Iron Dissolution of Dust Source Materials during Simulated Acidic

Article pubs.acs.org/est

Iron Dissolution of Dust Source Materials during Simulated Acidic Processing: The Effect of Sulfuric, Acetic, and Oxalic Acids Haihan Chen† and Vicki H. Grassian*,†,‡ Departments of †Chemical and Biochemical Engineering and ‡Chemistry, University of Iowa, Iowa City, Iowa 52242, United States S Supporting Information *

ABSTRACT: Atmospheric organic acids potentially display different capacities in iron (Fe) mobilization from atmospheric dust compared with inorganic acids, but few measurements have been made on this comparison. We report here a laboratory investigation of Fe mobilization of coal fly ash, a representative Fe-containing anthropogenic aerosol, and Arizona test dust, a reference source material for mineral dust, in pH 2 sulfuric acid, acetic acid, and oxalic acid, respectively. The effects of pH and solar radiation on Fe dissolution have also been explored. The relative capacities of these three acids in Fe dissolution are in the order of oxalic acid > sulfuric acid > acetic acid. Oxalate forms mononuclear bidentate ligand with surface Fe and promotes Fe dissolution to the greatest extent. Photolysis of Fe−oxalate complexes further enhances Fe dissolution with the concomitant degradation of oxalate. These results suggest that ligand-promoted dissolution of Fe may play a more significant role in mobilizing Fe from atmospheric dust compared with proton-assisted processing. The role of atmospheric organic acids should be taken into account in global-biogeochemical modeling to better access dissolved atmospheric Fe deposition flux at the ocean surface.

■

INTRODUCTION Iron (Fe) is an important limiting factor for phytoplankton growth in extensive regions of the ocean referred to as high nutrient low chlorophyll (HNLC) regions.1−3 While the transport and deposition of atmospheric dust is generally believed to be the main source of bioavailable Fe in open ocean waters,4−8 the flux of soluble Fe from various atmospheric dust remains as a critical uncertainty to evaluate global Fe connections in the Earth’s system.7,9 Fe solubility of atmospheric dust is closely related to intrinsic particle properties, such as size, surface area, and Fe speciation, as well as Fe distribution among single particles.10−23 Atmospheric chemical processing of aerosols during longrange transport enhances Fe solubility as suggested in many field, laboratory, and model studies.6,15,24−31 For example, the association of typical atmospheric inorganic acids onto aerosols can cause a dramatic decrease of surface pH along with a rapid release of dissolved iron into the adjacent aqueous phase.27,28 While this acidic processing by inorganic acids has been extensively explored,12,17,22,32 few studies have focused on the effect of organic acids. Indeed, a large number of organic acids such as acetic acid and oxalic acid that originate from anthropogenic and biogenic sources have been identified in the atmosphere and are often associated with aerosol particles.34,35 Low-molecular weight organic acids have been verified as major products in the photodegradation of volatile organic compounds in the atmosphere.89 Studies of precipitation chemistry have shown © 2013 American Chemical Society

that organic acids may account for a large fraction, up to 64%, of the total acidity in nonurban environments.36 Due to the polar nature of organic acids and the alkaline nature of dust particles, these organic acids are preferentially transferred to the aerosol phase.35,37 Field studies have widely reported the adsorption of organic acids on dust particles.37−41 These adsorbed organic acids potentially display different capacities in Fe mobilization compared to inorganic ones. Organic complexation has been suggested to promote Fe dissolution from typical components of dust and authentic dust.6,14,42−46 A positive correlation between soluble Fe and oxalate has been observed in ambient aerosol samples collected over the North Atlantic region.29,47,48 Laboratory experiments have also shown the enhancement of Fe dissolution from Fecontaining minerals and authentic dust in the presence of organic acids.14,44,76 In a model study to estimate soluble iron fluxes to the ocean, the incorporation of oxalate into the model showed a better correlation with the observations compared to the estimation with inorganic acids alone.33 Photoreduction of Fe(III) can also enhance Fe dissolution in atmospheric and ocean waters.49,50 Studies on photolysis of Feorganics complexes have shown the formation of dissolved Fe(II), a more bioavailable form of Fe, along with the Received: Revised: Accepted: Published: 10312

March 23, 2013 May 24, 2013 July 24, 2013 July 24, 2013 dx.doi.org/10.1021/es401285s | Environ. Sci. Technol. 2013, 47, 10312−10321

Environmental Science & Technology

Article

Figure 1. Dissolution of fly ash SRM 2689 as a function of time in pH 2 solutions acidified by sulfuric acid (6.6 mM), acetic acid (5.7 M), and oxalic acid (11.7 mM), respectively, at a solid loading of 2 g L−1. Measured dissolved Fe is shown as (a) dissolved Fe(II), (b) dissolved Fe(III), (c) total dissolved Fe, and (d) the fraction of total dissolved Fe present as Fe(II). When present, error bars represent one standard deviation from triplicate experiments.

■

decomposition of organics.14,25,26,51−58 The dissolved Fe(II) can be reoxidized to form a highly dispersed Fe(III) colloid, which is more soluble and reactive than more crystalline iron oxide.59 The Fe(II)−Fe(III) redox cycling driven by solar radiation is believed to enhance iron bioavailablitiy in open ocean waters.49,60 In our earlier study, we investigated Fe dissolution of coal fly ash (FA) samples, precursors of anthropogenic aerosols, and Arizona test dust (AZTD), a reference source material of mineral dust, in sulfuric acid solutions.21 Simulated atmospheric processing with sulfuric acid causes the disintegration of aluminosilicate glass that presents as a dominant material in FA, and thus enhances Fe dissolution.21 In contrast, AZTD mainly contains crystalline aluminosilicate mineral that is relatively stable during simulated atmospheric processing.21 In order to better understand the role organic acids play in Fe mobilization and compare the capacities of Fe mobilization between organic and inorganic acids, Fe dissolution of the three FA samples, as well as AZTD, is explored in aqueous environment acidified by various atmospherically relevant acids. Sulfuric acid, acetic acid, and oxalic acid are selected given that they are typical and abundant species present in the atmosphere and often associated with the aerosol phase.89 The aqueous environment at pH 2 and 3 is chosen to mimic the thin film of water, which is characteristic of low pH and high dust-to-liquid ratio, on the dust surface after cloud evaporation.61 The effects of ligand concentration, pH, and solar radiation on Fe mobilization by the organic acids are explored.

EXPERIMENTAL SECTION Reagents and Materials. All chemicals listed in the Supporting Information were reagent grade or better. Three FA standard reference materials (SRMs 2689, 2690, and 2691, National Institute of Standards and Technology), and Arizona fine test dust (ISO 12103-1 A2 test dust, Power Technology Inc.) were used as received. FA and AZTD samples have been characterized by various techniques discussed in previous publications.12,21 Details of the characterization methods are also provided in the Supporting Information. Dissolution Measurements. Dissolution experiments were performed in a glass vessel containing an appropriate aqueous solution. Solutions were either acidified with sulfuric acid to a desired pH with the concentration of acetate or oxalate, which ranged from 0 to 2 mM, or acidified with acetic acid or oxalic acid to a desired pH alone. The solution pH was controlled to be 2 ± 0.1 or 3 ± 0.1, which is within the low range found on the aerosol surface and in cloud droplets.61 Since minor variations in ionic strength did not significantly influence the results of our previous dissolution studies,12 the ionic strength for these dissolution experiments was not controlled. All solutions were stirred under oxygen exposure and kept at 298 K using a water jacket integrated into the vessel. The pH was monitored with a glass electrode standardized with pH buffer solutions. Experiments were performed at the desired pH and were adjusted with concentrated HCl if the pH value changes during the dissolution procedure. The solid loading was 2 g L−1. After addition of the particles into the solution, aliquots of suspension were periodically withdrawn for analyses 10313

dx.doi.org/10.1021/es401285s | Environ. Sci. Technol. 2013, 47, 10312−10321


Article

specific surface areas and Fe contents of FA and AZTD samples are tabulated in Table S1 in the Supporting Information. Influence of Different Acids on Fe Dissolution. Figure 1 shows dissolution of FA 2689 at pH 2 in sulfuric acid, acetic acid, and oxalic acid containing solutions. All acidic media show the increase of dissolved Fe species including both Fe(II) and Fe(III) with time but with different dissolution rate. Oxalic acid displayed the highest capacity in mobilizing Fe, with the total dissolved iron (Fedis) reaching 44 ± 3% after 45 h, followed by sulfuric acid with a Fedis value of 16 ± 1%. The Fedis in acetic acid is lowest, showing a value of 9 ± 1% after 45 h of dissolution. Proton-promoted and ligand-promoted dissolution are two main mechanisms of Fe dissolution from oxides and aluminosilicates.42,45,59,62,63 Dissolution depends on the surface interaction that takes place on the particles.42,45,64 Surface functional groups, such as H+ and OH−, are able to interact with H+ ions and ligands to form surface complexes. The formation of surface complexes results in the weakening of bonds in the proximity of a surface Fe ion center and causes detachment of Fe into the aqueous phase.42,45,64 All acidic media at the same pH (pH 2) should have similar H+-ion activity to promote Fe dissolution. The different capacities of the three acids in Fe mobilization shown here is due to the formation of different Fe−ligand complexes. Oxalate, which can form bidentate ligand with Lewis acid Fe centers, brings electron density into the coordination sphere of surface Fe, labilizes surface Fe−O bond, and eventually enhances Fe release into the adjacent solution. In addition, the complexation of oxalate with surface Fe(III) can act as an electron bridge to facilitate electron transfer between dissolved Fe(II) and surface Fe(III) and promote the dissolution of surface Fe(III).43 Sulfuric acid displays moderate capacity in iron mobilization given that HSO4− and SO42− can only form weak complexes with iron in the aqueous phase. Iron solubility in acetic acid solution was lowest and is attributed to the formation of even weaker complexes of acetic acid with iron.46 The monocarboxylic anion has a small effect on iron mobilization. It has been previously reported that dissolution of ripidolite, an Fe-containing mineral, by acetic acid appeared to be mainly related to proton-promoted dissolution, whereas sulfuric acid and oxalic acid significantly enhanced dissolution by forming metal−ligand complexes with framework metals.46 Dissolution of another Fe-containing mineral, serpentinite, also showed that sulfuric acid is more efficient at extracting Fe from serpentinite than acetic acid.65 Our dissolution results indicate that ligandpromoted dissolution could be as important as protonpromoted dissolution in controlling Fe mobilization from atmospheric dust. The results are in good agreement with a recent laboratory measurement, in which a positive correlation between oxalate concentration and Fe solubility from various dust sources was observed.76 Figure 1d displays the fractional Fe(II), i.e. Fe(II)/Fedis, in each dissolution experiment as a function of time. The Fe(II)/ Fedis stayed relatively constant at 27 ± 3% after 10 h of dissolution in sulfuric acid and acetic acid solutions, indicating that there was no redox reaction occurring in both systems. The initial difference of Fe(II)/Fedis can be explained by the differences in the dissolution rates of various Fe species in different acidic media. In contrast, the Fe(II)/Fedis in oxalic acid solution was lower and displayed an initial increase followed by a slightly decrease with time. Due to the high stability of Fe(III)−oxalate

of dissolved Fe(II) as well as total dissolved Fe, and oxalate if desirable. In order to explore the effect of solar radiation, a solar simulator with a 150 W xenon lamp (Oriel, model 67025) was mounted on the top of the vessel to allow irradiated experiments. For reference, dissolution experiments under dark conditions were conducted with the vessel wrapped in aluminum foil. Experiments were conducted in triplicate, and results represent the average and standard deviation of three measurements. The details of the characterization and analytical methods are provided in the Supporting Information. ATR-FTIR Spectroscopy. The adsorption of acetic acid and oxalic acid on FA 2689 were probed with attenuated total reflectance Fourier transform infrared (ATR-FTIR) spectroscopy using a Thermo Nicolet 670 FTIR spectrometer equipped with a MCT (mercury cadmium tellurium) detector. A thin, evenly coated FA 2689 film was deposited onto a germanium crystal mounted inside a horizontal ATR cell (Pike Technologies Inc.) by placing a suspension of FA 2689 (10 mg in 1 mL of Milli-Q water) onto the crystal surface and drying overnight. The deposited film was slowly flushed with Milli-Q water to remove any loosely bound particles. Acetic acid or oxalic acid solution at pH 2 was then introduced into the ATR cell, and the FTIR spectra were recorded. All spectra were collected in the range of 500−4000 cm−1 at a resolution of 4 cm−1, and 50 scans were averaged from each spectrum.

■

RESULTS AND DISCUSSION Characterization of FA and AZTD Particles. As discussed in previous publications,21 the characteristic morphologies of FA particles are mainly spherical with smaller particles typically attached to larger ones. Most of the particles ranged in size from less than 1 μm to greater than 10 μm. In contrast, AZTD contains particles that are more irregular in shape with particles in the size range extending from 0.1 to 3 μm. FA and AZTD samples are mainly composed of various oxides such as SiO2, Al2O3, Fe2O3, CaO, and MgO, as well as aluminosilicates. While well-crystalline aluminosilicate minerals are present as the dominant component of AZTD, FA samples were characterized to be mainly composed of aluminosilicate glass, which has a lower stability compared to the more crystalline aluminosilicate minerals. The Fe content of FA 2689, 2690, and 2691 certified by NIST are 9.32 ± 0.06%, 3.57 ± 0.06%, and 4.42 ± 0.03%, respectively. The Fe content of AZTD was determined to be 1.98 ± 0.08% using a digestion method that has previously been described.12 Among the four samples, AZTD has the highest specific surface area of 4.2 ± 1.0 m2 g−1, followed by FA 2690 and FA 2691, whose specific surface areas are 3.8 ± 0.1 and 2.2 ± 0.1 m2 g−1, respectively. FA 2689 exhibits the lowest specific surface area of 0.8 ± 0.1 m2 g−1. Fe speciation in FA and AZTD particles was determined in earlier studies using various techniques including Mössbauer spectroscopy.12,21 All three FA samples contain Fe(III) in oxides such as hematite or goethite and Fe(III) in aluminosilicate phases. In addition, a considerable amount of Fe(II) in aluminosilicate glass was identified for FA 2689 and FA 2690. AZTD was quantitatively determined to contain Fe(III)-substituted aluminosilicates and Fe(III) in oxides. An appreciable amount of Fe in AZTD exists as Fe(II)-substituted into aluminosilicates. Fe speciation plays a crucial role in controlling Fe dissolution of particle samples as discussed below and in previous studies.21 The speciation of Fe as well as 10314



Article

and dissolved Fe species after 45-h dissolution in each system are tabulated in Table S2 in the Supporting Information. While all of acetic acid solutions exhibited a lower capacity in Fe mobilization than sulfuric acid solution, the concentration of acetate only has small effect on Fe dissolution rate. Unlike acetic acid, the presence of oxalic acid greatly influences Fe dissolution, showing a positive correlation of oxalate concentration with Fe solubility. More interestingly, sulfuric acid containing 0.1 mM oxalate does not show a significant effect on Fe dissolution, and the total dissolved Fe concentration after 45-h experiment was approximately 0.7 mM, about seven times larger than the concentration of oxalate (0.1 mM). At low concentration of oxalate, dissolved Fe(III) effectively competes with the FA surface for oxalate through formation of Fe(III)−oxalate solution complexes, suppressing surface complexation. The similar competition between dissolved Al(III) and boehmite, corundum as well as aluminum−(oxy)hydroxides has previously been reported.67−69 It was proposed that the competition is more significant at low pH. Excess oxalate is therefore required to enhance Fe dissolution. Previous studies on the effect of oxalate on Fe dissolution from iron oxides and atmospheric dust were performed in the presence of excess oxalate and, therefore, did not observe the competition of oxalate between surface Fe and dissolved Fe.14,76 Influence of pH on Fe Dissolution and Dissolved Fe Speciation. Figure 3 shows the effect of pH on Fe dissolution in acetic acid and oxalic acid. As anticipated from established dissolution trends for Fe-containing atmospheric dust,12,21,70 increasing pH from 2 to 3 produced roughly 60% and 80% decrease in the concentrations of total dissolved Fe in acetic acid and oxalic acid solutions, respectively. Compared to acetic acid, Fe dissolution in oxalic acid showed a more significant decrease with increased pH. The concentrations of oxalate in pH 2 and 3 solutions are around 11.7 and 0.89 mM (Supporting Information Table S2), respectively. Given that Fe dissolution in oxalic acid solution is mainly driven by ligand complexation as discussed above, the drastic decrease of oxalate concentration with the increase of pH significantly suppressed Fe release from FA in oxalic acidic medium. The solution pH could also influence surface complexation of oxalate with Fe by controlling charge characteristics of both the surface and the ligand. Increasing pH makes the surface less positively charged, whereas the acid dissociates to become more negative, resulting in a less likelihood of the acid to bind to the surface due to unfavorable electrostatic interactions. 71 The Fe(II)/Fedis increased with increasing pH in both experiments as shown in Figure 3b. Such changes in dissolved Fe speciation are due to the greater solubility of Fe(II) relative to Fe(III) solution at high pH. Similar dissolution behavior of FA 2689 in sulfuric acid solution has previously been reported.21 Adsorption of Acetic Acid and Oxalic Acid on FA Particles. To better understand the interaction of organic acids with FA, ATR-FTIR spectroscopy studies were done. Figure 4 shows ATR-FTIR spectra collected after the exposure of FA 2689 particles to pH 2 acetic acid or oxalic acid. Also shown in Figure 4 are the spectra of acetic acid and oxalic acid solutions in the absence of FA particles. Acetic acid (Figure 4a) shows peaks at 1711, 1390, 1369, 1277, and 1015 cm−1, which are attributed to protonated acetic acid.72−74 The spectrum of adsorbed acetic acid on FA particles (Figure 4b) shows no additional peaks compared to aqueous phase acetic acid. The subtraction of Figure 4a from Figure 4b confirmed no

complexes, oxalate might preferentially mobilize Fe(III) other than Fe(II). The catalytic dissolution of Fe(III) in oxalate− Fe(II) solution can also contribute to the higher Fe(III) dissolution.43 Another plausible explanation for the lowest Fe(II)/Fedis in oxalic acid is that oxidation of Fe(II) into Fe(III) may occur in a considerable rate in aerated oxalic acid solution. Although the oxidation of Fe(II) is very slow at pH below 4 (e.g., t1/2 ∼ years in air-saturated water),66 Fe(II) oxidation can be accelerated by Fe−ligand complexation to change the reduction potential of Fe and create a labile coordination position capable of forming an inner-sphere complex with O2.64 Our previous research reported that FA particles can continually release dissolved Fe in pH 2 sulfuric acid as FA particles break up into irregular-shape fragments.21 The processed FA particles after dissolution in acetic and oxalic acid were collected in this study for TEM imaging as shown in Figure S1 in the Supporting Information. The disintegration of FA spheres also occurred during simulated acidic processing in both acetic and oxalic acid. Similar dissolution trends were observed for AZTD in the three acids investigated, as shown in Figure S2 in the Supporting Information. The total dissolved Fe after 45 h in pH 2 oxalic acid is 3.6 ± 0.3 and 4.4 ± 0.5 times greater than those in sulfuric acid and acetic acid, respectively. The results indicate that enhancement of Fe dissolution through Fe− oxalate complexation is generally applicable to mineral dust. Influence of Ligand Concentration on Fe Dissolution. The effect of ligand concentration on Fe dissolution of FA 2689 was also investigated. Sulfuric acid solutions with known concentrations of acetate or oxalate at pH 2 were used, and the results are shown in Figure 2. Also plotted in Figure 2 are the Fe dissolution of FA 2689 in pH 2 sulfuric acid, acetic acid, and oxalic acid for better comparison. The concentrations of anions

Figure 2. Total dissolved Fe production as a function of time from fly ash SRM 2689 in pH 2 sulfuric sulfuric acid (6.6 mM), acetic acid (5.7 M), and oxalic acid (11.7 mM), as well as sulfuric acid containing known concentrations of acetate and oxalate, at the solid loading of 2 g L−1. When present, error bars represent one standard deviation from triplicate experiments. 10315



Article

Figure 4. ATR-FTIR spectra of (a) pH 2 acetic acid (HAc) solution, (b) absorbed acetic acid on fly ash SRM 2689 particles, (c) pH 2 oxalic acid (H2Ox) solution, (d) absorbed oxalic acid on fly ash SRM 2689 after 1-min exposure, (e) absorbed oxalic acid on fly ash SRM 2689 after 10-min exposure, and (f) differential spectrum of parts e and d.

Figure 3. Influence of pH on Fe dissolution of fly ash SRM 2689 as a function of time in solutions acidified by acetic acid and oxalic acid, respectively, at a solid loading of 2 g L−1. The concentrations of acetic acid solution at pH 2 and 3 are 5.76 M and 58.5 mM, respectively. The concentrations of oxalic acid solution at pH 2 and 3 are 11.7 mM and 0.89 mM, respectively. Measured dissolved Fe is shown as (a) total dissolved Fe and (b) the fraction of total dissolved Fe present as Fe(II). When present, error bars represent one standard deviation from triplicate experiments.

oxalate with surface metal atoms forms chemically adsorbed species, significantly enhancing Fe dissolution from dust particles. Dissolution of FA 2690 and FA 2691 in Oxalic Acid. Dissolution experiments were also performed for two other FA samples, FA 2690 and FA 2691, in pH 2 oxalic acid to explore the impact of Fe speciation on Fe mobilization, and the results are shown in Figure S3 in the Supporting Information. FA 2691 yielded the highest total dissolved Fe among the three FA samples. Approximately 78% of the total Fe in FA 2691 dissolved into the solution after 45 h of simulated acidic processing, compared to 55% and 44% in FA 2690 and FA 2689, respectively. A portion of Fe in FA 2690 and FA 2691 is highly soluble, and thus, rapidly dissolved into the aqueous phase upon the addition of FA particles into the solution. Although there is no Fe(II) in FA 2691 particles according to Mössbauer spectroscopy analysis as discussed above and shown in Supporting Information Table S1, around 6% of total dissolved Fe present as dissolved Fe(II) in FA 2691 suspension. Redox reactions in the Fe(III)−oxalate system may occur with the formation of Fe(II). It is also possible that FA 2691 contained small amount of Fe(II), but low abundance of Fe(II) prevents its detection by Mössbauer spectroscopy. Linear correlation coefficients between dissolved Fe fractions, including Fe(II) (Fe(II)dis), Fe(III) (Fe(III)dis), and total dissolved Fe (Fedis), obtained from dissolution in oxalic acid, and the physicochemical properties, obtained from particle characterization, of three FA samples were calculated and tabulated in Table S3 in the Supporting Information. The Fe(III)dis is highly correlated, to the total Fe(III) species in FA particles with a correlation coefficient of 0.98. In contrast, the Fe(III)dis is moderately correlated (R = 0.74) to Fe(III) in aluminosilicate glass and is limitedly but positively correlated

difference except the absorbance bands of water, indicating no or very little adsorption of acetic acid on FA particles. In contrast, aqueous phase oxalic acid (Figure 4c) does not show any peaks at the low concentration of oxalate. However, adsorbed oxalic acid on FA particles shows several absorption bands with peaks at 1718, 1683, and 1407 cm−1 (Figure 4d), followed by bands arising at 1612 and 1312 cm−1 (Figure 4e). The difference spectrum of Figure 4d and 4e (Figure 4f) clearly shows the two sets of absorbance bands with different formation rates, indicating the formation of two different surface complexes of oxalate on FA particles. Previous studies of oxalate adsorption on oxides have assigned the absorbance bands at 1718, 1683, and 1407 cm−1 to inner-sphere complexes, most possibly a 5-member ring mononuclear bidentate structure.14,68,74,75 Absorbance bands at 1612 and 1312 cm−1 can be attributed to outer-sphere surface complexes of oxalate forming through electrostatic and H-bonding interactions after surface saturation of inner-sphere oxalate complexes.14,68,74,75 A negative band at 1015 cm−1 also shows along with oxalate adsorption and is due to ν(Si−O), indicating the detachment/ dissolution of FA particles during the course of the experiment. ATR-FTIR spectra revealed that acetic acid was not adsorbed or molecularly adsorbed on FA particles with no direct bonding to surface metal atoms. It therefore shows little effect on Fe dissolution as discussed earlier. In contrast, the complexation of 10316



Article

(R = 0.13) to Fe(III) in oxides. Although a large portion of dissolved Fe(III) originated from aluminosilicate Fe(III), iron oxides also played a role. Similarly, Fe(II)dis is mainly derived from the Fe(II) species in FA particles. The correlation coefficient between Fe(II)dis and the Fe(II) species in FA particles is 0.74. A correlation coefficient of 0.27 between specific surface area (SA) and Fedis indicates that specific surface area might also play a role in controlling Fe dissolution. Unlike FA−sulfuric acid solution, in which dissolved Fe was mainly released from Fe in aluminosilicate glass as previously reported,21 both Fe in oxides and aluminosilicates provied dissolved Fe in FA−oxalic acid solution. One plausible explanation is that the predominant mechanisms in promoting Fe dissolution from FA particles are different in sulfuric and oxalic acid solutions as discussed earlier. In addition, although Fe in aluminosilicates is more mobile than Fe in oxides, steric hindrance might limit the diffusion of oxalate ligand into the aluminosilicate layers to destabilize Fe−O bonds. It is noteworthy that the calculation of correction coefficient here might be ambiguous given the possible oxidation of Fe(II)dis into Fe(III)dis in oxalate-containing system. Future studies are needed to clarify the relative solubility of Fe in oxides and aluminosilicates. A recent study of oxalic acid in mobilizing Fe suggested that Fe in oxides plays a negligible role, and more than 95% dissolved Fe derived from clays.76 Influence of Solar Radiation on Fe Dissolution and Dissolved Fe Speciation. The dissolution experiment of FA 2689 was performed under irradiation in pH 2 oxalic acid solution to investigate the effect of solar radiation on Fe dissolution. Oxalate content and solution pH, as well as Fe species, were periodically determined during the whole experiment, and the results are shown in Figure 5. Oxalic acid under irradiation (Figure 5b) displayed a threestage decay: (i) initial slow decay, (ii) early fast decay, and (iii) late slow decay. The solution pH (Figure 5a) was controlled by oxalic acid, and was therefore negatively correlated with oxalic acid. While a controlled dissolution experiment under irradiation in the absence of FA particles showed no significant decay of oxalic acid (data not shown), the observed decay of oxalic acid in FA-containing system is attributed to Fe-catalyzed photodegradation of oxalate. The initial slow decay of oxalic acid indicates an induction period that can be attributed to the photodegradation of iron− oxalate complexes on the FA surface14,25,54,77−80

Figure 5. Dissolution profile of fly ash SRM 2689 as a function of time at pH 2 for oxalic acid solutions under dark and irradiation conditions, respectively. The results show the evolution of (a) pH, (b) the concentration of oxalic acid, (c) total dissolved Fe, and (d) the fraction of total dissolved Fe present as Fe(II). H+

C2O4 − · + O2 ⎯→ ⎯ HO2 ·/O2− · + 2CO2 H+

HO2 ·/O2− · + Fe2 +(aq) ⎯→ ⎯ H 2O2 + Fe3 +(aq)

(1)

>Fe IIC2O4 − · (ad) ↔ >Fe II + CO2 + CO2−

(2)

Irradiation promotes ligand-to-metal electron transfer in the Fe(III)−oxalate complexes, leading to the reduction of structural Fe(III) to Fe(II) along with the oxidation of oxalate. The surface-associated Fe(II) can then detach from the surface to yield dissolved Fe(II). The produced dissolve Fe species can further complex with oxalate in aqueous phase. The Fe(III)−oxalate complex strongly absorbed irradiation in the UV−vis region (290−570 nm) and thus are photochemically reactive54 [Fe3 +(C2O4 2 −)n ](3 − 2n) (aq) hv

↔ Fe2 +(aq) + (n − 1)C2O4 2 − + C2O4 −

(5)

Photolysis of ferrioxalate yields Fe(II) and oxalate radical (reaction 3).51 The oxalate radical then interacts with oxygen to form superoxide (O2−·) and its conjugated acid, hydroperoxide radicals (HO2·) (reaction 4). The HO2·/O2−· is a strong oxidant and can rapidly reoxidize Fe(II) to Fe(III), leading to the formation of H2O2 (reaction 5). Although other oxidants such as H2O2 and O2 can also reoxidize Fe(II), the reaction rate is much slower than that by HO2·/O2−·.51,54 The HO2·/O2−· oxidant can also react with Fe(III), but the reaction is suppressed in the presence of oxalate.81 With the accumulation of dissolved Fe species, the photodegradation of oxalic acid appeared to be autocatalytic with time, and its concentration therefore displayed a drastic decrease. With the consumption of oxalate and the increase of pH, Fe(III) started to precipitate, showing a rapid decrease of total dissolved Fe. Surface adsorption of oxalate on the particle surface is also suppressed.25,82 The less likelihood of surface complexation of oxalate, and the decrease of dissolved Fe(III) eventually led to a slow decay of oxalic acid after 7 h of dissolution. Compared to dark dissolution, photoreduction of Fe along with photodegradation of oxalate enhances the Fe release from FA 2689 (Figure 5c). Similar effects of irradiation in enhancing Fe dissolution has previously been observed by Cwiertny et al. in a study of photodissolution of α-FeOOH nanorods and microrods in oxalate.14

hv

>Fe IIIC2O4 −(ad) ↔ >Fe IIC2O4 − · (ad)

(4)

(3) 10317



Article

additional detail in future work. Global-biogeochemical modeling to assess dissolved atmospheric Fe deposition fluxes should be constrained by taking into account various atmospheric organic species and their capacities in forming soluble complexes with surface Fe on atmospheric dust.

The 10-h dissolution experiment under irradiation was also performed for AZTD (Figure S4 in the Supporting Information). Similar trends of oxalate decay and pH rising were also observed in an AZTD-containing system. Unlike FA 2689, which displayed an induction period of oxalate decomposition, oxalic acid in the AZTD-containing system initially showed a rapid decay. This can be explained by the highly soluble Fe in AZTD particles. The rapid release of Fe from AZTD particles led to photodegradation of oxalate in aqueous phase through reactions 3−5 and masked the induction period.

■

ASSOCIATED CONTENT

S Supporting Information *

Summary table of the physical and chemical properties of three FA samples and AZTD (Table S1); concentrations of anions and total dissolved Fe in different systems used in the study (Table S2); correlation coefficients between different parameters of FA samples (Table S3); TEM images of processed FA (Figure S1); Fe dissolution of AZTD in sulfuric, acetic, and oxalic acids (Figure S2); Fe dissolution of three FA samples in oxalic acid (Figure S3); Fe dissolution of AZTD in oxalic acid under irradiation (Figure S4). This material is available free of charge via the Internet at http://pubs.acs.org.

■

ATMOSPHERIC IMPLICATIONS The dissolution experiments of FA and AZTD dusts in atmospherically relevant acids, including sulfuric acid, acetic acid, and oxalic acid, were performed to investigate the effect of acidic processing on Fe dissolution of atmospheric dust. The capacities of the three acids on Fe mobilization are in the order of oxalic acid > sulfuric acid > acetic acid. Compared to protonpromoted dissolution, ligand-promoted dissolution may play a more important role in mobilizing Fe from atmospheric dust. The capacity of organic species in Fe mobilization is critically dependent on the ability of these anions to form complexes with surface Fe. A ligand, like oxalic acid, can form multiple bonds with surface Fe which can significantly enhance Fe mobilization. In contrast, monodentate ligands, like acetic acid, might only molecularly or not adsorb on the particle surface, and are less efficient for Fe mobilization. When both oxalate and sulfate present in aerosols, as usually observed in anthropogenic aerosols, the Fe−oxalate complexes are preferentially formed.33,48 Our results also suggest that the presence of oxalate in aerosol waters greatly facilitate Fe dissolution. However, under low concentration of oxalate, the competition of dissolved iron and surface iron for oxalate complexation suppress the Fe dissolution rate. Concentrations of oxalate and Fe in atmospheric waters are comparable, and both at micromolar levels.81,83−86 The influence of oxalate on Fe dissolution therefore can play a role, but this may be spatially and temporally dependent. Photochemical redox reactions of Fe− oxalate complexes can further enhance Fe dissolution with formation of more bioavailable Fe(II) and the decomposition of oxalate. The precipitation of Fe(III) into solid phase would occur after the consumption of oxalate, forming highly disperse and reactive Fe(III) colloids. Given that atmospheric oxalic acid is mainly derived from photochemical oxidation of hydrocarbons in the gas phase and in tropospheric liquids,87−89 formation of oxalic acid during atmospheric processing may continually provide oxalate to facilitate Fe mobilization. The photoreduction of Fe−oxalate complexes might also represent as a significant sink of atmospheric oxalate. While ligand-promoted Fe dissolution is highlighted in this study, a large number of atmospheric organic species is preferentially associated with atmospheric dust aerosols and can form soluble complexes with surface Fe of atmospheric dust.37−41 These organic compounds can also serve as electron donors to reduce Fe(III). It is also noteworthy that while aerosol pH is usually in the range of 3−5,22,61,90,91 and oxalate is at a micromolar level,81,83−85 further studies under more atmospherically relevant conditions are needed. To better understand the role of atmospheric dust in providing bioavailable Fe into the ocean waters, the effect of atmospheric organic species on Fe mobilization should be studied in

■

AUTHOR INFORMATION

Corresponding Author

*Phone: (319)335-1392. Fax: (319)353-1115. E-mail: [email protected]. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS This material is based on upon work supported by the National Science foundation under Grant No. CHE1012037. Any opinions, findings, and conclusions or recommendations expressed in this material are those of the authors and do not reflect the views of the National Science Foundation. The authors thank Professor Michelle Scherer for helpful discussions on the characterization of these different Fe-containing atmospheric dusts.

■

REFERENCES

(1) Martin, J. H. Glacial-interglacial CO2 change: The iron hypothesis. Paleoceanography 1990, 5 (1), 1−13. (2) Martin, J. H.; Coale, K. H.; Johnson, K. S.; Fitzwater, S. E.; Gordon, R. M.; Tanner, S. J.; Hunter, C. N.; Elrod, V. A.; Nowicki, J. L.; Coley, T. L.; Barber, R. T.; Lindley, S.; Watson, A. J.; Vanscoy, K.; Law, C. S.; Liddicoat, M. I.; Ling, R.; Stanton, T.; Stockel, J.; Collins, C.; Anderson, A.; Bidigare, R.; Ondrusek, M.; Latasa, M.; Millero, F. J.; Lee, K.; Yao, W.; Zhang, J. Z.; Friederich, G.; Sakamoto, C.; Chavez, F.; Buck, K.; Kolber, Z.; Greene, R.; Falkowski, P.; Chisholm, S. W.; Hoge, F.; Swift, R.; Yungel, J.; Turner, S.; Nightingale, P.; Hatton, A.; Liss, P.; Tindale, N. W. Testing the iron hypothesis in ecosystems of the equatorial pacific ocean. Nature 1994, 371 (6493), 123−129. (3) Moore, J. K.; Doney, S. C.; Glover, D. M.; Fung, I. Y. Iron cycling and nutrient-limitation patterns in surface waters of the World Ocean. Deep-Sea Res. Part II-Top. Stud. Oceanogr. 2002, 49 (1−3), 463−507. (4) Duce, R. A.; Tindale, N. W. Atmospheric transport of iron and its deposition in the ocean. Limnol. Oceanogr. 1991, 36 (8), 1715−1726. (5) Fung, I. Y.; Meyn, S. K.; Tegen, I.; Doney, S. C.; John, J. G.; Bishop, J. K. B. Iron supply and demand in the upper ocean. Glob. Biogeochem. Cycle 2000, 14 (1), 281−295. (6) Mahowald, N. M.; Baker, A. R.; Bergametti, G.; Brooks, N.; Duce, R. A.; Jickells, T. D.; Kubilay, N.; Prospero, J. M.; Tegen, I. Atmospheric global dust cycle and iron inputs to the ocean. Glob. Biogeochem. Cycle 2005, 19 (4), GB4025 DOI: 10.1029/ 2004GB002402. (7) Jickells, T. D.; An, Z. S.; Andersen, K. K.; Baker, A. R.; Bergametti, G.; Brooks, N.; Cao, J. J.; Boyd, P. W.; Duce, R. A.; Hunter, K. A.; Kawahata, H.; Kubilay, N.; laRoche, J.; Liss, P. S.; 10318



Article

Mahowald, N.; Prospero, J. M.; Ridgwell, A. J.; Tegen, I.; Torres, R. Global iron connections between desert dust, ocean biogeochemistry, and climate. Science 2005, 308 (5718), 67−71. (8) Mahowald, N. M.; Engelstaedter, S.; Luo, C.; Sealy, A.; Artaxo, P.; Benitez-Nelson, C.; Bonnet, S.; Chen, Y.; Chuang, P. Y.; Cohen, D. D.; Dulac, F.; Herut, B.; Johansen, A. M.; Kubilay, N.; Losno, R.; Maenhaut, W.; Paytan, A.; Prospero, J. A.; Shank, L. M.; Siefert, R. L. Atmospheric iron deposition: Global distribution, variability, and human perturbations. Annu. Rev. Mar. Sci. 2009, 1, 245−278. (9) Sholkovitz, E. R.; Sedwick, P. N.; Church, T. M.; Baker, A. R.; Powell, C. F. Fractional solubility of aerosol iron: Synthesis of a globalscale data set. Geochim. Cosmochim. Acta 2012, 89 (0), 173−189. (10) Veranth, J. M.; Smith, K. R.; Huggins, F.; Hu, A. A.; Lighty, J. S.; Aust, A. E. Mössbauer spectroscopy indicates that iron in an aluminosilicate glass phase is the source of the bioavailable iron from coal fly ash. Chem. Res. Toxicol. 2000, 13 (3), 161−164. (11) Baker, A. R.; Jickells, T. D. Mineral particle size as a control on aerosol iron solubility. Geophys. Res. Lett. 2006, 33 (17), L17608 DOI: 10.1029/2006GL026557. (12) Cwiertny, D. M.; Baltrusaitis, J.; Hunter, G. J.; Laskin, A.; Scherer, M. M.; Grassian, V. H. Characterization and acid-mobilization study of iron-containing mineral dust source materials. J. Geophys. Res. 2008, 113 (D5), D05202 DOI: 10.1029/2007JD009332. (13) Journet, E.; Desboeufs, K. V.; Caquineau, S.; Colin, J. L. Mineralogy as a critical factor of dust iron solubility. Geophys. Res. Lett. 2008, 35 (7), L07805 DOI: 10.1029/2007GL031589. (14) Cwiertny, D. M.; Hunter, G. J.; Pettibone, J. M.; Scherer, M. M.; Grassian, V. H. Surface chemistry and dissolution of α-FeOOH aanorods and microrods: Environmental implications of size-dependent interactions with oxalate. J. Phys. Chem. C 2009, 113 (6), 2175− 2186. (15) Shi, Z.; Krom, M. D.; Bonneville, S.; Baker, A. R.; Jickells, T. D.; Benning, L. G. Formation of iron nanoparticles and increase in iron reactivity in mineral dust during simulated cloud processing. Environ. Sci. Technol. 2009, 43 (17), 6592−6596. (16) Schroth, A. W.; Crusius, J.; Sholkovitz, E. R.; Bostick, B. C. Iron solubility driven by speciation in dust sources to the ocean. Nat. Geosci. 2009, 2 (5), 337−340. (17) Fu, H.; Cwiertny, D. M.; Carmichael, G. R.; Scherer, M. M.; Grassian, V. H. Photoreductive dissolution of Fe-containing mineral dust particles in acidic media. J. Geophys. Res. 2010, 115 (D11), D11304 DOI: 10.1029/2009JD012702. (18) Baker, A. R.; Croot, P. L. Atmospheric and marine controls on aerosol iron solubility in seawater. Mar. Chem. 2010, 120 (1−4), 4−13. (19) Rubasinghege, G.; Lentz, R. W.; Scherer, M. M.; Grassian, V. H. Simulated atmospheric processing of iron oxyhydroxide minerals at low pH: Roles of particle size and acid anion in iron dissolution. Proc. Natl. Acad. Sci. 2010, 107 (15), 6628−6633. (20) Shi, Z. B.; Krom, M. D.; Bonneville, S.; Baker, A. R.; Bristow, C.; Drake, N.; Mann, G.; Carslaw, K.; McQuaid, J. B.; Jickells, T.; Benning, L. G. Influence of chemical weathering and aging of iron oxides on the potential iron solubility of Saharan dust during simulated atmospheric processing. Glob. Biogeochem. Cycle 2011, 25, GB2010 DOI: 10.1029/2010GB003837. (21) Chen, H.; Laskin, A.; Baltrusaitis, J.; Gorski, C. A.; Scherer, M. M.; Grassian, V. H. Coal fly ash as a source of iron in atmospheric dust. Environ. Sci. Technol. 2012, 46 (4), 2112−2120. (22) Shi, Z.; Krom, M. D.; Jickells, T. D.; Bonneville, S.; Carslaw, K. S.; Mihalopoulos, N.; Baker, A. R.; Benning, L. G. Impacts on iron solubility in the mineral dust by processes in the source region and the atmosphere: A review. Aeolian Res. 2012, 5 (0), 21−42. (23) Fu, H.; Lin, J.; Shang, G.; Dong, W.; Grassian, V. H.; Carmichael, G. R.; Li, Y.; Chen, J. Solubility of iron from combustion source particles in acidic media linked to iron speciation. Environ. Sci. Technol. 2012, 46 (20), 11119−11127. (24) Zhuang, G. S.; Yi, Z.; Duce, R. A.; Brown, P. R. Link between iron and sulfur cycles suggested by detection of Fe(II) in remote marine aerosols. Nature 1992, 355 (6360), 537−539.

(25) Pehkonen, S. O.; Siefert, R.; Erel, Y.; Webb, S.; Hoffmann, M. R. Photoreduction of iron oxyhydroxides in the presence of important atmospheric organic-compounds. Environ. Sci. Technol. 1993, 27 (10), 2056−2062. (26) Zhu, X.; Prospero, J. M.; Savoie, D. L.; Millero, F. J.; Zika, R. G.; Saltzman, E. S. Photoreduction of iron(III) in marine mineral aerosol solutions. J. Geophys. Res. 1993, 98 (D5), 9039−9046. (27) Spokes, L. J.; Jickells, T. D.; Lim, B. Solubilization of aerosol trace-metals by cloud processing - A laboratory study. Geochim. Cosmochim. Acta 1994, 58 (15), 3281−3287. (28) Spokes, L. J.; Jickells, T. D. Factors controlling the solubility of aerosol trace metals in the atmosphere and on mixing into seawater. Aquat. Geochem. 1995, 1 (4), 355−374. (29) Johansen, A. M.; Siefert, R. L.; Hoffmann, M. R. Chemical composition of aerosols collected over the tropical North Atlantic Ocean. J. Geophys. Res. 2000, 105 (D12), 15277−15312. (30) Meskhidze, N.; Chameides, W. L.; Nenes, A.; Chen, G. Iron mobilization in mineral dust: Can anthropogenic SO2 emissions affect ocean productivity? Geophys. Res. Lett. 2003, 30 (21), 2085 DOI: 10.1029/2003GL018035. (31) Solmon, F.; Chuang, P. Y.; Meskhidze, N.; Chen, Y. Acidic processing of mineral dust iron by anthropogenic compounds over the north Pacific Ocean. J. Geophys. Res. 2009, 114, D02305 DOI: 10.1029/2008JD010417. (32) Shi, Z.; Bonneville, S.; Krom, M. D.; Carslaw, K. S.; Jickells, T. D.; Baker, A. R.; Benning, L. G. Iron dissolution kinetics of mineral dust at low pH during simulated atmospheric processing. Atmos. Chem. Phys. 2011, 11 (3), 995−1007. (33) Luo, C.; Gao, Y. Aeolian iron mobilisation by dust-acid interactions and their implications for soluble iron deposition to the ocean: a test involving potential anthropogenic organic acidic species. Environ. Chem. 2010, 7 (2), 153−161. (34) Kawamura, K.; Ng, L. L.; Kaplan, I. R. Determination of organic acids (C1-C10) in the atmosphere, motor exhausts, and engine oils. Environ. Sci. Technol. 1985, 19 (11), 1082−1086. (35) Falkovich, A.; Graber, E.; Schkolnik, G.; Rudich, Y.; Maenhaut, W.; Artaxo, P. Low molecular weight organic acids in aerosol particles from Rondonia, Brazil, during the biomass-burning, transition and wet periods. Atmos. Chem. Phys. 2005, 5 (3), 781−797. (36) Keene, W. C.; Galloway, J. N.; Holden, J. D. Measurement of weak organic acidity in precipitation from remote areas of the world. J. Geophys. Res. 1983, 88 (C9), 5122−5130. (37) Falkovich, A. H.; Schkolnik, G.; Ganor, E.; Rudich, Y. Adsorption of organic compounds pertinent to urban environments onto mineral dust particles. J. Geophys. Res. 2004, 109 (D2), D02208. (38) Falkovich, A. H.; Ganor, E.; Levin, Z.; Formenti, P.; Rudich, Y. Chemical and mineralogical analysis of individual mineral dust particles. J. Geophys. Res. 2001, 106 (D16), 18029−18036. (39) Russell, L. M.; Maria, S. F.; Myneni, S. C. B. Mapping organic coatings on atmospheric particles. Geophys. Res. Lett. 2002, 29 (16), 261−264. (40) Mochida, M.; Umemoto, N.; Kawamura, K.; Lim, H.-J.; Turpin, B. J. Bimodal size distributions of various organic acids and fatty acids in the marine atmosphere: Influence of anthropogenic aerosols, Asian dusts, and sea spray off the coast of East Asia. J. Geophys. Res. 2007, 112 (D15), D15209. (41) Sullivan, R. C.; Prather, K. A. Investigations of the diurnal cycle and mixing state of oxalic acid in individual particles in Asian aerosol outflow. Environ. Sci. Technol. 2007, 41 (23), 8062−8069. (42) Zinder, B.; Furrer, G.; Stumm, W. The coordination chemistry of weathering: II. Dissolution of Fe(III) oxides. Geochim. Cosmochim. Acta 1986, 50 (9), 1861−1869. (43) Suter, D.; Siffert, C.; Sulzberger, B. Catalytic dissoluion of iron(III)(hydr)oxides by oxalic acid in the presence of Fe(II). Naturwissenschaften 1988, 75 (11), 571−573. (44) Upadhyay, N.; Majestic, B. J.; Herckes, P. Solubility and speciation of atmospheric iron in buffer systems simulating cloud conditions. Atmos. Environ. 2011, 45 (10), 1858−1866. 10319



Article

(66) Stumm, W.; Lee, G. F. Oxygenation of ferrous iron. Ind. Eng. Chem. 1961, 53 (2), 143−146. (67) Fein, J. B.; Brady, P. V. Mineral surface controls on the diagenetic transport of oxalate and aluminum. Chem. Geol. 1995, 121 (1−4), 11−18. (68) Axe, K.; Persson, P. Time-dependent surface speciation of oxalate at the water-boehmite (γ-AlOOH) interface: Implications for dissolution. Geochim. Cosmochim. Acta 2001, 65 (24), 4481−4492. (69) Johnson, S. B.; Yoon, T. H.; Slowey, A. J.; Brown, G. E. Adsorption of organic matter at mineral/water interfaces: 3. Implications of surface dissolution for adsorption of oxalate. Langmuir 2004, 20 (26), 11480−11492. (70) Desboeufs, K. V.; Losno, R.; Vimeux, F.; Cholbi, S. The pHdependent dissolution of wind-transported Saharan dust. J. Geophys. Res. 1999, 104 (D17), 21287−21299. (71) Miller, W. P.; Zelazny, L. W.; Martens, D. C. Dissolution of synthetic crystalline and noncrystalline iron oxides by organic acids. Geoderma 1986, 37 (1), 1−13, DOI: 10.1029/2008JD011285. (72) Ito, K.; Bernstein, H. J. The vibrational spectra of the formate, acetate, and oxalate ions. Can. J. Chem. 1956, 34 (2), 170−178. (73) Berney, C. V.; Redington, R. L.; Lin, K. C. Infrared spectra of matrix-isolated acetic acid monomers. J. Chem. Phys. 1970, 53 (5), 1713−1721. (74) Hug, S. J.; Sulzberger, B. In situ Fourier transform infrared spectroscopic evidence for the formation of several different surface complexes of oxalate on TiO2 in the aqueous phase. Langmuir 1994, 10 (10), 3587−3597. (75) Hug, S. J.; Bahnemann, D. Infrared spectra of oxalate, malonate and succinate adsorbed on the aqueous surface of rutile, anatase and lepidocrocite measured with in situ ATR-FTIR. J. Electron Spectrosc. Relat. Phenom. 2006, 150 (2−3), 208−219. (76) Paris, R.; Desboeufs, K. V.; Journet, E. Variability of dust iron solubility in atmospheric waters: Investigation of the role of oxalate organic complexation. Atmos. Environ. 2011, 45 (36), 6510−6517. (77) Cornell, R. M.; Schindler, P. W. Photochemical dissolution of geothite in acid/oxalate solution. Clay Clay Min. 1987, 35 (5), 347− 352. (78) Kallay, N.; Matijević, E. Adsorption at solid/solution interfaces. 1. Interpretation of surface complexation of oxalic and citric acids with hematite. Langmuir 1985, 1 (2), 195−201. (79) Faust, B. C.; Zepp, R. G. Photochemistry of aqueous iron(III) polycarboxylate complexes: role in the chemistry of atmospheric and surface waters. Environ. Sci. Technol. 1993, 27 (12), 2517−2522. (80) Sulzberger, B.; Laubscher, H. Reactivity of various types of iron(III) (hydr)oxides towards light-induced dissolution. Mar. Chem. 1995, 50 (1−4), 103−115. (81) Sedlak, D. L.; Hoigné, J. The role of copper and oxalate in the redox cycling of iron in atmospheric waters. Atmos. Environ. 1993, 27 (14), 2173−2185. (82) Zhang, Y.; Kallay, N.; Matijević, E. Interaction of metal hydrous oxides with chelating agents. 7. Hematite-oxalic acid and -citric acid systems. Langmuir 1985, 1 (2), 201−206. (83) Grgić, I.; Dovžan, A.; Berčič, G.; Hudnik, V. The effect of atmospheric organic compounds on the Fe-catalyzed S(IV) autoxidation in aqueous solution. J. Atmos. Chem. 1998, 29 (3), 315−337. (84) Grgić, I.; Pozničc, M.; Bizjak, M. S(IV) autoxidation in atmospheric liquid water: The role of Fe(II) and the effect of oxalate. J. Atmos. Chem. 1999, 33 (1), 89−102. (85) Deguillaume, L.; Leriche, M.; Desboeufs, K.; Mailhot, G.; George, C.; Chaumerliac, N. Transition metals in atmospheric liquid phases: sources, reactivity, and sensitive parameters. Chem. Rev. 2005, 105 (9), 3388−3431. (86) Chen, Y.; Siefert, R. L., Determination of various types of labile atmospheric iron over remote oceans. J. Geophys. Res. 2003, 108 (D24). (87) Warneck, P. In-cloud chemistry opens pathway to the formation of oxalic acid in the marine atmosphere. Atmos. Environ. 2003, 37 (17), 2423−2427.

(45) Sulzberger, B.; Suter, D.; Siffert, C.; Banwart, S.; Stumm, W. Dissolution of Fe (III)(hydr) oxides in natural waters; laboratory assessment on the kinetics controlled by surface coordination. Mar. Chem. 1989, 28 (1), 127−144. (46) Hamer, M.; Graham, R.; Amrhein, C.; Bozhilov, K. Dissolution of ripidolite (Mg, Fe-chlorite) in organic and inorganic acid solutions. Soil Sci. Soc. Am. J. 2003, 67 (2), 654−661. (47) Xia, L.; Gao, Y. Chemical composition and size distributions of coastal aerosols observed on the US East Coast. Mar. Chem. 2010, 119 (1), 77−90. (48) Chen, Y.; Siefert, R. L. Seasonal and spatial distributions and dry deposition fluxes of atmospheric total and labile iron over the tropical and subtropical North Atlantic Ocean. J. Geophys. Res. 2004, 109 (D9), D09305. (49) Miller, W. L.; King, D. W.; Lin, J.; Kester, D. R. Photochemical redox cycling of iron in coastal seawater. Mar. Chem. 1995, 50 (1−4), 63−77. (50) Kieber, R. J.; Hardison, D. R.; Whitehead, R. F.; Willey, J. D. Photochemical production of Fe(II) in rainwater. Environ. Sci. Technol. 2003, 37 (20), 4610−4616. (51) Zuo, Y.; Hoigné, J. Formation of hydrogen peroxide and depletion of oxalic acid in atmospheric water by photolysis of iron(III)-oxalato complexes. Environ. Sci. Technol. 1992, 26 (5), 1014− 1022. (52) Siefert, R. L.; Pehkonen, S. O.; Erel, Y.; Hoffmann, M. R. Iron photochemistry of aqueous suspensions of ambient aerosol with added organic acids. Geochim. Cosmochim. Acta 1994, 58 (15), 3271−3279. (53) Zuo, Y.; Hoigné, J. Photochemical decomposition of oxalic, glyoxalic and pyruvic acid catalysed by iron in atmospheric waters. Atmos. Environ. 1994, 28 (7), 1231−1239. (54) Zuo, Y. G. Kinetics of photochemical/chemical cycling of iron coupled with organic-substances in cloud and fog droplets. Geochim. Cosmochim. Acta 1995, 59 (15), 3123−3130. (55) 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. (56) Wang, Z.; Chen, X.; Ji, H.; Ma, W.; Chen, C.; Zhao, J. Photochemical cycling of iron mediated by dicarboxylates: Special effect of malonate. Environ. Sci. Technol. 2010, 44 (1), 263−268. (57) Styler, S. A.; Donaldson, D. J. Heterogeneous photochemistry of oxalic acid on Mauritanian sand and Icelandic volcanic ash. Environ. Sci. Technol. 2012, 46 (16), 8756−8763. (58) Weller, C.; Horn, S.; Herrmann, H. Effects of Fe(III)concentration, speciation, excitation-wavelength and light intensity on the quantum yield of iron(III)-oxalato complex photolysis. J. Photochem. Photobiol. AChem. 2013, 255 (0), 41−49. (59) Stumm, W.; Morgan, J. J. Aquatic chemistry: chemical equilibria and rates in natural waters, 3rd ed.; John Wiley & Sons: New York, 1996. (60) Anderson, M.; Morel, F. Uptake of Fe (II) by a diatom in oxic culture medium. Mar. Biol. Lett. 1980, 1, 263−268. (61) Keene, W. C.; Sander, R.; Pszenny, A. A. P.; Vogt, R.; Crutzen, P. J.; Galloway, J. N. Aerosol pH in the marine boundary layer: A review and model evaluation. J. Aerosol. Sci. 1998, 29 (3), 339−356. (62) Kalinowski, B. E.; Schweda, P. Kinetics of muscovite, phlogopite, and biotite dissolution and alteration at pH 1-4, room temperature. Geochim. Cosmochim. Acta 1996, 60 (3), 367−385. (63) Werner, S. Reactivity at the mineral-water interface: Dissolution and inhibition. Colloid Surf. A-Physicochem. Eng. Asp. 1997, 120 (1−3), 143−166. (64) Remucal, C. K.; Sedlak, D. L. The role of iron coordination in the production of reactive oxidants from ferrous iron oxidation by oxygen and hydrogen peroxide. In Aquatic Redox Chemistry, Tratnyek, P. G., Grundl, T. J., Haderlein, S. B., Eds. American Chemical Society: Washington, DC, 2011; Vol 1071, p 177. (65) Teir, S.; Revitzer, H.; Eloneva, S.; Fogelholm, C.-J.; Zevenhoven, R. Dissolution of natural serpentinite in mineral and organic acids. Int. J. Miner. Process. 2007, 83 (1−2), 36−46. 10320



Article

(88) Crahan, K. K.; Hegg, D.; Covert, D. S.; Jonsson, H. An exploration of aqueous oxalic acid production in the coastal marine atmosphere. Atmos. Environ. 2004, 38 (23), 3757−3764. (89) Seinfeld, J. H.; Pandis, S. N. Atmospheric Chemistry and Physics: From Air Pollution to Climate Change. 2 ed.; John Wiley: New York, 2006. (90) Winkler, P. Observations on acidity in continental and in marine atmospheric aerosols and in precipitation. J. Geophys. Res.: Oceans 1980, 85 (C8), 4481−4486. (91) Ayers, G. P.; Larson, T. V. Numerical study of droplet size dependent chemistry in oceanic, wintertime stratus cloud at southern mid-latitudes. J. Atmos. Chem. 1990, 11 (1−2), 143−167.

10321


Iron Dissolution of Dust Source Materials during Simulated Acidic

Recommend Documents