Mechanism and Kinetics of Dark Iron Redox Transformations in

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Mechanism and Kinetics of Dark Iron Redox Transformations in Previously Photolyzed Acidic Natural Organic Matter Solutions Shikha Garg,† Hiroaki Ito,†,‡ Andrew L. Rose,†,§ and T. David Waite†,* †

School of Civil and Environmental Engineering, The University of New South Wales, Sydney, NSW 2052, Australia Tohoku University, Sendai, Japan § Southern Cross GeoScience, Southern Cross University, Lismore NSW 2480, Australia ‡

S Supporting Information *

ABSTRACT: Stable organic species produced on irradiation of Suwannee River Fulvic Acid (SRFA) are shown to be important oxidants of Fe(II) in aqueous solutions at acidic pH, with rate constants substantially larger than those for oxygenation of Fe(II) under the same conditions. These Fe(II)-oxidizing species, which are formed during photolysis by superoxide-mediated oxidation of reduced organic moieties that are present intrinsically in SRFA, are long-lived in the dark but prone to rapid oxidation by singlet oxygen (1O2) under irradiated conditions. The intrinsic reduced organic species are able to reduce Fe(III) at acidic pH. Although the exact identities of the organic Fe(II) oxidant and the organic Fe(III) reductant are unclear, their behavior is consistent with that expected of semiquinone and hydroquinone-like moieties respectively. A kinetic model is developed that adequately describes all aspects of the experimental data obtained, and which is capable of predicting dark Fe(II) oxidation rates and Fe(III) reduction rates in the presence of previously photolyzed natural organic matter.

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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. All organisms require iron for vital cellular functions including photosynthesis, respiration, and in some cases, nitrogen fixation. Despite its crustal abundance, labile iron is extremely scarce in oxygenated marine and freshwaters, with the inadequate availability of this critical element limiting growth of aquatic microorganisms in some instances.1,2 In natural waters, iron exists in two oxidation states. While the oxidized state of iron, Fe(III), is highly bioavailable in its dissolved inorganic form, the concentration of dissolved inorganic Fe(III) is typically extremely low due to its insolubility under circumneutral pH conditions. The reduced form of iron, Fe(II), is much more soluble and thus potentially more abundant in a bioavailable form, though the steady state concentration of Fe(II) will be determined by the relative rates of Fe(III) reduction and Fe(II) oxidation. Three major pathways are known to potentially account for abiotic photochemical reduction of dissolved Fe(III) in natural waters. These are (i) reduction of Fe(III) by photochemically produced superoxide/hydroperoxyl radical (O−2 /HO•2 ),3,4 that is, superoxide-mediated iron reduction (SMIR), (ii) ligand-tometal charge transfer (LMCT) in photoactive Fe(III) species,5−7 and (iii) direct reduction of Fe(III) by reduced organic species either initially present in natural organic matter (NOM) or formed during irradiation of NOM. In a previous study of the photochemical reduction kinetics of Fe(III), we showed that SMIR is the main pathway for photochemical © 2013 American Chemical Society

reduction of Fe(III) at pH 8 but is unimportant under acidic conditions due to the short-lifetime of O−2 /HO•2 under these conditions.8 The relative importance of LMCT or organic radical mediated Fe(III) reduction is not well established under acidic conditions. In the circumneutral pH range, Fe(II) oxidation is thought to occur primarily via its reactions with triplet dioxygen (O2) and, if present in sufficiently high concentration, hydrogen peroxide (H2O2).4,9−11 Other oxidants may include oxidizing organic radicals and/or other reactive oxygen species (ROS) such as O−2 /HO•2 and hydroxyl radicals (OH•), which are produced on photolysis of natural organic matter (NOM); however, these oxidants have not been shown to be important under conditions typically encountered in natural waters. For example, based on the results of kinetic calculations, Rose and Waite4 suggested that both OH• and oxidizing organic radicals play minor roles in Fe(II) oxidation in seawater. Furthermore, O−2 /HO•2 is more likely to behave as a Fe(III) reductant rather than oxidant of Fe(II) under circumneutral pH conditions.3 Under acidic conditions, oxygenation of Fe(II) is very slow, however the relative importance of other oxidants including organic species and/or ROS is unclear. In our recent study8 on Fe(II) generation from photolysis of Fe(III) in the presence of SRFA at pH 4 and 8, we concluded that, especially Received: Revised: Accepted: Published: 1861

September 4, 2012 January 18, 2013 January 19, 2013 January 19, 2013 dx.doi.org/10.1021/es3035889 | Environ. Sci. Technol. 2013, 47, 1861−1869

Environmental Science & Technology

Article

of the procedure for measurement of Fe(III) reduction kinetics in the presence of nonphotolyzed and previously photolyzed SRFA solutions and Fe(II) oxidation kinetics in the presence of previously photolyzed SRFA solution at pH 4 are provided in SI Section SI-1. Additional details of the various control experiments (including SOD addition and replacement of aqueous solution by D2O) performed to determine their effect on Fe redox transformations and H2O2 generation are also presented in SI Section SI-1. Fe(II) Determination. Concentrations of total Fe(II) were determined using a modified FZ method. FZ reacts rapidly with Fe(II) to form an Fe(FZ)3 complex which absorbs strongly at 562 nm. However, at low pH, FZ can facilitate reduction of Fe(III) which may result in over prediction of Fe(II) concentration in the samples. In order to minimize reduction of Fe(III) in the presence of FZ, DFB was also added to bind Fe(III) present in the samples. For determination of Fe(II) concentration, 60 μL of the FZ-DFB mix was added to 3 mL of sample and the resulting solution continuously circulated through a 1 m path length type II liquid waveguide capillary cell (World Precision Instruments). The absorbance of the solution was measured at 562 nm using an Ocean Optics fiber optic spectrophotometry system with correction for baseline drift by subtracting the absorbance at 690 nm (at which no components of the solution absorb significantly). Calibration of Fe(II) concentrations was performed immediately before undertaking experiments by standard addition of Fe(II) to the buffer solution containing the FZ-DFB mix. A molar absorption coefficient of 27 000 M−1cm−1 at 562 nm (close to the published value of 27 900 M−1cm−1)13 was obtained for Fe(FZ)3. Since a small amount of Fe(III) was reduced by FZ even in the presence of DFB and hence increased absorbance at 562 nm, calibration of Fe(III) was also performed using standard addition of Fe(III) to the buffer solution containing the FZ-DFB mix. The concentration of Fe(II) in the sample was deduced using the equation:

at pH 4, another important oxidant of Fe(II) was present in addition to dioxygen. We hypothesized that this oxidant could be either singlet oxygen (1O2) or a long-lived organic species generated on photolysis of SRFA solution.8 In this study, we investigate the kinetics of dark oxidation of Fe(II) and dark reduction of Fe(III) in previously photolyzed Suwannee River Fulvic Acid (SRFA) solutions under acidic conditions. On the basis of our experimental data, we propose a mechanism for the generation of an Fe(II) oxidant resulting from photolysis of SRFA and its subsequent involvement in controlling the steady state concentration of Fe(II) in previously photolyzed SRFA solutions.

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MATERIALS AND METHODS Reagents. Reagent solutions were prepared using 18.2 MΩ.cm resistivity Milli-Q water (MQ) unless stated otherwise. All experiments were performed at pH 4 in solutions containing 10 mM NaCl and 100 μM HCl. All experiments were performed at a room temperature of approximately 22 °C. The temperature of the solution increased slightly (1−3 °C) during irradiation for 10 min but returned to the room temperature 2 min after extinguishing the lamp. 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 was stable over the duration of the study (see SI-1.1 in the Supporting Information (SI) for 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 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.12 Stock solutions of 80 mM ferrozine (FZ; Sigma) and 20 mM desferrioxamine B (DFB; Sigma) were prepared in MQ water. A mixture containing FZ (50 mM) and DFB (5 mM), which was prepared weekly by dilution of the 80 mM FZ and 20 mM DFB stock solutions, was used for Fe(II) determination. A stock solution of 3 kU.mL−1 superoxide dismutase (SOD from bovine erythrocytes containing Cu and Zn; Sigma) was prepared in MQ and stored in 0.5 mL aliquots at −85 °C prior to use. A solution containing 10.0 mM NaCl and HCl at pD 4.0 ± 0.1 was prepared in D2O (99.9%; Sigma). Since all experiments were performed at pH 4 where O−2 /HO•2 mostly exists as HO•2 , we have used HO•2 to represent both O−2 and HO•2 from hereon. SRFA Photolysis. Irradiation of SRFA solutions was performed in a 1 cm path length quartz cuvette (volume ∼3.5 mL) using a ThermoOriel 150 W Xe lamp equipped with AM0 and AM1 filters to simulate solar radiation as the light source, positioned horizontally adjacent to the quartz cuvette. The spectral irradiance of the lamp and the absorbed photon irradiance (in μEinstein.m−2.s−1) of 10 mg·L−1 SRFA as functions of wavelength are shown in SI Figure SI-1. Details

[Fe(II)] = (A562 − εFe(III)[Fe]T )/(εFe(II) − εFe(III))

(1)

where A562 represents absorbance at 562 nm, εFe(II) represents molar absorption coefficient of Fe(II)FZ3 at 562 nm, εFe(III) represents the molar absorption coefficient of the small amount of Fe(II)FZ3 formed by reduction of Fe(III) in the presence of FZ, and [Fe]T represents the total Fe concentration. H2O2 Determination. The H2O2 concentration in the sample was quantified fluorometrically using the Amplex Red method14 in a Cary Eclipse spectrophotometer using settings and calibration procedures described previously.12 For measurement of H2O2 production in photolyzed SRFA solutions (both aqueous and D2O), 3 mL of solution containing SRFA was irradiated in a 1 cm quartz cuvette for 2, 5, or 10 min, then 1 mL of sample was mixed with 2 mL of 10 mM phosphate buffer and AR-HRP stock solution at final concentrations of 2.0 μM AR and 1 kU·L−1 HRP.

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RESULTS AND DISCUSSION Kinetics of Fe(III) Reduction and Fe(II) Oxidation in Nonphotolyzed SRFA Solution. When Fe(III) was added to SRFA in the dark, Fe(II) concentrations increased as a result of Fe(III) reduction for 10−30 min before reaching an asymptotic value (Figure 1), consistent with the results of earlier studies in which Fe(III) reduction kinetics in the presence of fulvic acid under acidic conditions have been investigated.15,16 The 1862

dx.doi.org/10.1021/es3035889 | Environ. Sci. Technol. 2013, 47, 1861−1869


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⎞ ⎛ k2 ⎜ − 1⎟[Fe(II)]2ss + (Fe0 + A 0)[Fe(II)]ss − Fe0A 0 = 0 ⎠ ⎝ k1 (3)

where Fe0 and A0 represent the initial concentrations of Fe(III) and A2− respectively. Based on the measured steady-state concentrations of Fe(II) produced from Fe(III) reduction in nonphotolyzed SRFA solutions, values of k1/k2 = 0.040 ± 0.005 and A0 = 35.4 ± 2.0 μmol.g−1 SRFA were calculated using eq 3 (see SI Section SI-3 for more details). The calculated concentration of stable reduced organic species in SRFA is reasonably consistent with the concentration (42 μmol·g−1) of stable reduced groups determined in nontreated humic acid by Aeschbacher and coworkers17 and with the concentration of reduced groups (25.4 μmol.g−1) in fulvic acid determined by Ratasuk and Nanny.18 However the concentration of the reduced group responsible for Fe(III) reduction determined here is much lower than the reported19 electron donating capacity of humic substances, which is mostly attributed to the presence of phenolic moieties, thereby suggesting that the functional groups responsible for Fe(III) reduction are distinct from phenolic moieties. The ratio k1/k2 also represents the equilibrium constant (K) for the reaction shown in eq 1. Based on the calculated equilibrium constant for reaction 1, and the calculated reduction potential of the Fe(III)/Fe(II) redox couple in our experimental matrix, we determine the standard reduction potential (E0H) of the reductant to be +0.36 V (see SI Section SI-3 for more details) which is in agreement with the standard reduction potential (E0H> +0.18 V) of reduced hydroquinonelike groups present in untreated humic acid determined by Aeschbacher and co-workers.17 The standard reduction potential of the reductant determined here is also in accord with the reported E0H range (0.178−0.448)20 for various hydroquinones. Kinetics of Fe(II) Oxidation and Fe(III) Reduction in Photolyzed SRFA Solution. Fe(II) oxidized when added to SRFA solutions that were irradiated for 10 min prior to ceasing illumination and adding Fe(II) (Figure 2). The rate of Fe(II) oxidation increased with increasing SRFA concentration, suggesting that the Fe(II) oxidant was generated as a result of photolysis of SRFA. Fe(II) was produced when Fe(III) was added to SRFA solutions that were irradiated for 10 min prior to cessation of irradiation and addition of Fe(III), but the steady-state Fe(II) concentration decreased with increasing SRFA concentration (Figure 3) and was lower than that observed when Fe(III) was added to nonphotolyzed SRFA solution (Figure 1). These observations suggest that the concentration of Fe(III) reductant decreased and/or the concentration of the Fe(II) oxidant increased when SRFA was irradiated. Fe(II) concentrations reached steady-state after addition of either Fe(III) or Fe(II) to photolyzed SRFA solutions (Figures 2 and 3), implying that a balance between the rates of Fe(III) reduction and Fe(II) oxidation was reached. These results are also consistent with the reaction scheme proposed under dark conditions (eq 1) if irradiation of SRFA results in a portion of A2− being oxidized to A−. Assuming this mechanism to be correct, the concentration of A− formed after irradiation of SRFA for 10 min (A0′ ) may be calculated from the relationship

Figure 1. (a) Generation of Fe(II) during reduction of 50 nM Fe(III) in nonphotolyzed 2.5 mg·L−1 (circles), 5 mg·L−1 (triangles) or 10 mg·L−1 (squares) SRFA solution. (b) Generation of Fe(II) in nonphotolyzed 10 mg·L−1 SRFA solution following addition of 25 nM (circles), 50 nM (triangles) or 100 nM Fe(III) (squares). Symbols represent experimental data (average of duplicate measurements); lines represent model values.

resulting steady-state concentration of Fe(II) increased nonlinearly with increasing initial SRFA and Fe(III) concentrations. If no back oxidation of Fe(II) was occurring, and the steadystate was thus due to complete consumption of the reductant, the steady-state Fe(II) concentration should have been linearly proportional to SRFA concentration, contrary to what was observed. Thus, we conclude that the steady-state concentration of Fe(II) is controlled by a balance between Fe(III) reduction and Fe(II) oxidation. Since no Fe(II) oxidation occurred in nonphotolyzed SRFA solutions for all SRFA concentrations investigated here (SI Figure SI-1), we further suggest that the Fe(II) oxidant must have been formed from the oxidation of reduced organic species by Fe(III). The experimental data were well described (Figure 1) by a simple reaction scheme in which reduced organic species (represented here by A2−) present in nonphotolyzed SRFA solution are oxidized by Fe(III) to form Fe(II) and oxidized organic radicals (A−), which may then reoxidize Fe(II): k1

A2 − + Fe(III) ⇌ A− + Fe(II) k2

(2)

This reaction scheme is also consistent with the observation that no Fe(II) oxidation occurs in the dark (SI Figure SI-1) provided that A− is absent in nonphotolyzed SRFA solution. The involvement of ROS (such as HO•2 and H2O2) or other organic species (peroxyl or phenoxyl radicals) is neglected here since the concentration of these species is expected to be very low in nonphotolyzed SRFA solution. The reaction scheme in eq 1 yields the following relationship: 1863



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Figure 2. (a) Oxidation of 100 nM total Fe(II) added after irradiation of solutions containing 1 mg·L−1 (diamonds), 2.5 mg·L−1 (squares), 5 mg·L−1 (circles) or 10 mg·L−1 (triangles) SRFA for 10 min. (b) Oxidation of Fe(II) in solutions containing 10 mg·L−1 SRFA that were irradiated for 10 min with subsequent addition of 25 nM (triangles), 50 nM (circles), 100 nM (squares), or 200 nM (diamonds) Fe(II). Symbols represent experimental data (average of duplicate measurements); lines represent model values.

Figure 3. (a) Generation of Fe(II) following addition of 50 nM Fe(III) to 5 mg·L−1 (triangles) or 10 mg·L−1 (squares) SRFA solutions that were irradiated for 10 min. (b) Generation of Fe(II), following addition of 25 nM (circles) 50 nM (triangles) or 100 nM Fe(III) (squares) to solutions containing 5 mg·L−1 SRFA that had previously been irradiated for 10 min. Symbols represent experimental data (average of duplicate measurements); lines represent model values.

A′0 = k1 Fe0A 0 k2

(

+ Fe0 −

k1 A k2 0

k

)

(

− 2 k1 Fe0 [Fe(II)]ss − 1 −

k1 Fe0 k2

2

(

+ 1−

k1 k2

k1 k2

lower total Fe to SRFA concentration ratios. This result suggests that Fe(III) is more reducible at higher Fe to SRFA concentration ratios, possibly as a result of weaker Fe(III) binding as the proportion of metal to organic increases. Involvement of other Fe(III) reductants such as HO•2 is also possible, with these processes increasing the steady-state concentration of Fe(II) produced and hence causing an error in estimated [A2−] and [A−] using eq 5. The concentration of Fe(III) reduced by HO•2 however will be 0.4 using a single tailed student ttest), while Fe(III) reduction in photolyzed SRFA solution increased in the presence of SOD (Figure 5b). This suggests that the Fe(II) oxidant is produced via an HO•2 -mediated pathway and/or HO•2 acts as a sink for the Fe(III) reductant. Both possibilities are consistent with the hypothesis that A2− reduces Fe(III) while A− oxidizes Fe(II) and are furthermore consistent with the expected reactions of a semiquinone/ hydroquinone couple with HO•2 .20 Role of Dioxygen. Oxidation of Fe(II) by 3O2 was unimportant at pH 4 on the time scale of our experiments (SI Figure SI-1), while oxidation of Fe(II) by 1O2 can be precluded in previously photolyzed solutions of SRFA based on the short lifetime (∼4 μs) of 1O2 in aqueous solution.26

Figure 4. Effect of time delay between extinguishing the lamp and adding 100 nM Fe(II) (squares) or 100 nM Fe(III) (circles) on the extent of Fe(II) oxidation and Fe(III) reduction respectively to 5 mg·L−1 SRFA solutions that were irradiated for 10 min and then stored in the dark before addition of Fe(II) or Fe(III). Symbols represent the measured steady-state concentration of Fe(II) achieved in each case. All measurements were performed in duplicates.

photolysis and Fe(III) addition (Figure 4), similarly suggesting that the Fe(III) reductant is also long-lived. Such longevity precludes the involvement of highly reactive species such as peroxyl radicals,21 HO•2 or OH•, which are known to oxidize Fe(II) but whose lifetime is much shorter than that of the oxidant generated in these studies. The involvement of H2O2 and organoperoxides (ROOH), which are stable byproducts of SRFA photolysis, also appears unlikely given that oxidation of Fe(II) by exogenous H 2 O 2 was negligible under the experimental conditions employed (SI Figure SI-6) and the reactivity of H2O2 and ROOH toward Fe(II) is similar.22 The involvement of organic peracids (RCOOOH) is a possibility as these compounds are recognized to be stronger oxidants than organoperoxides but they are likely to be present at low concentrations and direct production as a result of photolysis does not seem likely. It is clear from our observations of Fe(III) reduction in nonphotolyzed SRFA solution that stable reduced species (i.e., A2−) exist in SRFA solution; it is possible that the 1865



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Figure 6. (a): Oxidation of 100 nM Fe(II) added after irradiation of 10 mg·L−1 SRFA for 10 min in air saturated aqueous solution (squares), partially deoxygenated aqueous solution (triangles) or air saturated D2O solution (circles). (b) Generation of Fe(II) following addition of 100 nM Fe(III) to 10 mg·L−1 SRFA solution that was previously irradiated for 10 min in air saturated aqueous solution (squares), partially deoxygenated aqueous solution (triangles) or air saturated D2O solution (circles). Symbols represent the average of duplicate measurements; lines represent model values.

Figure 5. (a) Effect of SOD addition on Fe(II) concentration remaining 5 min after addition of 100 nM Fe(II) to solutions containing 5 mg·L−1 (circles) or 10 mg·L−1 (squares) SRFA that had been photolyzed for 10 min prior to adding Fe(II); (b) Effect of SOD addition on Fe(II) concentration generated after 5 min following addition of 100 nM Fe(III) to solutions containing 5 mg·L−1 (circles) or 10 mg·L−1 (squares) SRFA that had been photolyzed for 10 min prior to adding Fe(III). Symbols are the mean and error bars the standard deviation from duplicate experiments. SOD was added at final concentrations of 12.5 and 25 kU·L−1.

compared with 70% of added Fe(II) oxidized after the same period in H2O. Consistent with this observation, a much higher proportion of Fe(III) was reduced in D2O solution than in aqueous solution (Figure 6b). Thus, it appears that the presence of 1O2 enhances either (i) the generation of Fe(III)reducing A2− or (ii) the removal of Fe(II)-oxidizingA−. The possibility that 1O2 increases the generation rate of A2− is highly unlikely given that 1O2 is a powerful oxidant, strongly suggesting that 1O2 is capable of oxidizing either A− or the precursor organic moiety involved in the generation of A−. For the simple reaction mechanism where A− is formed by HO•2 mediated oxidation of A2−, the latter possibility is not feasible and hence we conclude that 1O2 oxidizes A−. Mechanism of Generation of the Oxidant Responsible for Fe(II) Oxidation in Photolyzed SRFA Solution. A proposed mechanism for the photomediated generation of A− is shown in Figure 7. Based on our earlier study,27 HO•2 is formed by reduction of O2 by the redox-active chromophore Q upon irradiation. Q appears to be similar to the electron accepting quinone-like moieties present in humic and fulvic acid which, upon reduction, can be reversibly oxidized in the presence of dioxygen.28 HO•2 formed as a result of this process either disproportionates to yield H2O2 or is oxidized to O2 on reaction with the light-generated organic radical R•. In addition to these two loss pathways, HO•2 can oxidize A2− to form A−, which is further oxidized to A on reaction with 1O2. Although we proposed that A is present in two redox states in our earlier

However, it is possible that 3O2 or 1O2 may be involved in generation or consumption of A2− and/or A− and hence may influence the kinetics of Fe redox transformations in photolyzed SRFA solutions. In order to probe the role of O2, we measured the rates of Fe(II) oxidation and Fe(III) reduction when Fe(II) or Fe(III) were added to irradiated solutions of 10 mg·L−1 SRFA that were partially (∼95%) deoxygenated prior to irradiation for 10 min. A decrease in O2 concentration decreased the Fe(II) oxidation rate (Figure 6a), suggesting that O2 is required for generation of the oxidant. Consistent with this observation, more Fe(II) generation from Fe(III) reduction in photolyzed SRFA solution occurred in partially deoxygenated solution compared to air-saturated solution, which we ascribe to decreased generation of the Fe(II) oxidant in these solutions (Figure 6b). In order to determine the involvement of 1O2, we measured rates of Fe(II) oxidation and Fe(III) reduction on addition of Fe(II) or Fe(III) to irradiated 10 mg·L−1 SRFA in D2O solution, in which 1O2 is longer-lived and hence persists at a higher steady-state concentration compared to that in aqueous solution. The results of the control experiments (see SI Section SI-1) and observation of similar H2O2 generation rates from irradiation of SRFA in D2O and aqueous solutions (SI Figure SI-5) ensured that any other effects unrelated to increasing 1O2 lifetime are unimportant in D2O systems. As shown in Figure 6a, only 14% of added Fe(II) was oxidized after 10 min in D2O 1866



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A kinetic model based on the reaction scheme shown in Figure 7 is displayed in Table 1 and a detailed description of each reaction and justification of the rate constants used is provided in SI Section SI-4. As shown (solid lines in Figures 1, 2, 3, and 6), the kinetic model simulates the Fe(III) reduction and Fe(II) oxidation kinetics under the various experimental conditions investigated here. The kinetic model is consistent with the data obtained in our earlier study on photochemical generation of O2− and H2O227,29 and also predicts the concentration of A− generated as function of irradiation time very well (SI Figure SI-9), as calculated from measured Fe(II) decay rates (SI Figure SI-8) in the presence of 5 mg·L−1 photolyzed SRFA solutions with increasing irradiation time. Furthermore, the kinetic model can be used to predict the effect of O2 concentration on rate of A− formation and H2O2 production (SI Figure SI-10) with the dependency of H2O2 production rate on O2 concentration similar to the rectangular hyperbolic form recently reported by Zhang and co-workers.29 The involvement of organic species A2− and A− as the only important Fe(III) reductant and Fe(II) oxidant respectively in photolyzed SRFA solution, as supported by our kinetic model (see Table 2), suggest that irrespective of the redox-state in which Fe is added, it essentially equilibrates with the A2−/A− redox couple in the dark. The presence of a reduced hydroquinone group in SRFA solution is consistent with the earlier work by Aeschbacher and co-workers on humic substances17 and is in agreement with the reported E0H range (0.178−0.448)20 for various hydroquinones. It should be noted that while we have developed the simplest model capable of describing all results obtained here, this does not exclude the possibility that other processes not described here may play some role. Implications of Findings. The findings described above demonstrate that stable reduced organic species exist in SRFA solution (in the dark) that reduce Fe(III) under acidic

Figure 7. Reaction scheme representing the mechanism of generation and decay of HO•2 and its involvement in redox transformations between A2−, A−, and A during irradiation of SRFA.

study27 (represented in that study by A− and A), invoking three redox states of A is essential to explain the experimental data obtained in this study. In our previous work, invoking two oxidation states for A was the simplest workable model, but this did not preclude the possibility of a third oxidation state. The existence of A in three oxidation states is consistent with hydroquinone-semiquinone-quinone type organic moieties where A2− and A− represent nonphotoreactive, redox active hydroquinone and semiquinone species respectively. On photolysis and oxidation of the semiquinone species by 1O2, quinone (A) is formed. Although we suggested in our earlier work that A2− (depicted in that study as A−) reacts with 1O2, reaction of A− (a semiquinone-like species) appears more reasonable given that oxidation of hydroquinone by 3O2 is spin restricted.20 The long-lived Fe(II) oxidant determined here thus exhibits behavior similar to that exhibited by semiquinone radicals. Such a result is consistent with earlier work23 showing the presence of stable semiquinone radicals in irradiated SRFA solution.

Table 2. Kinetic Model for Generation of Fe(II) Oxidant on Photolysis of SRFA at pH 4 no.

reaction

model value

Light-mediated Reactions 1 SRFA+hν→SRFA* SRFA* +3O2 → SRFA + 1O2 H2O 2 1 O ⎯⎯⎯⎯→ 3O 2

3

2

kd

kf

Q + hν → Q− → NRP +

4

Q− + 3O2 ⎯→ ⎯ Q + HO•2

5

HO•2 + HO•2 → H 2O2 + O2 R + hν → R• R• + HO•2 → R− + O2 + H+ R•+R•→R2

6 7 8 9 10 11

H

H+

A2 ‐ + HO•2 ⎯→ ⎯ A− + H 2O2 − • A + HO2 → A2− + O2 −

1

H

+

HO•2

published value

reference

calculated Φ ∼ 0.5% 2.4 × 105 s−1a

Φ ∼ 0.5% 2.4 × 105 s−1

32 33

kf = 6 × 10−5 s−1b; kd = 5.8 × 103

-

this work

−1 −1

−1 −1

∼1 × 10 M .s

1× 10 M .s

29

2 × 106 M−1.s−1

2 × 106 M−1.s−1

25

9

−6

−1c

9

−6

−1c

4 × 10 s 2.0 × 105 M−1s−1 1.0 × 103 M−1s−1 2 × 105 M−1.s−1d

7.5 × 10 s 104−109 M−1s−1 ∼1× 105 M−1.s−1

27 34 27 20

2.5 × 105 M−1.s−1d ∼1.5 × 108 M−1.s−1

∼ 10k9 diffusion limited

27 this work

A + O2 ⎯→ ⎯ A+ Fe(II) Oxidation and Fe(III) Reduction Reactions Occurring in Dark 1 × 105 M−1s−1e 12 Fe(II) + A− → Fe(III) + A2− 2− − 13 Fe(III) + A → Fe(II) + A k15/k14 = 0.04f

this work this work

a Rate constant is 1.6 × 104 s−1 in D2O solution.26 bPseudo-first order rate constant based on [Q]T = 1.7 mmol.g−1 SRFA;28 represents an apparent rate constant (see SI-4 for more details). cPseudo-first order rate constant based on [R]T = 44 mmol.g−1 SRFA.27 dBased on best-fit model results. e Based on best-fit model results to Fe(II) generation from Fe(III) reduction in nonphotolyzed SRFA solution and Fe(II) decay kinetics in irradiated SRFA solution. fRatio determined based on measured steady-state Fe(II) concentration generated from Fe(III) reduction in nonphotolyzed SRFA solution (see text for more details).

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conditions. These reduced species are oxidized on photolysis of SRFA by a HO•2 -mediated pathway to form semiquinone-like organic species, which are important Fe(II) oxidants under acidic conditions; oxidation half times resulting from the action of these oxidants are on the order of minutes compared to a half time for Fe(II) oxygenation on the order of hours to days (at pH 4). Although the Fe(II) oxidant is rapidly oxidized by 1 O2 under irradiated conditions, it is stable in the dark and does not appear to react with any other species, apart from Fe(II), in solution. The results of post-irradiation studies conducted here provide valuable insights into the critical role of relatively stable hydroquinone-like species present in NOM on Fe(III) reduction and of light-generated semiquinone-like species on Fe(II) oxidation, which are otherwise confounded in irradiated systems as a result of interference from other competing processes (such as ligand to metal charge transfer (LMCT) mediated Fe(III) reduction or ROS mediated Fe(II) oxidation). Additionally, based on the results obtained here, we can determine the importance of these competing processes. For example, Fe(II) generation rates resulting from Fe(III) reduction were substantially less in previously irradiated SRFA solutions compared to those observed in nonirradiated solutions; however, Fe(III) reduction rates when Fe(III) is present during photolysis could be higher or lower depending on the importance of LMCT at acidic pH. Furthermore, our results show that in natural sunlit environments where humic or fulvic-type NOM is present under mildly acidic conditions, steady state Fe(II) concentrations will be far less than would be the case if Fe(II) oxidation were controlled by the presence of either oxygen or H2O2 only. As the Fe(II) oxidant formed by SRFA photolysis is extremely long-lived, Fe(II) concentrations can be expected to decrease once sunlight-mediated Fe(II) generation ceases in a sunlit water body. Thus, significant diel variation in Fe(II) concentration is to be expected even in acidic waters since, in accord with earlier studies,30,31 time scales of light-mediated Fe(III) reduction and Fe(II) oxidation are expected to be similar. Overall, our work demonstrates the importance of stable organic quinone-like species in controlling Fe redox transformations in sunlit waters under acidic conditions. Although many investigators have suggested that organic species derived from fulvic and humic acids play a role in Fe redox transformations, this is the first study to provide clear evidence for the role of quinone-like moieties in both Fe(III) reduction and Fe(II) oxidation. While the results obtained here at pH 4 have direct relevance to atmospheric aerosols and acidic streams, preliminary results at higher pHs suggest similar rates of Fe(II) oxidation by the light induced semiquinone species and similar rates of Fe(III) reduction by the intrinsic hydroquinone species to those observed at pH 4. As such, the processes described here are expected to result in lower steady state Fe(II) concentrations across the range of pH typical of natural waters than would otherwise be the case. The effect of pH on the organic-mediated redox transformations of Fe in the presence of fulvic acid will be investigated in detail in future studies as will the importance of LMCT processes for Fe(III) reduction in sunlit waters. Additionally, studies of the lightmediated generation of long-lived semiquinone-like oxidants will be extended to natural organic compounds other than SRFA but which we suspect will exhibit similar redox behavior driven by the presence of quinone-like moieties.

Article

ASSOCIATED CONTENT

S Supporting Information *

Supporting Information contains details of (i) stability of SRFA solutions, (ii) Xe lamp used for SRFA irradiation, (iii) Fe(II) generation kinetics from Fe(III) reduction in nonphotolyzed SRFA solution, (iv) Fe(II) generation kinetics from Fe(III) reduction in photolyzed SRFA solution, (v) Fe(II) oxidation kinetics in photolyzed SRFA solution, and (vi) H 2 O 2 generation from photolysis of SRFA. In addition, further details are provided on the kinetic model as are calculations used to determine the initial Fe(III) reductant concentration, standard reduction potential of Fe(III) reductant, Fe(II) oxidant concentration and the rate constants for Fe(III) reduction and Fe(II) oxidation by organic species. Information is also provided on the kinetics of Fe(II) oxidant photogeneration and effect of dioxygen concentration on kinetics of H2O2 formation and Fe(II) oxidant photogeneration. This material is available free of charge via the Internet at http:// pubs.acs.org.

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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.

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REFERENCES

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Mechanism and Kinetics of Dark Iron Redox Transformations in

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