Environ. Sci. Technol. 2010, 44, 263–268
Photochemical Cycling of Iron Mediated by Dicarboxylates: Special Effect of Malonate ZHAOHUI WANG, XI CHEN, HONGWEI JI, WANHONG MA, CHUNCHENG CHEN, AND JINCAI ZHAO* Beijing National Laboratory for Molecular Sciences, Key Laboratory of Photochemistry, Institute of Chemistry, The Chinese Academy of Sciences, Beijing 100190, China
Received July 1, 2009. Revised manuscript received November 1, 2009. Accepted November 24, 2009.
Photochemical redox cycling of iron coupled with oxidation of malonate (Mal) ligand has been investigated under conditions that are representative of atmospheric waters. Malonate exhibited significantly different characteristics from oxalate and other dicarboxylates (or monocarboxylates). Both strong chelating ability with Fe(III) and strong molar absorptivities, but much low efficiency of Fe(II) formation (ΦFe(II) ) 0.0022 ( 0.0009, 300-366 nm) were observed for Fe(III)-Mal complexes (FMCs). Fe(III) speciation calculation indicated that Mal is capable of mediating the proportion between two photoactive species of Fe(III)-OH complexes and FMCs by changing the Mal concentration. Spin-trapping electron spin resonance (ESR) experiments proved the formation of both the · CH2COOH and · OH radicals at lower total Mal concentration ([Mal]T), but only · CH2COOH at higher concentrations of malonate, providing strong evidence for competition between malonate and OH- and subsequent different photoreaction pathways. Once FMCs dominate the Fe(III) speciation, both photoproduction and photocatalyzed oxidation of Fe(II) will be greatly decelerated. There exists an induction period for both formation and decay of Fe(II) until FeIII(OH)2+ species become the prevailing Fe(III) forms over FMCs as Mal ligand is depleted. A quenching mechanism of Mal in the Fe(II) photoproduction is proposed. The present study is meaningful to advance our understanding of iron cycling in acidified carbon-rich atmospheric waters.
Introduction Iron species have been identified as a ubiquitous component in atmospheric water droplets (i.e., cloud, rain, or fog droplets) in field measurements (1). The maximum concentration of dissolved iron species was up to 200 µM in fogwater (2). Photochemistry of atmospheric iron has a significant effect on numerous chemical processes in atmospheric waters, especially in redox and radical chain reactions such as natural fluctuation of reactive oxygen species (ROS) (3-5), oxidation of dissolved sulfur dioxide (SO2) (6) and organic substances (7), and redox cycling of other trace metals (e.g., Cu, Mn) (8). Furthermore, photolysis of iron species could contribute to iron input to open-ocean surface water via atmospheric deposition, thereby increasing the bioavailability of iron to aquatic biota (6). * Corresponding author fax: +86-10-8261-6495; e-mail: jczhao@ iccas.ac.cn. 10.1021/es901956x
2010 American Chemical Society
Published on Web 12/09/2009
In general, the monomeric ferric complex FeIII(OH)2+ is the most photoactive species in the absence of organic ligands in clouds and fog (3). The photolysis of FeIII(OH)2+ under UV irradiation leads to generation of Fe(II) and •OH(eq 1). hv
FeIII(OH)2+ 98 Fe(II) + •OH
(1)
DOM + •OH f oxidized products
(2)
Extensive field measurements have shown that dissolved organic matters (DOM) are very ubiquitous in atmospheric water droplets where the dissolved iron coexists at a comparable concentration (1). DOM that have relatively poor affinity to Fe(III), such as most monocarboxylates, may enhance the production of Fe(II) by scavenging •OH radical (eq 2), which decreases the rate of reoxidation of newly generated Fe(II) (9). Our recent work also revealed that influx of various DOM (10) or inorganic chromium species (11) can change Fe(II)/Fe(total) ratio in different ways. Among these Fe(III)-DOM species, Fe(III)-oxalato complexes are highlighted for their considerably high photoactivity under sunlight irradiation. The photolysis of ferrioxalate complex may proceed as follows (eqs 3-8) (4, 5): hv
[FeIII(C2O4)3]3- 98 [FeIII(C2O4)2]-+2CO•2 hv
[FeIII(C2O4)3]3- 98 [FeII(C2O4)2]2- + C2O•4
(3a)
(3b)
III 3CO•f [FeII(C2O4)2]2- + CO2 + C2O22 + [Fe (C2O4)3] 4 (4) •CO•2 + O2 f CO2 + O2
O•2
+
+H T
(5)
HO•2
(6)
2HO•2 f H2O2 + O2 •
Fe(II) + H2O2 f Fe(III) + OH + OH
(7) -
(8)
Ferrioxalate complex absorbs a photon and undergoes photodissociation without electron transfer from the oxalate to iron (eq 3a) (12) or with a ligand-to-metal charge transfer (LMCT) process, yielding Fe(II) and oxalate radical anion (C2O4•-) (eq 3b) (5, 13). CO2•- is a strong reducing agent (E0 ) -1.8 V (NHE)) and can react with another ferrioxalate molecule or can instead reduce O2 to superoxide anion (O2•-) at near-diffusion-controlled rate (k ) 2.4 × 109 M-1 s-1) (14). The photochemically generated Fe(II) can be reoxidized by O2, O2•-/HO2•, •OH, H2O2, and other oxidants. Malonic and succinic acids are also dominant dicarboxylic acids only inferior to oxalic acid in carbon-rich atmospheric waters (7, 15). Since oxalate and malonate have the strongest chelating capacity with Fe(III) among all dicarboxylates and monocarboxylates (see Table S1), the photochemistry of Fe-malonate complexes is expected to be important in iron cycling of atmospheric waters. However, there is little information available so far regarding the photochemistry of Fe(III)-dicarboxylate complexes except for that of ferrioxalate (4). The objective of this study is to investigate the photochemical behaviors of Fe(III)-malonate complexes (FMCs) comparing with those of Fe(III) complexes of oxalate and other carboxylates. The effect of malonate ligand on iron cycling and the photochemical reaction mechanism are also VOL. 44, NO. 1, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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discussed. One of the most important findings is that, in comparing with ferrioxalate/ferric hydroxo species, a high level of Fe(II)/Fe(t) can be maintained under UVA irradiation in the presence of excess Mal. This phenomenon is unique among the dicarboxylates (C2-C6) and monocarboxylates (C1-C3).
Experimental Section Chemicals. Iron(III) perchlorate hydrate, iron(II) perchlorate hydrate, and 2,2,6,6-tetramethylpiperidine-1-oxyl (TEMPO) were purchased from Aldrich. 5,5-Dimethyl-1-pyrroline-Noxide (DMPO) was from Sigma Chemical Co. Iron (III) tripotassium oxalate trihydrate was from Alfa Aesar. Oxalic acid, malonic acid, sodium hydroxide, perchloric acid, and 1,10-phenanthroline were of reagent grade and used as supplied. Barnstead UltraPure water (18.3 MΩ cm) was used for all experiments. Experimental Procedures. A 100-W Hg lamp (Toshiba SHL-100UVQ-2) was employed as the ultraviolet irradiation source. It mainly emits at 312 and 366 nm (16), overlapping the solar UVA spectrum (320-400 nm) (see Figure S1). All experiments were conducted in a 70-mL cylindrical Pyrex vial (Corning, Inc.) under continuous magnetic stirring. The Pyrex vial also served as a high-pass filter so that only light with wavelengths greater than 290 nm penetrated into the vessel. Unless otherwise specified, all experiments were performed under exposure to air at room temperature. An exhaust fan was employed during all the reaction processes to maintain the temperature below 35 °C. Fe(III)-organic acid solutions were freshly prepared by dilution of stock solutions of 0.01 M dicarboxylic acids (DAs), 5 mM Fe(III) at pH 1.5 (HClO4). The initial pH was adjusted with dilute HClO4 or NaOH. Under our experimental conditions ([Fe(III)], 100 µM; pH 3.0), the polynuclear iron complexes are negligible and are not considered in the photochemistry of iron cycling (see discussion in Supporting Information). For deaerated experiments, the solutions in the cap-sealed Pyrex vial were bubbled with high-purity Ar (O2 e 0.001%) for at least 20 min prior to UV irradiation and continuously purged throughout the experiment. During each kinetic experiment, a 1-mL aliquot was sampled with a new syringe each time and immediately disposed for the consequent analysis. Methods and Analysis. The concentration of Fe(II) was measured spectrophotometrically by a modified phenanthroline method (10, 17). Briefly, 0.5 mL of 1,10-phenanthroline solution (5.0 mM), 1 mL of sodium acetate/acetic acid buffer (pH 5.5), and 0.5 mL of ammonium fluoride solution (0.1 M) were premixed, followed by addition of 1 mL of the sample solution. The absorption of the resulting solution was read at 510 nm using a 1-cm quartz cell on a Hitachi U-3100 spectrophotometer. We eliminated the possibility that malonate would interfere with Fe(II) measurements by blank experiments (Figure S2). A DX-120 ion chromatograph (Dionex Co.) with conductivity detection was applied to quantitively determine organic acids. Five mM NaOH solution was chosen as an eluent. Six mM Ar-purged potassium ferrioxalate (0.05 M H2SO4) was used as a chemical actinometer to measure quantum yield (ΦFe(II)) for Fe(II) formation. An average actinometer quantum yield of 1.14 was used. All light at the photolysis wavelength (300-366 nm) was approximately absorbed by 6 mM ferrioxalate (18). Molar extinction coefficient of species involved in this study was calculated by spectral analysis method reported by Hug et al. (19). Electron spin resonance (ESR) spectra of spin-trapping radicals by DMPO were recorded at room temperature on a Bruker EPR ELEXSYS 500 spectrometer equipped with an in situ irradiation source (a Quanta-Ray ND:YAG laser system λ ) 355 nm). TEMPO (g ) 2.0051) was chosen as a standard for determination of g factors as recommended elsewhere 264
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FIGURE 1. UV-vis absorption spectra for solutions containing 100 µM Fe(III) and Mal ligand with different concentrations. Inset: Speciation (molar fraction) of 100 µM Fe(III) in solutions with total malonate concentration ([Mal]T) between 10 and 1000 µM. pH, 3.0. (20). Typical instrumental conditions were as follows: center field, 3480 G; sweep width, 100 G; resolution, 1024 pts; microwave frequency, ∼9.77 GHz; microwave power, 12.68 mW; modulation frequency, 100 kHz. To minimize experimental errors, the same quartz capillary was used for all the measurements. The simulations of ESR spectra were obtained with the use of WinSim EPR simulation software.
Results and Discussion Fe(III) Speciation. Table S1 (see Supporting Information) summarizes Fe(III) complexes with different carboxylates and the corresponding equilibrium constants. Malonate shows strong complexation ability with Fe(III) and is capable of forming mono-, di-, and trimalonato complexes with Fe(III), although the equilibrium constant for each stoichiometry of Fe(III)-Mal species is somewhat lower than that of the corresponding Fe(III)-oxalate complex. Due to their poor complexation capacity with Fe(III) (21, 22), the wellknown photochemistry of Fe(III)-OH complex is predominant in the presence of other mono- or dicarboxylates. Figure 1 shows the speciation of Fe(III) as a function of total malonate concentration ([Mal]T) at pH 3.0. The concentrations of all the hexaaquo and hydroxylated Fe(III) complexes decreased significantly with increasing [Mal]T, whereas FMCs predominated the speciation of Fe(III) at high [Mal]T. FeIII(Mal)+ and FeIII(Mal)2- are the major species of FMCs at pH 3.0 whereas FeIII(Mal)33- seems negligible under the present experimental conditions. Molar Absorptivities. Table S3 shows the values of molar extinction coefficients of individual complexes of FMCs within the range of 300-366 nm, which were obtained by multivariate linear regression method based on the known molar absorptivities of Fe(III)-OH complexes (Figure 1). Both Fe(Mal)+ and Fe(Mal)2- exhibited strong light-absorbing abilities, for example, εFe(Mal)+ and εFe(Mal)2- at 300 nm are 3512 and 3044 M-1 cm-1, respectively, larger than ferric hydroxyl chromophores. However, despite stronger light absorption of FMCs, their Fe(II) quantum yields were rather lower than that of monomeric Fe(OH)2+ (see discussion below). Photocatalyzed Oxidation of Fe(II). Dissolved Fe(II) is the predominant oxidation state in some atmospheric liquid waters and aerosol particles, accounting for 20-90% of the total Fe in fogwater samples from Zu ¨ rich (2), 40-72% of the dissolved Fe in cloudwater samples from Kleiner Feldberg (23), and 85 ( 13% of total dissolved Fe in the sunlit aerosol particles from Okinawa (24). Therefore, an initial ratio of Fe(II)/Fe(t) ) 63% was chosen to examine the photocatalyzed oxidation of Fe(II) in the presence of different carboxylates. The photocatalyzed oxidation of Fe(II) proceeded rather
FIGURE 2. Photocatalyzed oxidation of Fe(II) in the presence of malonate and oxalate. Ultraviolet irradiation source: 100-W Hg lamp (Toshiba SHL-100UVQ-2, average 3.6 W m-2 at 320-400 nm). Fe(II):Fe(III) ) 63:37; [Fe(t)], 100 µM; dicarboxylate, 500 µM; pH, 3.0.
FIGURE 3. Effect of [Mal]T (marked beside the lines) on photocatalyzed oxidation of Fe(II). [Fe(t)], 100 µM; Fe(III):Fe(II) ) 37:63; pH 3.0. slowly in the Mal system during the time scale of the experiment (Figure 2), whereas Fe(II) was rapidly oxidized to Fe(III) in the presence of oxalate within 15 min of irradiation. However, oxalate might be quickly mineralized (11) and then Fe(II) would revive from Fe(III) and further approach the photosteady state equilibrium of Fe(III)/Fe(II) (10). In contrast, for acetate and other di- or monocarboxylate systems (data not shown), photocatalyzed oxidation of Fe(II) showed behaviors similar to that without any carboxylates. The unique photoreaction behaviors of Ox and Mal among di- or monocarboxylates should be attributed to the different photoreaction mechanisms of their ferric complexes. Addition of Ox renders the acceleration of photocatalyzed oxidation of Fe(II), because photogenerated superoxide/ hydroperoxide radicals (eqs 5-6) may act as the oxidants sink of Fe(II) species in acidic solution (5). However, the reason photocatalyzed oxidation of Fe(II) was greatly retarded in the Fe/Mal system has not been reported. Therefore, it is of interest to examine the factors controlling the slow photocatalyzed oxidation of Fe(II) in the presence of Mal ligand. Figure 3 shows the photocatalyzed oxidation of Fe(II) at different [Mal]T. The control experiments (curves a, b) indicated that both light exposure and presence of O2 are indispensable for the Fe(II) oxidation even in the presence of Mal. In addition, [Mal]T in solutions should be closely associated with the photocatalyzed oxidation of Fe(II). Fe(II) was rapidly oxidized after 20 min of irradiation in the absence of Mal, whereas quite a long induction period for the photocatalyzed oxidation of Fe(II) was observed in irradiated
FIGURE 4. Rates of Fe(II) quantum yield (ΦFe(II)) and molar fraction of Fe(III) species as a function of [Mal]T. [Fe(III)], 100 µM; pH 3.0. Irradiation source: 100 W Hg lamp; irradiation time: 8 min. Six mM Ar-purged potassium ferrioxalate (0.05 M H2SO4) was used as a chemical actinometer. aerobic Mal-containing solutions. Higher initial [Mal]T led to the longer induction period. As more Fe(III) was introduced to the Fe(III)-Mal-Fe(II) solution at the beginning of the reaction, the induction period was shortened considerably (Figure S3). Our recent study proved that the duration of induction period for the Fe(II) oxidation is closely related to the content of FeIII(OH)2+ species in Fe(III)-Fe(II) systems (10), that is, the photocatalyzed oxidation of Fe(II) should be initiated and then accelerated by photolysis of FeIII(OH)2+. In the present Fe(III)/Mal systems, Mal can substitute OHligand of FeIII(OH)2+ and further influence the Fe(III)photocatalyzed oxidation of Fe(II), which has been verified by the following experiments. Fe(II) Photoproduction. According to Fe(III) speciation calculation, it is expected that Mal is able to control the proportion between Fe(III)-OH complexes and FMCs, which should further affect the photochemical reactions in irradiated solutions. Figure S4 shows the photoproduction of Fe(II) from Fe(III) at different [Mal]T (0-1000 µM). The addition of Mal led to an induction period for the Fe(II) production, which increased with [Mal]T. However, there was no induction period observed for Mal degradation (Figure S5), implying that slow Fe(II) production should be mainly attributed to the secondary (photo) chemical reactions involving reoxidation of Fe(II) but not primary photolytic reaction of FMCs. Quantum Yield (ΦFe(II)). The average ΦFe(II) at different concentrations of Mal was calculated according to the literature (eqs 9-11) (19). Where fi is the photon fraction absorbed by a specific species i; ci and εi are concentration and molar extinction coefficient of species i, respectively; Kai is the rate of light absorption by species i; I0 is the incident photon flux from the Hg lamp; l is optical path length; Φi is individual quantum yield of species i for Fe(II) formation; x(λ) is x (x ) fi or εi or I0) at certain wavelength λ. fi(λ) )
εi(λ)ci
∑ ε (λ)c i
(9) i
i
Kai )
∑ f (λ)I (λ)[1 - 10 ∑ i
0
λ
ΦFe(II) )
-εi(λ)cil
]∆λ
(10)
i
d[Fe(II)]/dt
∑K i
ai
∑ΦK
i ai
)
i
∑K
(11) ai
i
Figure 4 plots ΦFe(II) and molar fraction of Fe(III)-OH species against total malonate concentration (0-1000 µM). VOL. 44, NO. 1, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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SCHEME 1. Proposed pathways for photodecomposition of FMCs. For simplification, water ligand is omitted from (H2O)4Fe(III)Mal complex
FIGURE 5. ESR spectra of DMPO radical-adducts formed during the photolysis of Fe(III)/Mal solution under argon-purged conditions. DMPO, 0.04 M; Fe(III), 100 µM; pH, 3.0. (1a) experimental data, Mal, 1000 µM; (1b) simulated data of 1a; (2a) experimental data, Mal, 200 µM; (2b) simulated data of 2a. (*) DMPO- · CH2COOH, (V) DMPO- · OH. Since FeIII(OH)2+ was expected to have photoreduction reactivity roughly close to that of FeIII(OH)2+ due to their similar nLpOH-dLMCT absorption (25), FeIII(OH)2+ was regarded equivalent to FeIII(OH)2+ in this study. In the absence of Mal, the average ΦFe(II) (300-366 nm) was 0.028 ( 0.0014, which is comparable to the reported values by Faust and Hoigne´ (3) (Φ313 ) 0.14, Φ360 ) 0.017). However, with the increase of [Mal]T, molar fraction of Fe(III)-OH complexes and ΦFe(II) decayed rapidly. This suggests that contribution of the photolysis of FMCs to Fe(II) generation was minor. Assuming that the calculated quantum yield for FeIII(OH)2+ (0.028, 300-366 nm) is constant irrespective of the presence of Mal, then the quantum yield for individual FMCs can be further calculated by eq 11. The obtained wavelengthaveraged ΦFe(II) of FMCs was 0.0022 ( 0.0009, about 1 order of magnitude lower than that in Fe(III) control system (0.028 ( 0.0014). Since ΦFe(II) is the ratio of total Fe(II) formed to photons absorbed, all contributions from direct and indirect photochemical reactions are incorporated into the Fe(II) quantum yield. There was a significant difference between quantum yields for Fe(II) generation (0.0074 ( 0.0008) and Mal degradation (0.036 ( 0.005) in anaerobic solution containing 100 µM Fe(III) and 500 µM Mal. This indicates that low yield of Fe(II) generation is not due to poor photoactivity of Fe(III)-Mal complexes but rapid reoxidation of the photoproduced Fe(II). Under aerobic conditions, even lower quantum yield for Fe(II) formation (0.0016 ( 0.0007) was observed because of O2-involved reoxidation of Fe(II). Free Radicals Involved. Spin-trapping ESR technique was employed to identify the possible short-lived radicals involved in the reaction systems. It was carried out under anaerobic conditions to avoid the interference of quenching effect of dioxygen on carbon-centered radicals. Our results indicated the generation of two major kinds of radicals during the photoreaction of Fe(III)-Mal solutions at a lower concentration of Mal (200 µM) (2a in Figure 5). They were assigned to be DMPO- · CH2COOH (RH ) 23.01 G, RN ) 15.34 G) (26) and DMPO- · OH (RN ) RH ) 14.8 G) (27), respectively. The current ESR measurements provided direct evidence for the generation of · CH2COOH radical during the photodecarboxylation of FMCs. It is interesting to note that when 1000 µM of Mal was added, the DMPO- · OH signal was not observed. · OH radical is derived from the photolysis of FeIII(OH)2+, so the ESR signal intensity of DMPO- · OH adducts can be regarded as an indicative of concentration of photoactive FeIII(OH)2+. The generation of DMPO- · OH adduct in 2a (Mal, 200 µM) but not in 1a of Figure 5 (Mal, 1000 µM) supported our 266
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presumption that Mal can significantly decrease molar fraction of FeIII(OH)2+ by forming FMCs. Under this condition, · OH radical should be hardly detected. DMPO- · CH2COOH adduct increased with the irradiation time, with the rate of radical formation for curve a (Mal, 1000 µM) and b (Mal, 200 µM) in Figure S6, 35.2 and 5.9 au s-1, respectively. Mechanism Discussion. FeIII(OH)2+ and Fe(III)-Ox complexes have been considered as the predominant photoactive species in DOM-free (3) and DOM-rich atmospheric liquids (4, 5), respectively, and could be switched mutually with the influx and depletion of organic ligands (10, 11). The involved photochemical reactions have been extensively studied. Here we present another photochemical reaction pathway of iron redox cycling in the presence of malonate ligand (Scheme 1). Photoproduction of Fe(II). Fe(II) is formed through the Fe(III) photoreduction of major light-absorbing species in solutions: FMCs (process I) and Fe(III)-OH complexes (process II), respectively. Similar to photolysis of FeIII(OH)2+ species, FMCs with higher molar absorptivities transfer to an electronically excited state upon UV irradiation, followed by two competitive reactions: (1) return to ground state; (2) LMCT process with forming Fe(II) and a carboxylate radical in solvent cage (species a). The radical undergoes decarboxylation to yield CO2 and a Fe(II)/ · CH2COOH radical pair (species b). In this study, · CH2COOH radical has been detected by ESR-trapping technology (Figure 5), proving that the common LMCT process of the irradiated FMCs actually takes place. However, it seems difficult for Fe(II) to diffuse out of the solvent cage since the reoxidation of Fe(II) by · CH2COOH readily occurs (28-30) especially in the presence of uncomplexed Mal ligand. This quenching mechanism of Mal in Fe(II) formation is similar to that proposed in Cu(II)/Mal systems (31-34), which is further evidenced by measuring acetate as an intermediate (Figure S5). We observed the Mal degradation (Figure S7), providing a direct evidence that this quenching mechanism operates even in the absence of dioxygen. Mal degradation coupled with the Fe(II) reoxidation can explain the quite low Fe(II) quantum yield of FMCs. Therefore, once FMCs control the Fe(III) speciation, few Fe(II) are accumulated in solutions. FeIII(OH)2+ species gradually prevail over FMCs as Mal ligand is depleted (Figure 1). Once the photolysis of Fe(III)-OH complexes revives, Fe(II) would be produced rapidly (Figure S4). Photocatalyzed Oxidation of Free Fe(II). Fe(II) is very stable at pH 3 in the dark. The half-life for the Fe(II) oxidation by O2 is at least longer than 285 days (35). The presence of Mal ligand did not obviously accelerate the thermal/dark oxidation of Fe(II). Both light irradiation and presence of O2 are
indispensable for the Fe(II) oxidation (Figure 3). Since Fe(II) species at pH 3 do not absorb UV light (>300 nm) (3), it is Fe(III)-OH complex or FMCs that probably acts as a main chromophore to photocatalyze Fe(II) oxidation by O2. Our previous study has proved that the photocatalyzed oxidation of Fe(II) in aerated solutions is closely related to the content of FeIII(OH)2+ species in Fe(III)-Fe(II) systems (10). Although photogenerated Fe(II)/ · CH2COOH radical pair may also react with O2 to form a relatively stable six-membered chelating ring (species c) that can further decompose or oxidize another proximate ferrous ion (k ) (5.0 ( 1.0) × 106 M-1 s-1) (30), this process only give minor contribution to catalytic oxidation of Fe(II) since the free Fe(II) in solutions was oxidized more slowly when FMCs dominated the Fe(III) speciation (Figure 3). Therefore, it should be Fe(III)-OH complex but not FMCs that is pivotal for Fe(II) oxidation by O2. In conclusion, the presence of Mal greatly limits the formation of Fe(III)-OH complex, thereby indirectly preventing Fe(II) against oxidation. This phenomenon is unique among all di- or monocarboxylates studied since photolytic reaction of ferrioxalate • accelerates the formation of O•2 /HO2 oxidants (eqs 5-6) and Fe(III)-OH complex is always the prevailing Fe(III) species when other di- or monocarboxylates besides oxalate and malonate are present. Environmental Implications. Our study reveals that malonic acid plays a quite different role in the photoredox cycling of iron relative to oxalic acid and other di- or monocarboxylic acids. Both oxalate and malonate are possible subproducts from the chain reactions of dicarboxylic acids (35) and both can form stable Fe(III) complexes with high molar absorptivities, but their fates are quite different. Fe(III)-oxalato complexes are very photosensitive and can easily generate reducing CO•2 radical upon irradiation, which can further reduce another Fe(III)-oxalato complex (eq 4) and thereby enhance the Fe(II) formation (5). However, FMCs have much lower Fe(II) quantum yield because an oxidizing · CH2COOH radical is generated and reoxidizes the photogenerated Fe(II) in the presence of excess Mal. Therefore, the different photoreaction pathways and nature of radicals derived from oxalate and malonate result in their opposite roles in iron cycle. Zuo et al. reported that the Fe(III)-Fe(II)-Fe(III) cycling time for Fe(III)-Ox complex was on the order of only 100 s (7). The half-life of Fe(III)-aquo complexes was approximately 9 min at pH 3 (36). Our study reveals that high level of Fe(II)/Fe(t) (Figure 3) can be maintained for a long time in the presence of excess Mal. Many field measurements indicated that Fe(II) accounts for most of the dissolved iron in rainwaters (23, 24, 37), but the mechanism involved for (photo)stability of Fe(II) is unclear (38). It is of interest to totally identify the dissolved organic matters in rainwaters, particularly dicarboxylic acids, to understand the speciation of dissolved iron in carbon-rich troposphere, although the correlation of Mal with Fe(II) against (photo)oxidation in rainwater remains obscure so far because rainwater may have high concentrations of other oxidants generated via non metal redox pathways. Despite that the concentration of Fe(III) and Mal in this study are higher than those found in most atmospheric waters, which usually contain micromolar to nanomolar concentrations of Fe and Mal, the ratios of Fe to Mal in our study (1:1-1:10) should be representative of those in rainwater, fog, snow, and cloud waters (see Table 3 in ref (1) and Table S4). Concentrations of dissolved iron up to 200 µM (greater than these in the present study) have also been deemed environmentally relevant in some cases (2). Therefore, our findings in this study would be of great significance in evaluating photochemical cycling of iron in the troposphere.
Acknowledgments This work was financially supported by 973 project (2007CB613306 and 2010CB933503), NSFC (20537010, 20677062 and 20777076), and CAS/SAFEA.
Supporting Information Available Additional experimental evidence. This material is available free of charge via the Internet at http://pubs.acs.org.
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