected]. † School of Civil and Environmental Engineering. ‡ Centre for Water and Waste Technology. 10.1021/es0617892 CCC: $37.00 Published on Web 03/30/2007

 2007 American Chemical Society

of iron uptake by the benthic cyanobacterium Lyngbya majuscula. Two major mechanisms have been proposed for the production of iron(II) by superoxide mediated Fe(III)L reduction, which has been measured in earlier studies (1, 2). In the first mechanism, which we term Dissociative Reduction (DR), superoxide reduces inorganic iron(III) (Fe(III)′) to inorganic iron(II) (Fe(II)′) following dissociation of the Fe(III)′ from an organic complex (5, 9) (Figure 1, pathway A). In the second mechanism, which we term Non-Dissociative Reduction (NDR), direct reduction of the Fe(III)L by superoxide occurs (1), initially forming Fe(II)L that may then dissociate to give Fe(II)′ (Figure 1, pathway B). The steady-state concentration of total iron(II) (Fe(II)′ and Fe(II)L) formed by superoxide-mediated Fe(III)L reduction depends on the oxidation kinetics of either the Fe(II)′ (in the case of DR) or the Fe(II)L (in the case of NDR). As the oxidation kinetics of Fe(II)L can be either retarded or accelerated depending on the type of ligand (10-13), the nature of the ligand affects the oxidation kinetics and, as a result, the steady-state concentration of total iron(II). From a biological perspective NDR may be more important as it leads to conversion of non-labile Fe(III)L to labile Fe(II)′ (which occurs via formation of Fe(II)L) whereas DR results in conversion of labile Fe(III)′ to Fe(II)′. Thus, if we are to fully understand the factors controlling the steady-state concentration of total iron(II) and bioavailability of iron, it is critical to also understand the mechanism of superoxidemediated reduction of Fe(III)L and the effect of ligand type and concentration on the kinetics of this process. To investigate these issues, we have conducted a series of kinetic experiments designed to distinguish between DR and NDR by examination of the effect of the presence of superoxide on the rate of complexation of inorganic iron (both Fe(III)′ and Fe(II)′) by the fungal siderophore desferrioxamine (DFB), and complexati1on of Fe(II)′ by ferrozine (FZ).

Experimental Section Reagents. Reagents were prepared using 18 MΩ cm resistivity Milli-Q water unless stated otherwise. All experiments were performed in 2 mM NaHCO3 solution containing 10 mM NaCl at pH 8 (referred to as NaHCO3/NaCl hereafter). All pH measurements were made using a Hanna HI9025 pH meter calibrated daily using pH 7.01 and 10.01 buffers. Stock solutions of 2% HCl and 2% NaOH, prepared by dilution of high purity 30% w/v HCl (Sigma) and 30% w/v NaOH (Fluka puriss p.a. plus) respectively, were used for pH adjustment. A maximum pH variation of ( 0.05 units was allowed during experiments. An 8 mM stock solution of FeCl3‚6H2O (Ajax Chemicals) was prepared in 0.1 M HCl. A 4 mM stock solution of Fe(II) in 0.1 M HCl was prepared from ferrous ammonium sulfate hexahydrate. Experiments were performed with iron(III) that was complexed with the relatively strong ligands ethylenediaminetetraacetic acid (EDTA), 3,4-dihydroxy benzoic acid (DHB), Suwannee River fulvic acid (SRFA), and the relatively weak ligands citrate and sulfosalicylic acid (SSA). For EDTA and citrate, stock solutions of the iron(III) complexes were prepared at final iron concentrations of 200 µM, while for DHB, SRFA, and SSA 50 µM Fe stock solutions were prepared. The appropriate mass of ligand was dissolved in Milli-Q water along with the appropriate volume of iron(III) stock (250 or 62.5 µL for final Fe(III) concentrations of 200 µM and 50 µM, respectively) and the solution volume was made up to 10 mL in a volumetric flask. The pH of the stock solutions of each complex was then adjusted to 8.0. Stock solutions of 0.1 M VOL. 41, NO. 9, 2007 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 1. Reaction scheme showing the superoxide-mediated transformations of iron, as well as the various pathways of FeDFB production (when DFB is present) and Fe(FZ)3 production (when FZ is present). “A” represents the dissociative reduction pathway and “B” represents the non-dissociative reduction pathway for Fe(II)′ formation. FZ (3-(2-pyridyl)-5,6-diphenyl-1,2,4-triazine-4′,4′′-disulfonic acid sodium salt; Fluka) and 50 mM DFB (Sigma) were prepared in Milli-Q water and the pH was adjusted to 8.0. A 13.4 mM stock of xanthine (X; Sigma) was prepared in Milli-Q water and the pH was adjusted to 9. Xanthine oxidase (XO; Sigma) on receipt was dissolved in Milli-Q water and was stored at -85 °C in 1 mL aliquots. All other stock solutions were stored at 4 °C when not in use. Fe(II)′ Production Measurement. The superoxide generating X/XO system (14) ([X] ) 50 µM and [XO] ) 2 U L-1) was used to continuously generate superoxide in NaHCO3/ NaCl solution containing 200 nM Fe(III)L complex, and Fe(II)′ production was measured over time using FZ to trap Fe(II)′ as it was formed (1). The sample solution was pumped continuously through a 1 m path length Type-II Liquid Waveguide Capillary Cell (World Precision Instruments) and the absorbance of the solution at 562 nm (the wavelength at which Fe(FZ)3 absorbs most strongly) was monitored continuously using an Ocean Optics fiber optic spectrophotometry system consisting of a DH-2000 UV-Visible lamp coupled to USB2000 spectrometer as the detector. A molar absorptivity of 28 000 M-1 cm-1 at 562 nm (close to the published value of 27 900 M-1 cm-1 (15)) was obtained for Fe(FZ)3. To account for Fe(II)′ formation occurring due to processes other than superoxide-mediated reduction of Fe(III)L, the Fe(II)′ formation rate was measured in the absence of 2 U L-1 XO and subtracted from the total measured rate. All experiments were performed in triplicate in a waterjacketed reactor maintained at 25 °C and encased in foil to avoid interference from light. FeDFB Production Measurement. The production rate of the FeDFB complex was determined at pH 8.0 in the absence and presence of the superoxide generation source (i.e., the X/XO system) by measuring the absorbance of FeDFB at 470 nm, employing the same Ocean Optics spectrophotometry system described above. A final concentration of 500 µM DFB was added to NaHCO3/NaCl solution containing 1 µM Fe(III)L, 50 µM xanthine, and 0-2 U L-1 of XO and the formation of FeDFB was measured over time by pumping the samples continuously through the 1 m path length TypeII Liquid Waveguide Capillary Cell. A molar absorptivity of 2200 M-1 cm-1 at 470 nm (close to the reported value of 2280 M-1 cm-1 (16)) was obtained for the FeDFB complex. All experiments were performed in triplicate in a water-jacketed dark reactor maintained at 25 °C. The total iron concentrations used in our experiments measuring Fe(FZ)3 and FeDFB production are higher than the expected iron concentrations in many natural waters. These high concentrations of iron 3206

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were based on the detection limits of the methods used for measuring Fe(FZ)3 and FeDFB. As the molar absorptivity of FeDFB is lower than that of Fe(FZ)3, it was necessary to increase the total iron concentration from 200 nM to 1 µM in the case of FeDFB measurement to ensure that the FeDFB formed could be detected easily by 1 m path length spectrophotometry. Superoxide Production and Disproportionation Rate Measurement. Superoxide concentrations were determined by measuring the chemiluminescence of the product of the reaction of superoxide with the reagent methyl Cypridina luciferin analogue (MCLA) using the FeLume chemiluminescence system (Waterville Analytical) (1). Samples were continuously withdrawn from the reaction vessel using a peristaltic pump operating at 25 rpm and pumped directly into the FeLume for analysis. Superoxide for calibration purposes was generated photochemically by irradiation of 5 mM acetone (BDH) and 1 M ethanol (Fluka) in 1 mM borate buffer at pH 12 and was standardized by UV spectrometry at 240 nm, as described previously (1). The disproportionation rate constant of superoxide at pH 8.0 and its production rate from the X/XO system were determined as described earlier (1).

Results and Discussion The theoretical background to analysis of the results obtained in all experiments is presented in the Supporting Information, Sections A and B. It is recommended that this theoretical analysis be examined in consort with the results and discussion presented below. Fe(FZ)3 Production Rate. The effect of ligand concentration on Fe(FZ)3 production rates (reported as an average rate over 20 min) for various ligands from superoxide-mediated reduction of iron(III) complexes in the presence of excess FZ is shown in Figure 2. The increase in Fe(FZ)3 was linear for t > 2 min in almost all cases, thus confirming that steadystate is achieved for intermediates including Fe(III)′, Fe(II)′, and Fe(II)L within 2 min. Decreasing Fe(FZ)3 production rates with increasing citrate and SSA concentrations suggests that reduction occurs mostly via DR (i.e., pathway A in Figure 1) when iron is complexed to weak ligands such as citrate and SSA. The increase in ligand concentration decreases the effective Fe(III)′ concentration available for reduction, thereby decreasing the Fe(FZ)3 production rate (see Supporting Information, Section A). An alternative explanation for the above results might be that the decrease in Fe(FZ)3 production rate with increasing ligand concentration is associated with a change in the speciation of iron(III)-organic complexes. This possibility was rejected however as speciation calculations (see Supporting Information, Section C) indicated that there was no notable variation in the speciation of the iron(III)-citrate and iron(III)-SSA complexes over the range of iron and ligand concentrations examined in this study. In the presence of the strong ligands EDTA and DHB, as well as SRFA (which is often considered to contain a suite of ligands exhibiting a range of iron-binding strengths), the production rate of Fe(FZ)3 was not affected by increase in ligand concentration. Thus, these results suggest that NDR dominates in the presence of these ligands (see Supporting Information, Section A). The low Fe(FZ)3 production rate in the case of FeEDTA compared to FeSRFA can be explained by the rapid rate of oxygenation of the iron(II) complex formed on reduction. Earlier reports show that the rate constant for oxidation of Fe(II)EDTA by dissolved oxygen (2270 M-1 s-1) (11) is about 20 times higher than that for the Fe(II)SRFA complex (100 M-1 s-1) (13). Additionally, as FZ can trap only inorganic iron(II), the rate of dissociation of the iron(II) complex formed on reduction also affects the Fe(FZ)3 production rate. Our ongoing studies suggest that the dissociation rate of the Fe(II)EDTA complex is less than

FIGURE 2. Average initial Fe(FZ)3 production rate from superoxide-mediated reduction of (a) Fe(III)citrate, (b) Fe(III)SSA, (c) Fe(III)SRFA, (d) Fe(III)DHB, and (e) Fe(III)EDTA in the presence of 1 mM FZ. Points represent the experimentally measured rate and lines represent the rate predicted by kinetic modeling. Experimentally measured and model-predicted rates are based on the slope of Fe(FZ)3 concentration versus time over a 20 min period. The increase in Fe(FZ)3 concentration was linear for t > 2 min in almost all cases. Concentrations of 200 nM total iron, 2 U L-1 XO, and 50 µM xanthine were used in all the experiments. Error bars represent the standard deviation from triplicate experiments. that for Fe(II)SRFA, thus further decreasing the rate of Fe(FZ)3 production in the FeEDTA system. FeDFB Formation Rate. Increase in the rate of FeDFB formation on addition of superoxide (Figure 3a) is in agreement with our previous results (Figure 2e) and supports the earlier conclusion that NDR dominates in the presence of strong iron(III) complexes such as Fe(III)EDTA (see Supporting Information, Section A). The formation rate of FeDFB from dissociation of Fe(III)Citrate remains unchanged (Figure 3b) on addition of superoxide, supporting the conclusion that, in this case, DR dominates (see Supporting Information, Section A). In the case of both EDTA and citrate complexes, addition of DFB resulted in a sudden initial increase in the absorbance of the solution, which may be attributable to trace amounts of ferrioxamine (the Fe(III)DFB complex) that were already present as contaminants in the DFB stock and/or rapid complexation of very weakly bound Fe(III)′ in the solution. This initial increase in absorbance probably occurs through processes other than

those described in Figure 1, however it does not affect our conclusions based on these experiments since: (i) in the Fe(III)EDTA system only a small (7%) increase (compared to the total iron concentration) in absorbance was observed, and thus the overall effect of the sudden increase in absorbance was relatively unimportant; and (ii) in the Fe(III)citrate system a similar increase in absorbance was observed in the absence and presence of superoxide, and thus these results are not inconsistent with DR being the dominant mechanism in this system. These experiments could not be performed with iron(III) complexes of SSA, SRFA, and DHB as these complexes absorb strongly at wavelengths similar to the peak absorbance of FeDFB (17) resulting in interference with the measurement of FeDFB. Rate of Superoxide Production and Decay in the X/XO System. Superoxide disproportionates readily in the absence of trace metals to yield hydrogen peroxide and oxygen. The second-order rate constant for disproportionation (kdisp) of VOL. 41, NO. 9, 2007 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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by FZ. The rate law for superoxide is given by

d[Fe(FZ)3] d[O•2 ] 2 ) PO2•- - 2kdisp[O•2 ] dt dt

(3)

and so at steady-state, 2 •2kdisp[O•2 ]ss ) PO2

FIGURE 3. Kinetics of formation of FeDFB in the presence and absence of superoxide generated from 50 µM xanthine and varying concentrations of xanthine oxidase in the presence of 500 µM DFB and (a) FeEDTA (1 µM Fe; 20 µM EDTA) and (b) FeCitrate (1 µM Fe; 100 µM citrate). superoxide in the NaHCO3/NaCl system at pH 8.0 was calculated to be (4.0 ( 1.6) × 104 M-1 s-1 (see Supporting Information, Section D) which is close to the value reported by other investigators (18, 19). The production rate of superoxide from the X/XO system as described earlier (1) (details of the derivation are provided in Supporting Information, Section D) is given by

PO2•- ) k1[X][O2]

(1)

where k1 is the second-order rate constant for superoxide generation from X/XO. In the absence of trace metals, the rate of superoxide production is eventually balanced by its rate of decay due to disproportionation, resulting in a steady-state superoxide concentration described by the equation 2 k1[X][O2] ) 2kdisp[O•2 ]ss

(2)

Based on a measured steady-state superoxide concentration of 34.8 ( 2.1 nM in the absence of trace metals and a disproportionation rate constant of (4.0 ( 1.6) × 104 M-1 s-1 (as above), the rate of superoxide production in the presence of 50 µM xanthine and 2 U L-1 xanthine oxidase was calculated to be 93.5 ( 11.2 pM s-1. In order to account for any decay in superoxide due to its interaction with excess ligand present, we measured the steady-state superoxide concentration (produced from X/XO) in the absence as well as the presence of ligands at the maximum concentration used in our experiments. Similar steady-state superoxide concentrations were observed in both the presence and absence of ligand confirming that excess ligand does not interact with superoxide (Supporting Information, Section E). When iron is present in the system along with an excess of FZ, superoxide may be consumed by either disproportionation or by reduction of iron(III) to iron(II) with the subsequent trapping and removal of Fe(II)′ from the system 3208

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d[Fe(FZ)3] dt

(4)

Thus we are able to use the measured values of kdisp, PO2•-, and Fe(FZ)3 formation rate to calculate the steady-state superoxide concentration under each of the experimental conditions in the presence of ferrozine. Comparison of DR and NDR. Our experimental results suggest that in the presence of weak iron(III) complexes (Fe(III)Citrate and Fe(III)SSA), DR of the iron(III)-organic complex occurs primarily, resulting in more labile iron(II) than in the case of NDR. In the presence of strong iron(III) complexes, NDR dominates. The best-fit model results using an approach in which both DR and NDR contribute to Fe(FZ)3 production are shown in Figure 2. The various rate constants used and deduced rate constants for modeling are shown in Table 1. The best-fit model results were calculated using kinetic modeling software ACHUCHEM (20). ACHUCHEM was used for model fitting rather than the analytic solution derived earlier to account for the non-steady behavior of intermediates like Fe(II)L and superoxide during the initial few minutes of the experiment. However, in natural waters where various species are often present at steady-state, the analytic solution derived earlier will be valid and can be used for predicting the Fe(II)′ formation rates. As seen in Figure 2, the model accurately describes the rate of Fe(FZ)3 formation from the superoxidemediated reduction of Fe(III)L for all five of the ligands examined. The slight discrepancy between the predicted and measured Fe(FZ)3 production rate at the lowest ligand concentration in the case of weak citrate and SSA complexes may be due to precipitation of some iron from the solution at this ligand concentration. The rate of Fe(FZ)3 production would be slightly lower under these circumstances, as the iron precipitate would be reduced more slowly under either a DR or NDR pathway. Precipitation of Fe(III)′ released from Fe(III)L will occur when the rate of Fe(III)′ precipitation (given by kppt[Fe(III)′]2 where kppt ) 1.6 × 107 M-1 s-1 at pH 8 (21)) exceeds the rate of recomplexation by L (given by kf1[Fe(III)′][L]), i.e. when

kf1 [Fe(III)′] . k [L] ppt

(5)

Naturally occurring organic ligands in marine systems typically have kf values in the range 105 to 107 M-1 s-1 (17, 22), implying that iron/ligand ratios of less than 0.1 (or perhaps even higher for stronger ligands) would result in very little precipitation in natural waters. This suggests that the precipitation phenomenon observed in our experiments is unlikely to be of concern in most natural systems. We have also used the model rate constants to calculate the relative contribution of DR and NDR to the Fe(FZ)3 formation rate for each of the ligands and at each of the concentrations used (see Table 2). Our analysis suggests that about 60% of Fe(FZ)3 formed with ligand concentrations of 500 µM citrate and 20 mM SSA is produced by NDR, whereas nearly 100% of the Fe(FZ)3 formed with EDTA, DHB, and SRFA as ligands is produced by NDR at all ligand concentrations examined. Since the DR and NDR pathways are independent, the overall Fe(FZ)3 production rate via both NDR and DR at

TABLE 1. Rate Constants Used for Best Fit Model ligand

kd1 (s-1)

kf1 (M-1s-1)

citrate SSA

1.4 × 10-4 b 1.0 × 10-3 b

4.4 × 105 c 3.5 × 104 c

weak ligands 1.5 × 108 1.5 × 108

EDTA

1.0 × 10-6 i

4.6 × 107 c

strong ligands 1.5 × 108

DHB

1.5 × 10-6 k

7.5 × 106 c

1.5 × 108

2.0 × 10-6 d

6 × 106 c

1.5 × 108

SRFA

kr2 (M-1s-1)a

kox1 kd2

kd2 (s-1)

kox1 (M-1s-1)

kr1 (M-1s-1)

2 × 10-3 d 8 × 10-3 b

2.6 e 1.1 × 104 g

8 × 102 f 1.6 × 105 h

2 × 10-4 b

2270 j

2.3 × 105 h

≈ 5 × 10

6b

2.6 × 105 k

(a)

8 × 10-4 d

2.8 × 105 h

100 e

a Ref 3. b Based on best fit model line shown in Figure 2. c Unpublished data. The values of the complexation rate constants were determined in NaHCO3/NaCl solution using a method reported earlier (17). d Ref 17. e Ref 13. f Based on best fit model line. This value is different from that reported earlier (1) probably due to change in speciation of Fe(III)citrate complex, as the ligand concentration used here is greater (5-50 times) than that used in the earlier study. g Reported value for Fe(II)-salicylate (12). h Determined in NaHCO3/NaCl solution using the method reported earlier (1). These are close to reported values (1). i Ref 29. j Ref 11. k Average of reported values for EDTA and SRFA.

TABLE 2. Calculated Contribution of NDR to Fe(FZ)3 Formation Rate dFe(FZ)3

ligand concentration (µM)

from NDR dt (pM s-1) (% of total)

20 40 100 200 500

Fe-citrate system 0.9 (14%) 0.9 (17%) 0.9 (22%) 0.9 (33%) 0.9 (58%)

[O•2 ]ss (nM) 32.9 33.1 33.3 33.6 33.8

kr2kd1 kf1[L] +

500 1000 2000 4000 10000 20000

Fe-SSA system 2.3 (12%) 2.3 (13%) 2.3 (17%) 2.3 (25%) 2.3 (42%) 2.3 (60%)

30.3 30.8 31.5 32.4 33.1 33.4

4-100

Fe-EDTA system 0.37 (100%)

34.0

2-60 a

Fe-SRFA 15.8 (100%)

31.1

20-400

FeDHB system 0.83 (100%)

34.0

a

kr2[O•2 ]ss

)

kr1kd2 kox1 + kd2

(7)

From eq 7, we can deduce that, when NDR and DR coexist with equal contributions to the rate of Fe(II)′ formation, the relationship between the conditional stability constant of the iron(III) complex (K′) and ligand concentration is described by:

K′ )

Units in mg L-1.

steady-state superoxide concentration can be calculated using the sum of the DR and NDR rates which, as shown in Supporting Information, Section A (eqs S3 and S7 respectively), can be written as

[

pathways for superoxide removal (e.g., uncatalysed disproportionation and scavenging by copper) under environmental conditions. The particular case where both DR and NDR coexist and contribute equally to Fe(II)′ production depends on the strength of the iron complex (both iron(II) and iron(III)), ligand concentration, and superoxide production rate. The ligand concentration corresponding to this situation can be determined using the following equation (noting that each side of the equation corresponds to the DR and NDR terms in eq 6):

]

dFe(FZ)3 kr2kd1 kr1kd2 + ) [Fe(III)L][O•2 ] •dt kf1[L] + kr2[O2 ] kox1 + kd2 (6) The relative contribution of each pathway to the Fe(FZ)3, and hence Fe(II)′, production rate depends on the nature and concentration of the organic ligand present. In the presence of either strong ligands or high concentrations of weaker ligands, the fraction of iron(III) present as Fe(III)′ and available for reduction in this form is very small, and consequently the Fe(III)L complex is mostly reduced directly to the corresponding Fe(II)L complex. However, at lower concentrations of the weaker ligands, DR dominates over the direct reduction of the Fe(III)L complex. According to our model, we would expect the pathway for reduction of Fe(III)L (DR and/or NDR) by superoxide to be independent of total iron concentration, although the overall amount of iron reduced by superoxide will depend on both the total iron concentration and the relative importance of competing

kf1 kr2(kox1/kd2 + 1) ) kd1 kr2[O•2 ] kr1 [L] + kf1

(

)

(8)

This relationship divides the regions in which NDR and DR each contribute the majority fraction of Fe(II)′ formation during the superoxide-mediated reduction of iron(III). Whereas NDR converts iron(III)-organic complexes into more labile Fe(II)′, DR results in conversion of only Fe(III)′, which is already highly labile. The calculated overall increase in total inorganic iron (i.e., Fe(III)′ + Fe(II)′) from DR ranges between 5 and 15 times while the increase is about 102-106 times in case of NDR. Thus, from a biogeochemical perspective it is perhaps more important to establish for which values of K′ and [L] NDR makes any significant contribution to superoxide-mediated reduction of iron(III) since NDR will result in a remarkable increase in the overall lability of iron. If we nominate a value of 10% as a “significant” contribution of NDR to superoxide-mediated reduction, then the relationship in eq 8 is modified to give an equation describing the regions in which NDR contributes more or less than 10% to the fraction of Fe(II)′ formation during the superoxidemediated reduction of iron(III):

K′ )

kf1 0.1 × kr2(kox1/kd2 + 1) ) kd1 kr2[O•2 ] kr1 [L] + kf1

(

)

(9)

To apply this relationship across a range of values of K′ and [L], it is necessary to assume general values for the rate VOL. 41, NO. 9, 2007 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 4. Predominance of DR and NDR pathway in production of Fe(II)′ as function of ligand concentration and complex conditional stability constant. The solid points represent our experimental data. Plus symbols represent data from coastal ligands (17) and crosses are a compilation of open ocean data (22-26). Dashes represent the conditions at which NDR contributes 10% to the Fe(II)′ production rate for the ligands used in this study. The lines represent upper and lower bounds where NDR contributes 10% to the Fe(II)′ production rate for a steady-state superoxide concentration of 30 nM (solid line) and 0.1 nM (dotted line). The rate constants used for estimating the upper and lower bound are described in the text.

constants kr1, kf1, kox1, and kd2, which may actually vary depending on the specific complex involved. The relationship also varies depending on the concentration of superoxide. To consider this relationship in the context of environmentally relevant ligand concentrations and complex conditional stability constants, we have established upper and lower bounds for the line based on known variability of the major rate constants kf1, kox1, and kd2 (kr1 and kr2 remain nearly constant with variation in ligand type). The rate constant kf1 is known to vary from 1.0 × 104 to 1 × 108 M-1 s-1, based on published data, whereas the ratio kox1/kd2 is observed to vary from 1.2 × 105 and 1.2 × 107 M-1 for the ligands used in this study (based on the parameters listed in Table 1, used to fit the data in Figure 2). The upper and lower bound values for which NDR is calculated to make a significant contribution to Fe(II)′ formation based on a superoxide concentration of 30 nM (around the steady-state superoxide concentration in all our experiments) are shown in Figure 4 and are plotted together with our experimental data. We have additionally calculated the concentrations of citrate and SSA for which NDR contributes 10% to the Fe(II)′ formation rate. Our experimental data are consistent with eq 9 with the calculated values for citrate and SSA lying within the bounds of the two extreme conditions. The calculated ligand concentrations for EDTA, DHB, and SRFA at which NDR contributes 10% to the Fe(II)′ formation rate is negative, which implies that the contribution of NDR will always be greater than 10% when the steadystate superoxide concentration is >1 nM. Since superoxide concentration affects the position of the dividing line, we also calculated the upper and lower bounds of NDR and DR at a superoxide concentration of 100 pM (dotted line, Figure 4). Applying eq 9 with a lower superoxide concentration extends the linear portion of the relationship to higher values of K′. This is the case since eq 9 simplifies to

K′ ) 3210

9

0.1 × kr2(kox1/kd2 + 1) kr1[L]

(10)

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 41, NO. 9, 2007

i.e., log K′ ) -log [L] - log (kr2(kox1/kd2 + 1)/kr1)

(11)

when [O•2 ] , kf1[L]/kr2. At lower superoxide concentrations, this condition will therefore be met over a wider range of ligand concentrations. In addition to our experimental data, literature values of L and K′ collected in marine systems (22-26) are also sown in Figure 4. Interestingly, most of these data points lie within the upper and lower bound lines calculated from eq 9 for both concentrations of superoxide (0.1 and 30 nM). This suggests that NDR may only be important for some iron(III)-organic complexes. If these complexes have high values of kf1 and kox1/kd2 and the superoxide concentration is at sub-nanomolar levels, DR will dominate in most cases. However, it is possible that when superoxide is being generated biologically, the superoxide concentration near the cell surface (where reduction takes place) may be high (possibly in the nanomolar range). In this case, NDR would occur to a significant extent for complexes with characteristics similar to the complexes used in this study. Overall, the dominance plot developed in Figure 4 can be used to estimate the major pathway of superoxide-mediated reduction for a range of iron(III)-organic complexes found in natural waters. The dominance plot developed here extends a plot developed earlier by Shaked and co-workers (9) but also reflects an important difference. The dividing line between NDR and DR proposed by Shaked and coworkers (9) was estimated using data for one ligand only (EDTA). However, as explained above, the dividing line between NDR and DR cannot necessarily be generalized by extrapolation of values for a single complex as K and L are not the only parameters affecting the position of the dividing line. The effect of stability of the iron(II) complex (reflected by kox1 and kd2) and superoxide concentration also play a critical role in deciding the position of the dividing line. Implications of the Findings. The results of this study confirm that superoxide-mediated reduction of organically complexed iron may result in an increase in the concentration of potentially bioavailable iron(II). The efficiency of the

reduction process (calculated as the ratio of moles of Fe(II)′ produced for each mole of superoxide produced) varied from 0.4% to 27.5% and depends on the concentration and nature of the organic ligand present. Fe(II)′ production from superoxide-mediated reduction of Fe(III)L can occur via both DR and NDR pathways. In the presence of a weak iron(III) complex and at relatively low ligand concentration, DR is important. However, due to a decrease in the concentration of Fe(III)′ available for reduction, NDR becomes important when either a strong ligand or a high concentration of weaker ligand is present. Overall, the relative contribution of DR and NDR to Fe(II)′ production rate is determined by the stability of the iron(II) and iron(III) complexes, the superoxide production rate and the ligand concentration. The concentration of labile iron increases on reduction of Fe(III)L by superoxide irrespective of the pathway of reduction, however the increase in bioavailability is much more significant (∼102-106 times) in the case of NDR compared to DR (∼5-15 times). Thus, superoxide-mediated iron reduction would be expected to be most useful for biological iron acquisition in the presence of strong iron(III) complexes when superoxide concentrations are high as may potentially be the case near the surface of cells actively producing this reductant. In natural waters, other reductive processes such as biological reduction by cell surface reductases, photochemical reduction through ligand-to-metal charge transfer (LMCT), and reduction by organic radicals produced photochemically or otherwise can also occur and may be more efficient than superoxide-mediated reduction of Fe(III)L. Reduction of Fe(III)L by cell-surface reductases or organic radicals most likely involves an outer-sphere electron-transfer mechanism as in the case of superoxide. Thus we expect that similar mechanisms of reduction (DR or NDR) would occur in these systems, resulting in efficiency of iron(III) reduction comparable to that in the presence of superoxide. In the case of LMCT, the mechanism of Fe(III)L reduction is somewhat different because oxidation of the ligand results in production of highly labile iron(II) complexes which might be expected to rapidly dissociate to Fe(II)′. Thus, the efficiency of LMCT in terms of moles of Fe(II)′ produced per mole of reductant (in this case, photoexcited ligands) may well be considerably higher than for the outer-sphere processes described above. In addition, in the presence of other species which can compete with iron for superoxide, the overall rate of superoxide-mediate Fe(II)′ production may decrease and other reductive processes may become more important. Despite this, superoxide-mediated reduction of iron is likely to be an important process under many environmental conditions. A recent study showed that reduction of iron(III) to iron(II) by photochemically produced superoxide was considerably more important than LMCT in a simulated system typical of coastal waters (4). Modeling of photochemical iron reduction in oceanic waters by Miller and coworkers (in which superoxide was assumed to participate solely by DR) similarly implied that superoxide was at least as important as LMCT for iron reduction (27). In the absence of light, biologically important concentrations of iron(II) have been shown to exist even in marine waters with total iron concentrations of less than a few nanomoles per liter (28), implying that thermal reductive mechanisms must be operating. The relative importance of superoxide-mediated reduction versus other known mechanisms for iron reduction (e.g., direct reduction at cell surfaces and reduction by organic radicals) under such conditions is not known. However, the results of this study would suggest that it could well be an important contributor to Fe(II)′ production, especially at or in the vicinity of locations where this reactive oxygen species is being actively produced.

Supporting Information Available Detailed mathematical analysis of the rates of Fe(II)′ production under both DR and NDR pathways (Section A), additional details for justification of steady-state assumptions (Section B), complex speciation calculations (Section C), superoxide production rate and disproportionation rate constant determination (Section D), and examination of the effect of ligand concentration on steady-state superoxide concentration (Section E). This material is available free of charge via the Internet at http://pubs.acs.org.

Literature Cited (1) Rose, A. L.; Waite, T. D. Reduction of organically complexed ferric iron by superoxide in a simulated natural water. Environ. Sci. Technol. 2005, 39, 2645-2650. (2) Voelker, B. M.; Sedlak, D. L. Iron reduction by photoproduced superoxide in seawater. Mar. Chem. 1995, 50, 93-102. (3) Rush, J. D.; Bielski, B. H. J. Pulse radiolytic studies of the reactions of HO2/O2- with Fe(II)/Fe(III) ions. The reactivity of HO2/O2with ferric ions and its implication on the occurrence of the Haber-Weiss reaction. J. Phys. Chem. 1985, 89, 5062-5066. (4) Rose, A. L.; Waite, T. D. Role of superoxide in the photochemical reduction of iron in seawater. Geochim. Cosmochim. Acta 2006, 70, 3869-3882. (5) Kustka, A.; Shaked, Y.; Milligan, A.; King, D. W.; Morel, F. M. M. Extracellular production of superoxide by marine diatoms: Contrasting effects on iron redox chemistry and bioavailability. Limnol. Oceanogr. 2005, 50, 1172-1180. (6) Marshall, J.-A.; Ross, T.; Pyecroft, S.; Hallegraeff, G. M. Superoxide production by marine microalgae. II. Towards understanding ecological consequences and possible functions. Mar. Biol. 2005, 147, 541-549. (7) Oda, T.; Nakamura, A.; Shikayama, M.; Kawano, I.; Ishimatsu, A.; Muramatsu, T. Generation of reactive oxygen species by raphidophycean phytoplankton. Biosci., Biotechnol. Biochem. 1997, 61, 1658-1662. (8) Rose, A. L.; Salmon, T. P.; Lukondeh, T.; Neilan, B. A.; Waite, T. D. Use of superoxide as an electron shuttle by the marine cyanobacterium Lyngbya majuscula. Environ. Sci. Technol. 2005, 39, 3708-3715. (9) Shaked, Y.; Kustka, A. B.; Morel, F. M. M. A general kinetic model for iron acquisition by eukaryotic phytoplankton. Limnol. Oceanogr. 2005, 50, 872-882. (10) Santana-Casiano, J. M.; Gonza´lez-Da´vila, M.; Millero, F. J. Oxidation of nanomolar levels of Fe(II) with oxygen in natural waters. Environ. Sci. Technol. 2005, 39, 2073-2079. (11) Santana-Casiano, J. M.; Gonza´lez-Da´vila, M.; Rodrı´guez, M. J.; Millero, F. J. The effect of organic compounds in the oxidation kinetics of Fe(II). Mar. Chem. 2000, 70, 211-222. (12) Santana-Casiano, J. M.; Gonza´lez-Da´vila, M.; Millero, F. J. The oxidation of Fe(II) in NaCl-HCO3- and seawater solutions in the presence of phthalate and salicylate ions: a kinetic model. Mar. Chem. 2004, 85, 27-40. (13) Rose, A. L.; Waite, T. D. Effect of dissolved natural organic matter on the kinetics of ferrous iron oxygenation in seawater. Environ. Sci. Technol. 2003, 37, 4877-4886. (14) Fridovich, I. Quantitative Aspects of the Production of Superoxide Anion Radical by Milk Xanthine Oxidase. J. Biol. Chem. 1970, 245, 4053-4057. (15) Stookey, L. L. Ferrozine: a new spectrophotometric reagent for iron. Anal. Chem. 1970, 42, 779-781. (16) Rose, A. L.; Waite, T. D. Kinetics of hydrolysis and precipitation of ferric iron in seawater. Environ. Sci. Technol. 2003, 37, 38973903. (17) Rose, A. L.; Waite, T. D. Kinetics of iron complexation by dissolved natural organic matter in coastal waters. Mar. Chem. 2003, 84, 85-103. (18) Zafiriou, O. C. Chemistry of superoxide ion-radical (O2-) in seawater. I. pKasw* (HOO) and uncatalyzed dismutation kinetics studied by pulse radiolysis. Mar. Chem. 1990, 30, 31-43. (19) Bielski, B. H. J.; Cabelli, D. E.; Arudi, R. L.; Ross, A. B. Reactivity of HO2/O2- radicals in aqueous solution. J. Phys. Chem. Ref. Data 1985, 14, 1041-1100. (20) Braun, W.; Herron, J. T.; Kahaner, D. K. ACUCHEM: A computer program for modeling complex chemical reaction systems. Int. J. Chem. Kinet. 1988, 20, 51-60. (21) Pham, A. N.; Rose, A. L.; Feitz, A. J.; Waite, T. D. Kinetics of Fe(III) precipitation in aqueous solutions at pH 6.0-9.5 and 25 °C. Geochim. Cosmochim. Acta 2006, 70, 640-650. VOL. 41, NO. 9, 2007 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

3211

(22) Witter, A. E.; Hutchins, D. A.; Butler, A.; Luther, G. W. I. Determination of conditional stability constants and kinetic constants for strong model Fe-binding ligands in seawater. Mar. Chem. 2000, 69, 1-17. (23) Boye, M.; van den Berg, C. M. G.; de Jong, J. T. M.; Leach, H.; Croot, P.; de Baar, H. J. W. Organic complexation of iron in the Southern Ocean. Deep Sea Res. I 2001, 48, 1477-1497. (24) Croot, P. L.; Andresson, K.; Ozturk, M. The distribution and speciation of iron along 6°E in the Southern Ocean. Deep Sea Res. II 2004, 51, 2857-2879. (25) Rue, E. L.; Bruland, K. W. Complexation of iron(III) by natural organic ligands in the Central North Pacific as determined by a new competitive ligand equilibration/adsorptive cathodic stripping voltammetric method. Mar. Chem. 1995, 50, 117138. (26) Witter, A. E.; Luther, G. W. I. Variation in Fe-organic complexation with depth in the Northwestern Atlantic Ocean as

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determined using a kinetic approach. Mar. Chem. 1998, 62, 241-258. (27) Miller, W. L.; King, D. W.; Lin, J.; Kester, D. R. Photochemical redox cycling of iron in coastal seawater. Mar. Chem. 1995, 50, 63-77. (28) Hopkinson, B. M.; Barbeau, K. A. Organic and redox speciation of iron in the eastern tropical North Pacific suboxic zone. Mar. Chem. In Press. (29) Hudson, R. J. M.; Covault, D. T.; Morel, F. M. M. Investigations of iron coordination and redox reactions in seawater using 59Fe radiometry and ion-pair solvent extraction of amphiphilic iron complexes. Mar. Chem. 1992, 38, 209-235.

Received for review July 27, 2006. Revised manuscript received January 30, 2007. Accepted February 9, 2007. ES0617892