Article pubs.acs.org/est
Iron Redox Transformations in Continuously Photolyzed Acidic Solutions Containing Natural Organic Matter: Kinetic and Mechanistic Insights Shikha Garg,† Chao Jiang,† Christopher J. Miller,† Andrew L. Rose,‡ and T. David Waite†,* †
School of Civil and Environmental Engineering, The University of New South Wales, Sydney, NSW 2052, Australia Southern Cross GeoScience, Southern Cross University, Lismore NSW 2480, Australia
‡
S Supporting Information *
ABSTRACT: In this work, the various pathways contributing to the formation and decay of Fe(II) in photolyzed acidic solutions containing Suwannee River fulvic acid (SRFA) are investigated. Results of experimental and computational studies suggest that ligand to metal charge transfer (LMCT), superoxide-mediated iron reduction and interaction with reduced organic species that are present intrinsically in SRFA each contribute to Fe(III) reduction with LMCT the most likely dominant pathway under these conditions. Fe(II) oxidation occurs as a result of its interaction with a variety of light-generated species including (i) short-lived organic species, (ii) relatively stable semiquinone-like organic species, and (iii) hydroperoxy radicals. While not definitive, a hypothesis that the short-lived organic species are similar to peroxyl radicals appears most consistent with our experimental and modeling results. The semiquinone-like organic species formed during photolysis by superoxide-mediated oxidation of reduced organic moieties are long-lived in the dark but prone to rapid oxidation by singlet oxygen (1O2) under irradiated conditions and thus play a minor role in Fe(II) oxidation in the light. A kinetic model is developed that adequately describes all aspects of the experimental data obtained and which is capable of predicting Fe(II) oxidation rates and Fe(III) reduction rates in the presence of natural organic matter and light.
1. INTRODUCTION Iron is the fourth most abundant element in the Earth’s crust and is regarded as an essential micronutrient in both aquatic and terrestrial environments. Photoreduction of Fe(III) (hydr)oxides and/or iron-organic complexes is an important control on iron bioavailability in many sunlit surface and atmospheric waters. The major pathways for abiotic photochemical reduction of Fe(III) in natural waters include (i) reduction of Fe(III) by photochemically produced superoxide/ hydroperoxyl radicals (O−2 /HO•2 ), i.e., superoxide-mediated iron reduction (SMIR);1,2 (ii) ligand-to-metal charge transfer (LMCT) in photoactive Fe(III) species;3,4 (iii) direct excited state electron transfer from organic donor chromophores (not complexed with Fe); and (iv) direct reduction of Fe(III) by reduced organic species (e.g., hydroquinone and/or semiquinone-type species)5 either initially present in natural organic matter (NOM) or formed during irradiation of NOM. In our earlier investigation6 of the light-mediated reduction kinetics of Fe(III) bound to Suwannee River fulvic acid (SRFA), we showed that SMIR was the main pathway for photochemical reduction of Fe(III) at pH 8 but played a minor role under acidic conditions due to the short lifetime of (O−2 /HO•2 ). In a more recent study of light-mediated transformations of iron in acidic solutions of SRFA,5 we further showed that reduced organic species, which are intrinsically present in SRFA, are important reductants of Fe(III) at acidic pH. Direct excited © 2013 American Chemical Society
state electron transfer from organic donor chromophores (not complexed with Fe) appears unlikely since these organic chromphore are more likely to react with dioxygen which is present in far more excess than Fe(III) in oxygenated natural waters. However, the importance of LMCT or short-lived organic intermediates in Fe(III) reduction, although suggested previously,7 has not been well elucidated under acidic conditions. The persistence of photochemically produced Fe(II) in natural waters will also depend on its oxidation kinetics. Oxidation of Fe(II) in seawater is thought to occur primarily via its reactions with triplet dioxygen (3O2) and hydrogen peroxide (H2O2), if present at sufficiently high concentrations.2,8−10 In our recent work, we showed that long-lived semiquinone-like organic species, produced on irradiation of SRFA, are important oxidants of Fe(II) in aqueous solutions at acidic pH.5 While other oxidants may include reactive oxygen species (ROS: O−2 /HO•2 ,11 hydroxyl radical,12 and 1O26), shortlived oxidizing organic species (such as phenoxyl and peroxyl radicals,13 excited triplets, and organo-peroxides14), their Received: Revised: Accepted: Published: 9190
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complexed form. A ThermoOriel 150 W Xe lamp equipped with AM0 and AM1 filters to simulate solar radiation was used as the light source and was positioned horizontally adjacent to, and ∼5 mm from, the quartz cuvette. The spectral irradiance of the lamp and the absorbed photon irradiance (in μEinstein·m−2·s−1) by 10 mg·L−1 SRFA as a function of wavelength was the same as reported in our earlier study.5 As only a very small fraction (∼ 5%) of the incident light is absorbed over the 1 cm optical path length used in experiments, we have neglected the light screening effect in this study. 2.3. Fe(II) Determination. Concentrations of total Fe(II) were determined spectrophotometrically using a modified FZ method as described previously.5 For measurement of Fe(II) generated by photochemical Fe(III) reduction, 3 mL of a pH 4 solution containing SRFA and Fe(III) was irradiated in a 1 cm quartz cuvette for 0.5, 1, 2, 5, or 10 min followed by addition of 60 μL FZ-DFB mix. The concentration of Fe(II) in the sample was then determined using the modified FZ method. For measurement of Fe(II) oxidation kinetics, 3 mL of a pH 4 solution containing SRFA and Fe(II) was irradiated in a 1 cm quartz cuvette for 0.5, 1, 2, 5, or 10 min, then the lamp was extinguished and the concentration of Fe(II) remaining in the sample was determined using the modified FZ method. Additional details of the various control experiments (e.g., SOD and DMSO addition and substitution of H2O by D2O as the solvent) that were performed to determine the major intermediates involved in Fe redox transformations are provided in SI-1. 2.3. H2O2 Determination. H2O2 concentrations were quantified fluorometrically by the Amplex Red method16 using a Cary Eclipse spectrophotometer with settings and calibration procedures described previously.15 For measurement of H2O2 production in photolyzed SRFA solution, 3 mL of solution containing SRFA was irradiated in a 1 cm quartz cuvette for 0.5, 1, 2, 5, or 10 min, then 1 mL of sample was mixed with 2 mL of 10 mM phosphate buffer (pH 7) and ARHRP stock solution at final concentrations of 2.0 μM AR and 1 kU.L−1 HRP. 2.4. Data Analysis and Numerical Modeling. The pseudo first order rate constant for Fe(II) oxidation was calculated using the forward-difference method from the slope of a plot of log transformed Fe(II) concentration versus time. Kinetic modeling was undertaken using ACUCHEM.17
importance under conditions typically encountered in nature is not well-known. In this study, we investigate the kinetics of Fe(III) reduction in SRFA solution under acidic conditions (pH 4), and evaluate the importance of several possible Fe(III) reduction mechanisms under these conditions. As the rate of Fe(II) oxidation is a critical determinant of steady state Fe(II) concentration, we also examine the importance of the various Fe(II) oxidants (i.e., ROS, phenoxyl and peroxyl radicals, excited triplets, and organo-peroxides) generated in irradiated SRFA solution under acidic conditions where oxidation by triplet dioxygen is relatively unimportant. On the basis of our experimental data, we propose a mechanism for generation of Fe(II) via photochemical Fe(III) reduction in the presence of SRFA. Since all experiments were performed at pH 4, where superoxide is present principally as the hydroperoxy radical, HO•2 is used to represent both O−2 and HO•2 from here on.
2. EXPERIMENTAL SECTION 2.1. Reagents. Reagent solutions were prepared using 18.2 MΩ·cm resistivity Milli-Q water (MQ) unless stated otherwise. All experiments were performed at room temperature close to 22 °C in solutions at pH 4 containing 10 mM NaCl and 100 μM HCl. A solution of 0.5 M HCl for flushing of the experimental apparatus was prepared from 30% HCl (Sigma, reagent grade). A 2.0 g.L−1 stock solution of standard SRFA (International Humic Substances Society) was prepared in MQ and stored in the dark at 4 °C when not in use. The stock solution of SRFA was observed to be stable over the study duration (see SI SI-1 for more details). A working 4 μM Fe(II) stock in 0.2 mM HCl was prepared weekly by 1000-fold dilution with MQ of a primary 4.0 mM Fe(II) stock solution in 0.2 M HCl. The working stock pH was 3.5, which was sufficiently low to prevent significant Fe(II) oxidation over a week but sufficiently high to prevent significant pH change when added to pH 4 solutions. A 20 μM Fe(III) stock solution in 2 mM HCl was prepared every week by dilution of a primary 2.0 mM Fe(III) stock solution in 0.2 M HCl. The solution pH was sufficiently low to avoid polymerization or precipitation of iron. Stock solutions of 100 μM Amplex Red (AR; Invitrogen) mixed with 50 kU·L −1 horseradish peroxidase (HRP; Sigma) for H2O2 determination were prepared and stored as described previously.15 Stock solutions of 80 mM ferrozine (FZ; Sigma) and 20 mM desferrioxamine B (DFB; Sigma) were prepared in MQ water. For Fe(II) determination, a mixture containing FZ (50 mM) and DFB (5 mM) was prepared weekly by dilution of the 80 mM FZ and 20 mM DFB stock solutions. A stock solution of 3 kU.mL−1 superoxide dismutase (SOD; Sigma) was prepared in MQ and stored in 0.5 mL aliquots at −85 °C prior to use. A 10 mM sodium phosphate solution was prepared in MQ and its pH adjusted to 7 by addition of NaOH. A solution containing 10.0 mM NaCl and HCl at pD 4.0 ± 0.1 was prepared in D2O (99.9%; Sigma). Dimethyl sulfoxide (DMSO) was used as received from Sigma. 2.2. Photochemical Experimental Setup. Photochemistry experiments were performed in a 1 cm quartz cuvette (volume ∼3.5 mL) containing solutions of SRFA (in the range 5−10 mg·L−1) in which concentrations of Fe(II) and/or H2O2 were monitored over time during irradiation. At these concentrations of SRFA, the hydrolysis and subsequent precipitation of iron oxyhydroxides is expected to be minor with the large majority of Fe(III) likely to be present in
3. RESULTS AND DISCUSSION 3.1. Photochemical Reduction Kinetics of Fe(III) in Presence of SRFA. During photochemical reduction of Fe(III) in the presence of SRFA, Fe(II) concentrations reached a maximum within 2 min of commencing irradiation before declining over time at all three SRFA concentrations examined (Figure 1a). Both the maximum Fe(II) concentration achieved on Fe(III) reduction and the rate of Fe(II) oxidation observed after the maximum concentration was reached varied in a nonlinear manner with increase in SRFA concentration, with these observations supporting the conclusion that both Fe(III) reduction rate and Fe(II) oxidation rate are dependent on SRFA concentration. Recently, we proposed a mechanism for Fe(III) reduction by reduced hydroquinone-like organic species present in SRFA (A2−) and subsequent oxidation of Fe(II) by semiquinone-like organic species (A−) produced by HO•2 -mediated oxidation of A2− at pH 4.5 These long-lived organic species were shown to be important in controlling the redox transformations of Fe in 9191
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Figure 2. (a) Photochemical oxidation of 100 nM Fe(II) in the presence of 5 mg·L−1 (triangles), 7.5 mg.L−1 (squares) and 10 mg·L−1 (circles) SRFA. (b) Photochemical oxidation of 100 nM (triangles), 50 nM (squares) and 25 nM (circles) Fe(II) in the presence of 5 mg·L−1 SRFA. Symbols represent experimental data (average of duplicate measurements); lines represent model values.
Figure 1. (a) Fe(II) generation on photochemical reduction of 100 nM Fe(III) in the presence of 5 mg·L−1 (triangles), 7.5 mg·L−1 (squares) and 10 mg·L−1 (circles) SRFA. (b) Generation of Fe(II) on photochemical reduction of 100 nM (circles), 50 nM (squares) and 25 nM (triangles) Fe(III) in the presence of 5 mg·L−1 SRFA. Symbols represent experimental data (average of duplicate measurements); lines represent model values.
Fe(II) oxidation rate increased slightly with increasing SRFA concentration (Figure 2). As shown, the amount of Fe(II) remaining after 5 min of irradiation decreased from 35% to 15% on increasing the SRFA concentration from 5 to 10 mg.L−1. This suggests that the Fe(II) oxidant was generated as a result of SRFA photolysis and that the concentration of this oxidant increased with increasing SRFA concentration. The Fe(II) oxidation rate observed during the later stages of irradiation (>5 min; see SI Figure SI-1) was very slow suggesting that either the Fe(II) oxidant is exhausted or the Fe(III) reduction rate increases during later stages of irradiation. Given that A− is continuously generated on irradiation,5 the decreasing rate of Fe(II) oxidation over time cannot result from exhaustion of Fe(II) oxidant but rather is likely due to existence of an additional Fe(III) reduction pathway that becomes important after a significant amount of Fe(III) is generated via Fe(II) oxidation. Thus, we suggest that at least two pathways exist for both photochemical Fe(II) oxidation and Fe(III) reduction in the presence of SRFA, one of which involves A2− and A−.5 Possible additional pathways for Fe(III) reduction include SMIR, LMCT and excited state electron transfer from (non-Fe bound) organic chromophores while possible additional pathways for Fe(II) oxidation include ROS-mediated oxidation and oxidation by short-lived organic species. The possible involvement of H2O2 and organo-peroxides, which are also byproducts of SRFA photolysis and have similar reactivity
the dark in the presence of nonphotolyzed and previously photolyzed SRFA solutions. The nature and concentration of the A2−species initially present in SRFA were in agreement with properties of the reduced hydroquinone-like group that is present in untreated humic acid determined by Aeschbacher and co-workers.18 If A2− was the only important reductant of Fe(III) and A− the only important oxidant of Fe(II) in the irradiated system investigated here, then the Fe(II) concentration should reach steady-state within a few minutes of irradiation, which is inconsistent with our observations (Figure 1). Furthermore, the observed rate of Fe(III) reduction (0.60 nM·s−1 for 100 nM Fe(III) in the presence of 10 mg·L−1 SRFA, calculated from the 30 s data point shown in Figure 1) was much faster than the maximum calculated rate of Fe(III) reduction by A2− (∼0.14 nM·s−1 under the same conditions based on the previously reported rate constant).5 Thus, the nonsteady-state behavior of the Fe(II) concentration profile along with the fast reduction kinetics of Fe(III) suggests that, in irradiated SRFA solutions, Fe(III) reduction, and Fe(II) oxidation occur by pathways other than those involving A2− and A− alone. 3.2. Photochemical Fe(II) Oxidation Kinetics in Presence of SRFA. No Fe(II) oxidation was observed in the dark (see Figure SI-1 in Garg et al.5); however, when Fe(II) was added to SRFA solution and subsequently photolyzed, the 9192
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toward Fe(II) as H2O2,14 was excluded since addition of 2 μM H2O2 did not change Fe(II) oxidation kinetics noticeably (SI Figure SI-2). We attempt to identify the additional Fe(III) reductants and Fe(II) oxidants below. 3.3. Role of 1O2. In order to investigate the involvement of 1 O2, we measured rates of photochemical Fe(II) oxidation and Fe(III) reduction on addition of Fe(II) or Fe(III) to 5 mg.L−1 SRFA in D2O solution in which 1O2 is longer-lived and hence persists at a higher steady-state concentration than in aqueous solution. As shown in SI Figure SI-3, neither Fe(III) reduction nor Fe(II) oxidation kinetics are significantly different in D2O solution compared to aqueous solution (p > 0.5 single tailed student’s t-test), thus rejecting direct involvement of 1O2 in Fe(II) oxidation. This contrasts with results from our earlier study5 in which Fe(III) reduction rates increased and Fe(II) oxidation rates decreased when Fe(III) or Fe(II) was added to previously photolyzed D2O rather than aqueous solutions containing SRFA. We previously attributed these effects to a decrease in the concentration of the Fe(II) oxidant (A−) resulting from the more rapid oxidation of A− by the substantially higher concentration of 1O2 in D2O compared to that in aqueous solution. Given that we observed no change in Fe(III) reduction kinetics or Fe(II) oxidation kinetics in D2O solution compared to aqueous solution, it appears that A− is a minor Fe(II) oxidant in irradiated SRFA solution under the experimental conditions used here. We furthermore conclude that other Fe(II) oxidant(s) present in our experimental matrix are not generated via a 1O2-mediated pathway. 3.4. Role of Superoxide. In order to determine the role of HO•2 in Fe redox transformations, both Fe(III) reduction and Fe(II) oxidation rates were measured in solutions containing SOD which catalyzes the decay of HO•2 to dioxygen and H2O2. Addition of 25 kU·L−1 SOD increased the concentration of Fe(II) generated by Fe(III) reduction and decreased the proportion of Fe(II) oxidized (Figure 3). A similar effect was observed in our earlier study5 and was shown to occur because semiquinone-like Fe(II) oxidants are produced by the HO•2 mediated oxidation of Fe(III)-reducing, hydroquinone-like species. However, addition of SOD was observed to cause complete inhibition of Fe(II) oxidation in previously photolyzed SRFA solution in our earlier study5 whereas the Fe(II) oxidation rate decreased by ∼30% in irradiated SRFA solution here. This observation suggests that other Fe(II) oxidants must be generated via non-HO•2 mediated pathways, and hence induce Fe(II) oxidation even when SOD is present. Since the addition of SOD resulted in an increase in the concentration of Fe(II) generated by Fe(III) reduction after 10 min of irradiation, we can disregard SMIR as the main pathway for Fe(III) reduction under acidic conditions. Although SOD is known to also scavenge 1O2, this effect is neglected here given that 1O2 was found to play a negligible role in Fe redox transformation. 3.5. Effect of DMSO. As shown in Figure 4, addition of 1 mM DMSO, a scavenger of hydroxyl radical or hydroxylating intermediates,19 resulted in a significant (p < 0.08 using singletailed t-test) decrease in Fe(II) oxidation rate and, correspondingly, a substantial (p < 0.07 using single-tailed t-test) increase in net Fe(III) reduction rate. No interaction between DMSO and either Fe(II) or Fe(III) was observed in the absence of SRFA (data not shown) confirming that the observed effect was due to interaction between DMSO and SRFA. Consistent with the observed DMSO effect, direct involvement of HO• in Fe(II) oxidation is neglected here since the steady state
Figure 3. (a) Fe(II) generation on photochemical reduction of 100 nM Fe(III) in solutions containing 5 mg.L−1 SRFA in the presence (closed circles) and absence (open circles) of 25 kU·L−1 SOD. (b) Photochemical Fe(II) oxidation following addition of 100 nM Fe(II) to solutions containing 5 mg·L−1 SRFA in the presence (closed circles) and absence (open circles) of 25 kU·L−1 SOD. Symbols are the mean and error bars the range from duplicate experiments; lines represent model values.
concentration of HO• is expected to be very low.20 Applying previously reported values for HO• production rate of 0.38 nM·(mg·L−1)−1min−1 21 and rate constant for scavenging of HO• by SRFA of 7000 L·mg−1·s−1 22 yields a steady-state HO• concentration of 9 × 10−16 M in a solution containing 5 mg·L−1 SRFA. This steady-state HO• concentration yields a pseudo first order Fe(II) oxidation rate constant of 3.9 × 10−7 s−1 (using a second-order rate constant of 4.3 × 108 M−1s−1 for reaction of Fe(II) with HO• 23) which is 1000-fold lower than the measured apparent overall oxidation rate constant (see SI Figure SI-1). The radical intermediates formed from the reaction of HO• with SRFA are expected to be unstable22 and will either be rapidly oxidized by O2 yielding HO2• or will incorporate O2 to yield peroxyl radicals (RO2•) with the probability of either pathway dependent on the local chemical environment of the radical formed.24,25 Both HO2• and RO2• are recognized to be effective Fe(II) oxidants13,26 and hence could account for the effects of DMSO on Fe redox transformation observed here. Given that SOD addition decreased the extent of Fe(II) oxidation by ∼30% only, the possibility that HO2• is the only oxidant appears unlikely. Furthermore, peroxyl radical generation on photolysis of Aldrich humic acid (HA) at previously reported rates of 6−10 nM·min−1(mg HA·L−1)−1 27 would be expected to result in significant Fe(II) oxidation. The fact that 9193
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the observation that SOD addition did not completely suppress Fe(II) oxidation (Figure 3). It should be recognized that methyl radicals may well be generated on reaction of DMSO with hydroxylating intermediates, which could subsequently form methylperoxy radicals on reaction with oxygen that are capable of oxidizing Fe(II). If termination pathways for these methylperoxy radicals other than Fe(II) oxidation exist, then the argument for the role of peroxyl radicals in Fe(II) oxidation presented earlier still holds. If however these methylperoxy radicals are effective Fe(II) oxidants, alternate pathways for Fe(II) oxidation, then possibly involving other light generated oxidizing organic radicals, would need to be considered. As such, we are unable to distinguish conclusively between the various potential Fe(II) oxidants (i.e., peroxyl radicals or other oxidizing organic species produced on irradiation) based on the experimental evidence presented here. 3.6. Mechanism of Photochemical Redox Transformation of Fe. On the basis of the discussion presented above, Fe(III) reduction mainly occurs via a non-superoxide dependent pathway that may include LMCT or reduction by excited state electron transfer from non-Fe bound organic chromophores. Although LMCT within photoactive Fe(III)SRFA complexes has not been directly demonstrated here, its role in Fe(III) reduction on irradiation of acidic SRFA solutions is strongly supported by evidence from the literature.28−30 Long-range excited state electron transfer from organic chromophores, such as phenolic groups which are not complexed with Fe(III), is improbable as these chromophores are more likely to react with O2 or quinone-like species which are present at concentrations far in excess of Fe(III). Due to the difficulty in isolating these processes experimentally, we are unable to determine the exact mechanism of Fe(III) reduction, however, we can conclude that reaction with A2− (which is the key Fe(III) reductant in the dark) and reaction with HO•2 (the major Fe(III) reductant under alkaline conditions) play minor roles in Fe(III) reduction in irradiated SRFA solution at pH 4. On the basis of the data presented above and results of our earlier studies,5,31 Fe(II) oxidation by reaction with both longlived semiquinone species (A−) and the hydroperoxyl radical (HO•2 ) is minor. Potentially significant oxidants include excited singlet or triplet species and/or other short-lived organic species including phenoxyl and peroxyl radicals. Due to the
Figure 4. (a) Photochemical generation of Fe(II) following addition of 100 nM Fe(III) to 5 mg·L−1 SRFA solution in the absence (open circles) and presence (closed circles) of 1 mM DMSO. (b) Photochemical Fe(II) oxidation following addition of 100 nM Fe(II) to 5 mg·L−1 SRFA in the absence (open circles) and presence (closed circles) of 1 mM DMSO. The model results in the presence of DMSO were obtained by assuming that DMSO-derived methylperoxy radicals are ~100 times (or more) less reactive with Fe(II) under the conditions used than the peroxyl radicals generated on SRFA irradiation. Symbols represent experimental data (average of duplicate measurements); lines represent model values.
peroxyl radicals are generated via a non-HO2• mediated pathway24 in irradiated SRFA solution is also consistent with
Figure 5. Reaction scheme representing the mechanism of redox transformations of Fe and generation of various Fe(III) reductants and Fe(II) oxidants during irradiation of SRFA at pH 4. 9194
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Table 1. Kinetic Model for Fe Redox Cycling on Continuous Photolysis of pH 4 SRFA Solutions. no. 1
reaction
SRFA + hν → SRFA* SRFA* + 3O2 → SRFA + 1O2
2
1
3
Q + hν → Q− → NRPb
4 5 6 7 8
H2O
O2 ⎯⎯⎯⎯→ 3O2 kd
kf
H+
H+
A2 − + HO•2 ⎯→ ⎯ A− + H 2O2 A− + HO•2 → A2 − + O2 A− + 1O2 ⎯→ ⎯ A + HO•2 HO•2 + HO•2 → O2 + H 2O2
2.4 × 105 s−1
R + hν → R
10
R• + HO•2 → R− + O2 + H+ •
R +R →R−R O2,HO•
∼1 × 10 M s
3 × 105 M−1s−1
2.5 × 105 M−1s−1
5
−1 −1
1.5 × 10 M s
5
1.1 × 107 M−1s−1
1.1 × 107 M−1s−1
11
s
1.0 × 10 M s
Fe(III) + A2 − → Fe(II) + A−
4 × 103 M−1s−1
Fe(II) + RO•2 → Fe(III) + RO2 H
1 × 10 M s
20
Fe(II) + HO•2 → Fe(III) + H 2O2 Fe(III) +
HO•2
→ Fe(II) + O2
s
−1 −1
1.0 × 10 M s 3
1 × 10−6s−1e,f
16
19
7.5 × 10
−1
104−109 M−1s−1
−1 −1
3
Fe(II) + A− → Fe(III) + A2 −
hν , L
−6
1.5 × 105 M−1s−1
15
Fe(III) ⎯⎯⎯→ Fe(II) + Lox
8
−1e
14
18
−1 −1
1.0 × 10 M s
RO•2
17
5 37
RO•2 + RO•2 → RO4 R
+
1 × 10 M s
∼1 × 105 M−1s−1
13
→R
−1 −1
9
2 × 105 M−1s−1d
R + hν ⎯⎯⎯⎯⎯⎯⎯→ RO•2
HO•2
Dalrymple, et al.36 5
−1 −1
12
•+
reference
Paul et al.35
kf = 10 × 10−5 s−1c;kd = 5.8 × 103 s−1
6 × 10
9
11
2.4 × 105 s−1a
−6
•
•
Φ ≈ 0.5%
8
+
published value
Φ ≈ 0.5%
9
Q− + 3O2 ⎯→ ⎯ Q + HO•2
H
model value light-mediated reactions calculated
1 × 10
s
105−106 M−1s−1
−1f
75% (Figure 6; see SI-4 for
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ASSOCIATED CONTENT
S Supporting Information *
Additional detail on experimental methods, manipulative experiments, and kinetic model. This material is available free of charge via the Internet at http://pubs.acs.org.
■
AUTHOR INFORMATION
Corresponding Author
*Phone +61-2-9385 5059; fax +61-2-9385 6139; e-mail d.
[email protected] Notes
The authors declare no competing financial interest.
■
ACKNOWLEDGMENTS Funding provided by the Australian Research Council Discovery Grant Scheme (DP0987188) is gratefully acknowledged.
■
Figure 6. Relative contribution of various Fe(III) reduction pathways as a function of SRFA concentration during irradiation.
REFERENCES
(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) Rose, A. L.; Waite, T. D. Role of superoxide in photochemical reduction of iron in seawater. Geochim. Cosmochim. Acta 2006, 70, 3869−3882. (3) Faust, B. C.; Zepp, R. G. Photochemistry of aqueous iron(III)polycarboxylate complexes: Roles in the chemistry of atmospheric and surface waters. Environ. Sci. Technol. 1993, 27, 2517−2522. (4) Sima, J.; Makánová, J. Photochemistry of iron(III) complexes. Coord. Chem. Rev. 1997, 160, 161−189. (5) Garg, S.; Ito, H.; Rose, A. L.; Waite, T. D. Mechanism and kinetics of dark iron redox transformations in acidic previously photolyzed acidic natural organic matter solutions. Environ. Sci. Technol. 2013, 47, 1861−1867. (6) Garg, S.; Rose, A. L.; Waite, T. D. pathways contributing to the formation and decay of ferrous iron in sunlit natural waters. In Aquatic Redox Chemistry, ACS Symposium Series: 2011; Vol. 1071. (7) Borer, P.; Kraemer, S. M.; Sulzberger, B.; Hug, S. J.; Kretzschmar, R. Photodissolution of lepidocrocite (γ-FeOOH) in the presence of desferrioxamine B and aerobactin. Geochim. Cosmochim. Acta 2009, 73, 4673−4687.
detailed calculation). It should be recognized however that considerable uncertainty remains as to the exact processes contributing to the rate and extent of Fe(III) reduction and to the rates of these various processes. Additionally, the relative contributions of the various processes contributing to Fe(III) reduction are expected to vary with varying environmental conditions such as pH, organic concentration, total Fe concentration and radiation intensity. In particular, the contribution of SMIR will increase at higher pH since superoxide is longer-lived and present predominantly as the more reducing superoxide anion rather than the hydroperoxyl radical at higher pH. Diel variation in Fe(III) reduction rate is also expected, with LMCT and SMIR likely to be the most important pathways during the daytime while reaction with reduced organic species will be most important in the dark. Our results also show that stable semiquinone-like organic species, oxidizing radicals (such as peroxyl radicals), and HO•2 are all important oxidants of Fe(II) under acidic conditions. These results confirm that, in natural sunlit environments 9196
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(8) González-Dávila, M.; Santana-Casiano, J. M.; Millero, F. J. Oxidation of iron (II) nanomolar with H2O2 in seawater. Geochim. Cosmochim. Acta 2005, 69 (1), 83−93. (9) King, D. W.; Farlow, R. Role of carbonate speciation on the oxidation of Fe(II) by H2O2. Mar Chem 2000, 70, 201−209. (10) Miller, C. J.; Lee, S. M. V.; Rose, A. L.; Waite, T. D. Impact of natural organic matter on H2O2-mediated oxidation of Fe(II) in coastal seawaters. Environ. Sci. Technol. 2012, 46, 11078−11085. (11) 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 (4), 1041−1100. (12) Buxton, J. D.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. Critical review of rate constants for reactions of hydrated electrons, hydrogen atoms and hydroxyl radicals (•OH/O−•) in aqueous solution. J. Phys. Chem. Ref. Data 1988, 17, 513−531. (13) Khaiken, G. I.; Alfassi, Z. B.; Huie, R. E.; Neta, P. Oxidation of ferrous and ferrocyanide ions by peroxyl radicals. J. Phys. Chem. 1996, 100, 7072−7077. (14) Chevallier, E.; Joliboisa, R. D.; Meuniera, N.; Carliera, P.; Monodb, A. “Fenton-like” reactions of methylhydroperoxide and ethylhydroperoxide with Fe2+ in liquid aerosols under tropospheric conditions. Atmos. Environ. 2004, 38, 921−933. (15) Garg, S.; Rose, A. L.; Waite, T. D. Production of reactive oxygen species on photolysis of dilute aqueous solutions of quinones. Photochem. Photobiol. 2007, 83, 904−913. (16) Zhou, M.; Diwu, Z.; Panchuk-Voloshina, N.; Haugland, R. P. A stable nonfluorescent derivative of resorufin for the fluorometric determination of trace hydrogen peroxide: Applications in detecting the activity of phagocyte NADPH oxidase and other oxidases. Anal. Biochem. 1997, 253 (2), 162−168. (17) 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. (18) Aeschbacher, M.; Sander, M.; Schwarzenbach, R. P. Novel electrochemical approach to assess the redox properties of humic substances. Environ. Sci. Technol. 2010, 44, 87−93. (19) Pochon, A.; Vaughan, P. P.; Gan, D.; Vath, P.; Blough, N. V.; Falvey, D. E. Photochemical oxidation of water by 2-methyl-1,4benzoquinone: Evidence against the formation of free hydroxyl radical. J. Phys. Chem. A 2002, 106, 2889−2894. (20) Gan, D. Aqueous photochemistry of 1,4-benzoquinones and their possible role in the photochemistry of natural organic matter. University of Maryland, 2005. (21) Page, S. E.; Arnold, W. A.; McNeill, K. Terephthalate as a probe for photochemically generated hydroxyl radical. J. Environ. Monit. 2010, 12 (9), 1658−1665. (22) Westerhoff, P.; Mezyk, S. P.; Cooper, W. J.; Minakata, D. Electron pulse radiolysis determination of hydroxyl radical rate constants with Suwannee River fulvic acid and other dissolved organic matter isolates. Environ. Sci. Technol. 2007, 41 (13), 4640−4646. (23) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. Critical review of rate constants for reactions of hydrated electrons, hydrogen atoms and hydroxyl radicals (•OH/•O−) in aqueous solution. J. Phys. Chem. Ref. Data 1988, 17 (2), 513−886. (24) von Sonntag, C.; Dowideit, P.; Fang, X.; Mertens, R.; Pan, X.; Schuchmann, M. N.; Schuchmann, H.-P. The fate of peroxyl radicals in aqueous solution. Water Sci. Technol. 1997, 35 (4), 9−15. (25) von Sonntag, C.; Schuchmann, H.-P. The elucidation of peroxyl radical reactions in aqueous solution with the help of radiationchemical methods. Angew. Chim., Int. Ed. Eng. 1991, 30, 1229−1253. (26) 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/O2− with ferric ions and its implication on the occurrence of the Haber-Weiss reaction. J. Phys. Chem. 1985, 89, 5062−5066. (27) Blough, N. V. Electron paramagnetic resonance measurements of photochemical radical production in humic substances. 1. Effects of oxygen and charge on radical scavenging by nitroxides. Environ. Sci. Technol. 1988, 22, 77−82.
(28) Waite, T. D.; Morel, F. M. M. Photoreductive dissolution of colloidal iron oxides in natural waters. Environ. Sci. Technol. 1984, 18, 860−868. (29) Voelker, B. M.; Morel, F. M. M.; Sulzberger, B. Iron redox cycling in surface waters: Effects of humic substances and light. Environ. Sci. Technol. 1997, 31 (4), 1004−1011. (30) Southworth, B. A.; Voelker, B. M. Hydroxyl radical production via the photo-Fenton reaction in the presence of fulvic acid. Environ. Sci. Technol. 2003, 37 (6), 1130−1136. (31) Garg, S.; Rose, A. L.; Waite, T. D. Photochemical production of superoxide and hydrogen peroxide from natural organic matter. Geochim. Cosmochim. Acta 2011, 74, 4310−4320. (32) Zhang, Y.; Del Vecchio, R.; Blough, N. V. Investigating the mechanism of hydrogen peroxide photoproduction by humic substances. Environ. Sci. Technol. 2012, 46 (21), 11836−11843. (33) McKnight, D. M.; Kimball, B. A.; Bencala, K. E. Iron photoreduction and oxidation in an acidic mountain stream. Science 1998, 240, 637−640. (34) Scott, D. T.; Runkell, R. L.; McKnight, D. M.; Voelker, B. M.; Kimball, B. A.; Carraway, E. R., Transport and cycling of iron and hydrogen peroxide in a freshwater stream: Influence of organic acids. Water Resour. Res. 2003, 39. (35) Paul, A.; Hackbarth, S.; Vogt, R. D.; Röder, B.; Burnison, B. K.; Steinberg, C. E. W. Photogeneration of singlet oxygen by humic substances: Comparison of humic substances of aquatic and terrestrial origin. Photochem. Photobiol. Sci. 2004, 3, 273−280. (36) Dalrymple, R. M.; Carfagno, A. K.; Sharpless, C. M. Correlations between dissolved organic matter optical properties and quantum yields of singlet oxygen and hydrogen peroxide. Environ. Sci. Technol. 2010, 44 (15), 5824−5829. (37) Roginsky, V.; Barsukova, T. Kinetics of oxidation of hydroquinones by molecular oxygen. Effect of superoxide dismutase. J. Chem. Soc.−Perkin Trans. II 2000, 2000, 1575−1582. (38) Neta, P.; Huie, R. E.; Ross, A. B. Rate Constants for Reactions of Peroxyl Radicals in Fluid Solutions. J. Phys. Chem. Ref. Data 1990, 19, 413. (39) Garg, S.; Rose, A. L.; Waite, T. D. Superoxide mediated reduction of organically complexed iron(III): Comparison of nondissociative and dissociative reduction pathways. Environ. Sci. Technol. 2007, 41 (9), 3205−3212. (40) Wilkinson, F.; Helman, W. P.; Ross, A. B. Rate constants for the decay and reactions of the lowest electronically excited singlet state of molecular oxygen in solution. An expanded and revised compilation. J. Phys. Chem. Ref. Data 1995, 24, 663. (41) Aeschbacher, M.; Vergari, D.; Schwarzenbach, R. P.; Sander, M. Electrochemical analysis of proton and electron transfer equilibria of the reducible moieties in humic acids. Environ. Sci. Technol. 2011, 45, 8385−8394. (42) Aeschbacher, M.; Graf, C.; Schwarzenbach, R. P.; Sander, M. Antioxidant properties of humic substances. Environ. Sci. Technol. 2012, 46, 4916−4925.
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dx.doi.org/10.1021/es401087q | Environ. Sci. Technol. 2013, 47, 9190−9197
