Surface Chemistry and Dissolution of Î±-FeOOH Nanorods and

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Surface Chemistry and Dissolution of r-FeOOH Nanorods and Microrods: Environmental Implications of Size-Dependent Interactions with Oxalate† David M. Cwiertny,*,‡,§,# Gordon J. Hunter,‡ John M. Pettibone,⊥ Michelle M. Scherer,§ and Vicki H. Grassian*,‡,⊥ Department of Chemistry, Department of CiVil and EnVironmental Engineering, and Department of Chemical and Biochemical Engineering, UniVersity of Iowa, Iowa City, Iowa 52242 ReceiVed: August 16, 2008; ReVised Manuscript ReceiVed: October 07, 2008

Although recent evidence suggests that particle size plays an important role in the dissolution of iron from mineral dust aerosol, a fundamental understanding of how particle size influences the rate and extent of iron oxide dissolution processes remains unclear. In this study, surface spectroscopic methods are combined with solution phase measurements to explore ligand-promoted dissolution and photochemical reductive dissolution of goethite (R-FeOOH) of different particle sizes in the presence of oxalate at pH 3 and 298 K. Both X-ray photoelectron spectroscopy and attenuated total reflectance Fourier transform infrared spectroscopy (ATRFTIR) revealed differences between R-FeOOH particles in the nanometer-size range as compared to R-FeOOH particles in the micrometer-size range (nanorods and microrods, respectively). ATR-FTIR spectra showed a significant presence of surface hydroxyl groups as well as differences in surface complexes formed on nanorod surfaces. Furthermore, the saturation coverage of oxalate adsorbed on nanorods relative to microrods is ∼30% less as determined from solution phase batch adsorption isotherms. Despite less oxalate uptake per unit surface area, the surface-area-normalized rate of oxalate-promoted dissolution was ∼4 times greater in nanorod suspensions, suggesting this process is particle-size-dependent. Photochemical dissolution experiments revealed only a moderate increase in the rate of oxalate oxidation per gram of R-FeOOH with decreasing particle size. However, concentration profiles of photochemically generated Fe(II) and Fe(III) suggest differences in the dominant mechanisms controlling nanorod and microrod dissolution. Although loss of reactive surface area arising from oxalate-induced particle aggregation can contribute to size-dependent reactivity trends toward oxalate, our data, taken collectively, suggest unique surface chemistry of nanorods as compared to larger microrods. Results from ligand-promoted and photochemical dissolution experiments also highlight the important, and sometimes dominant, role that iron oxides on the nanoscale may play in iron mobilization relative to the larger oxide phases present in mineral dust aerosol. Introduction Iron-containing phases present in mineral dust aerosol play a key role in the global biogeochemical cycling of iron, providing a link between processes in air, water, and sea.1 Roughly 450 Tg of dust is deposited into the ocean each year.2 Iron-containing phases in these deposited dusts represent the dominant source of nutrient iron to remote ocean surface waters,3,4 where low iron concentrations have been shown to limit primary productivity.5,6 Iron comprises only ∼5% of these aerosol particles by mass, however, and only 3-5% of this iron content is capable of being solubilized via dissolution processes in marine environments.2 Evidence suggests that these soluble forms of iron are more bioavailable as a nutrient source than iron that remains trapped in solid phases.7-9 Because mineral dust aerosol has terrestrial origins, the iron phases it contains generally reflect iron’s crustal abundance.10,11 Accordingly, iron, the major transition metal in dust, is most often found in its +3 valence state (i.e., ferric iron), either as † Part of the special section “Physical Chemistry of Environmental Interfaces”. * Corresponding authors. (V.H.G.) Phone: 319-335-1392. Fax: 319-3351270. E-mail: [email protected]. (D.M.C.) Phone: 951-827-7959. Fax: 951-827-5696. E-mail: [email protected]. ‡ Department of Chemistry. § Department of Civil and Environmental Engineering. ⊥ Department of Chemical and Biochemical Engineering. # Current address: Department of Chemical and Environmental Engineering, University of California, Riverside, Riverside, CA 92521

Fe(III) oxides or as ionic Fe(III) trapped in aluminosilicate clay minerals.12 Of these Fe(III)-containing solid phases, (oxyhydr)oxides (or simply oxides) such as hematite, goethite, and ferrihydrite are often considered the forms most likely to be solubilized.13 Whereas dissolution of Fe-containing clays may be important in extremely acidic environments,14 a variety of physical and chemical processes can take place on oxide surfaces either during atmospheric transport or in marine surface waters to facilitate their conversion to soluble iron forms.15,16 For iron oxides present in mineral dust aerosol, nonreductive dissolution processes are promoted by low pH (“protonpromoted”)orthepresenceoforganicligands(“ligand-promoted”).12,15 Proton-promoted dissolution is important only at low pH regimes, where high proton concentrations can destabilize Fe-O bonds in the crystal lattice, thereby facilitating dissolution.17 At near-neutral pH values, oxide dissolution typically requires an organic ligand because the formation of metal-organic surface complexes also destabilizes the Fe-O bond and lowers the energy barrier for dissolution.18 Both types of dissolution ultimately increase Fe bioavailability through the generation of dissolved forms of Fe(III). Another important solubilization mechanism is reductive dissolution, which in marine and atmospheric waters (e.g. clouds, fog, and deliquescent layers on aerosol particles) is usually photochemical in nature. Reductive dissolution yields Fe(II) (i.e., ferrous iron), which exhibits greater solubility than Fe(III), but it is rapidly reoxidized to Fe(III) at the near-neutral

10.1021/jp807336t CCC: $40.75  2009 American Chemical Society Published on Web 11/14/2008

2176 J. Phys. Chem. C, Vol. 113, No. 6, 2009 pH values characteristic of most natural waters.18 In many natural environments, continuous Fe(II)-Fe(III) redox cycling is driven by the photochemical reduction of Fe(III) followed by Fe(II) reoxidation.16 Such redox cycling is believed to increase iron bioavailability in marine waters; dissolved Fe(II) is a labile form of iron to diatoms,7 and reoxidation of Fe(II) is expected to yield highly dispersed, insoluble Fe(III) colloids that will exhibit greater solubility and reactivity than more crystalline Fe(III) oxides.18,19 Several variables are known to influence the rate and extent of iron oxide dissolution. For nonreductive processes, important factors include the nature and degree of crystallinity of the iron oxide, solution pH, and the types of organic ligands present in solution.17 For reductive dissolution mechanisms, these same factors merit consideration, in addition to the nature of the reductant. Another important consideration that has yet to be fully explored is the influence that iron oxide particle size exerts on aerosol solubilization processes. Several recent laboratory investigations have demonstrated that Fe(III) oxide nanoparticles exhibit enhanced rates of reductive dissolution relative to largerscale iron oxides.20-23 In each case, size-dependent trends in dissolution rates could not be explained simply by measured differences between the specific surface areas of the oxide particle sizes considered. Rather, results have been most consistent with an increase in the inherent reactivity of iron oxides as particle sizes become increasingly smaller, behavior frequently associated with nanoparticles.24 In air, water, and soil systems, naturally occurring iron oxides are ubiquitous and exhibit a broad range of particle sizes that extends into the nanoscale.25 Nanoscale iron oxides therefore represent likely components of mineral dust aerosol.26 Accordingly, previous reports of size-dependent oxide reactivity have significant implications for the dissolution processes that control the bioavailability of iron-containing solid phases present in mineral dust aerosol. For example, it was recently reported that the amount of leachable iron present in aerosol increased with decreasing aerosol particle size,27 leading to the hypothesis that aerosol particle size was the controlling factor for the solubilization of iron from aerosols. Although Baker and Jickells27 attributed this relationship to the greater surface area available on small aerosol particles, it is also certainly possible that their result is a reflection of size-dependent reactivity of the iron oxide phases present within the aerosol they investigated. In the current study, complementary spectroscopic methods and solution-phase measurements are used to explore the influence of particle size on the dissolution of goethite (RFeOOH) by oxalic acid (pKa1 ) 0.97, pKa2 ) 3.57),28 hereafter simply referred to as oxalate, in aqueous solutions at pH 3.0. The R-FeOOH-oxalate system is of great relevance to the global cycling of iron. Oxalate is considered the most abundant water-soluble organic compound present in rural and urban aerosol,29 and goethite is one of the dominant forms of iron present in soils30 and has frequently been detected in aerosol particles of various origins.31-34 The reaction between goethite and oxalate can proceed via both nonreductive (ligand-promoted) and reductive (photochemical) dissolution mechanisms. These reactions can therefore take place either in atmospheric waters during aerosol transport or in the photic zone of the marine water column. There also are an extensive number of prior investigations focusing on various aspects of the interactions between oxalate and goethite.35-43 These previous studies serve as a fundamental basis for the results presented herein, which focus exclusively on elucidating the role that particle size plays in oxalate-goethite interactions.

Cwiertny et al. Solution phase measurements were used to quantify the extent of oxalate adsorption and the rate of oxalate-promoted dissolution in suspensions with two sizes of R-FeOOH particles, referred to as microrods and nanorods.20 Similar experiments were performed to quantify the influence of particle size on the rate of photochemical reductive dissolution. Results from these solution-phase macroscopic reactivity studies were then combined with the results from various characterization methods to provide molecular-level insights into the size-dependent interactions of goethite and oxalate. These different methods of characterization include attenuated total reflectance Fourier transform infrared spectroscopy (ATR-FTIR), transmission electron microscopy (TEM), and light-scattering techniques (e.g., dynamic light-scattering and sedimentation rates from transmitted UV/visible light) to better understand the fundamental physicochemical properties of R-FeOOH nanorods and microrods and their interactions with an environmentally prevalent dicarboxylic acid. Experimental Methods Reagents. All chemicals were reagent grade or better and were used as received. Ferric nitrate nonahydrate (Fe(NO3)3 · 9H2O; Sigma Aldrich; g98%), sodium bicarbonate (NaHCO3, Sigma Aldrich, g99.5%), and potassium hydroxide (KOH, Sigma Aldrich; ACS reagent) were used for goethite synthesis. Oxalate solutions were prepared from sodium oxalate (Sigma Aldrich). Sodium perchlorate (Sigma Alrich) was used to poise ionic strength. When necessary, solutions were adjusted to the appropriate pH with hydrochloric acid (HCl; Fisher). Measurements of dissolved Fe(II) and total dissolved iron were performed with 1,10-phenanthroline (Sigma-Aldrich, g 99%), hydroxylamine hydrochloride (Sigma Aldrich, 98%), and a buffer from ammonium acetate (Fisher, 98.5%) and glacial acetic acid (EMD, 99.7%). The mobile phase used during HPLC analysis was prepared from 1 N sulfuric acid (H2SO4; Fisher). Goethite Synthesis. Nanoparticles were synthesized according to the method of Anschutz and Penn20 as previously described.44 Briefly, this method involved the dropwise addition of NaHCO3 to a solution of Fe(NO3) · 9H2O, resulting in the formation of ferrihydrite nanoparticles. Ferrihydrite nanoparticles were subsequently aged at pH 12 for 24 h at 90 °C to yield the nanoscale goethite product referred to as nanorods. Larger goethite particles, hereafter referred to as microrods, were synthesized using the method of Schwertmann and Cornell.45 Goethite Characterization. Nanorods and microrods were characterized using powder X-ray diffraction (XRD) performed on a Bruker D-5000 diffractometer with a Cu KR source. Nanorod and microrod dimensions were obtained from single particle analysis with TEM. Suspensions (∼0.2 g/L) of nanorods and microrods were prepared in deionized water, and a single drop was applied to a 40 mesh Cu TEM grid. Images of goethite particles were collected on a JEOL JEM-1230 transmission electron microscope operated at a 100 keV accelerating voltage. Digital images were acquired using a Gatan UltraScan CCD camera with Gatan imaging software. The size of goethite particles was then determined by analyzing TEM images in the software package Image J. Particle size distributions were determined from the analysis of roughly 500 nanorods and 300 microrods. Surface areas were determined from a seven-point N2-BET isotherm using a Quantachrome Nova 1200 surface area analyzer. Samples for BET were outgassed overnight (∼12 h) at a temperature of 60 °C. To avoid the possibility of goethite transformation to hematite, higher outgassing temperatures were not used.

R-FeOOH Nanorods and Microrods The near-surface region of nanorods and microrods was probed using X-ray photoelectron spectroscopy (XPS). Analyses were performed using a custom-designed Kratos Axis Ultra X-ray photoelectron spectroscopy system. The ultrahigh vacuum surface analysis chamber has been described in detail elsewhere.46 Briefly, the chamber is equipped with monochromatic radiation at 1486.6 eV from an aluminum KR source using a 500 mm Rowland circle silicon single crystal monochromator. The X-ray gun was operated using a 15 mA emission current at an accelerating voltage of 15 kV. Low-energy electrons were used for charge compensation to neutralize the sample. Survey scans were collected using the following instrument parameters: energy scan range of 1200 to -5 eV; pass energy of 160 eV; step size of 1 eV; dwell time of 200 ms, and an X-ray spot size of 700 × 300 µm. High-resolution spectra were acquired in regions of interest (e.g., Fe 2p) using the following experimental parameters: 20-40 eV energy window; pass energy of 20 eV; step size of 0.1 eV, and dwell time of 1000 ms. Three sweeps were used to acquire the spectra for all regions. The absolute energy scale was calibrated to the Cu 2p2/3 peak binding energy of 932.9 eV using an etched copper plate. Additional information regarding the electronic properties of microrods and nanorods was determined from UV/visible spectra collected for aqueous suspensions of each particle size. Stock suspensions of nanorods and microrods (0.5 g/L) were prepared in pH 3 solution. They were then diluted to the following goethite loadings: 0.005, 0.01, 0.05, and 0.1 g/L. Absorbance spectra were collected at each solids concentration on a PerkinElmer Lambda 20 UV/visible spectrophotometer over 200-800 nm at a scan rate of 480 nm/min using a quartz cuvette with 1 cm path length. Sedimentation Studies of Goethite Suspension Stability. The relative stability of nanorod and microrod suspensions at pH 3.0 was investigated using UV/visible spectrophotometry to measure rates of particle sedimentation. Goethite suspensions were added to a 1 cm path length cuvette, and the change in transmitted light (λ ) 510 nm) was monitored as a function of time. The suspensions contained a goethite concentration of 0.2 g/L and 5 mM NaClO4. Additional studies exploring the influence of oxalate on suspension stability also included either 0.2 or 2 mM oxalate. All suspensions were prepared 24 h prior to conducting sedimentation measurements. Adsorption of Oxalate on Goethite. Isotherms for oxalate adsorption on nanorods and microrods were conducted at pH 3 in suspensions containing 5 mM NaClO4, a goethite concentration of either 1 or 2 g/L, and an initial oxalate concentration ranging from 0.1 to 2 mM. Reactors were prepared by diluting appropriate volumes of a 5 g/L goethite stock suspension in 5 mM NaClO4 and a 50 mM oxalate stock solution in 5 mM NaClO4. The total volume was adjusted to 6 mL with 5 mM NaClO4. Each stock solution was adjusted to pH 3 prior to reactor assembly. Once assembled, reactors were sealed and wrapped in Al foil to prevent light from entering the reactors and were rotated end-over-end at 60 RPM in the dark for 12 h. After this time interval, approximately 1 mL of sample was removed from the reactor via a disposable plastic syringe and passed through a 0.2 µm syringe-driven PTFE filter, and the filtrate was subsequently analyzed with HPLC to determine the concentration of oxalate as described below. The concentration of adsorbed oxalate was then determined from the difference in initial and final concentrations measured in the experimental systems, for which initial concentrations were determined from control reactors prepared without goethite. The remainder of the suspension was used for pH measurements with a pH

J. Phys. Chem. C, Vol. 113, No. 6, 2009 2177 microelectrode (Orion) and an Accumet pH meter. Generally, little pH drift was observed over the course of these experiments; although the final pH values of the nanorod suspensions were all within 0.1 pH unit of the initial pH, the pH of microrod reactors increased from 3.0 to 3.4 over the oxalate concentration range of 0.1 to 2 mM. To obtain molecular-level details pertaining to the surface complexes of oxalate adsorbed on nanorods and microrods, IR measurements were collected in a Thermo Nicolet FTIR equipped with a mercury cadmium telluride A (MCTA) detector using a horizontal ZnSe or AMTIR ATR cell (Pike Technologies, Inc.). For each spectrum, 150 scans were collected. The crystal was coated with a hydrosol of goethite rods (approximately 1.5 mL of a 2 g/L goethite suspension) and allowed to air-dry for 24 h. This process left behind a thin layer of goethite that uniformly coats the crystal. The entire horizontal ATR cell was placed inside the internal compartment of the FTIR spectrometer that contains optics for the ATR cell. Optima water (Fisher) that had been adjusted to pH 3 with dilute HCl was introduced to the cell to obtain a reference spectrum, and the cell was sealed quickly to alleviate problems of evaporation. An oxalate stock solution at pH 3 was added and allowed to equilibrate with the sample for at least 1 h. Increasing concentrations were examined with the addition of the stock solution ranging from 0.01 to 1 mM. After a solution was added to the cell, spectral scans were collected in 15 min intervals for up to 1.5 h until no change in the spectra was observed, suggesting equilibrium had been achieved. Solution-phase experiments over a range of dissolved oxalate concentrations (0.1-20 mM) were also conducted on the bare crystal to examine the contribution of oxalate in the aqueous phase, thereby allowing differentiation between adsorbed and dissolved oxalate species in IR spectra. Proton- and Oxalate-Promoted Goethite Dissolution. Experiments examining the proton- and ligand-promoted dissolution of nanorods and microrods were conducted in 20 mL serum bottles, crimp-sealed and mixed end-over-end at 60 rpm. Reactors for proton-promoted dissolution contained 0.5 g/L of R-FeOOH and 5 mM NaClO4 and were at pH 3.0. For ligandpromoted dissolution experiments, reactors also contained 1 mM of oxalate. All reactors were constructed in a manner similar to our isotherm experiments. First, the appropriate volumes of R-FeOOH and oxalate stock solutions necessary to achieve the stated initial concentrations were combined in a 25 mL beaker, and the beaker’s contents were subsequently mixed with a stir bar. The suspension pH was then rapidly adjusted to pH 3.0 using dilute HCl, and the contents of the beaker were transferred to a 20 mL serum bottle, crimp-sealed, and placed on a mixer. Over the next 24 h, 1 mL samples were removed periodically from the reactors in a manner identical to that described for our adsorption isotherm experiments, and the filtrate was then analyzed for dissolved Fe(II) and total dissolved iron. Finally, a select number of long-term dissolution experiments were also conducted. Once constructed, these reactors were allowed to react for 200 h, at which point a sample of the suspension was removed for TEM analysis to identify possible morphological changes resulting from the dissolution reaction. Photochemical Reductive Dissolution of Goethite by Oxalate. A solar simulator with a 150 W xenon lamp (Oriel Corp.) was used as the irradiation source for all photolysis experiments. The photochemical reaction vessel had a capacity of 70 mL and was stirred using a magnetic stir bar. The vessel had a quartz window with an area of 12.5 cm2 mounted on top that permitted deoxygenated experiments and was removable to allow for

2178 J. Phys. Chem. C, Vol. 113, No. 6, 2009 reactions to proceed under atmospheric conditions. In experiments performed under deoxygenated conditions, suspensions were purged with N2 for ∼10 min prior to irradiation, and during irradiation the headspace of the reaction vessel was purged to maintain positive N2 pressure. The ionic medium used for photolysis experiments was 5 mM NaClO4. The solid concentration used was 0.5 g L-1. The temperature was kept constant (298 ( 1 K) through use of a water jacket integrated to the reaction vessel. The pH was monitored with a glass electrode standardized with pH buffer solutions. Experiments were performed at pH ) 3.0 ( 0.1 and were adjusted with HCl as necessary. Over time, samples were periodically removed from the reactor using a disposable syringe that was connected to 10 cm of Teflon tubing. Aliquots (1 mL) were collected after passing through a 0.2 µm PTFE filter (Expertek) and immediately acidified with 40 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 dissolved Fe(II), total dissolved iron, and oxalate concentration. All photochemical reactions were conducted in triplicate. TEM Analysis of Reacted Particles. To examine how proton-promoted, ligand-promoted, and photoassisted dissolution reactions influenced the morphology of goethite particles, samples were taken from each of these experimental systems at the conclusion of experiments for TEM analysis. Approximately 200 µL of sample was removed from each reactor and diluted in 1 mL of deionized water. A drop of this suspension was then added to a Cu TEM grid as described previously. For photochemical experiments conducted under deoxygenated conditions, samples were removed using an airtight syringe that was capped and then transferred to an anaerobic chamber (95% N2, 5% H2) for TEM sample preparation. Dimensions of reacted particles were determined from TEM analysis of individual particles as described previously by our group.44 Analytical Methods. Ferrous iron was measured colorimetrically with 1,10-phenanthroline, which forms a complex with Fe(II) that absorbs at 510 nm.47,48 For Fe(II) analysis, 200 µL of a 5 mM 1,10-phenanthroline solution and 200 µL of an ammonium acetate buffer were added to 1 mL of sample. The mixture sat in the absence of light for 30 min prior to UV/ visible analysis, during which time a reddish-orange color developed if Fe(II) was present. Total dissolved iron was determined via the same protocol, except that 20 µL of 1.5 M hydroquinone, which reduces Fe(III) to Fe(II), was added to the sample. Absorbances measured at 510 nm were converted to concentrations using aqueous standards prepared from anhydrous beads of ferrous chloride (Sigma Aldrich). Oxalate concentrations were determined using an Agilent 1200 Series HPLC that was equipped with a UV/visible diode array detector (HPLC-DAD) and a Prevail organic acid column (Grace Davison Discovery Sciences). A flow rate of 1 mL/min and a mobile phase of 0.01 N H2SO4 were used. Oxalate was analyzed at 254 nm, and instrument response was converted to aqueous phase concentrations using standards of oxalate that were prepared daily in an appropriate aqueous solution. Results and Discussion Characterization of Nanorods and Microrods of r-FeOOH. Details regarding the purity, size distribution and specific surface area of the R-FeOOH nanorods and microrods have been described elsewhere.44 Briefly, powder X-ray diffraction patterns were consistent with that expected for goethite, with patterns

Cwiertny et al. for nanorods exhibiting line broadening as is typically observed with decreasing particle size. As in previous work with similarly synthesized materials,20,49,50 goethite rods were elongated along the crystallographic c-axis [001] and modeled as rhomboidal prisms bounded by (110) with ends capped by (021) surfaces. This model yields rods with a diamond-like cross section,30 and does not account for minor faces such as the (100) and (010), which only develop to a limited extent, if at all, during synthesis.51-53 Nanorods were 81 ((27) nm by 7 ((2) nm, and microrods were 670 ((370) nm by 25 ((9) nm (uncertainties represent one standard deviation). Note that widths correspond to the 110-type faces; absolute widths of nanorods and microrods measured with TEM were corrected to account for the particles’ preferred orientation on the grid according to the procedure outlined by Anschutz and Penn.20 The BET specific surface areas of nanorods and microrods were 110 ((7) and 40 ((3) m2/g, respectively. These values are somewhat lower than estimates of geometric surface area for nanorods (∼200 m2/g) and microrods (∼50 m2/g) determined using the aforementioned particle dimensions and structural model of goethite. As shown in Figure 1, XPS spectra were similar for both particle sizes, with the survey spectra and spectra of the Fe 2p and O 1s regions essentially being independent of particle size. However, an additional feature was observed at ∼3 eV in the valence band region of the nanorods, indicating possible differences in the band gap of goethite as a function of particle size. UV/visible absorption spectra were used to explore further possible differences in the band gap of nanorods and microrods. Absorption spectra in Figure 2 were collected for 0.05 g/L suspensions of goethite at pH 3 and have been normalized to the peak near 380 nm. Although collected over a range of R-FeOOH loadings (0.005-0.1 g/L), normalized spectra were independent of goethite concentration. In nanorod suspensions, local absorption maxima around ∼300 and 380 nm were slightly shifted by ∼4 nm toward lower wavelengths relative to microrods. Although small, this shift was consistent and reproducible, suggesting a possible shift in the band gap energy of nanorods, behavior that is frequently observed for semiconducting nanomaterials.24 Characterization of nanorods and microrods with ATR-FTIR spectroscopy showed more pronounced differences as a function of particle size (Figure 3). Most clearly observed are distinct absorption bands at 3489 and 3661 cm-1 in the spectra for nanorods, which are much stronger than the corresponding bands in the spectra of microrods. These bands arise from hydroxyl groups on the surface of the goethite nanorods.54 On the basis of adsorption studies with phosphate, Russell et al.55 originally assigned the band near 3660 cm-1 to singly coordinated surface hydroxyl groups and the band near 3489 cm-1 to surface O-H groups of higher coordination (doubly and triply coordinated). A more recent investigation by Boily and Felmy, 54 however, assigned both bands to singly coordinated O-H groups arising from the protonation of -O groups on the (110) and (021) faces of the goethite surface (i.e., -Oz- + H+ T -OH1-z), which are suspected to play an important, if not dominant, role in adsorption reactions. Normalization to the bulk hydroxyl region revealed a roughly 3-4 fold increase in the intensity of these surface O-H bands on nanorods relative to microrods (Table 1), suggesting that the density of singly coordinated surface hydroxyl groups is greater on nanorods than on microrods. More subtle differences were observed for the other absorption bands characteristic of goethite. From previous investigations by Cambier,56,57 the band near 670 cm-1 is due to the

R-FeOOH Nanorods and Microrods

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Figure 1. Characterization of goethite nanorods and microrods by XPS. Shown are (a) survey spectra collected over a broad range of binding energies, as well as higher resolution spectra collected in the (b) Fe 2p, (c) O 1s, and (d) the valence band (VB) regions.

Figure 2. UV/visible absorbance spectra of 0.05 g/L nanorod and microrod suspensions at pH 3. Each spectrum was normalized to its respective maximum near 380 nm (i.e., 379 nm for nanorods and 383 nm for microrods). These peaks are shown in the inset.

symmetric Fe-O stretch, and the broad band near 3120 cm-1 is attributable to the stretching mode of the bulk hydroxyl groups in the goethite structure (νOH). For nanorods, the bands at 897 and 796 cm-1 can be assigned to the bending modes of the bulk hydroxyl groups in (δOH) and out (γOH) of the (001) plane. These absorption bands are slightly shifted to lower wavenumbers for microrods, at 892 and 794 cm-1, respectively. Traditionally, the location of the O-H bending bands can be used as a diagnostic for the relative crystallinity of goethite particles.30 In a study with a suite of synthetic goethites, Cambier56,57 found that the O-H bending bands broaden, values of δOH and γOH decrease, and the peak-to-peak separation between the bending bands narrows with decreasing particle size and crystallinity. It should be emphasized that we observed opposite trends with respect to particle size; the O-H bending modes for nanorods are located at higher frequencies and display

Figure 3. (a) Comparison of ATR-FTIR spectra collected for dry powders of nanorods and microrods. (b) Enlarged region of spectra between 3400 and 3800 cm-1, in which features indicative of surface hydroxyl groups are shown to be more prominent for nanorods.

greater peak-to-peak separation relative to microrods. Accordingly, these diagnostics seem to suggest that nanorods exhibit a higher degree of crystallinity than the microrods. This suggestion runs counter to conventional wisdoms associated with nanoparticle crystallinity, in which structural disorder increases with decreasing particle size.25,58 For example, Waychunas et al.25 used EXAFS measurements to infer sub-

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Cwiertny et al.

TABLE 1: Comparison of the Absorption Band Intensities Associated with Surface Hydroxyl Groups on Microrods and Nanorods of r-FeOOH frequency band (cm-1) 3660 3490 a

normalized intensitya microrods nanorods 0.105 0.111

0.407 0.402

intensity ratio, nanorods/microrods 3.9 3.6

Integrated intensities normalized to the internal hydroxyl region

Figure 4. ATR-FTIR spectra comparing oxalate adsorption on nanorods and microrods (middle and bottom spectra, respectively). Experiments were conducted at pH 3.0 with an equilibrium dissolved oxalate concentration of 0.25 mM, although similar results were obtained for all dissolved oxalate concentrations investigated (0.1-1 mM). Also shown is the spectrum for an aqueous solution of 0.25 mM oxalate at pH 3.0, collected in the absence of goethite (top spectrum). The top spectrum shows that at the low concentration of 0.25 mM, absorption bands are very weak for dissolved oxalate. These ATRFTIR data show that oxalate adsorbs readily on goethite and indicate the oxalate complexes on goethite are influenced by particle size (see text for further details).

stantial distortion of the Fe(O,OH)6 octahedra in 5 nm goethite nanoparticles, presumably arising from the high degree of surface curvature in their oblate particles. However, similar structural disorder was not observed for larger, rodlike goethite nanoparticles (i.e., with lengths of 24 and 75 nm), which are most comparable to the nanorods used in the present work. Thus, it is unclear whether the low degree of crystallinity typically associated with particles of only a few nanometers in diameter should also be anticipated for nanoparticles of larger dimensions. In light of these ATR-FTIR results, further work comparing the surface and bulk structural properties of nanocrystals to larger crystallites may be warranted because such factors will ultimately impact interfacial reactivity. Adsorption of Oxalate on r-FeOOH Nanorods and Microrods. Figure 4 shows ATR-FTIR spectra collected after the exposure of nanorods and microrods to a pH 3 solution of 0.25 mM oxalate. Also shown is the spectrum for a pH 3 solution of 0.25 mM oxalate in the absence of goethite, illustrating that dissolved oxalate species do not contribute to the measured absorbances in goethite-containing experiments. Rather, in the presence of nanorods and microrods, the absorption bands

resulted entirely from oxalate species adsorbed on the goethite surface. We note that peaks attributable to aqueous phase oxalate species were detected only at much higher concentrations (>10 mM) than that shown in Figure 4. ATR-FTIR spectra revealed a distinct difference between the adsorbed oxalate species on nanorod and microrod surfaces. Although several of the absorption bands were observed for both particle sizes, the absorption band at 1607 cm-1 in the nanorod system was absent in the spectrum for microrods. This result was obtained at each of the dissolved oxalate concentrations used in ATR-FTIR experiments to probe adsorption on goethite (0.1-1 mM). To better understand the difference between nanorod and microrod systems, results from the current study were compared to previous ATR-FTIR investigations of oxalate adsorption on goethite. The spectra in Figure 4 agree reasonably well with IR results obtained at pH 2.7 by Persson and Axe,37 who examined the pH-dependent uptake of oxalate on a high surface area goethite (84 m2/g) most comparable to the nanorods considered herein (110 m2/g). Through EXAFS and complementary IR analysis of solution-phase Fe(III)-oxalate species, they proposed surface complexes responsible for each of the IR bands arising from oxalate adsorption on goethite. Absorption bands at 1307 and 1582 cm-1 were assigned to outer-sphere or so-called hydrogen-bonded surface complexes, whereas bands near 1710, 1690, 1400, and 1255 cm-1 were credited to inner-sphere complexes. According to the assignments of Persson and Axe, the absorption band at 1607 cm-1 likely corresponds to an outersphere surface complex, implying that such complexes form more readily or to a greater extent on nanorods relative to microrods. We note that this peak is shifted toward higher wavenumbers in our nanorod system (1607 cm-1) relative to the value reported by Persson and Axe (1587 cm-1).37 This may be the result of differences in the type of goethite used in each study or due to slight variations in solution pH, which in turn will impact the degree of oxalate protonation. Nevertheless, ATR-FTIR results suggest that the oxalate complexes formed on the surface of goethite are, indeed, influenced by particle size. To assess the impact of these molecular-level differences on goethite reactivity, complementary quantitative adsorption experiments were conducted in aqueous batch systems at pH 3. On the basis of goethite mass (Figure 5a), isotherms for oxalate adsorption revealed a greater degree of uptake on nanorods relative to microrods, behavior that is frequently observed in experimental investigations of size-dependent adsorption on nanoparticles.25,59,60 Normalization of adsorbed oxalate concentrations by goethite specific surface area from BET, however, indicates that there is less adsorbed oxalate per unit area on nanorod surfaces (Figure 5b). Similarly, use of geometric surface area estimates results in adsorption maxima of ∼1.3 and ∼2.6 µmol/m2 for nanorods and microrods, respectively. This sizedependent trend for oxalate adsorption is the same as what we previously observed for Fe(II) sorption on nanorods and microrods; smaller particles have a lower adsorption capacity on the basis of surface area.44 In our prior investigation,44 we concluded that aggregation likely influenced the relative reactivity of nanorods and microrods toward dissolved Fe(II). Scanning electron microscopy images revealed that goethite aggregates as large as 10 µm were produced in the pH 7.5 suspensions used for Fe(II) sorption experiments. Ultimately, such extensive aggregation, particularly in nanorod suspensions, raised questions as to the amount of



Figure 7. Dissolved Fe(III) production via ligand-promoted dissolution in suspensions containing 0.5 g/L goethite and 1 mM oxalate. Experiments were conducted in the absence of light in pH 3 solutions containing 5 mM NaClO4. Reactors were mixed end-over-end at 60 RPM.

Figure 5. Isotherms for the uptake of oxalate on nanorods and microrods in pH 3 suspensions of 5 mM NaClO4. Data are presented for adsorbed oxalate normalized by (a) mass and (b) surface area in 1 g/L suspensions of goethite.

Figure 6. Sedimentation plots comparing the behavior of 0.2 g/L suspensions of nanorods (NR) and microrods (MR) in pH 3.0 solutions of 5 mM NaClO4. Data are shown for nanorod and microrod suspensions in the presence and absence of oxalate. Two different oxalate concentrations (0.2 and 2 mM) were examined.

reactive surface area available for Fe(II) uptake and whether surface areas determined from characterization of dry powders were an appropriate basis for comparing nanorod and microrod reactivity. In the current work, aggregation was not expected to play an important role in oxalate adsorption because experiments were conducted at pH 3, a value far from the reported pHzpc for goethite.30 Not surprisingly, therefore, sedimentation experiments (Figure 6) demonstrated that nanorod and microrod suspensions were relatively stable under such conditions, with

minimal settling in a 0.2 g/L suspension observed over ∼ 2.5 h. However, extensive aggregation was observed upon the addition of oxalate to goethite suspensions. The adsorption of negatively charged oxalate on the positively charged surface in turn minimized electrostatic repulsion between particles. This destabilized nanorod and microrod suspensions, as evidenced by the considerable increase in the sedimentation rates observed in the presence of oxalate. Interestingly, the rate of sedimentation was essentially equivalent for nanorods and microrods in the presence of oxalate. These rates of sedimentation were less than those previously determined for goethite aggregates at pH 7.5,44 suggesting that aggregates formed via oxalate adsorption likely are not as large as those observed previously. As in our previous work,44 aggregation introduces uncertainty as to the amount of surface area in goethite suspensions available for oxalate adsorption. Oxalate-induced aggregation will cause the amount of reactive surface area to change in response to adsorbate-adsorbent interactions. Thus, the ambiguity regarding the available surface area in each suspension makes it difficult to determine whether the size-dependent trend in oxalate adsorption (Figure 5) reflects molecular-level differences in nanorod and microrod surface reactivity or differences in their aggregation state upon oxalate addition. This highlights the need to consider suspension stability both before and after addition of the adsorbate when exploring size-dependent trends in adsorption. Oxalate-Promoted Dissolution of r-FeOOH Nanorods and Microrods. At pH 3, proton-promoted goethite dissolution was not observed over 30 h (data not shown). In the presence of oxalate, however, the rate of ligand-promoted dissolution was appreciable over this time scale and was considerably greater for nanorods compared to microrods (Figure 7). From linear regression analysis, the initial rate of dissolved Fe(III) production in nanorod suspensions was 52 ((10) µmol g-1 h-1. For microrods, this rate was 4.6 ((1.4) µmol g-1 h-1, indicating roughly an 11 ((4)-fold difference in the initial rate of oxalatepromoted dissolution on the basis of goethite mass. Surfacearea-normalized rates of goethite dissolution were 0.47 ((0.09) and 0.12 ((0.04) µmol m-2 h-1 for nanorods and microrods, respectively. On the basis of particle surface area, therefore, nanorods exhibit a 4.1 ((1.3)-fold increase in their rate of dissolution relative to microrods. Although there are previous reports of increasing surface-area-normalized rate constants with decreasing particle size for proton-promoted dissolution61,62 and


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Figure 8. (a) TEM images of acicular goethite rods prior to reaction with oxalate. Inset illustrates the dominant crystal faces of goethite nanorods, whereas the arrow highlights their well-defined ends terminated by (021) faces. (b) TEM images of goethite nanorods collected 200 h after the addition of 1 mM oxalate. The arrows highlight the narrow, rounded ends resulting from dissolution. Similar images were collected for nanorods after photochemical reductive dissolution mediated by oxalate. TEM analysis was used to construct distributions for the (c) length and (d) width of nanorods before and after reaction with oxalate. For unreacted nanorods (t ) 0 h), distributions were constructed from sizing of n ) 530 nanorods. Distributions for reacted nanorods (t ) 200 h) were constructed after analysis of n ) 100 particles. Note that the widths in part d represent the width of the 110-type faces (see text for further details).

reductive dissolution20-23 of iron oxides, to the best of our knowledge, this is one of the first demonstrating such sizedependent behavior for ligand-promoted dissolution of iron oxides. We note that the difference in the inherent surface reactivity of nanorods and microrods is likely greater than that estimated simply from consideration of specific surface area. Specifically, the production rate of dissolved Fe(III), which is equal to the rate of goethite dissolution, is described by a second-order rate law, in which the rate is proportional to the surface concentration of adsorbed oxalate ([Ox]ads) in the form of a bidentate mononuclear complex and the concentration of Fe(III) surface sites ([Fe(III)]surf) (eq 1).63

dissolution rate ) d[Fe(III)] ⁄ dt ) k[Ox]ads[Fe(III)]surf (1) Recall that adsorption isotherms (see Figure 5) indicate that nanorods adsorbed ∼30% less oxalate per unit BET surface area than microrods at pH 3. Thus, the rate of goethite dissolution per unit of adsorbed oxalate is roughly 6-fold greater for nanorods relative to microrods. It appears, therefore, that the oxalate complexes responsible for goethite dissolution are more reactive and promote dissolution to a greater extent when formed on nanorod surfaces. The different rates of nanorod and microrod ligand-promoted dissolution are reminiscent of early work focusing on the sizedependent, proton-promoted dissolution of goethite. Cornell et al.61,62 observed that surface-area-normalized rate constants for the proton-promoted dissolution of goethite increased with decreasing particle size. Moreover, TEM analysis revealed that

the ends of their goethite rods became rounded over the course of the reaction, suggesting that the rate of proton-promoted dissolution was greatest on these crystal faces. Using a rectangular model for the morphology of goethite and particle dimensions from TEM, they calculated that the relative amount of surface area available on the ends of acicular goethite increased with decreasing particle size. Accordingly, they proposed that rates of dissolution were greatest on small goethite particles because they possessed the largest fraction of the reactive faces primarily responsible for dissolution. We, too, observed changes in nanorod morphology upon exposure to oxalate-containing solutions at pH 3 (Figure 8). TEM images of nanorods prior to oxalate addition (Figure 8a) show sharp, well-defined ends characteristic of their (021) faces. In contrast, images collected 24 (not shown) and 200 h (Figure 8b) after the addition of 1 mM oxalate revealed changes in particle morphology; over time, nanorods became cigar-shaped, developing narrow, rounded ends. After 200 h, the average length of nanorods also decreased from an initial value of 81 ((27; n ) 530) nm to 66 ((33; n ) 100) nm. Although the standard deviations associated with these averages are relatively large, the particle size distribution clearly illustrates that oxalatepromoted dissolution yields a greater percentage of nanorods with smaller lengths relative to unreacted particles (Figure 8c). In contrast, little change was observed in nanorod width, measured at the rod’s center, over long-term dissolution experiments (Figure 8d). Similar morphological changes were

R-FeOOH Nanorods and Microrods far more difficult to discern on microrods, presumably due to the slower rate of oxalate-promoted dissolution observed in these systems. The observed morphological changes indicate that rates of oxalate-promoted dissolution are greatest on the (021) faces of goethite, as well as at their intersection with 110-type surfaces. Thus, the anisotropic dissolution of goethite previously observed for proton-promoted dissolution61 also takes place during the ligand-promoted reaction. Cornell et al.61 proposed that this anisotropy results from differences in the coordination and concentration of the surface hydroxyl groups present on each crystal face. Consistent with this hypothesis and our results, Baro´n and Torrent64 calculated that the highest density of singly and doubly coordinated surface hydroxyl groups occurs on the (021) faces of goethite (8.2 groups/nm2), whereas significantly lower densities (3 groups/nm2) were calculated for the (110) faces. It is possible, therefore, that the greater rate of dissolution on (021) surfaces results from their higher density of surface hydroxyl groups assumed to be the reactive site. This characteristic would likely impact other interfacial processes, as well, and several other reports have noted unique reactivity of the goethite (021) faces. For example, Weidler et al. 65 and, more recently, Chun et al.50 identified these surfaces as the most reactive with respect to particle growth. There are several possible explanations for the enhanced reactivity toward oxalate of nanorods relative to microrods (see Figure 7). On the basis of the similarities between our results and those of Cornell et al.,61 the size-dependent rates of dissolution may reflect differences in the relative amount of crystal faces present on nanorods and microrods. In fact, recent atomic force microscopy measurements have shown that the relative area of goethite crystal faces can change substantially as a result of decreasing particle size.66 Additional factors may also be at play. For instance, our ATR-FTIR characterization suggests a greater density of surface hydroxyl groups on nanorods relative to microrods. It is also possible that the surface hydroxyl groups on the (021) faces of nanorods are inherently more reactive than their counterparts on microrod surfaces. For example, molecular simulations suggest that the degree of protonation of the (021) and (110) faces is influenced by goethite particle size,67 behavior that would be expected to produce differences in the interfacial reactivity of nanorods and microrods. Photochemical Reductive Dissolution of Goethite by Oxalate. Results of photochemical dissolution experiments are presented in Figure 9, which shows concentration profiles for oxalate decay (Figure 9a), and the production of dissolved Fe(III) (Figure 9b) and Fe(II) (Figure 9c) in irradiated microrod and nanorod suspensions. Data are shown for experiments conducted in the presence and absence of O2. As anticipated, rates of photochemical reductive dissolution were considerably greater than ligand-promoted dissolution, consistent with results of previous investigations.40,41,68 Initial (t < 1 h) rates of oxalate decay were essentially equivalent in all systems, with the exception that oxalate decay in the aerated nanorod suspension was initially more rapid than the other three systems (Figure 9a). In contrast, a much clearer reactivity trend emerged at longer time scales (t > 1 h); on the basis of mass, nanorods were more reactive than microrods, and the rate of oxalate decay for each goethite particle size was greatest in the presence of oxygen. The role of O2 in promoting oxalate transformation is consistent with the work of Sulzberger and co-workers,40,41 who extensively investigated iron oxide photochemical dissolution in the presence of oxalate. Sulzberger and Laubscher41 first observed that the rate of oxalate decay


Figure 9. Results of photochemical reduction of goethite by oxalate. Data are shown as a function of time for (a) normalized oxalate concentration, (b) dissolved Fe(III), and (c) dissolved Fe(II) both in the presence and absence of O2. Uncertainties represent one standard deviation of triplicate experiments. Experiments were conducted in 0.5 g/L suspensions of nanorods or microrods in pH 3 solutions of 5 mM NaClO4.

was greatest in aerated suspensions. They hypothesized that O2 could reoxidize Fe(II) produced from the reaction with oxalate to yield a Fe(III) precipitate more prone to photochemical reduction than goethite. For both nanorod and microrod systems, dissolved Fe(III) profiles (Figure 9b) were independent of the presence of oxygen. In all cases, an initial burst of Fe(III) production was followed by a gradual decline over longer time scales. This decrease in Fe(III) is not due to homogeneous precipitation because maximum values of dissolved Fe(III) were well below the solubility limit for ferric hydroxide at pH 3 (Fe(OH)3(s) T Fe3+ + 3OH-, log K ) 10-37.11).69 Rather, it appears that dissolved Fe(III) is consumed via some other mechanism, likely the photochemical reduction of a dissolved Fe(III)-oxalate complex, which is photoactive.40,41 The initial rate of dissolved Fe(III) production was far greater than anticipated from the dark, oxalate-promoted dissolution experiments (see Figure 7). This behavior was most pronounced in microrod systems, for which initial rates of dissolved Fe(III)


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production determined from linear regression analysis were ∼500 times greater in the presence of light than in dark dissolution experiments. In contrast, only a 6-fold increase was observed in photochemical experiments with nanorods. Accordingly, whereas oxalate-promoted dissolution yielded higher concentrations of dissolved Fe(III) in nanorod suspensions, dissolved Fe(III) production was more rapid for microrods in the photochemical systems. Although it is clear that the formation of dissolved Fe(III) shown in Figure 9b is predominantly photochemical in nature, a molecular-level explanation for the enhanced rate of its production is less obvious. We note that enhanced photochemical production of dissolved Fe(III) was also reported by Siffert and Sulzberger,40 but a mechanistic explanation for its occurrence was not provided. Dissolved Fe(II) was observed in all systems, regardless of the presence or absence of oxygen (Figure 9c). This result is in marked contrast to that of Sulzberger and Laubscher,41 who observed Fe(II) production only in deoxygenated goethite suspensions. They interpreted their results as evidence that the rate-determining step in dissolved Fe(II) production was its detachment from the oxide surface and that in the presence of O2, reoxidation of surface-associated Fe(II) by O2 was fast relative to the detachment rate. Sulzberger and Laubscher41 did report dissolved Fe(II) production in oxygenated lepidocrocite suspensions, a result they attributed to its lower degree of crystallinity relative to goethite. Notably, their goethite had a BET specific surface area of 14 m2/g, suggesting their particles were larger than both the microrods and nanorods used herein. It is possible, therefore, that our observation of dissolved Fe(II) in the presence of O2 reflects differences in the crystallinity of goethite used in each study. Although O2 had no influence on the amount of Fe(II) generated in microrod suspensions, dissolved Fe(II) concentrations were depressed in aerated nanorod suspensions (Figure 9c). The lower Fe(II) concentrations measured in aerated nanorod suspensions is consistent with a greater rate of Fe(II) reoxidation on the surface of nanorods relative to microrods. According to the aforementioned hypothesis of Sulzberger and co-workers,40,41 faster rates of Fe(II) reoxidation are most likely to occur on particles with a greater degree of crystallinity. In the context of this hypothesis, our dissolved Fe(II) data could be interpreted as indirect evidence that nanorods are more crystalline than microrods, a scenario consistent with diagnostics observed from ATR-FTIR characterization. Trends in the production of Fe(II) lend insights into the dominant mechanism of photochemical dissolution of nanorods and microrods by oxalate. From their work, Sulzberger and co-workers40,41 proposed two pathways for the reductive dissolution of goethite in the presence of light and oxalate (see Figure 5 of Sulzberger and Laubscher41). The first is a direct photochemical reduction mechanism in which the oxalate-goethite surface complex serves as the chromophore (eqs 2-4):

>FeIIIOH + HC2O4- f > FeIIIC2O4- + H2O

(2)

>FeIIIC2O4- + hν f > FeIIC2O4-•

(3)

>FeIIC2O4-• f > FeII + CO2 + CO2-

(4)

After formation of the Fe(III)-oxalate surface complex (eq 2), light absorption followed by ligand-to-metal charge transfer results in the reduction of structural Fe(III) to Fe(II) (eq 3) and the formation of oxalate oxidation products (eq 4). The surfaceassociated Fe(II) can then either detach from the oxide surface to generate dissolved Fe(II) or be reoxidized by O2 (in aerated systems) to form a surface Fe(III) precipitate.

The second pathway for dissolved Fe(II) production proposed by Sulzberger and co-workers40,41 involves Fe(II)-promoted dissolution of the goethite surface (eqs 5 and 6):

>FeIIIC2O4- + Fe2+(aq) f > FeIIIC2O4FeII+

(5)

>FeIIIC2O4FeII+ f > FeIIFeIIIC2O4+(aq)

(6)

Dissolved Fe2+ generated from the direct photochemical reduction pathway can form a complex on the goethite surface using oxalate as a bridging ligand (eq 5). This Fe(II) complex can then reduce Fe(III) at the goethite surface to yield a surfaceassociated Fe(II) species and a dissolved Fe(III)-oxalate complex (eq 6), which will be readily photolyzed to yield an additional dissolved Fe(II) species. Ultimately, this second pathway causes Fe(II) production and oxalate oxidation to be autocatalytic in nature; for every mole of Fe(II) generated from the heterogeneous photochemical mechanism, another mole of Fe(II) can be generated as a result of the homogeneous reduction of the Fe(III)-oxalate complex. From the Fe(II) concentration profiles shown in Figure 9c, we can speculate as to the relative importance of each of these mechanisms as a function of goethite size. Initially, rates of Fe(II) production are approximately equal for each particle size, when direct photochemical reduction is expected to be the dominant mechanism for Fe(II) production. This suggests that in systems with equivalent goethite mass loading, nanorods and microrods exhibit comparable photochemical reactivity despite nanorods adsorbing more oxalate per gram than microrods (see Figure 5). A potential explanation for this behavior relates to the different complexes formed on nanorod and microrod surfaces (see Figure 4). Because inner sphere complexes between oxalate and surface Fe(III) are proposed as the dominant chromophore in these systems,40,41 the observation that outer sphere complexes form more readily on nanorod surfaces could explain their lower photochemical reactivity. Another possible explanation may be related to the differences in the electronic states of nanorods and microrods suggested from Figures 1 and 2. At longer time scales, the rate of Fe(II) production in nanorod systems is considerably greater than that observed for microrods. For example, after 100 min of light exposure, the concentration of Fe(II) remains nearly constant in microrod suspensions. In contrast, the dissolved Fe(II) concentration more than doubles over the same time interval in nanorod systems. At these greater time scales, aqueous Fe(II) concentrations are appreciable, and the second dissolution pathway catalyzed by dissolved Fe(II) is more likely to contribute to goethite dissolution. We contend that the enhanced rate of Fe(II) production in nanorod systems at longer time scales is evidence that the rate of the Fe(II)promoted dissolution pathway is far greater on nanorods, consistent with the size-dependent trend we observed for nonphotochemical, ligand-promoted goethite dissolution (see Figure 7). Particle aggregation and uncertainty in reactive surface area prevents more quantitative comparisons of rates of oxalate decay and dissolved Fe(II) and Fe(III) production. However, insights into the nature of the surface site at which dissolution occurs were obtained from TEM analysis of reacted suspensions; changes in particle morphology were identical to those observed after dark dissolution experiments with oxalate (see Figure 8). This suggests that proton-promoted, ligand-promoted, and photochemical reductive dissolution occur at the same location on the goethite surface namely, the (021) crystal faces.


Figure 10. Comparison of total dissolved iron concentration generated in aerated nanorod and microrod suspensions in the presence of oxalate and light. Experiments were conducted in 0.5 g/L suspensions in pH 3 solutions of 5 mM NaClO4.

Implications for Iron Mobilization from Mineral Dust Aerosol. To explore the implications of our results for bioavailable iron formation, the most logical data set to consider is how total dissolved iron concentration () dissolved Fe(III) + dissolved Fe(II)) varies as a function of time in oxygenated systems (Figure 10). Initially, microrods release more total dissolved Fe than nanorods, although the amount of Fe released from nanorods eventually surpasses that of microrods. The behavior at short time scales is primarily attributable to the sharp initial rate of Fe(III) release in microrod systems (Figure 9b), because the production of Fe(II) is roughly equivalent for both particle sizes over this same time interval (Figure 9c). At longer time scales, on the other hand, the dominant contribution to total dissolved Fe is dissolved Fe(II) production, which is eventually liberated from the surface of nanorods in much greater quantities relative to microrods. Due to the long time scales associated with transport of mineral dust, the data suggest that nanorods may be a more important long-term source of bioavailable Fe than microrods. Thus, they could play a dominant role in iron mobilization for iron-containing aerosols with long atmospheric residence times or deposited aerosol particles with relatively long residence times in the marine photic zone. Conclusions Results from the current study suggest that molecular-level differences exist in the surface chemistries of goethite nanorods and microrods and that these differences have implications for their interactions with oxalate. Most notably, ATR-FTIR analysis suggests that nanorods exhibit a greater density of surface hydroxyl groups relative to microrods and that differences likely exist in the oxalate complexes formed on the surfaces of different sized particles. It is challenging to quantitatively assess the impact of these molecular-level differences on macroscopic reactivity because oxalate induces extensive aggregation, introducing much uncertainty as to the available surface area in each suspension. However, results of ligand-promoted reactions suggest that nanorods tend to be more labile than microrods, a result with fairly important implications for the role that nanoscale iron oxides may play in iron solubilization processes. Specifically, ligand-promoted, nonreductive dissolution reactions occur at a much greater rate on nanorod surfaces, and the difference in reactivity far exceeds that expected from simple considerations of nanorod surface area.

J. Phys. Chem. C, Vol. 113, No. 6, 2009 2185 In terms of photochemical dissolution processes, nanorods appear to be a more important, long-term source of dissolved iron production in the presence of light and oxalate. Although microrods produce more total dissolved iron initially, their rate of dissolved iron production curtails over time. In contrast, the rate of total dissolved Fe production steadily increases in nanorod systems. This difference is likely linked to the dominant mechanism of reductive dissolution taking place in these systems at low pH. We speculate that the enhanced rate of Fe(II) production observed for nanorods at longer time scales results from their greater susceptibility to Fe(II)-promoted dissolution relative to microrods. On the other hand, the relatively comparable rates of Fe(II) production observed initially may be an indication that the rate of direct photochemical reduction of nanorods and microrods is roughly comparable at equivalent goethite mass loadings. Given the strong evidence provided by Sullivan and Prather70 using single particle mass spectrometry that oxalate is preferentially associated with mineral dust aerosol compared to other atmospheric aerosols (e.g., carbonaceous aerosols), further studies that extend the results of the current work to consider dissolution reactions at higher pH values would be valuable. These higher pH values are more relevant in some atmospheric waters, as well as marine environments in which such ligandpromoted and photochemical reductive dissolution mechanisms will also be important. Acknowledgment. The authors acknowledge Dr. Jonas Baltrusaitis for his assistance with XPS, Drew Latta and Robert Handler for their insightful discussions, and the three anonymous reviewers whose comments and suggestions greatly improved the quality of this manuscript. This material is based upon work supported by the National Science Foundation under NSF EAR0506679 and CHE-0503854. Any opinions, findings, and conclusions or recommendations expressed in this material are those of the authors and do not necessarily reflect the views of the National Science Foundation. References and Notes (1) Ridgwell, A. J. Philos. Trans. R. Soc. A 2002, 360, 2905–2924. (2) Jickells, T. D.; et al. Science 2005, 308, 67–71. (3) Duce, R. A.; Tindale, N. W. Limnol. Oceanogr. 1991, 36, 1715– 1726. (4) Jickells, T. D.; Spokes, L. Atmospheric Iron Inputs to the Oceans. In The Biogeochemistry of Iron in Seawater; Turner, D. R., Hunter, K. A., Eds.; John Wiley & Sons: Chichester, England, 2001; pp 85-121. (5) Boyd, P. W.; et al. Science 2007, 315, 612–617. (6) Martin, J. H.; et al. Nature 1994, 371, 123–129. (7) Anderson, M. A.; Morel, F. M. M. Mar. Biol. Lett. 1980, 1, 263– 268. (8) Anderson, M. A.; Morel, F. M. M. Limnol. Oceanogr. 1982, 27, 789–813. (9) Rich, H. W.; Morel, F. M. M. Limnol. Oceanogr. 1990, 35, 652– 662. (10) Goudie, A. S.; Middleton, N. J. Earth-Sci. ReV. 2001, 56, 179– 204. (11) Usher, C. R.; Michel, A.; Grassian, V. H. Chem. ReV. 2003, 103, 4883–4939. (12) Deguillaume, L.; Leriche, M.; Desboeufs, K.; Mailhot, G.; George, C.; Chaumerliac, N. Chem. ReV. 2005, 105, 3388–3431. (13) Claquin, T.; Schulz, M.; Balkanski, Y. J. J. Geophys. Res. 1999, 104, 22243–22256. (14) May, H. M.; Kinniburgh, D. G.; Helmke, P. A.; Jackson, M. L. Geochim. Cosmochim. Acta 1986, 50, 1667–1677. (15) Behra, P.; Sigg, L. Nature 1990, 344, 419–421. (16) Stumm, W.; Sulzberger, B. Geochim. Cosmochim. Acta 1992, 56, 3233–3257. (17) Martin, S. T. Precipitation and Dissolution of Iron and Manganese Oxides. In EnVironmental Catalysis; Grassian, V. H., Ed.; CRC Press: Boca Raton, FL, 2005; pp 61-81.

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