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Ligand-Promoted Photoreductive Dissolution of Goethite by Atmospheric Low-Molecular Dicarboxylates Zhenzhen Wang,† Hongbo Fu,*,†,‡ Liwu Zhang,† Weihua Song,† and Jianmin Chen*,† †

Shanghai Key Laboratory Atmospheric Particle Pollution and Prevention, Department of Environmental Science and Engineering, Fudan University, Shanghai 200433, China ‡ Collaborative Innovation Center of Atmospheric Environment and Equipment Technology, Nanjing University of Information Science and Technology, Nanjing 210044, China S Supporting Information *

ABSTRACT: Recent evidence suggested that organic ligands in atmosphere water play an important role in the mobilization of iron from mineral aerosol. In this study, the dissolution of goethite (α-FeOOH) was investigated in the presence of three lowmolecular dicarboxylates enriched in the atmosphere, as well as a reference organic acid of methanesulfonate (MSA), which is especially abundant in marine atmospheric boundary layer. Iron mobilized from α-FeOOH was promoted under the irradiation and the deaerated condition, and the soluble Fe(II) concentration was enhanced greatly in the ligand-containing suspensions exposed to light. Irrespective of the reaction conditions, the capacities of the dicarboxylates on Fe mobilization were in the following order: oxalate > malonate > succinate, which were closely correlated with the carbon chain length of dicarboxylates: n = 2 > 3 > 4. The space barrier action of carbon atoms inhibited ligand-promoted Fe dissolution by affecting the structure and stability of the complexes. MSA also acted as an organic ligand to mobilize iron and showed weak capacity to reduce Fe(III) under the irradiation. The reactive oxygen species (ROS) analysis indicated that · OH, O2·−, and H2O2 could be involved in the Fe(II)−Fe(III) redox circle, and the ligand-promoted photoreductive dissolution process could be an important source of ROS in atmosphere water. Both transmission electron microscopy analysis and zeta potential data supported that the adsorption of oxalate molecules onto the surface would change the aggregation state of goethite nanoparticles, which increased the effective surface area, and therefore facilitated Fe mobilization from the oxide. The data shown herein deepens our understanding on the ligand-promoted dissolution mechanisms, which could be an important formation pathway of bioavailable Fe in the atmosphere.

1. INTRODUCTION Given the accumulated evidence of Fe cycle feedback on the global carbon, nitrogen, and sulfur cycle, iron bio-geochemistry is of great interest to oceanographers and environmental chemists.1,2 Fe solubilities in remote marine aerosols are significantly higher than that in soil source materials.3−6 During long transport from source regions to open oceans, dust undergoes atmospheric processes, such as gravitational settling, cloud processing, photoreduction, and acid processing, all of which could greatly increase Fe solubility via complex physical− chemical processes.6−11 Cornell and Schwertmann have reviewed the dissolution mechanisms of iron oxides, including protonation, complexation, reduction, and biological dissolution.12 Hereafter, the kinetics of iron dissolution from mineral aerosols under the conditions of simulating atmospheric aerosol processes have been studied.3,6−8,13,14 The reaction of natural organic ligands with mineral surfaces is important in environmental systems, which has been studied extensively over the last two decades.15−26 Field measurements have reported that dust aerosols tend to complex with organic species, mostly of short-chain oxygenated hydrocarbons and carboxylic acids during aging in the atmosphere.26 Further, most © 2017 American Chemical Society

of dissolved iron (>99%) in the seawater is bound by organic ligands.16 Organic complexation enables both enhancing the iron mobilization from aerosol particles, as well as stabilizing iron in dissolved form after wet deposition in seawater.25,26 Two mechanisms have been proposed to account potentially for iron mobilization by ligands in an aquatic environment.26 One mechanism is the nonreductive ligand-promoted dissolution, in which the organic compounds acting as ligands form relatively strong metal−organic surface complexes. The complexes destabilize the Fe−O bond and lower the energy barrier for dissolution, followed by the dissociation of complex at the surface in the form of dissolved Fe(III) ligand, for example, Fe(C2O4)n, in the presence of oxalate. The other is the reductive ligandpromoted dissolution, in which process organic compounds play a role on iron dissolution, as ligands, for example, complexing agent, and as electron donors, for example, reducing agent.27 The formation of surface iron organic complexes serves as the chromophores for photolysis, leading to a photoinduced ligandReceived: September 10, 2016 Revised: January 5, 2017 Published: February 1, 2017 1647

DOI: 10.1021/acs.jpca.6b09160 J. Phys. Chem. A 2017, 121, 1647−1656

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The Journal of Physical Chemistry A

nium ester 10-methyl-9-(p-formylphenyl)acridinium carboxylate trifluoromethane-sulfonate (AE), and catalase were from Sigma. All of the solutions were prepared using 18 MΩ Milli-Q water yielded from Mingche Q-Gard A2. Synthesis and Characterization of Goethite. Goethite was synthesized from ferrihydrite according to a modified method reported by Cwiertny and co-workers.8 X-ray diffraction (XRD) was performed to identify mineral phases in synthesized products using a Bruker D5000 diffractometer equipped with a Cu source. Infrared spectra (4000− 625 cm−1) of the synthesized goethite surface was collected on a PerkinElmer model 682 dispersive infrared spectrometer with a PerkinElmer model 3600 data station. The specific surface area of goethite was determined from seven-point N2 Brunauer−Emmett− Teller adsorption isotherms performed on a Quantachrome Nova 1200 surface area analyzer. The size and morphology of the synthesized goethite were observed with an FEI TECNAI G2 S-TWIN F20 transmission electron microscopy (TEM) operated in a bright field mode at 120 kV. For TEM analysis, suspensions (0.2 g L−1) of goethite were prepared in ethanol and sonicated for 1 h. The prepared suspension (50 μL) was then dropped to a carbon-coated 400 mesh Cu TEM grid (Electron Microscopy Science Inc.) and allowed to air-dry. Dissolution Experiments. All of dissolution experiments were performed in goethite (0.2 g L−1) suspensions at constant temperature (298 ± 1 K) by using a Pyrex glass vessel with a water jacket. A quartz window was placed on top of the vessel to prevent evaporation. The experimental setting is shown in Figure S1. Suspensions were prepared by goethite and 2 mM oxalate (pKa1 = 1.25, pKa2 = 4.27), malonate (pKa1 = 2.84, pKa2 = 5.79), succinate (pKa1 = 4.42, pKa2 = 5.42), and MSA (pKa = 1),40,41 and were stirred at 60 rpm using a magnetic stir bar. The pH of the suspension was rapidly adjusted to pH 3.0 by adding dilute HCl or NaOH. Experiment was also conducted in HCl solution at pH 3 as a control. Previous works have found that the minor variations in ionic strength did not significantly influence Fe dissolution;8,42 therefore, the ionic strength was not controlled in the present work. For dark experiments, the reactors were wrapped in aluminum foil to prevent exposure to light. For irradiation experiments, a 300 W xenon lamp (OSRAM) with AM 1.5 filter was used as a solar simulator. For deoxygenated experiments, suspensions were purged with N2 for 15 min prior to irradiation, and during irradiation the headspace of the reaction vessel was purged to maintain positive N2 pressure. All of dissolution reactions were conducted in triplicate. Over time, aliquots were periodically removed from the reactor using a syringe. Aliquots were passed through 0.2 μm poly(tetrafluoroethylene) (PTFE) filters (Millipore) and immediately acidified with 30 μL of 5 M HCl to preserve the sample for iron analysis. At each sampling event, enough sample volume was taken to allow for analysis of the total dissolved iron (FeT) and dissolved Fe(II).8,42 Determination of Fe Species. Ferrous iron was measured colorimetrically with 1,10-phenanthroline, which was reported in previous study.42 For Fe(II) analysis, 200 μL of a 5 mM 1,10phenanthroline solution and 200 μL of an ammonium acetate buffer were added to 1 mL of sample. To avoid possible interference from Fe(III), which can also form a complex with 1,10-phenanthroline when present at high concentrations, 50 μL of 0.43 M ammonium fluoride was added to the sample prior to 1,10-phenanthroline. The mixture was allowed to sit in the dark for 30 min prior to ultraviolet−visible (UV− vis) spectroscopy analysis, during which time a reddish-orange color developed if Fe(II) was present. FeT was determined via the same protocol, except that 30 μL of 1.44 M hydroxylamine hydrochloride, which reduces Fe(III) to Fe(II), was added to the sample rather than ammonium fluoride. Absorbance measured at 510 nm was converted to concentrations using aqueous standards prepared from anhydrous beads of ferrous chloride. Standards were prepared in each acid used in dissolution studies, and no matrix effects were observed. These conditions resulted in a detection limit of 1 μM. The concentration of dissolved Fe(III) was calculated from the difference in experimentally measured concentrations of dissolved FeT and dissolved Fe(II). Total Organic Carbon Analysis. The concentration changes of organic components in the suspensions were determined by a Jena multi N/C 2100 TOC. Suspensions were prepared using 0.05 g L−1 of

to-metal charge transfer (LMCT) reaction by reducing Fe(III) to Fe(II), along with oxidation of ligands.17,23,26,28−30 On the basis of strong evidence, the following pathways have been proposed: (i) the photolysis of Fe(III)-hydroxo groups and/or Fe(III)ligands complex,16,27,29 (ii) dark iron redox by previously photolyzed acidic natural organic matter,21 and (iii) Fe(II)catalyzed (photo)reduction dissolution.31−34 In general, the photoreductive ligand-promoted dissolution of iron oxides is an order of magnitude faster than nonreductive ligand-promoted dissolution, because release of Fe(II) from the surface of the crystal lattice is energetically more favorable than release of Fe(III).22,34,35 Beyond, the total rates driven by synergistic effect of both ligand and reductant on goethite dissolution were greater than the sum of the individual rates.35 It has been proposed that monocarboxylic acids (e.g., formic and acetic) do not complex substantial amounts of metal under typical atmospheric conditions. On the contrary, dicarboxylic oxalic acid and malonic acid are strong acids with high Henry’s law constants and large stability constants, leading to substantial iron complexation even at very low concentrations (approximately ppt).26 Three low-molecular species of oxalate, malonate, and succinate accounted for 70% of total dicarboxylates, all of which were ubiquitously acted as ligands associated with particles in the atmosphere.36,37 Several previous studies have employed the dicarboxylates to study the complexation structures formed on the surface of Fe(III) (hydro)oxides. Oxalate and malonate were found to be effective ligands and were expected to form bidentate chelating or bidentate bridging structure,38,39 whereas succinate would form monodentate complexation structure.38 Both oxalate and malonate were indicated to cause an enhancement of iron solubility associated with an increase of dissolved Fe(II) concentration under the irradiation.26,30,38,39 It is well-known mineral aerosols during clouds process will generate Fe rich nanoparticles,3,16 and iron oxides in atmospheric mineral aerosols are generally as goethite (α-FeOOH) and hematite (α-Fe2O3).16 Goethite nanoparticles have been used as a model to understand various effects in iron dissolution.8 In the present work, goethite nanoparticles was used as a surrogate iron phase to investigate the kinetics of ligand-promoted dissolution in the presence of atmospherically rich low-molecular organic ligands, including oxalate, malonate, and succinate. Methanesulfonate (MSA) is abundant in the marine boundary layer, as an important dimethylsulfide oxidation product. Currently, the capacity of MSA on ligand-promoted iron dissolution has not been studied. For comparison, MSA (CH3S(4+)O3H) was also embodied into this study. The aqueous environment at pH 3 was chosen to mimic the thin film of acidic water in the atmosphere, which was characteristic of low pH at water−mineral interfaces after cloud evaporation. The results from the laboratory experiments will provide the important insights on the processes of Fe dissolution mobilized from dust particles by typical organic ligands in the atmosphere.

2. EXPERIMENTAL SECTION Reagents. All of the chemicals were reagent grade or better. Ferric nitrate nonahydrate (Sigma), sodium bicarbonate (Sigma), and potassium hydroxide (Sigma) were used for goethite synthesis. Ligand solutions were prepared with oxalate, malonate, succinate, and MSA (Sigma). Solutions were adjusted to the appropriate pH with HCl (Sinopharm) and sodium hydroxide (Sinopharm). Measurements of dissolved iron were performed with 1,10-phenanthroline (Sigma), hydroxylamine hydrochloride (Sigma), ammonium fluoride (Sigma), ammonium acetate (Fisher), and glacial acetic acid (Fisher). 5,5Dimethyl-1-pyrroline-N-oxide (DMPO), chromium(III) oxide, acridi1648


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Figure 1. Comparison of Fe dissolution as a function of time in the presence of oxalate, malonate, succinate, and MSA under different experimental conditions (dark or irradiation, aerated or deaerated). Dissolution experiment in HCl solution at pH 3 was set as control. The data are shown for the dissolved FeT and the dissolved Fe(II) as a function of time. Uncertainties represent one standard deviation from triplicate experiments. goethite and 500 μM of oxalate, malonate, or succinate. The pH of the suspension was adjusted to pH 3.0. The experiments were conducted in the dark and under the irradiation as previously described. After reaction, suspension was passed through 0.2 μm PTFE filters and immediately acidified with 30 μL of 5 M HCl. Chromium(III) oxide (60 mg) was used as catalyst. Measurements of Hydrogen Peroxide (H2O2). A flow injection analysis instrument with acridinium ester chemiluminescence (AE-CL) detection was used for the determination of H2O2.43 The CL detector

was a Hamamatsu HC-135 photon counting photomultiplier tube (Hamamatsu Corp., Bridgewater, NJ) operated at the manufacturer’s recommended voltage (700 V) for optimal signal/noise ratio for nanomolar concentrations of H2O2. The AE reagent (10 μM) was prepared by adding AE to Milli-Q water and was buffered to pH 3 with 1 mM phosphate buffer to inhibit base hydrolysis of AE. The Milli-Q water used for preparation of reagent, and buffer solutions were treated with 3 mg L−1 catalase for at least 30 min to remove trace amounts of H2O2. 1649


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Table 1. Species Composition Matrix: Components and Their Corresponding Stoichiometric Coefficients and Equilibrium Constant for Each Species components species 2+

Fe(OH) Fe(OH)2+ Fe(OH)3 Fe2(OH)24+ Fe(OH)4− Fe(OH)+ Fe(ox)+ Fe(ox)2− Fe(ox)33− Fe(ox)0 Fe(ox)22− Fe(mal)+ Fe(mal)2− Fe(mal)3− Fe(mal)0 Fe(mal)22− Fe(suc)+ a

Fe3+

Fe2+

ox2−

mal2−

suc2−

1 1 1 1 1 1 1 1 1 1 1

1 2 3 1 2

1 1 1 1 1

1 2 3 1 2

1

1

H+

log10 β

reference

−1 −2 −3 −2 −4 −1

−2.2 −5.8 −3.2 −2.9 −21.6 −9.5 9.4 16.2 20.8 4.3 6.4 9.3 15.5 18.4 3.4 4.4 8.8

47 47 47 47 47 47 37 37 37 47 47 37 37 37 47 47 26

β is cumulative equilibrium constant for formation of species. bAll thermodynamic data are for 298 K, 1 atm, with ionic strength→0.

The concentrations of H2O2 were determined spectro-photometrically using the H2O2 molar absorptivity of 38.1 ± 1.4 M−1 cm−1 at 240 nm. Electron Spin Resonance Analysis. The electron spin resonance (ESR) technique using DMPO as the spin-trap reagent was applied to identify free radical intermediates involved in Fe mobilization under the irradiation. Suspensions were prepared using 100 μM ferric sulfate and 200 μM oxalate, malonate, or succinate. The pH of the suspension was adjusted to pH 3.0. The ESR spectra of radicals were recorded on a Bruker EPR ELEXSYS 500 spectrometer after 0, 30, 90, and 160 s under the irradiation (λ = 355 nm) at room temperature. The setting for the ESR spectrometer was as follows: center field, 3486.70 G; sweep width, 100 G; microwave frequency, 9.78 GHz; modulation frequency, 100 kHz; microwave power 12.7 mW. To minimize experimental errors, the same quartz capillary tube was used for all of the ESR measurements. TEM Analysis. To examine the morphological changes and aggregation state of goethite during dissolution, aliquot of the reaction mixture (0.2 g L−1) with 2 mM oxalate was collected for TEM analysis at 7 h after the onset of the irradiation experiment. To obviously observe the morphological changes, a higher concentration 100 mM of oxalate was also used to enforce the dissolution. High-resolution TEM (HRTEM) images were also collected to observe structural features of the samples at the nanoscale. Zeta Potentials of the Suspensions. The colloidal properties of the suspensions were characterized by measuring zeta potential. Zeta potential of the suspensions were obtained as a function of pH by Horiba SZ-100-z voltmeter measurement at 298 K. Suspensions were prepared using 100 μM of goethite and 200 μM of oxalate, malonate, or succinate and were sonicated for 20 min. The pH values of the suspensions were adjusted to a wide pH range from 2.0 to 12.0 by the stepwise addition of titrant (1 mol L−1 HCl or NaOH).

Ligand-Promoted Dissolution of Goethite. The release of soluble iron species as a function of time in the dark is shown in Figure 1, in which soluble iron species present as the dissolved FeT (Figure 1a) and the dissolved Fe(II) (Figure 1b). One can see that iron dissolved faster in the presence of organic ligands as compared to one in the case of HCl. The dissolved FeT concentrations in the presence of organic ligands were followed in the order of oxalate, malonate, succinate, and MSA. As is shown in Figure 1a, the dissolved FeT concentration was lowest (10.8 μM) in the HCl solution after 50 h. However, the FeT concentration displayed the highest value of 126.5 μM in the oxalate solution compared to the other suspensions, in which the FeT concentration was ∼11 times of the value in the HCl solution. For the malonate and succinate ligands, 83.2 μM and 32.5 μM of the dissolved FeT concentrations were measured at the same time, respectively. In the presence of MSA, the FeT was relatively low in concentration, only reaching 24.3 μM. The mobilization of Fe from the oxides could be composed of a proton-promoted process and a ligand-promoted dissolution process.38 Since the pH values of all the solutions were not significantly changed during the dark reaction (Figure S5), all of the acidic media at pH 3.0 would contain the similar H+ ion concentration to proton-promoted Fe dissolution. Although Cl− could complex with Fe(III) at pH 3, the formed complexation was so poor that the dissolution rate of the Fe−Cl complexes often was thought of as the proton-promoted rate.38 It was thus proposed that the ligand-promoted processes accounted for ∼91.5%, 87.0%, 66.8%, and 55.6% of dissolved FeT in the oxalate, malonate, succinate, and MSA solutions, respectively. The different capacities of organic ligands for Fe mobilization shown here were due to the formation of different Fe(III)ligand complexes on the surface of the nanoparticles. Dissolved Fe existed mainly as ferric iron in the various ligand solutions. The relative proportion of dissolved Fe(III) to dissolved FeT in oxalate, malonate, succinate, and MSA solutions were 87.4%, 82.7%, 79.2%, and 76.1% on average at every sampling point, suggesting that the Fe dissolution was mainly contributed to a nonreductive ligand-promoted dissolution. The leaching of some dissolved Fe(II) in the dark indicated that there

3. RESULTS AND DISCUSSION Characterization of Goethite Nanoparticles. On the basis of the powder XRD pattern (Figure S2) and the IR spectrogram (Figure S3), the as-prepared particle was identified to be goethite and did not contain any other crystalline phase, such as hematite or magnetite. The specific surface area of the goethite was 110 (±7) m2 g−1. The TEM images showed that goethite was nanosized acicular or rodlike crystal (Figure S4). On the basis of measurements of 200 particles, the nanosized particles were ∼81 nm in the length and 7 nm in the width on average. 1650


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the irradiation to investigate the effect of the simulated sunlight on the ligand-promoted iron dissolution. The generation of the soluble Fe species as a function of time is presented in Figure 1c,d. Goethite released considerably higher dissolved FeT in the organic ligand solutions compared to the HCl solution under the irradiation. Similarly with the reaction in the dark, the highest dissolved FeT concentration was measured in the presence of oxalate, followed by malonate, succinate, and MSA. After 36 h, the concentration of FeT under the irradiation reached 220.6, 176.8, 84.9, and 57.2 μM, in the oxalate, malonate, succinate, and MSA solution, respectively, which were enhanced by 1.9, 2.7, 3.3, and 2.7 times compared to the dark reaction. The dissolved Fe(II) concentration enhanced markedly under the irradiation, which increased by 6.6, 5.9, 3.0, and 2.0 times in the oxalate, malonate, succinate, and MSA solutions, respectively, as compared to that in the HCl. The average Fe(II)/FeT ratios determined at each sampling point were 88.8%, 87.3%, 83.0%, and 85.0% in the presence of oxalate, malonate, succinate, and MSA, respectively. The significant increase of the dissolved Fe(II) indicated that the iron mobilization was mostly contributed via the photoreductive pathway.8,13,14,17−19 It has been reported that many organic acids, such as salicylic and gluconic acid, could increase greatly photochemical production of Fe(II) in the solution, depending on the nature and concentration of the ligand.22,48 The photoreduction mechanism in the ferrioxalate system has been subjected to numerous investigations.14,19,21,22,49 As for the LMCT reaction involved in FeIII(C2O4)n, a radical complex of FeII(C2O4)n−1C2O4·− or a C2O4·− was first formed in the aqueous phase. An escape of C2O4·− to the solvent bulk followed instantly by its decarboxylation, then formed CO2 and/or CO2·−. Following this step, C2O4·− and/or CO2·− would further reduce the unphotolyzed Fe(III)-oxalato complexes to facilitate Fe leaching.20 The chain length of aliphatic ligands would affect electron transfer, because the carbon chain increased the electronegativity of CO2− cluster, inhibiting the electron donor capacity.26 Hence the decrease of the dissolved iron in concentration with the increase of carbon chains in length was observed: oxalate (C2) > malonate (C3) > succinate (C4). Because of the different sizes of three dicarboxylates, the difference in the photoreduction rates of iron in the presence of oxalate, malonate, and succinate could also be attributed to the steric hindrance of these molecules limiting the access of −COOH group to the surface site.26 Succinate could block some of the adjacent surface hydroxyl sites, therefore reducing the rate of Fe mobilization from the surface. In the case of MSA, the LMCT of [Fe(III) (CH3SO3)] complex also resulted in the reduction of Fe(III) to Fe(II). According to the previous work, the aqueous SO32− could react with Fe(III) to yield Fe(II), as well as SO42− and H2O2 via the formation of SO3·−, SO4·−, SO5·−, and HO2·.37,49 MSA enhanced the Fe(II) formation by 2.0 times that in the presence of HCl under the irradiation, suggesting that the marine biological production of MSA may be a potential participant to reduce the Fe(III)-containing aerosol in the marine atmospheric boundary layer. Influence of Dissolved Oxygen on Ligand-Promoted Dissolution. The photochemical dissolution experiments were also performed in the deaerated solution, and the results are shown in Figure 1e,f. Regardless of the ligand type, the dissolved Fe in concentration was substantially higher in the deoxygenated solution than those in the air-saturated solutions. After 7 h, FeT in the deaerated solutions reached 170.2, 135.2, 55.0, and 35.2 μM in the cases of oxalate, malonate, succinate, and MSA,

was a thermal reductive ligand-promoted dissolution occurring in these systems (Figure 1b). Previous study has recognized that humic substances were capable of reducing Fe(III) to Fe(II) at pH 3 and 5 in the dark.24 The data shown herein supported that oxalate, malonate, succinate, and MSA presented the reduction capability of Fe(III) to Fe(II), most likely via the redox active of stronger complexes, but generally thermal reduction pathway was too slow to produce large amounts of dissolved Fe(II). Ligand-promoted dissolution was mainly contributed to strong interactions between the ligand and the Fe atom on the surface, which induced polarization of the bonds between the Fe atom and the oxygen atom, thus reducing the activation energy of Fe leaching. The rate coefficients for goethite dissolution by ligands were closely correlated with ligand binding strength.35,38 Ligand binding strength depended upon the complexation structure formed between the functional groups of ligand and the surface hydroxyls groups of iron oxides.30,38 Dicarboxylates may form three complexation structures, including bidentate mononuclear, bidentate binuclear, and monodentate mononuclear on the crystal surface (Figure S6). Oxalate and malonate often formed inner-sphere bidentate mononuclear complexation structures, with one oxygen of each carboxylic group coordinated to the surface sites, as five- and six-membered ring chelate structures, respectively, which could accelerate the dissolution.38 Furthermore, the structure with the five-membered ring of oxalate yielded more stable complexes to facilitate Fe leaching from the surface as compared to the six-membered ring of malonate.38,44,45 It was worth noting that the formed innersphere bidentate binuclear complexation on the surface also facilitated Fe mobilization, but much more energy was needed to remove simultaneously two center Fe atoms from the crystalline lattice than to remove solely one center.46 Succinate always binds through one moiety to form monodentate complex with the surface iron atoms,38 having only one oxygen bond to the surface, donating less electron density to a surficial iron atom than the bidentate surface complexation structure, resulting in less contribution to Fe mobilization from the surface. The ligand-promoted dissolution rate was proportional to the binding strength of the ligand,38 in good agreement with the dissolution experiments. Apart from the ligand binding strength, the stability of the Fe(III)-ligand complexes also played an important role on Fe mobilization from the oxides.35,38 The equilibrium constants of the related complexation species are shown in Table 1.26,37,47 The equilibrium constants of oxalate with ferric iron were higher than those of the malonate and succinate ligands, and thus oxalate imposed the stronger influence on the surface complexation, therefore facilitating Fe solubility. One can see that the dissolved FeT concentrations in different ligand solutions decreased with decreasing the equilibrium constants of the complexes: oxalate > malonate > succinate, correlated with carbon number in the chain: n = 2 > 3 > 4, which was consistent with the conclusion of Duckworth and Martin.38 The space barrier action of carbon atoms affected the structure and stability of the complexes, thereby inhibiting Fe mobilization. Currently, the equilibrium constant of the Fe(III)MSA complex has still not been reported, but MSA might bind with iron to form poor monodentate complex [Fe(III) (CH3SO3)]2+. The FeT concentration in the MSA solution was 2.4 times that in the HCl solution, suggesting that the Fe(III)MSA complexes were more capable to mobilize iron as compared to ones of the Fe(III)-Cl complexes. Influence of Irradiation on Ligand-Promoted Dissolution. The dissolution experiments were then performed under 1651


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The Journal of Physical Chemistry A respectively, which were ∼1.6, 2.6, 2.0, and 2.4 times the ones for 7 h in the corresponding aerated solutions. The dissolved Fe(II) in the deaerated solutions reached 157.5, 123.1, 48.4, and 31.2 μM for oxalate, malonate, succinate, and MSA after 7 h. Furthermore, the dissolved Fe(II)/FeT ratios in the four deaerated suspensions were apparently higher than those in the aerated solutions, the average values of which reached 92.2%, 97.3%, 88.3%, and 89.3%, respectively. It was well-known that dissolved oxygen could rapidly reoxidize Fe(II) back to Fe(III) before Fe(II) can be detached from the surface, resulting in dissolution inhibition, although the previous studies have demonstrated that the Fe(II)−Fe(III) reoxidation reaction rate is very slow at pH below 4 in the presence of dissolved oxygen.23,29 Fe(II) oxidation could be accelerated by the organic ligands. The presence of the organic ligands could induce a series of free-radical chain reactions under the irradiation. Some oxidants, such as O2·− and ·OH, could form in the presence of O2.10 There are competitive reactions between O2 and Fe(III) with the carboxylate radicals. Taking oxalate as an example, the reactive fragments (C2O4·− and/or CO2·−) react with O2 to form reactive oxygen species (ROS), such as O2·−/ HO2·, H2O2, and ·OH in the air-saturated solution,18,20 which could inhibit Fe mobilization by reoxidizing the photogenerated Fe(II), as denoted by the formula of R1−R4, whereas in the deaerated solution, C2O4·− and/or CO2·− react with the unphotolyzed Fe(III)-oxalato complexes to promote reductive dissolution as expressed by R5. The removal of O2 could minimize the production of ROS in the solution, therefore facilitating the photoinduced formation of Fe(II), which was in good agreement with the dissolved data under the irradiation (Figure 1e,f). C2O4 ·− +O2 + H+ → O2 ·/HO2 ·+2CO2

(R1)

O2 ·/HO2 ·+Fe II → H 2O2 + Fe III

(R2)

Fe II + H 2O2 → Fe III + ·OH

(R3)

Fe II + ·OH → Fe III + H 2O

(R4)

Figure 2. TOC changes of the suspensions as a function of irradiation time in the presence of organic ligands.

surface complexes, which may occupy active sites to inhibit dissolution.52 Formation of ROS in the Ligand-Containing Suspensions. The H2O2 formations in the suspensions under the irradiation are shown in Figure 3. The H2O2 concentration increased as a function of the irradiation time, while almost no H2O2 could be measured in the dark. Under the irradiation, the formation and decomposition of H2O2 occurred simultaneously.18,19,53 On the one hand, it was well-known that H2O2 yielded through the dismutation of O2− and/or bimolecular reaction of ·OH via radical chain reaction.18 On the other hand, H2O2 could be consumed by oxidizing Fe(II) to Fe(III).18,19,53 The accumulation of H2O2 herein suggested the formation of H2O2 out-competed its decomposition. The H2O2 concentration was 4.27, 7.83, and 2.28 μM in the deaerated suspension of oxalate, malonate, and succinate after 7 h under the irradiation, respectively. Maximal production of H2O2 was 7.23, 10.45, and 4.87 μM in the aerated oxalate, malonate, and succinate solutions, respectively, which were 1.7, 1.3, and 2.1 times those in the deaerated solution. In the first 60 min, the accumulation rate of H2O2 under the simulated sunlight was 1.68 nM s−1 in the presence of goethite (2.2 mM) and oxalate (2 mM) at pH 3, which was slightly lower than that reported by Zuo and Hoigné.19 This difference was due to the concentration of Fe(III)(C2O4)n and/or the wavelengths of irradiation, because the rate of H2O2 formation depended greatly on light intensity, oxalate concentration, and Fe(III) concentration.19,20 As the composition of oxalate, the formation rate of H2O2 decreased by the reaction of H2O2 with surficial and dissolved Fe(II). An initial rapid rise and subsequently steady growth in the concentration of H2O2 were thus observed in the oxalate solution. The ESR signal intensities of intermediate reactive radicals as a time function are shown in Figure 4. No ESR signals were observed in the dark in the presence of organic ligands. Under the irradiation, the characteristic quartet peaks of the DMPO-·OH adduct with a 1:2:2:1 intensity were observed after 30 s, consistent with the similar spectra reported by others for the ·OH adduct.54 The peak intensities further increased after 160 s. This indicated that the free ·OH radical could be a main active oxygen species in this photochemical dissolution process. Note that different spectrum shapes could be observed in the oxalate solution, which was characterized by six peaks. It was speculated

C2O4 ·− +[Fe III(C2O4 )n(3 − 2n) + ](aq) + hv → Fe II + 2CO2 + nC2O−4

(R5)

As shown in Figure 1f, the dissolved Fe(II) in the goethite− oxalate system exhibited an initially sharp increase, followed by a steady increase with the prolonged time. The change of dissolution rate was found to be correlated with the concentration of oxalate. The total organic carbon (TOC) measurement (Figure 2) showed that the concentrations of malonate and succinate were not significantly changed, but oxalate decomposed obviously. It was due to the shortest carbon chain of oxalate, which could photodegrade and decarboxylate more easily than those of malonate and succinate. With a drastic decrease in the concentration of the organic ligand, the surface Fe-oxalato complexation was greatly suppressed,10 leading to slowing the dissolution rate. Moreover, it was observed that the rate of iron photoreduction decreased with an increase of pH in the goethite−oxalate system. Compared to malonate and succinate, the pH of oxalate solution showed a slight increase after 4 h under the irradiation (Figure S5). The increase of pH would weaken the proton-promoted Fe mobilization and would decrease the adsorption of oxalate on the particle surface.51 In addition, CO2 could react with surface Fe(III) to form carbonate 1652


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Figure 3. The H2O2 formation as a function of time in the deaerated and aerated suspensions. The H2O2 concentration in the aerated solution in the dark was determined as control.

Figure 4. ESR signal of DMPO-·OH (left, water solution) and DMPO- O2·− (right, ethanol solution) under the irradiation for (a) the Fe2(SO4)3/ oxalate system, (b) the Fe2(SO4)3/malonate system, and (c) the Fe2(SO4)3/succinate system. Reaction conditions: Fe2(SO4)3 loading, 100 μM; the dicarboxylates concentration, 200 μM; DMPO concentration, 40 mM.

that the spectrum was the signals of ·OH radical, coupling with carbon radical CO2·−, which originated from the photoinduced decarboxylation. Wang et al. have provided the evidence that both ·CH2COOH and ·OH radicals were formed in the presence of lower concentrations of organic ligand.55 It is well-documented that the superoxide radical anions O2·− tended to be unstable, which readily converted to H2O2 and O2 in a H2O solvent system.54 Therefore, the ESR spectra of the DMPO-·OOH/O2·− spin adducts were then recorded in ethanol media to probe O2·−. One can see that the six characteristic peaks of the ·OOH/O2·− adducts appeared under the irradiation, and the signal intensity increased slightly with the irradiation time. No such signals were observed in the dark; that is, generation of O2·− anions in the goethite−dicarboxylate suspensions inherently implicated irradiation. The ROS analysis gave the direct evidence that the active species (·OH, O2·−, and H2O2) were involved in the Fe(II)−Fe(III) circle during the Fe dissolution. Morphological Changes of Goethite after Dissolution. TEM images of nanosized goethite before and after the dissolution are shown in Figure 5. The original goethite nanorods (Figure 5a) showed clear lattice fringes, well-defined edges, and faceted shapes characteristic of the (021) and (110) surface planes as described in the previous research.33,56 The goethite nanorods were present as both dispersed particles and

aggregates in aqueous solution, and the former particles dissolved initially before the latter. Most of the goethite nanorods existed as dense aggregates, resulting in transformation of the nature of actual exposure concentrations (on mass and surface area) of goethite.56 After 7 h of dissolution in the presence of oxalate under the irradiation, some isolated goethite nanorods were observed under the TEM (Figure 5b), suggesting fewer aggregates as compared to the original material in the suspension. Reacted goethite became cigar-shaped, developing narrow, rounded ends. The particle size distribution showed an anisotropic change, and the average length of nanorods decreased from an initial value of 81 (±33; n = 100) nm to 70 (±31; n = 100) nm, whereas the width measured at the rod’s center was changed little. The observed morphological changes indicated that the dissolution preferentially happened on the both ends (021) faces of αFeOOH, which was in good agreement with the previous research.56,57 A high concentration of 100 mM oxalate was used to inspect the morphological change after dissolution. After 7 h under the irradiation (Figure 5c), goethite showed noticeable dissolution characteristics. Both the length and width of nanorods displayed significantly decreased. The average length of nanorods decreased to 55 (±41; n = 100) nm. HRTEM showed the disorder shapes on 110 facets after etching, 1653


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Figure 5. TEM images of the goethite nanorods before and after dissolution (a) original goethite before dissolution; (b) reacted goethite with 2 mM oxalate under the irradiation after 7 h; (c) reacted goethite with 100 mM oxalate under the irradiation after 7 h.

dicarboxylate would result in a polyanionic organic coating on mineral particles altering essentially the surface properties of minerals.59,60 In the presence of succinate and malonate, the pHiep value of goethite shifted to 5.1 and 2.0, respectively. It was noted that the amount of complexed ligand on the particle surface increased with the lowering of pH.51 In the presence of oxalate, the goethite surface’s zeta potentials shifted to the negative values throughout the pH spectrum, even reaching the stability value of −30 mV (Malvern Instruments technical note on Zeta Potential theory), suggesting oxalate has a “stabilizing” effect for the particle suspension.16 This “stabilizing effect” made the nanoparticles keep dispersed state in the suspension, which increased the effective surface area, therefore facilitating Fe mobilization from the oxides. As shown in Figure 6, oxalate could change the goethite surface’s charge distribution to a much larger degree. The zeta potential of goethite adsorbed by oxalate was much more negative than those in the case of malonate and succinate. At pH 3.0, the zeta potential of goethite decreased from 19.9 in the HCl solution to 14.5 in the succinate solution, to −2.5 in the malonate solution, to −28.5 in the oxalate solution. The results demonstrated that the oxalate molecule with small size was adsorbed preferentially onto the surface of goethite, which was more likely to form the Fe-oxalato complexes on goethite surface,39,44 which facilitated LMCT and thus iron mobilization from the surface.

suggesting that goethite displayed irregular surface defects rather than being perfect acicular nanorods. Zeta Potentials of Suspensions. The zeta potentials of goethite in the suspensions were measured as a function of pH, and the results are shown in Figure 6. Zeta potential as a property

Figure 6. Zeta potentials of goethite as a function of pH in the different suspensions.

of particle surface charge is responsible for particle distribution status in the suspensions.50,58 At the isoelectric point (pHiep), the crystal surfaces tend to be attractive to each other, and the particles are ready to flocculate and then settle. The pHiep value of goethite nanorods was 7.3 in water. The aggregated morphologies under the TEM was easily observed in the sample prepared using the Milli-Q water (pH = 7.1). The adsorption of

4. CONCLUSIONS The dissolution experiments of goethite in the presence of four atmospherically relevant acids were performed to investigate ligand-promoted Fe mobilization. Compared to the protonpromoted dissolution, ligand-promoted dissolution may play a more important role in mobilizing Fe from the oxides. The 1654


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(4) Zhuang, G. S.; Yi, Z.; Duce, R. A.; Brown, P. R. Link between iron and sulphur cycles suggested by detection of Fe(II) in remote marine aerosols. Nature 1992, 355, 537−539. (5) Bonnet, S.; Guieu, C. Dissolution of atmospheric iron in seawater. Geophys. Res. Lett. 2004, 31, L03303. (6) Hand, J. L.; Mahowald, N. M.; Chen, Y.; Siefert, R. L.; Luo, C.; Subramaniam, A.; Fung, I. Estimates of atmospheric-processed soluble iron from observations and a global mineral aerosol model: Biogeochemical implications. J. Geophys. Res. 2004, 109, 5305. (7) Baker, A. R.; Croot, P. L. Atmospheric and marine controls on aerosol iron solubility in seawater. Mar. Chem. 2010, 120, 4−13. (8) 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, D05202. (9) Maters, E. C.; Delmelle, P.; Bonneville, S. Atmospheric processing of volcanic glass: Effects on iron solubility and redox speciation. Environ. Sci. Technol. 2016, 50, 5033−5040. (10) 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, 2056− 2062. (11) Zhang, T. R.; Shi, J. H.; Gao, H. W.; Zhang, J.; Yao, X. H. Impact of source and atmospheric processing on Fe solubility in aerosols over the Yellow Sea, China. Atmos. Environ. 2013, 75, 249−256. (12) Cornell, R. M.; Schwertmann, U. The Iron Oxides: Structure, Properties, Reactions, Occurrences and Uses, 2nd ed.; Wiley-VCH: Weinheim, Germany, 2003. (13) Fu, H. B.; Lin, J.; Shang, G. F.; 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, 11119−11127. (14) Chen, H.; Grassian, V. H. Iron dissolution of dust source materials during simulated acidic processing: the effect of sulfuric, acetic, and oxalic acids. Environ. Sci. Technol. 2013, 47, 10312−10321. (15) Kadar, E.; Cunliffe, M.; Fisher, A.; Stolpe, B.; Lead, J.; Shi, Z. Chemical interaction of atmospheric mineral dust-derived nanoparticles with natural seawater−EPS and sunlight-mediated changes. Sci. Total Environ. 2014, 468, 265−271. (16) Rue, E. L.; Bruland, K. W. Complexation of iron(III) by natural organic-ligands in the Central North Pacific as determined by a new competitive ligand equilibration adsorptive cathodic stripping voltammetric method. Mar. Chem. 1995, 50, 117−138. (17) Sulzberger, B.; Laubscher, H. Reactivity of various types of iron(III) (hydr)oxides towards light-induced dissolution. Mar. Chem. 1995, 50, 103−115. (18) Borer, P.; Sulzberger, B.; Hug, S. J.; Kraemer, S. M.; Kretzschmar, R. Photoreductive dissolution of iron(III) (hydr)oxides in the absence and presence of organic ligands: Experimental studies and kinetic modeling. Environ. Sci. Technol. 2009, 43, 1864−1870. (19) 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, 1014−1022. (20) 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 2013, 255, 41−49. (21) Garg, S.; Ito, H.; Rose, A. L.; Waite, T. D. Mechanism and kinetics of dark iron redox transformations in previously photolyzed acidic natural organic matter solutions. Environ. Sci. Technol. 2013, 47, 1861− 1869. (22) Kuma, K.; Nakabayashi, S.; Matsunaga, K. Photoreduction of Fe(III) by hydroxycarboxylic acids in seawater. Water Res. 1995, 29, 1559−1569. (23) Siffert, C.; Sulzberger, B. Light-induced dissolution of hematite in the presence of oxalate. A case-study. Langmuir 1991, 7, 1627−1634. (24) Voelker, B. M.; Sulzberger, B. Effects of fulvic acid on Fe(II) oxidation by hydrogen peroxide. Environ. Sci. Technol. 1996, 30, 1106− 1114.

capacities of three dicarboxylates on Fe mobilization were in the order of oxalate > malonate > succinate, which were correlated with carbon chain length: n = 2 > 3 > 4. The space barrier action of carbon atoms affected the structure and stability of the complexes, thereby inhibiting Fe mobilization. The effect of ligand-promoted dissolution was closely related with the binding strength of the ligand. The stability of the Fe(III)-ligand complexes also played an important role on Fe mobilization from the particle. In the case of MSA, the soluble FeT and Fe(II) concentrations were higher than that in the HCl solutions, suggesting that MSA may be a potential contributor to mobilize Fe from aerosol in marine atmospheric boundary layer. ESR analysis and H2O2 measurements suggested that the ROS species, including ·OH, O2·−, and H2O2, were involved in the photochemical Fe(II)−Fe(III) circles when Fe was leached from the particle surface. TEM showed that the dissolution preferentially produced on both ends (021) faces of αFeOOH. One can see that nanosized goethite under the TEM was more well-dispersed in the oxalate solution as compared to ones in the presence of malonate and succinate, which could be attributed to surface charge change due to the formation of Feoxalato complexes, as confirmed by the zeta potential data. The results shown herein were helpful to deepen our understanding of Fe mobilization from dust aerosols, especially organic ligand promoted process in the atmosphere.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpca.6b09160. Illustrated experimental setting, characterization results, variation of pH during dissolution experiments, graphics of three common complexation structures (PDF)

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AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]. (H.F.) *E-mail: [email protected]. (J.C.) ORCID

Zhenzhen Wang: 0000-0003-3796-4826 Weihua Song: 0000-0001-7633-7919 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS Financial support was provided by National Natural Science Foundation of China (Nos. 21577022, 21190053, and 40975074), Ministry of Science and Technology of China (2016YFC0203700), and International Cooperation Project of Shanghai Municipal Government (15520711200).

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