Atmospheric Processing and Iron Mobilization of Ilmenite: Iron

Subscriber access provided by READING UNIV

Article

Atmospheric Processing and Iron Mobilization of Ilmenite: An Iron Containing Ternary Oxide in Mineral Dust Aerosol Eshani Hettiarachchi, Omar Hurab, and Gayan Rubasinghege J. Phys. Chem. A, Just Accepted Manuscript • Publication Date (Web): 16 Jan 2018 Downloaded from http://pubs.acs.org on January 16, 2018

Just Accepted “Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted online prior to technical editing, formatting for publication and author proofing. The American Chemical Society provides “Just Accepted” as a free service to the research community to expedite the dissemination of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully peer reviewed, but should not be considered the official version of record. They are accessible to all readers and citable by the Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore, the “Just Accepted” Web site may not include all articles that will be published in the journal. After a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web site and published as an ASAP article. Note that technical editing may introduce minor changes to the manuscript text and/or graphics which could affect content, and all legal disclaimers and ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or consequences arising from the use of information contained in these “Just Accepted” manuscripts.

The Journal of Physical Chemistry A is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

Page 1 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

The Journal of Physical Chemistry

Atmospheric Processing and Iron Mobilization of Ilmenite: An Iron Containing Ternary Oxide in Mineral Dust Aerosol Eshani Hettiarachchi, Omar Hurab and Gayan Rubasinghege* Department of Chemistry, New Mexico Institute of Mining and Technology, Socorro, NM 87801 [email protected] Abstract Over the last several decades, iron has been identified as a limiting nutrient in about half of the world’s oceans. Its most significant source is identified as deposited iron-containing mineral dust that has been processed during atmospheric transportation. The current work focuses on chemical and photochemical processing of iron-containing mineral dust particles in the presence of nitric acid, and an organic pollutant dimethyl sulfide under atmospherically relevant conditions. More importantly, ilmenite (FeTiO3) is evaluated as a proxy for the iron-containing mineral dust. The presence of titanium in its lattice structure provides higher complexity to mimic mineral dust, yet it is simple enough to study reaction pathways and mechanisms. Here, spectroscopic methods are combined with dissolution measurements to investigate atmospheric processing of iron in mineral dust, with specific focus on particle mineralogy, particle size, and their environmental conditions, i.e., pH and solar flux. Our results indicate that the presence of titanium elemental composition enhances iron dissolution from mineral dust, at least by two fold comparison with its non-titanium containing counterparts. The extent of iron dissolution and speciation is further influenced by the above factors. Thus, our work highlights these important, yet unconsidered, factors in the atmospheric processing of iron-containing mineral dust aerosol. *Author to whom correspondence should be addressed

1 ACS Paragon Plus Environment

The Journal of Physical Chemistry 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 2 of 35

Introduction Iron is the fourth most abundant element in the Earth’s crust. However, about 30% of the oceans show a limited supply of dissolved, and bioavailable iron. Given that the iron is an essential element for all living organisms including those in marine environments, primary productivity of phytoplankton in these regions is highly constrained by the amount of bioavailable iron.1-3 The oxic pH of seawaters (~pH 8.3) limits the amount of dissolved iron in the ocean to 0.1 nM from the direct dissolution of the Earth’s crust.4 However, the total iron concentration of the ocean reaches as high as 2 nM, in spite of this limitation. Therefore, it is hypothesized that the addition of iron to ocean waters from Fe-containing atmospheric aerosol, due to the acidic aerosol waters, elevates the total bioavailable iron.5,6 More recent satellite studies on algal populations in Atlantic Ocean have further confirmed enhanced blooming of algae during dust storms.7 Moreover, changes in the supply of dissolved iron from the atmosphere to the ocean may alter the oceanic carbon uptake. Nonetheless, significant uncertainties regarding the mineralogy of the mineral dust persist. The total Fe concentration in aerosol particulate matter (PM) ranges from 5 to 17,000 ng.m-3 where the dominant source is the mineral dust particles that had lifted into the atmosphere by wind action. However, a number of recent studies have reported other significant sources of soluble iron. Journet, et al. reported that Fe-containing clays accounted for about 90% of the soluble iron.8 Anthropogenic sources of soluble iron include combustion sources, i.e., biomass burning that releases fly ash and industrial activities, and represent up to 30% of total iron deposited into ocean waters.3,9,10,11 The average iron concentration in the marine atmosphere ranges from 200 to 500 ng.m-3 that could easily reach 10,000 ng.m-3 during Asian dust storms. It has also been estimated that annually 14-16 Tg of Fe-containing aerosol is deposited into the oceans, with approximately 50% of the total deposition occurring in the North Atlantic Ocean.3,9,12


Page 3 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Given that iron is only marginally soluble in seawater, the higher fraction of bioavailable iron in deposited mineral dust could be as high as 50% of iron content, suggesting that dissolution occurs during atmospheric transport.10,13 Among several explanations for the above increased solubility of iron, aerosol acidification and surface reactions, in particular with those involving anthropogenic gases (HNO3, SO2), have been prominent. Numerous previous laboratory and field studies have shown that the particle size, different acid types, and photochemical reactions affect proton-promoted dissolution rates of mineral dust.3, 5, 14-16 Jickells et al. observed higher iron solubility for the fine mode aerosols compared to the coarse mode.12 Because smaller particles are transported farther from the source, there is a greater probability that the deposited Fe-containing particles are smaller in size with increased surface reaction sites. However, Rubasinghege et al. found that at higher ionic strengths, aggregation of smaller particles dominates over surface area, leading to quenching of dissolution reactions.17 Apart from the particle size, it has been shown that the extent of crystallinity and weathering influences iron dissolution. Higher iron dissolution has been reported for the amorphous phase of atmospheric aerosol compared to that of the crystalline phases.18, 19,20, and 21 Given that iron dissolution is greatly influenced by the elemental composition and energetics of mineral surface, mineralogy of atmospheric dust may play a vital role on solubility. However, most of the previous work on iron dissolution has been mainly focused on single iron oxides, i.e., hematite (-Fe2O3), goethite (-FeOOH).8, 10, 20 Mineral dusts are complex mixtures of various metals and metal oxides with significant differences in their surface and bulk propoerties.22, 23 Hence, single iron oxide proxies do not describe the complexity in mineralogy or synergetic effects among various metals in the mineral dust. On the other hand, studies using only authentic mineral dust are too complex and do not provide molecular-level insights to understand the details of the reaction mechanisms.



Page 4 of 35

The work presented here, for the first time, discusses the influence of mineralogy on iron dissolution using a new model system, ilmenite (FeTiO3) that exhibits a higher complexity to mimic mineral dust; it is also simple enough to reveal the details of the reaction pathways. Ilmenite has been reported as a component of Fe-containing mineral in Saharan dust, the largest sources of mineral dust deposited to the Atlantic Ocean.13,24 Kandler et al. reported the mineralogy of the dust plumes collected at Izana; a Saharan desert reaching at Spain to have a Ti:Fe ratio of 1:6 suggesting that 10% of iron-rich minerals are ilmenite while 40% of it is hematite.25 Moreover, rutile (TiO2), another major reactive component of mineral dust, has shown higher uptake of acidic gases (e.g., NO2, SO2) that could potentially enhance the metal dissolution.22, 24, 26 Thus, a mix of iron and titanium would make a perfect model system to study the effects of mineralogy, in particular the metal-metal synergistic effects. The current study evaluates the dissolution of ilmenite for iron, under different environmental conditions, and compares against wellstudied simple Fe sources, hematite and maghemite (-Fe2O3). Experimental Materials & Methods Reagents. All chemicals were reagent grade or better and were used as received. Ilmenite (FeTiO3, American elements, 99.5%) was used as the proxy for iron-containing mineral dust. This sample is referred as “large” ilmenite. The dissolution results are compared with other commonly used single component proxies, i.e., hematite (-Fe2O3, US nanomaterials, 99.5%), maghemite (-Fe2O3, US nanomaterials, 99.5%). Heterogeneous uptake experiments were carried out using concentrated nitric acid (HNO3, Fischer Scientific, 67-70%) and DMS ((CH3SCH3, Sigma Aldrich, 99%). Sodium sulfate (Na2SO4, Sigma Aldrich, 99%), and sodium nitrate (NaNO3, Acros, crystals) were used in solution phase adsorption studies. All the solutions were prepared in Milli-Q water (18Ω, Milli-Q Advance 10). Hydroxylamine hydrochloride (NH2·OH·HCl, Acros Organic, 99%), ammonium fluoride (NH4F, Baker Chemicals, 99%), concentrated hydrochloric acid (HCl, Scholar chemicals, 36%), ammonium acetate 4 ACS Paragon Plus Environment

Page 5 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(CH3COONH4, Mallinckrodt, 99%), acetic acid (CH3COOH, VWR International, glacial),

1,10-

phenanthroline (C12H8N2, Acros organics, 99+%), and ferrous ammonium sulfate hexahydrate (Fe(NH4)2(SO4)2·6H2O, Fischer Scientific, 98.5%) were used during the analysis of dissolved iron based on the procedure previously described in Stucki et al.27 Experimental Methods Preparation of “small” Ilmenite Prospero and others have reported that the mineral dust is subjected to physical and chemical weathering during their long range transportation. This may even result distorted structures.28 In this work, physically weathered ilmenite with smaller particle size was prepared by high energy milling of large ilmenite. The preparation was carried out using Fritsch Pulverisette Mill in a zirconia jar. During the milling process, approximately 200 g of 10 mm beads were used. ~ 30 g of large ilmenite was first milled at 400 rpm for 16 repeats of 15 mins with the total time of 240 mins followed by a 30 min pause; the milling time was doubled to 480 mins by 32 repeats of 15 min millings. The milled samples are referred to as “small” ilmenite. Particle Characterization The shape and size of mineral particles were determined from single particles analysis with scanning electron microscope (SEM). The size distribution was determined by analyzing ~800 particles using the software package ImageJ. Surface areas of mineral samples were measured in a seven-point N2-BrunauerEmmet-Teller (BET) isotherm using a Quantachrome Autosorb-1 surface area analyzer. Samples were outgassed overnight (~24 h) at a temperature of 105°C prior to the BET analysis. X-Ray diffraction (XRD) of mineral proxies was measured in a Pananalytical XPert Pro Diffractometer equipped with Cu Source. Sample spectra were compared with reference XRD using ICDD database provided by International Centre for Diffraction Data. 5 ACS Paragon Plus Environment


Page 6 of 35

Attenuated total reflectance-Fourier transform infrared (ATR)-FTIR spectra of samples were obtained using a Thermo Scientific Nicolet IS50 spectrophotometer, equipped with a liquid-nitrogen cooled mercury-cadmium-telluride (MCT) detector and Ge ATR element. In ATR–FTIR experiments, ~20 mg of the solid sample was suspended in 1 mL of Milli-Q water followed by sonication for 20 min to break aggregates. The suspension was evenly spread onto the Ge ATR element using a transfer pipette. The sample was then allowed to completely dry in the presence of dry air (CO2 and H2O free) flow that resulted a thin film on the Ge ATR element for further analysis. Heterogeneous Uptake of Gas-phase HNO3 and DMS on Ilmenite The adsorption of HNO3 and DMS on ilmenite surfaces was studied using a custom-built flow reactor system coupled with an ATR–FTIR spectrometer. A schematic diagram of the complete flow reactor system is shown in Figure 1. A thin layer of the sample was prepared on the ATR crystal as described above. Here, dry air (%RH < 5, CO2 free) was used as the carrier gas. The air flow was controlled with flow meters (FM) and mass flow controllers (MFC). Two water bubblers were used to obtain the desired humidity while the acid bubbler was used to introduce the HNO3 or DMS vapor to the gas phase. A 1:3 mixture of conc. HNO3 and H2SO4 was used to generate HNO3 vapor. DMS solution, without any additives, was used to produce gas-phase DMS. Carrier gas was mixed with other reactants of interest in the mixing chamber prior to being introduced into the sample cell. ATR-FTIR spectra were collected at certain time intervals while the deposited sample was exposed to the desired gas mixture. After 120 mins, the acid vapor flow was stopped and the sample was evacuated under dry air flow overnight. The “exposed” and “evacuated” spectra were used to study reversible and irreversible uptake of the gases onto mineral surfaces. All the gas phase adsorption studies were carried out in the dark at 25 °C. These experiments were triplicated to confirm the reproducibility and consistency of adsorption.


Page 7 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Figure 1. Schematic diagram of flow reactor system combined with Furrier Transform Infrared Spectrometer (FTIR). AF – moisture air filter; MV – manual valve; MFC – Mass Flow Controller; FM – Flow Meter; RGA – Residual Gas Analyzer. Dissolution Measurements-Batch Reactor Studies Batch reactor studies were carried out to measure dissolved iron using custom-built glass reactors. The reaction vessel has a capacity of 100 mL with a removable airtight top. The particle loading was ~0.2 g/L of Fe-containing mineral in acid solutions at low pH (pH 1 and 2). Prior to the dissolution experiment, the acid solution was purged with nitrogen gas at 5 sccm for 5 mins to obtain a reduced atmosphere. The experiments were performed in the absence and presence of a solar simulator (150 W xenon lamp, New Port Corp). The quartz window (12.5 cm2) mounted on the top permitted light during the solar experiments. The glass reactor is also equipped with a temperature probe and a standardized pH electrode to measure these parameters throughout the dissolution experiment. The temperature was kept constant



Page 8 of 35

through the use of a water jacket. Over time, samples were periodically removed from the reactor using a disposable syringe that was connected to 12 cm Teflon tubing. The collected samples were filtered through 0.2 m filters and analyzed using the 1,10-phenanthroline method.27 All the batch reactor studies were triplicated. All dark and light experiments were conducted in triplicate with average measurements reported. Reported errors represent one standard deviation. Molecular Level Study of Solution Phase Adsorption Solution phase adsorption studies were performed using an ART-FTIR flow-through system to better understand the molecular level adsorption of acid anion on Fe-containing mineral surface. A thin layer of the mineral sample was prepared on the ATR crystal as described above. The dried mineral surface was exposed to a continuous flow of acid anion solution (1 mL min-1; 0.01 M Na2SO4 or HNO3) for a total period of 60 mins. The surface was allowed to equilibrate before collecting the IR spectra of the mineral surface. The solution spectrum for anion was collected exposing the bare crystal to the acid anions solution. Results and Discussion Characterization of Fe-containing Mineral Particles Surface and bulk characterization of the mineral particles provides important information on their particle size, shapes as well as crystal phases and purity. SEM images for representative samples of Fe-containing minerals are provided in Figure 2(a-d). As illustrated, the mineral particles are different from each other in their shapes and sizes. Given that spherical particles have a shape factor (χ) closer to 1, hematite and maghemite samples were more spherically shaped.29 However, both smaller and larger ilmenite samples were more irregular in shape and randomly oriented with χ ranging from 1.2 to 1.5. The size distribution; determined from SEM images, revealed that these particles range from nanoscale to microscale. Hematite, maghemite, and small ilmenite particle sizes were less than a micron with 32 (±4) nm, 35 (±3) 8 ACS Paragon Plus Environment

Page 9 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


nm, and 858 (±199) nm, respectively, while large ilmenite was 3400 (±300) nm. The broader size distribution of small ilmenite resulted from the method of preparation.

Figure 2. SEM images of a) -Fe2O3, b) -Fe2O3, c) Large FeTiO3, d) Small FeTiO3 and XRD patterns of e) - Fe2O3, f)  - Fe2O3, g) FeTiO3. The BET surface areas of the samples correspond to their particle sizes. Hematite, maghemite, small ilmenite, and large ilmenite showed surface areas of 100(±4) m2 g-1, 65(±3) m2 g-1, 6.8(±0.1) m2 g-1, and 0.51(±0.01) m2 g-1, respectively. The small ilmenite thus exhibits a 13-fold higher surface area than large 9 ACS Paragon Plus Environment


Page 10 of 35

ilmenite. No pronounced differences between small and large ilmenite was identified in ATR–FTIR spectroscopy. (Results are not shown here). The absorption bands at ~3500 cm-1 were observed for surface hydroxyl groups. As shown in Figure 2(e-g), XRD patterns of mineral samples are in good agreement with their respective reference XRD patterns confirming no phase change or contamination. However, the relative intensity of small ilmenite is less compared to large ilmenite, indicating a reduction in crystallinity of ilmenite during the milling. The mechanical milling may also introduce defects to the ilmenite lattice structure. Li et al. reported similar defects, especially along with the c-axis of the unit cell.19. Having Fe2+ and Ti4+ ions alternatively aligned, the modifications along the c-axis of ilmenite crystal would be rather important in detaching any of these ions from its lattice.

Gas Phase Adsorption Studies-Flow Reactor Measurements Heterogeneous uptake of acidic gases and water vapor on Fe-containing mineral surface yields an acidic deliquescence layer that promotes iron dissolution, and thus the bioavailable iron production.5 The adsorption of HNO3 and DMS onto ilmenite particles, small vs. large, was studied using a flow reactor system coupled with the ATR-FTIR. Both large and small ilmenite showed stronger adsorption of HNO3 and DMS via various oxide-coordinated surface complexes, and thus acidifying the particle surface. Uptake of HNO3 on Ilmenite ATR-FTIR measurements of HNO3 uptake on large and small ilmenite at a flow of 5 sccm are shown in Figure 3. FTIR spectra provide information on vibrational modes of different surface-adsorbed species, molecular-reversible and dissociative-irreversible adsorption. As can be seen in Figure 3(a) and (b), exposure of both large and small ilmenite to gas phase HNO3 results in adsorbed nitric acid species on the surface (ca. 900–1400 cm–1) with a concomitant loss of surface -OH, as indicated by the negative intensities in the -OH stretching region (ca. 3000–3700 cm–1).5,26, 30, 31 In the presence of HNO3 flow, the spectra in Figure 3 are labeled “total HNO3 uptake”. When the nitric acid reacted-ilmenite surface is 10 ACS Paragon Plus Environment

Page 11 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


exposed to dry air without nitric acid vapor, several bands disappear from the spectrum, indicating a loss of weakly or molecularly bound nitric acid that can be referred to as the “reversible uptake” of nitric acid. The surface exposed to dry air results in only the strongly bound chemisorbed nitrate on the ilmenite. These spectra are labeled as “irreversible HNO3 uptake” in Figure 3. The spectral assignments for the dissociative irreversible adsorption of nitric acid on ilmenite, and more specifically to oxide-coordinated and adsorbed water-solvated nitrate, are given in Table 1. The cartoon in Figure 3(c) illustrates the three modes of oxide-coordinated nitrate on ilmenite surfaces. It should be noted that both large and small ilmenite surfaces show the characteristic peaks of adsorbed nitrate onto the mineral surface forming monodentate (Cs), bidentate (C2V), and bridged complexes. The peak at 1038 cm-1 on large ilmenite and the peak at 1011 cm-1 on small ilmenite can be assigned to the 1 mode of the bridged nitrate. The 3 (high) vibrational mode of bridged complex was assigned to the peak at 1611 cm-1 on large ilmenite. The same vibrational mode was attributed to the peak at 1606 cm-1 on small ilmenite. The peaks around 1240 cm-1 were assigned to the 3(low) vibrational mode of the bidentate complexes while peaks at 1579 cm-1 and at 1563 cm-1 on large and small ilmenite were assigned to 3 (high) of the same nitrate complex. Peaks at 1292 cm-1 and 1301 cm-1 on large and small ilmenite were assigned to the 3 (low) of the monodentate complexes, respectively. The 3 (high) of the monodentate complex were assigned to the peak at 1495 cm-1 and 1494 cm-1 on large and small ilmenite. On many mineral surfaces, nitrate can co-adsorb with water as solvated nitrate.26, 31 In the current work, the peaks obtained around 1345 cm-1 and 1358 cm-1 were assigned to solvated nitrate on large and small ilmenite, respectively.



Page 12 of 35

Figure 3. FTIR spectra of adsorbed HNO3 on (a) large ilmenite and (b) small ilmenite at a flow of 5 SCCM of the acid to show reversible and irreversible uptake of HNO3. “Total HNO3 uptake” are the spectra collected in the presence of HNO3 acid in the air stream. “Irreversible HNO3 uptake” is in the absence of HNO3 in the air stream. See text for further details. (c) Illustration of the surface complexes form on Ilmenite surface with nitric acid. 12 ACS Paragon Plus Environment

Page 13 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Table 1. Assignment of vibrational frequencies for adsorbed HNO3 on large and small ilmenite Experimental Frequency (cm-1) Vibrational Mode Large Ilmenite Molecular adsorbed HNO3 1 (NO2)

Small Ilmenite

Reference Frequency for -Fe2O3 (cm-1)31

Reference Frequency for TiO2 (cm-1)26

1674 n.o. n.o.

1675 n.o. n.o.

1679 1337 1297

1683 1336 1305

Oxide-coordinated nitrate monodendate 3 (high) 3 (low) 1

1495 1238 n.o.

1494 1301 n.o.

1555 1281

1509 1282

bidendate 3 (high) 3 (low) 1

1579 1292 n.o.

1563 1240 n.o.

1588 1229

1581 1243

bridged 3 (high) 3 (low) 1

1611 n.o. 1038

1606 n.o. 1011

1622 1203 993

1636 1230 1005

1358 n.o. n.o.

1399 1346 814

1406 1331 1046

δ (OH)

s (NO2)

Water-solvated adsorbed nitrate 1345 3 (high) n.o. 3 (low) n.o. 1 n.o. – not observed in the current work

Uptake of DMS on Ilmenite As shown in Figure 4 and Table 2, adsorption of DMS vapor onto both large and small ilmenite surfaces showed characteristic spectral bands.32-35 The peaks at 956 cm-1 and 1029 cm-1 in large ilmenite and peaks at 957 cm-1 and 1014 cm-1 in small ilmenite can be assigned to the rocking modes of CH3 group of



Page 14 of 35

adsorbed DMS. The absorption at 1405 cm-1 on large ilmenite was assigned to the bending mode of the C-H bond in the methyl group. The bands for the asymmetric deformation of CH3 group were observed around 1435 cm-1 on large ilmenite and 1441 cm-1 on small ilmenite. The C–H stretching bands were observed around 3000 cm-1 for the both large and small ilmenite. However, the absorption bands for large ilmenite were less intense than the corresponding bands for small ilmenite, indicating greater adsorption on smaller particles. On both the ilmenite surfaces, the predominant bands appeared were due to various vibrational modes of -CH3 suggesting that the CH3 group of adsorbed DMS is free to move.

Table 2. Assignment of vibrational frequencies for adsorbed DMS on large and small ilmenite Vibrational Mode  (C-H) – stretching mode of CH3

Experimental Frequency (cm-1) Large Ilmenite Small Ilmenite

Reference Frequency (cm-1)32,35 ca. 3000 for CH3 stretching

2846

2858

2923 2976

2930 2985

ρr (CH3) – rocking mode of CH3

956 1029

957 1014

δ (CH3) - bending mode of CH3

1405

1405 (br.)

1037

Asymmetric deformation of CH3

1435

1441

1432

3350

3714

3395

 (O-H) – stretching mode of Hbonded –OH br.- broad peak

972.5 1030.0

Thus, DMS chemisorbed on the ilmenite surface via sulfur-iron bond. Moreover, a spectral band appeared at 3350 cm-1 for large ilmenite and 3714 cm-1 for small ilmenite; these can be attributed to O-H stretching that involved in H-bonding. This feature suggests that DMS could also H-bonded to the ilmenite surface. The proposed DMS adsorbed complexes are illustrated in the cartoon shown in Figure 4(c). 14 ACS Paragon Plus Environment

Page 15 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Figure 4. FTIR spectra of adsorbed DMS on (a) large ilmenite and (b) small ilmenite at a flow of 5 SCCM of the acid to show reversible and irreversible uptake of DMS. “Total DMS uptake” are the spectra collected in the presence of DMS acid in the air stream. “Irreversible DMS uptake” is in the absence of DMS in the air stream. See text for further details. (c) Illustration of the surface complexes form on Ilmenite surface with nitric acid.



Page 16 of 35

Batch Reactor Studies on Dissolved Iron Exposed to HNO3 and H2SO4 Role of Mineralogy The mineralogy of the source materials influences the formation of surface complexes, and thus dissolution chemistry of the surface species. The mechanisms of iron dissolution from mineral dust, especially in proton-promoted dissolution at low pH conditions, further highlight the importance of surface complexation.5, 36-38 A comparison of total dissolved iron for the three model systems: hematite, maghemite and small ilmenite at pH 2 is shown in Figure 5. These data have been normalized to the respective total mass and surface area of the mineral sample. As seen in Figure 5(a), the extents of total iron dissolution under dark conditions for hematite and maghemite are about ten- and seven-fold higher than that of small ilmenite on mass basis. Both hematite and maghemite have higher surface areas compared to small ilmenite, and thus showed higher iron dissolution on mass basis. Given that high surface area of smaller particles masks mineralogy effects, it is more important to compare total dissolved iron normalized to their respective surface areas. According to Figure 5(b), small ilmenite showed a higher extent of dissolution compared to hematite or maghemite on surface area basis. Here, total iron dissolution from small ilmenite after 48 hrs is ~2-fold higher than that of hematite and maghemite. It is also worth highlighting that even with lower % Fe by mass in ilmenite compared to the other two proxies, 37% vs. 70%, small ilmenite produced more total dissolved iron compared to single component iron oxides. Moreover, it is also important to mention that similar trend was observed with large ilmenite, suggesting that the observed enhancement is primarily due to minerology and crystallinity effects rather than surface modifications by the milling process. These data are provided in Figure S1 in Supporting Information. In dark, both hematite and maghemite did not leach any detectable Fe(II) concentration. However, small ilmenite and large ilmenite leached 75% and 68% Fe(II) fractions, respectively. One possible reaction


Page 17 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


mechanism for the dissolution of ilmenite in acidic media, proposed in previous work, is given in Equation 1.39,40 FeTiO3(s) + 2 HA (aq)  Fe2+ (aq) + TiO2+ (aq) + 2OH-(aq) + 2A-(aq)

Equation 1

Where; HA is the acid used. Here, simultaneous dissolution of Ti4+ and Fe2+ via a surface reaction has been proposed. Ti and Fe are alternatively aligned along the c-axis of the ilmenite surface, and thus, removal of one metal atom alters the electronic properties of the surface and further enhances the dissolution.19,26 Therefore, it can be speculated that the presence of Ti4+ in the ilmenite lattice may enhance the dissolution of iron in nitric acid medium under conditions.

Figure 5. Dissolution of α-Fe2O3, -Fe2O3, and small FeTiO3 is monitored by the formation of total soluble Fe and Fe(II) in nitric acid solution at pH 2.0 under light and dark conditions. These plots compare iron dissolution on per mass (a, c, and e) and per surface area (b, d, and f) basis. The data has fitted to Langmuir model. 17 ACS Paragon Plus Environment


Page 18 of 35

As seen in Figure 5(c) and (d), total iron dissolution is enhanced in the presence of solar radiation for all three minerals, regardless of their total iron percentages. On surface area basis, as shown in Figure 5(d), the extent of total iron dissolution for irradiated small ilmenite was at least an order of higher compared to that of hematite and maghemite under the same conditions. Presence of titanium in the ilmenite lattice can influence the photochemistry of these semiconductor minerals.41,42 According to Xiong et al., Ti(III) surface defects (TSDs) can be found on TiO2, and are proposed to be generated in large quantities under irradiation, thereby increasing the surface activity of TiO2.43 Given similarities in the structural and electronic features between TiO2 and ilmenite, such Ti(III) surface defects could also be expected on ilmenite upon irradiation. Thus, the reduction of surface Ti(IV) to Ti(III) can introduce a photo-reductive mechanism increasing the dissolution of ilmenite. The dissolved Fe(II) concentrations, on mass and surface area basis, are shown in Figure 5(e) and (f), respectively. The Fe(II) production for ilmenite and single component iron oxides did not show a significant difference on mass basis. However, the surface area normalized data highlight that dissolved Fe(II) concentration for small ilmenite was at least 10-fold higher than that of hematite and maghemite. Since hematite and maghemite contain Fe(III) in their lattice, the production of Fe(II) is due to photochemical reactions initiated by electron-hole pairs.5,36 Upon irradiation, photoexcited electrons in the conduction band reduces Fe(III) to Fe(II) whereas positively charged holes oxidizes adsorbed water to H+ and oxygen. 5,36 While acidic conditions enhances the dissolution, produced oxygen could oxidize some of the Fe(II) to Fe(III). Further, excess nitrate could oxidize Fe(II) to Fe(III) under acidic conditions, lowering the net production of Fe(II). (Discuss in detail below). On the other hand, having Fe(II) in the ilmenite lattice, high dissolved Fe(II) fraction is expected. However, dissolved Fe(II) fraction is about 25% of total dissolved iron, similar to other iron oxides, highlighting that the oxidation of Fe(II) to Fe(III) is dominant even under irradiation. Thus, ilmenite demonstrates very interesting iron dissolution properties, especially synergetic effects


Page 19 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


based on its elemental composition. In the current work, this new model system was further studied to better understand the roles of particle size, pH, and solar radiation on its iron dissolution. Role of Particle Size and pH of the Medium Smaller the particle size corresponds to higher the surface area available for surface complexation and greater the detachment of iron from the surface. As seen in Figure 6(a), small ilmenite showed higher total iron dissolution, on mass basis, for both pH 1 and pH 2 in nitric acid medium, in the dark conditions. The initial rates of iron dissolution, determined using Langmuir isotherm, are given in Table 3. At pH 2, the rate of total iron dissolution for large ilmenite was 35±1 M.g-1.hr-1 and that of the small ilmenite was 69±1 M.g-1.hr-1. At pH 1, the rate of total iron dissolution increased to 46±2 M.g-1.hr-1 for large ilmenite while small ilmenite showed no significant difference with 70±2 M.g-1.hr-1. Thus, the total iron dissolution for small ilmenite, at both pH 1 and pH 2, were ~2-fold higher compared to large ilmenite under the same conditions. Table 3. Initial rates of the total iron dissolution, determined from Langmuir model, for small and large ilmenite under different experimental conditions. Both mass normalized, M.g-1.hr-1, and surface area normalized, M.m-2.hr-1, rates are presented. Surface area normalized rates are given in parenthesis. HNO3 medium Sample

Small Ilmenite

Large Ilmenite

pH 1

H2SO4 medium

pH 2

pH 1

pH 2

Dark

Dark

Light

Dark

Dark

Light

70±2

69±1

197±2

218±2

104±1

37±2

(12±1)

(10±1)

(29±1)

(32±1)

(15±1)

(5±1)

46±2

35±1

42±2

40±2

48±2

28±3

(91±4)

(69±2)

(82±4)

(78±4)

(93±5)

(54±6)



Page 20 of 35

Figure 6. A comparison of total iron dissolution and Fe(II) formation between small and large FeTiO3 in nitric acid medium at pH 1 and pH 2 under dark conditions. The data are presented on per mass (a & c) and per surface area (b & d) basis. The data has fitted to Langmuir model.

Nonetheless, on surface area basis, large ilmenite showed a higher iron dissolution compared to small ilmenite, as shown in Figure 6(b) and Table 3. Dissolution may take place more rapidly in certain surface planes of ilmenite over the others. According to Li et al. and several others, a rapid dissolution has been reported from basal (0001) and (1011) crystal planes of “non-milled” ilmenite.19 However, during the preparation of small ilmenite using high energy milling, the crystallinity of particles decreases, as 20 ACS Paragon Plus Environment

Page 21 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


confirmed by the XRD data.19, 44 Therefore, the lower Fe dissolution for smaller particles, on surface area basis, could be attributed to the loss of basal (0001) and (1011) crystal planes. These results further confirm that the higher surface area of smaller ilmenite particles compensates this loss yielding higher solubility, on mass basis. Similar to HNO3, the studies carried out using H2SO4 medium showed a higher extent of total iron dissolution at pH 1 compared to that of pH 2. These results are presented in Figure S2 in Supporting Information. As given in Table 3, the rate of total iron dissolution in sulfuric acid at pH 1 is higher than that of pH 2, especially for small ilmenite with a ~2-fold increase. Thus, these results are in a good agreement with the established proton-promoted mechanism for Fe solubilzation at low pH environments.18, 36 The production of dissolved Fe(II) under dark conditions is shown in Figure 6(c) and (d) on a per mass and per surface area basis, respectively. At both pH 1 and pH 2, large ilmenite showed a higher dissolved Fe(II) fraction (= [dissolved Fe(II)]/[Total dissolved Fe]), ~75%, indicating no or very little oxidation of Fe(II) to Fe(III) occurs under these conditions. Even though small ilmenite showed a higher Fe(II) fraction at pH 2, ~75%, a very low dissolved Fe(II) concentration was observed at pH 1. These results suggest that the oxidation of Fe(II) to Fe(III) is more noticeable in small ilmenite in the presence of high nitric acid concentrations under dark conditions. Several recent studies have reported that in the presence of high nitrate concentrations at lower pH, dissolved Fe(II) oxidizes to Fe(III) while NO3- reduce to NO2, according to Equation 2.37 NO3- + 2Fe(II) + 2H+  NO2- + H2O + 2Fe(III)

Equation 2

Given that the free energy change for the above process is estimated to be -15 kJ mol-1, oxidation of Fe(II) to Fe(III) is thermodynamically feasible under highly acidic conditions. However, in sulfuric acid medium, dissolved Fe(II) from small ilmenite did not show any oxidation at pH 1. The redox reaction



Page 22 of 35

with sulfate anion, shown in Equation 3, has a free energy change of about +170 kJ mol-1 and thus, is not thermodynamically feasible. SO42- + 2Fe(II) + 2H+  SO32- + H2O + 2Fe(III)

Equation 3

Role of Acid Anion and Sunlight Different acid anions in the deliquescence layer form surface complexes on the mineral surface that are distinguishable from each other, resulting in significant differences in the production of bioavailable Fe.18 The effect of acid anion on iron dissolution may be further enhanced by the availability of solar radiation.45,46 In the current study, the dissolution of iron in the presence of two anions, nitrate and sulfate, was studied under light and dark conditions. These results are shown in Figure 7. Some of these differences between the anions are already discussed in the previous section under particle size and pH effects. As seen in Figure 7(a) and (b), regardless of the particle size, the rate of total iron dissolution is higher in the presence of sulfate compared to that of nitrate, at pH 2 under dark conditions. The rate of total iron dissolution for large and small ilmenite in sulfuric acid were 48±2 M.g-1.hr-1 and 104±1 M.g1.hr-1, respectively, indicating ~1.5-fold increase in the presence of sulfate compared to that of nitrate. A recent study by Fu et al. on iron dissolution from authentic mineral dust reported a similar enhancement in the presence of H2SO4 compared to HNO3.18 However, Rubasinghege et al. recently reported an opposite effect in their study on iron dissolution from goethite particles.5 Therefore, these previous works suggest that the effect of acid anion highly depends on the mineralogy of the mineral particle. Moreover, Fu et al. propose that the acid anion further influences the solubility by compressing electrical double layer, and consequently, destabilizing mineral particles colloids.18 Highlighting significant differences in their surface electrical properties, different minerals may thus exhibit significant variations in iron solubility based on their surface mineralogy. However, the extent of total iron dissolution showed a more noticeable enhancement in small ilmenite at pH 2, a ~2-fold increase; compared to that that of large 22 ACS Paragon Plus Environment

Page 23 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


ilmenite. These differences can be attributed to the higher surface charge introduced to small ilmenite during its preparation.

Figure 7. Dissolution of small and large ilmenite particles is monitored by the formation of total soluble Fe and Fe(II) in solution at pH 2.0. These plots compare the dissolution of ilmenite in the presence of two acid anions, SO42- vs. NO3-, under dark and light conditions. (a) and (b) Total iron, (c) and (d) Fe(II) production. The data has fitted to Langmuir model. In the presence of sunlight, adsorbed anion complex on the mineral surface undergoes dissociation via different redox reaction pathways. As shown in Figure 7(a) and (b), the rate of total iron dissolution is 23 ACS Paragon Plus Environment


Page 24 of 35

lower in the presence of solar radiation compared to its dark counterpart, in the presence of sulfate. The rate of total iron dissolution for large ilmenite in the presence of sulfate and solar radiation at pH 2 was 28±3 M.g-1.hr-1 showing a 1.7 times lower dissolution compared to that of the dark condition. Nevertheless, large ilmenite did not show a significant difference in the extent of total iron dissolution for both sulfate and nitrate anions. According to Figure 7(b), small ilmenite showed a more evident decline in both rate and extent of total iron dissolution in the presence of sulfate under solar radiation. The rate of total iron dissolution for small ilmenite under these conditions was 37±2 M.g-1.hr-1 which was about a 3-fold decrease compared to dark conditions. On the other hand, small ilmenite in HNO3 showed an enhancement in total iron dissolution upon irradiation. The rate of total iron dissolution for small ilmenite in the presence of nitrate and solar radiation was 197±2 M.g-1.hr-1 indicating a 3-fold increase compared to its dark condition counterpart. As discussed previously, the enhancement in total iron dissolution for ilmenite in the presence of solar radiation, at least partly, due to the formation of Ti(III) surface defects that introduces additional photo-reductive mechanisms. Moreover, in HNO3 solutions, adsorbed nitrate could also act as a chromophore and initiate additional photochemical reaction pathways, further enhancing total iron dissolution.47 However, it is apparent that in the presence of sulfate anion, the formation of surface defects no longer increases the total iron dissolution. Surface passivation partially explains the observed decrease of solubility in the presence of sulfate under the irradiation. A recent study done by Silveira et al. on catalytic activity of natural ilmenite revealed that, under irradiation, adsorbed sulfate layers passivate the surface, thereby decreasing the catalytic activity of ilmenite. The authors further suggested that the loss of iron from the ilmenite surface forms FeSO4 that deposits on the surface.41 However; it is unclear why the effect is not sufficient to passivate the surface without irradiation. Being an antiferromagnetic semiconductor oxide with TiO32- anion, ilmenite has a band gap varying from 2.4eV to 2.9eV,41, 48 which can photo-cleave water molecules in 24 ACS Paragon Plus Environment

Page 25 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


order to produce reactive oxygen species such as hydroxyl radicals in the case of TiO2 where Ti is present in the form of Ti4+. 22,49 Radical reactivity can convert Fe(II) to Fe(III) ions whose sulfate solubility is even less in the sulfuric acid medium.41, 49 In the presence of nitric acid, the possible products, Fe(NO3)2 and Fe(NO3)3, are both soluble in HNO3, therefore surface passivation can be overcame by diffusion of the products away from the mineral surface. Figure 7(c) and (d) illustrates a comparison of dissolved Fe(II) for small and large ilmenite under dark and light conditions at pH 2. As discussed above, both small and large ilmenite showed higher Fe(II) fractions, 75% and 68% respectively, in the nitric acid medium. These data are in good agreement with the previously reported value for the Fe(II)/Fe(III) ratio, 0.66.41 Even though large ilmenite in sulfuric acid medium showed a higher Fe(II) fraction (72%), that of small ilmenite was only about 1%. The reasons for the very low dissolved Fe(II) are unclear, and further studies are suggested here to investigate interactions between surface defects and acid anions, and the influence on redox reactions. Upon irradiation, both small and large ilmenite showed a lower Fe(II) fraction in the presence of nitrate or sulfate anions. The Fe(II) fraction for small and large ilmenite in HNO3 was 23% and 38%, respectively, under light conditions. In H2SO4, the dissolved Fe(II) for small and large ilmenite were only 31% and 14% for small ilmenite and large ilmenite, respectively. The lower Fe(II) fractions could be due to the oxidation of Fe(II) by free radicals generated in the presence of ilmenite under irradiation, as shown in Equation 4.22, 41, 49 Fe(II) + OH  Fe(III) + OH-

Equation 4

Solution Phase Adsorption Studies - Molecular Level Insights As discussed above, variations in surface-adsorbed complexes of acid anion result in significant differences in total iron dissolution and iron speciation. To better understand at the molecular level of surface–anion complexes formed on the ilmenite surfaces, the solution phase adsorption studies were 25 ACS Paragon Plus Environment


Page 26 of 35

performed. Those studies were carried out for both sulfate and nitrate media. However, in the nitrate medium, it is hard to distinguish the absorption bands in the aqueous phase, and thus, only the results from sulfate adsorption are discussed below. Figure 8 presents several ATR-FTIR spectra for sulfate adsorption on both large and small ilmenite at pH 2. For comparison, a solution-phase spectrum of sulfate anion in the absence of ilmenite is also shown.

Figure 8. ATR-FTIR surface spectra illustrating adsorption of aqueous phase sulfate on large and small ilmenite at pH 2. Spectra shown here are average of six successive additions of sulfate in 15 min intervals. The solution-phase spectrum, collected in the absence of ilmenite, is also shown. Illustrations of the surface complexes form on Ilmenite surface with aqueous sulfate is shown on the right. Sulfate in aqueous solution possesses a tetrahedral symmetry that showed only one peak ~1,100 cm-1 for the triply degenerate 3 asymmetric stretch. However, adsorbed sulfate showed a reduced symmetry that


Page 27 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


yields additional absorption peaks in the IR spectrum. During these experiments, small ilmenite surface showed immediate saturation with the sulfate solution, suggesting a higher rate of surface complexation. The peaks at 972 cm-1 and ~1010 cm-1 on large and small ilmenite, respectively, can be assigned to 1 vibrational mode of the iron–sulfate monodentate complex.50-52 Nevertheless, the peaks on small ilmenite are less intense, suggesting weak adsorption and formation of fewer surface complexes. Similarly, the peaks at 1141 cm-1 and 1144 cm-1 on large and small ilmenite are in good agreement with the 3 mode of iron-sulfate monodentate complex.50-52 Thus, adsorption of sulfate on large and small ilmenite yields ilmenite-sulfate complexes with monodentate coordination. However, the peaks on small ilmenite surface are broad and low in intensity. This suggests the surface of small ilmenite is less prone to make surface complexes with sulfate and further supports the lower dissolution rates obtained on the surface area normalized basis. Conclusions and Environmental Implications In the current study, ilmenite has been used as a proxy for Fe-containing mineral dust aerosol to investigate several factors that affect iron mobilization in low pH environments with a specific focus on the mineralogy of the particle. The effect of particle size and other environmental factors such as sunlight, pH, and type of acid anion were further studied. Ilmenite dissolution is about 2-fold higher than that of hematite or maghemite, on the surface area basis, suggesting the enhanced synergetic effects of iron and titanium mineralogy in ilmenite. Small ilmenite showed higher total iron dissolution than large ilmenite, on mass basis. However, on surface area basis, large ilmenite showed a higher total iron dissolution that goes beyond the surface area effect, which suggests that the high extent of dissolution of small ilmenite arose from its increased surface area due to smaller particle size. This can be inferred as the basal (0001) and (1011) planes of crystalline ilmenite are more prone to dissolution and the loss of crystallinity resulting from the preparation method of small ilmenite decreases its iron solubility on the surface area



Page 28 of 35

basis under dark conditions. Therefore, more crystalline ilmenite can contribute to more bioavailable iron due to the abundance of specific crystal planes. The total iron dissolution at lower pH is higher than that in the higher pH conditions. However, for small ilmenite, at lower pH, the amount of Fe (II) dissolved has been drastically lowered in the nitric acid medium, possibly due to nitrate redox cycling occurring on the ilmenite surface. Dissolution of iron under atmospherically relevant conditions showed significant differences in the presence of different acid anions. In dark conditions, usually sulfuric acid favors total iron dissolution from the ilmenite particles. However, the light conditions showed a suppression of the iron dissolution in the presence of sulfate ions, whereas nitric acid enhances iron dissolution under the same conditions. This suggests that the production of the soluble iron in the acidic conditions during the daytime conditions is not always enhanced. Depending on the mineralogy of the mineral dust particle, different anions behave differently under irradiation. The suppression in the presence of sulfate under the irradiation can be attributed to the surface passivation of ilmenite due to the formation of Fe(SO4) and Fe2(SO4)3. Nitric and sulfuric acids are generated in the atmosphere from precursor atmospheric gases such as nitrogen oxides and sulfur dioxide respectively. Any increase of such acidic gases in the atmosphere will yield higher adsorption of the nitrates and sulfates on Fe-containing mineral particles such as ilmenite. Quantifying the iron mobilization in different environmental conditions using single component iron oxides such as hematite and maghemite is useful yet does not provide knowledge on the effect of mineralogy of the particle. The atmospheric gases synergistically interact with other metal cations in the mineral dust surface yielding synergetic effect on overall iron mobilization. The current study suggests having just Ti4+ in the lattice of the ilmenite demonstrates interesting iron mobilization properties. As mineral dust being a complex mixture of many minerals, considering the effect solely of pure iron oxides will not provide the perfect model for iron mobilization. Therefore, it is necessary to study the synergetic 28 ACS Paragon Plus Environment

Page 29 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


effects of other metal cations in order to understand the atmospheric processing of iron-containing mineral dusts. Supporting Information 01. Figure S1 - A comparison of total iron dissolution in large ilmenite, hematite and maghemite at pH 2 nitric acid under dark conditions. 02. Figure S2 – Effects of particle size and pH on dissolution of Fe from Ilmenite in the sulfuric acid medium in the dark condition. Acknowledgement This work is supported by New Mexico Institute of Mining and Technology. Authors also would like to thank the Center for Micro-engineered Materials at University of New Mexico and Dr. Paul Furrier of the Department of Materials Engineering, New Mexico Tech and Dr. Virgil Leuth at New Mexico Bureau of Geology.

References (1) Martin, J. H.; Gordon, R. M.; Fitzwater, S. E. Iron in Antarctic Waters. Nature 1990, 345, 156– 158 (2)

De Baar, H. J. W.; Boyd, P. W. The Role of Iron in Plankton Ecology and Carbon Dioxide Transfer of the Global Oceans. Cambridge Univ Press, New York 2000

(3)

Mahowald, N. M.; Engelstaedter, S.; Luo, C.; Sealy, A.; Artaxo, P.; Benitez-Nelson, C.; Bonnet, S.; Chen, Y.; Chuang, P. Y.; Cohen, D. D.; et al. Atmospheric Iron Deposition: Global Distribution, Variability, and Human Perturbations . Annu. Rev Mar Sci. 2008, 245-278.

(4)

Rose, A. L; Waite, T. D. Kinetic Model for Fe(II) Oxidation in Seawater in the Absence and Presence of Natural Organic Matter. Environ Sci Technol. 2002, 36, 433– 444.



(5)

Page 30 of 35

Rubasinghege, G.; Lentz, R. W.; Scherer, M. M.; Grassian, V. H.; Finlayson-Pitts, B. J. Simulated Atmospheric Processing of Iron Oxyhydroxide Minerals at Low pH: Roles of Particle Size and Acid Anion in Iron Dissolution. PNAS 2010, 107, 6628–6633.

(6)

Anastasio, C.; Martin, S. T. 8 Atmospheric Nanoparticles. Mineral. Geochemistry 2001.

(7)

Gassó, S.; Grassian, V. H.; Miller, R. L. Interactions between Mineral Dust, Climate, and Ocean Ecosystems. Elements 2010.

(8)

Journet, E.; Desboeufs, K. V.; Caquineau, S.; Colin, J. L. Mineralogy as a Critical Factor of Dust Iron Solubility. Geophys. Res. Lett. 2008, 35, L07805.

(9)

Guieu, C.; Bonnet, S.; Wagener, T.; Loye-Pilot, M. D.; Biomass Burning as a Source of Dissolved Iron to the Open Ocean? Geophys. Res. Lett. 2005, 32, L19608, doi:101029/2005GL022962

(10) Borgatta, J.; Paskavitz, A.; Kim, D.; Navea, J. G. Comparative Evaluation of Iron Leach from Different Sources of Fly Ash under Atmospherically Relevant Conditions. Environ. Chem. 2016, 13, DOI: 10.1071/EN16046. (11) Sedwick, P. N.; Sholkovitz, E. R.; Church, T. M. Impact of Anthropogenic Combustion Emissions on the Fractional Solubility of Aerosol Iron: Evidence from the Sargasso Sea. Geochemistry, Geophys. Geosystems 2007, 8, 1–21. (12) Jickells, T. D. S.; An, Z. S.; Andersen, K. K.; Baker A R; Bergametti G; Brooks, N.; Cao, J. J.; Boyd, P. W.; Duce, R. A.; Hunter, K. A.; et al. Global Iron Connections Between Desert Dust, Ocean Biogeochemistry, and Climate. Science 2005, 308, 68–71 DOI: 10.1126/science.1105959. (13) Cornell, R. M.; Posner, A. M.; Quirk, J. P.; Crystal Morphology and Dissolution of Goethite. .J Inorg Nucl Chem 1974, 36, 1937– 1946 (14) Duggen, S.; Croot, P.; Schacht, U.; Hoffmann, L. Subduction Zone Volcanic Ash Can Fertilize the Surface Ocean and Stimulate Phytoplankton Growth: Evidence from Biogeochemical Experiments and Satellite Data. Geophys. Res. Lett. 2007, 34, L0161. 30 ACS Paragon Plus Environment

Page 31 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(15) Stefánsson, A. Iron(III) Hydrolysis and Solubility at 25 °C. Environ. Sci. Technol., 2007, 17, 6117– 6123. (16) Nadtochenko, V. A.; Kiwi, J. Photolysis of FeOH2+ and FeCl2+ in Aqueous SolutionPhotodissociation Kinetics and Quantum Yields. Inorg Chem 1998, 37, 5233– 5238 (17) Rubasinghege, G.; Elzey, S.; Baltrusaitis, J.; Jayaweera, P. M.; Grassian, V. H. Reactions on Atmospheric Dust Particles: Surface Photochemistry and Size-Dependent Nanoscale Redox Chemistry. J. Phys. Chem. Lett. 2010, 1729–1737. (18) Fu, H.; Cwiertny, D. M.; Carmichael, G. R.; Scherer, M. M.; Grassian, V. H. Photoreductive Dissolution of Fe-Containing Mineral Dust Particles in Acidic Media. J. Geophys. Res. Atmos. 2010, 115. (19) Li, C.; Liang, B.; Guo, L. hong; Wu, Z. b. Effect of Mechanical Activation on the Dissolution of Panzhihua Ilmenite. Miner. Eng. 2006, 19, 1430–1438. (20) Zarzycki, P.; Kerisit, S.; Rosso, K. M. Molecular Dynamics Study of Fe(II) Adsorption, Electron Exchange, and Mobility at Goethite (-FeOOH) Surfaces. J. Phys. Chem. C 2015, 119, 3111–3123 DOI: 10.1021/jp511086r. (21) Alexandrov, V.; Rosso, K. M. Insights into the Mechanism of Fe(II) Adsorption and Oxidation at Fe-clay Mineral Surfaces from First-principles Calculations. J. Phys. Chem. C 2013, 117, 22880– 22886 DOI: 10.1021/jp4073125. (22) Gankanda, A.; Grassian, V. H. Nitrate Photochemistry on Laboratory Proxies of Mineral Dust Aerosol: Wavelength Dependence and Action Spectra. J. Phys. Chem. C 2014, 118, 29117–29125 DOI: 10.1021/jp504399a. (23) Gankanda, A.; Grassian, V. H. Nitrate Photochemistry in NaY Zeolite: Product Formation and Product Stability under Different Environmental Conditions. J. Phys. Chem. A 2013, 117, 2205– 2212 DOI: 10.1021/jp312247m. 31 ACS Paragon Plus Environment


Page 32 of 35

(24) Cwiertny, D. M.; Hunter, G. J.; Pettibone, J. M.; Scherer, M. M.; Grassian, V. H. Surface Chemistry and Dissolution of alpha-FeOOH Nanorods and Microrods: Environmental Implications of Sizedependent Interactions with Oxalate. J.

Phys.

Chem. C 2009, 113, 2175– 2186.

(25) Kandler, K.; Benker, N.; Bundke, U.; Cuevas, E.; Ebert, M.; Knippertz, P.; Rodríguez, S.; Schütz, L.; Weinbruch, S. Chemical Composition and Complex Refractive Index of Saharan Mineral Dust at Izaña, Tenerife (Spain) Derived by Electron Microscopy. Atmos. Environ. 2007, 41, 8058–8074 DOI: 10.1016/j.atmosenv.2007.06.047. (26) Goodman, A. L.; Bernard, E. T.; Grassian, V. H. Spectroscopic Study of Nitric Acid and Water Adsorption on Oxide Particles: Enhanced Nitric Acid Uptake Kinetics in the Presence of Adsorbed Water. J. Phys. Chem. A 2001, 105, 6443–6457. (27) Stucki, J. W. The Quantitative Assay of Minerals for Fe2+ and Fe3+ Using 1, 10-Phenanthroline; Sources of Variability. Soil Sci. Soc. Am. J. 1981, 45, 638–641. (28) Prospero, J. M. Long-range Transport of Mineral Dust in the Global Atmosphere: Impact of African Dust on the Environment of the Southeastern United States. Proc. Natl. Acad. Sci. 1999, 96, 3396– 3403 DOI: 10.1073/pnas.96.7.3396. (29) Podczeck, F. A Shape Factor to Assess the Shape of Particles using Image Analysis. Powder Technol. 1997, 93, 47–53. (30) Schuttlefield, J.; Rubasinghege, G.; El-Maazawi, M.; Bone, J.; Grassian, V. H. Photochemistry of Adsorbed Nitrate. J. Am. Chem. Soc. 2008, 130, 12210–12211. (31) Wijenayaka, L. A.; Rubasinghege, G.; Baltrusaitis, J.; Grassian, V. H. Surface Chemistry of FeOOH Nanorods and Microrods with Gas-Phase Nitric Acid and Water Vapor: Insights into the Role of Particle Size, Surface Structure, and Surface Hydroxyl Groups in the Adsorption and Reactivity of -FeOOH with Atmospheric Gases. J. Phys. Chem. C 2012, 116, 12566–12577.


Page 33 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(32) Wierzejewska, M. FTIR Matrix Isolation Studies of Complexes of Dimethyl Sulfide, Dimethyl Disulfide and Hydrogen Sulfide with Nitric Acid. Vib. Spectrosc. 2000, 23, 253–262. (33) Key, J. M.; Paulk, N.; Johansen, A. M. Photochemistry of Iron in Simulated Crustal Aerosols with Dimethyl Sulfide Oxidation Products. Environ. Sci. Technol. 2008, 42, 133–139. (34) Andreae, M. O.; Ferek, R. J.; Bermond, F.; Byrd, K. P.; Engstrom, R. T.; Hardin, S.; Houmere, P. D.; Lemarrec, F.; Raemdonck, H.; Chatfield, R. B. Dimethyl Sulfide in the Marine Atmosphere. J. Geophys. Res. 1985, 90900 (20), 891–12. (35) National

Institute

of

Standards

&

Technology (NIST)

–

Dimethyl

Sulfide

Page.

http://webbook.nist.gov/cgi/cbook.cgi?ID=C75183&Units=SI&Mask=80#Refs (accessed on Jan 3rd 2017) (36) Wiederhold, J. G.; Kraemer, S. M.; Teutsch, N.; Borer, P. M.; Halliday, A. N.; Kretzschmar, R. Iron Isotope Fractionation during Proton-Promoted, Ligand-Controlled, and Reductive Dissolution of Goethite. Environ. Sci. Technol. 2006, 40, 3787–3793. (37) 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. Atmos. 2008, 113. (38) Shi, Z.; Bonneville, S.; Krom, M. D.; Carslaw, K. S.; Jickells, T. D.; Baker, A. R.; Benning, L. G. Iron Dissolution Kinetics of Mineral Dust at Low pH During Simulated Atmospheric Processing. Atmos. Chem. Phys. 2011, 11, 995–2011. (39) Zhang, S.; Nicol, M. J. Kinetics of the Dissolution of Ilmenite in Sulfuric Acid Solutions under Reducing

Conditions.

Hydrometallurgy

2010,

103,

196–204

DOI:

10.1016/j.hydromet.2010.03.019. (40) Van Dyk, J. P.; Vegter, N. M.; Pistorius, P. C. Kinetics of Ilmenite Dissolution in Hydrochloric Acid. Hydrometallurgy 2002, 31–36. 33 ACS Paragon Plus Environment


Page 34 of 35

(41) Silveira, J. E.; Paz, W. S.; Garcia-Muñoz, P.; Zazo, J. A.; Casas, J. A. UV-LED/ilmenite/persulfate for Azo Dye Mineralization: The Role of Sulfate in the Catalyst Deactivation. Appl. Catal. B Environ. 2017, 219, 314–321 DOI: 10.1016/j.apcatb.2017.07.072. (42) Singh, J.; Uma, S. Efficient Photocatalytic Degradation of Organic Compounds by Ilmenite AgSbO3 under Visible and UV Light Irradiation. J. Phys. Chem. C 2009, 113, 12483–12488 DOI: 10.1021/jp901729v. (43) Xiong, L. B.; Li, J. L.; Yang, B.; Yu, Y. Ti3+ in the Surface of Titanium Dioxide: Generation, Properties and Photocatalytic Application. J. Nanomater. 2012, 1–13 DOI: 10.1155/2012/831524. (44) Sasikumar, C.; Srikanth, S.; Mukhopadhyay, N. K.; Mehrotra, S. P. Energetics of Mechanical Activation

-

Application

to

Ilmenite.

Miner.

Eng.

2009,

22,

572–574

DOI:

10.1016/j.mineng.2009.01.013. (45) Chen, H.; Navea, J. G.; Young, M. A.; Grassian, V. H. Heterogeneous Photochemistry of Trace Atmospheric Gases with Components of Mineral Dust Aerosol. J. Phys. Chem. A 2011, No. 115, 490–499 DOI: 10.1021/jp110164j. (46) Wang, Z.; Fu, H.; Zhang, L.; Song, W.; Chen, J. Ligand-Promoted Photoreductive Dissolution of Goethite by Atmospheric Low-Molecular Dicarboxylates. J. Phys. Chem. A 2017, 121, 1647–1656 DOI: 10.1021/acs.jpca.6b09160. (47) Nanayakkara, C. E.; Jayaweera, P. M.; Rubasinghege, G.; Baltrusaitis, J.; Grassian, V. H. Surface Photochemistry of Adsorbed Nitrate: The Role of Adsorbed Water in the Formation of Reduced Nitrogen Species on -Fe2O3 Particle Surfaces. J. Phys. Chem. A 2014, 118, 158–166 DOI: 10.1021/jp409017m. (48) Yoshimatsu, K.; Mashiko, H.; Umezawa, N.; Horiba, K.; Kumigashira, H.; Ohtomo, A. Electronic Structures and Photoanodic Properties of Ilmenite-type M TiO3 Epitaxial Films ( M = Mn, Fe, Co, Ni). J. Phys. Chem. C 2017, 121, 18717–18724 DOI: 10.1021/acs.jpcc.7b06076. 34 ACS Paragon Plus Environment

Page 35 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(49) Nakamura, R.; Nakato, Y. Primary Intermediates of Oxygen Photoevolution Reaction on TiO2 (Rutile) Particles, Revealed by in Situ FTIR Absorption and Photoluminescence Measurements. J. Am. Chem. Soc. 2004, 126, 1290–1298. (50) Peak, D.; Ford, R. G.; Sparks, D. L. An in Situ ATR-FTIR Investigation of Sulfate Bonding Mechanisms on Goethite. J. Colloid Interface Sci. 1999, 218, 289–299. (51) Rietra, R. P. J. J.; Hiemstra, T.; Van Riemsdijk, W. H. Sulfate Adsorption on Goethite. J. Colloid Interface Sci. 1999, 218, 511–521. (52) Hug, S. J. In Situ Fourier Transform Infrared Measurements of Sulfate Adsorption on Hematite in Aqueous Solutions. J. Colloid Interface Sci. 1997, 188, 415–422.

TOC Graphic


Atmospheric Processing and Iron Mobilization of Ilmenite: Iron

Recommend Documents