Different Arsenate and Phosphate Incorporation Effects on the

Sep 18, 2014 - Heterogeneous Nucleation and Growth of Nanoparticles at Environmental Interfaces. Young-Shin Jun , Doyoon Kim , and Chelsea W. Neil...
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Different Arsenate and Phosphate Incorporation Effects on the Nucleation and Growth of Iron(III) (Hydr)oxides on Quartz Chelsea W. Neil,† Byeongdu Lee,‡ and Young-Shin Jun*,† †

Department of Energy, Environmental and Chemical Engineering, Washington University, St. Louis, Missouri 63130, United States X-ray Science Division, Argonne National Laboratory, Argonne, Illinois 60439, United States

‡

S Supporting Information *

ABSTRACT: Iron(III) (hydr)oxides play an important role in the geochemical cycling of contaminants in natural and engineered aquatic systems. The ability of iron(III) (hydr)oxides to immobilize contaminants can be related to whether the precipitates form heterogeneously (e.g., at mineral surfaces) or homogeneously in solution. Utilizing grazing incidence small-angle X-ray scattering (GISAXS), we studied heterogeneous iron(III) (hydr)oxide nucleation and growth on quartz substrates for systems containing arsenate and phosphate anions. For the iron(III) only system, the radius of gyration (Rg) of heterogeneously formed precipitates grew from 1.5 to 2.5 (±1.0) nm within 1 h. For the system containing 10−5 M arsenate, Rg grew from 3.6 to 6.1 (±0.5) nm, and for the system containing 10−5 M phosphate, Rg grew from 2.0 to 4.0 (±0.2) nm. While the systems containing these oxyanions had more growth, the system containing only iron(III) had the most nucleation events on substrates. Ex situ analyses of homogeneously and heterogeneously formed precipitates indicated that precipitates in the arsenate system had the highest water content and that oxyanions may bridge iron(III) hydroxide polymeric embryos to form a structure similar to ferric arsenate or ferric phosphate. These new findings are important because differences in nucleation and growth rates and particle sizes will impact the number of available reactive sites and the reactivity of newly formed particles toward aqueous contaminants.

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iron(III) (hydr)oxides.14−17 Arsenate is particular interesting due to its chronic toxic effects on humans. The U.S. EPA maximum contaminant level (MCL) for arsenic in drinking water is 10 μg/L.18 According to the World Health Organization, arsenic contamination of drinking water affects over 137 million people worldwide.19 This arsenic frequently comes from naturally occurring minerals, such as arsenicbearing pyrites (e.g., arsenian pyrite or arsenopyrite), in groundwater aquifers.4,20−22 Arsenic can be mobilized through oxidation of these minerals, simultaneously releasing iron (II, III) and arsenate ions and subsequently precipitating arsenicbearing iron(III) (hydr)oxides.1,4,23,24 Due to its structural similarity to arsenate, phosphate can compete with arsenate in sorption and coprecipitation with iron(III) (hydr)oxides.14,15,25,26 Furthermore, phosphate is often introduced to natural aquatic systems by human sources, such as fertilizer runoff and groundwater infiltration of sewage and industrial discharge.27−30 Decades of studies have investigated arsenate removal through sorption and coprecipitation with iron(III) (hydr)oxides,1,14,26,31,32 as well as phosphate competition over available sorption sites.17,25,26 However, only limited studies

INTRODUCTION Iron(III) (hydr)oxides play a central role in the geochemical cycling of both natural and anthropogenic aqueous contaminants.1−3 Iron(III) (hydr)oxides can form during the oxidative dissolution of Fe(II) minerals, such as pyrite and arsenopyrite, and in acidic systems, such as acid mine drainage, where the dissolution of iron minerals results in supersaturation with respect to iron(III) (hydr)oxides.4−6 Initial precipitates can be amorphous and have a high specific surface area, making them powerful sorbents for water-borne trace metal contaminants (e.g., copper(II), chromium(III), and lead(II))7,8 and organic pollutants.9,10 The ability of iron(III) (hydr)oxides to immobilize these contaminants can be related to their formation location.2 Heterogeneously formed precipitates (e.g., formed at at mineral surfaces) will be immobilized, while homogeneously formed precipitates in solution can be transported in aqueous systems. In addition, the kinetics, morphology, composition, and formation location of iron(III) (hydr)oxides on mineral surfaces are significantly affected by water chemistry such as pH, ionic strength, and the presence of different aqueous ions.11−13 Hence, to better understand immobilization mechanisms of toxic metal and organic contaminants in aquatic environments, we have focused in this study on the heterogeneous nucleation and growth of iron(III) (hydr)oxides. Among environmentally important anions, arsenate and phosphate are of interest owing to their strong interactions with © XXXX American Chemical Society

Received: December 27, 2013 Revised: September 12, 2014 Accepted: September 18, 2014

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Table 1. Conditions for Iron(III) (Hydr)oxide Precipitation Experimentsa System

NaNO3 (mM)

Fe(NO3)3 (mM)

H2AsO4− (mM)

H2PO4− (mM)

ISb (mM)

pHc

SId Fe(OH)3

ζe (mV)

ζ*,f (mV)

1 2 3

10 10 10

0.1 0.1 0.1

0 0.01 0

0 0 0.01

10.19 10.26 10.18

3.6 ± 0.2 3.6 ± 0.2 3.6 ± 0.2

0.31 0.35 0.33

39.9 ± 1.9 25.9 ± 2.3 30.1 ± 4.1

51.3 ± 2.1 44.2 ± 2.4 39.3 ± 1.9

a

Comparison of measured pHs, zeta potentials with or without quartz powder, and saturation indices (SI = log(Q/K)) calculated using Geochemist’s Workbench’s thermo.dat database file (GWB, release 8.0, RockWare, Inc.). bIS = Ionic strength, calculated using GWB. cpH values measured after solution mixing. They are consistent with GWB calculations for the low concentrations of arsenate/phosphate. dSI = Saturation Index = log(IAP/ Ksp), with respect to ferrihydrite (simplified as Fe(OH)3) at 20 °C calculated with GWB using thermo.dat database. IAP: Ion activity product and Ksp: solubility product. eZeta potential (ζ) of homogeneously formed precipitates measured without quartz powder. Measurements taken every 1 min until values stabilized (20 min to 1 h). fZeta potential measurements of heterogeneously formed precipitates on suspended quartz powder (ζ*).

systems (Table 1). Because the first pKa’s for arsenic and phosphoric acid are 2.3 and 2.2, respectively, both oxyanions will be doubly protonated at the system pH.40,41 Based on thermodynamic calculations, the speciation of arsenate was calculated to be 96.5% H2AsO4− and 3.4% H3AsO4, and the speciation for phosphate was calculated to be 97.3% H2PO4− and 2.7% H3PO4. Reported saturation indices (SIs) and aqueous speciation percentages were calculated using Geochemist’s Workbench software (GWB, Release 8.0, RockWare, Inc., Urbana, IL) using the thermo.dat database file. Because sodium abounds in natural aqueous systems and nitrate is not expected to interact with iron(III) (hydr)oxides, 10 mM sodium nitrate provided the background ionic strength. Reaction systems are outlined in Table 1. In Situ GISAXS Measurements. To conduct GISAXS, a clean quartz substrate was first placed in a cleaned, specially designed GISAXS fluid cell. Ultrapure water (resistivity >18.2 MΩ·cm) was injected and the surface was aligned with the Xray beam. The water was then removed, and the reaction solution (Table 1) was introduced. Then, 1 mL of the solution was immediately injected into the cell and in situ GISAXS measurements began. Approximately 2 min elapsed between creating the solution, which started the precipitation reaction, and the first GISAXS data recording. The reaction time was defined as starting when the solution was created. During the in situ reaction period of 1 h, GISAXS measurements were taken at 1 min intervals. An incidence angle (αi) of 0.11° was chosen to obtain 98% reflectivity at the beam energy of 14 keV. The beam size was 40 μm (vertical) × 400 μm (horizontal). Supplementary analysis using AFM showed insignificant homogeneous particle settling during the 1 h reaction period (Supporting Information, Figure S2). The scattering vector range (q range) was 0.007−0.300 Å−1. Particle sizing for heterogeneously formed particles was carried out by fitting the shape of the 1D scattering intensities over the analyzed q range. More details on the GISAXS experimental setup can be found in the Supporting Information. Experiments were conducted at the Advanced Photon Source (APS), beamline 12-ID-B, at Argonne National Laboratory (ANL) (Argonne, IL). GISAXS Scattering Data Analysis. GISAXS measurements produced time-resolved 2D scattering images of nanoparticles at the mineral surface. The first scattering image was used as a background and subtracted from subsequent images. The selected time series for the 2D scattering data can be found in Figure S5 in the Supporting Information. The 2D images were reduced to 1D by cutting along the Yoneda wing, where the scattering is enhanced by the grazing incidence effect (Vineyard effect).33 1D intensities (I) were plotted versus q for different time points to show the evolution in scattering intensities due to iron(III) (hydr)oxide

systematically investigate in situ the effects of these oxyanions on iron(III) (hydr)oxide precipitation, particularly during the early stages of nucleation and growth. We have previously used a unique fluid cell setup to observe in situ the precipitation of iron(III) (hydr)oxide in the presence of environmentally abundant anions (nitrate, chloride, and sulfate) on quartz, using grazing incidence small-angle X-ray scattering (GISAXS). We found that chloride inhibited nucleation, while sulfate greatly increased precipitate growth.2,33 Based on our previous research,2,33−36 the main objective of our current study was to investigate the in situ nucleation and growth of iron(III) (hydr)oxides in the presence of arsenate or phosphate oxyanions, utilizing our established GISAXS approach. By adding arsenate and phosphate anions to the iron(III) (hydr)oxide system, we established a unique scenario with important implications for natural and engineered aquatic systems. Nucleation and growth kinetics of iron(III) (hydr)oxide were investigated on quartz, an environmentally abundant substrate in natural aquatic systems. The mechanisms behind iron(III) (hydr)oxide precipitation in these systems were then explored by complementary ex situ approaches, including atomic force microscopy (AFM), high resolution X-ray diffraction (HRXRD), and thermal gravimetric analysis (TGA), to explain observed differences in precipitation kinetics and precipitate morphology (size, volume, and, therefore, surface area evolution). The new information gained in this study can be useful for modeling the reactive transport of arsenate and phosphate in natural and engineered aqueous systems.

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EXPERIMENTAL SECTION Substrate and Solution Preparation. Substrates used for heterogeneous precipitation studies were high quality single crystal quartz wafers with a surface roughness of less than 5 Å (MTI Corporation, Richmond, CA). Quartz was x-cut, exposing the (110) surface to contact with the aqueous solution. Quartz lacks a distinctive cleavage plane, thus the abundance of the (110) surface in natural systems would be approximately equal to the abundance of any other surface with similar surface energies (e.g., 1̅01 and 101̅ surfaces).37−39 Quartz wafers were cut to 1 cm squares to fit in GISAXS reaction cells, then cleaned with Nochromix and sulfuric acid as outlined in the (Supporting Information, S1). Reaction solutions were created using ultrapure water and reagent grade Fe(NO3)3·9H2O, NaNO3, Na2HAsO4·7H2O, and Na2HPO4·7H2O. Fe(III) concentrations were kept at 10−4 M for all systems, while concentrations of 10−5 M arsenate and phosphate were used for in situ GISAXS and ex situ experiments. The pH of these systems was 3.6 ± 0.2, and the saturation index (SI) did not vary significantly between the B

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Figure 1. GISAXS 1D scattering intensity for iron(III) hydroxide precipitation on quartz in the presence of (A) 10 mM sodium nitrate only, (B) 10 mM sodium nitrate with 10−5 M arsenate, and (C) 10 mM sodium nitrate with 10−5 M phosphate. Scattering curves were produced by cutting along the Yoneda wing. Experiments were conducted for 1 h. No significant water evaporation occurred during this experimental period.

Atomic Force Microscopy (AFM) Measurement. Ex situ AFM measurements of GISAXS samples were conducted to complement in situ GISAXS measurements of heterogeneously formed particle sizes. This approach provided consistent evidence of newly formed particle size trends. No conclusions on particle sizes or trends are drawn from the ex situ AFM data only. Tapping mode AFM (AFM, Veeco Inc.) was used to measure the height, amplitude, and phase of precipitates on reacted substrates. AFM tapping mode probes were 125 μm long, with phosphorus (n) doped silicon tips (nominal tip radius of 10 nm, MPP-11100-10, Bruker). Imaging used a scanning rate of 0.988 Hz and drive frequencies between 312 and 320 kHz. Images were processed using Nanoscope 7.20. Quantifying Arsenate and Phosphate Incorporation in Iron(III) (Hydr)oxides. Arsenate and phosphate incorporation into heterogeneously and homogeneously precipitated iron(III) (hydr)oxides was quantified. For homogeneously formed precipitates, a large batch of the reaction solution (Table 1) was created and precipitates were concentrated at 5000 rpm using Millipore Amicon ultra-15 centrifugal filter units (Millipore Corporation, Billerica, MA, U.S.A.). The nominal molecular weight limit (NMWL) of this membrane was 100,000. The particle-rich solution was acidified with 2% nitric acid to dissolve iron(III) (hydr)oxides. The pH of the 2% nitric acid was 100 precipitates at different locations on the substrate surfaces.

iron(III) hydroxide monomers and polymeric embryos due to lower electrostatic driving forces, leading to slower growth. Particle numbers were also calculated (N = V/Rg3) for the three systems (Figure 2C). While the systems containing arsenate and phosphate oxyanions had approximately 2 and 1.5 times more growth (based on Rg), respectively, compared to the iron(III) only system, the system containing iron(III) only had approximately 7 times more nucleation than the arsenate system and 4 times more nucleation than the phosphate system (based on particle number). Ex Situ Zeta Potential, Composition, and Phase Identification of Iron(III) (Hydr)oxides. To explain the observed trends in heterogeneous precipitation rates for our three experimental systems, we used ex situ approaches to determine the composition, zeta potential, and phase of both E

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which may influence the size and growth of these heterogeneous precipitates. One such factor was the oxyanion content of iron(III) (hydr)oxide precipitates. For the 10−4 M Fe(III)−10−5 M arsenate system, based on ICP-MS results, the oxyanion content was 8.1 ± 2.3 mol % for homogeneously formed particles and 6.5 ± 1.3 mol % for heterogeneously formed particles. For the 10−4 M Fe(III)−10−5 M phosphate systems, the oxyanion content was 13.1 ± 1.8 mol % for homogeneously formed particles and 12.2 ± 1.1 mol % for heterogeneously formed particles. The close agreement between precipitate compositions indicates that results from analyses carried out on homogeneously formed precipitates, such as phase identification, likely hold for heterogeneously formed precipitates as well. Next, samples of homogeneously formed precipitates in the reactions systems in Table 1 were analyzed using HRXRD for ex situ phase identification (Figure 4B). In the 10−4 M Fe only system, the observed peaks were characteristic of 6-line ferrihydrite. For systems containing arsenate and phosphate oxyanions, the characteristic peaks were much less defined, indicating that the incorporation of oxyanions during iron(III) (hydr)oxide nucleation and growth resulted in more amorphous precipitates for these systems. Furthermore, although the peak locations in the spectra resemble ferric arsenate and ferric phosphate, they are off by a few degrees from the ferric arsenate45,31 and ferric phosphate46 peak centers. This indicates that the phases can be amorphous or poorly crystalline ferric arsenate and phosphate, or at least amorphous or poorly crystalline iron(III) (hydr)oxide containing a significant portion of ferric arsenate/phosphate type bonds. Note that the phases identified in HRXRD analysis can present as more crystallized than the in situ newly formed phases. Ferric phosphate and ferric arsenate contain cornerlinked FeO6 tetrahedra bridged by PO4 or AsO4 tetrahedra, respectively.47,48 This bridging by arsenate and phosphate anions may account for the larger observed sizes of iron(III) (hydr)oxides. The observed decreased crystallinity is also consistent with previous reports published on the effects of arsenate on the aging of iron(III) (hydr)oxide precipitates. Waychunas et al.31 found that the incorporation of high quantities of arsenate into iron(III) hydroxide precipitates slowed the transformation of ferrihydrite into hematite by preventing FeOFe polymerization. Pedersen et al.32 found that trace amounts of arsenate (up to 0.5 mol % As) have no effect on crystallization rates. It has also been found that for lower arsenate loadings (8416 mg As/kg iron(III) (hydr)oxide ≈ 1.2 mol % As), the crystallization of iron(III) (hydr)oxides can lead to stable, irreversible arsenic attenuation.49 ICP-MS showed that heterogeneously formed precipitates in the phosphate system had a higher oxyanion content, suggesting that more bridging occurred in the phosphate system than the arsenate system. However, nanoparticles were smaller in the phosphate system. Thus, we hypothesized that differences in the hydrated radii of incorporated oxyanions and water content of precipitates could also contribute to larger precipitates for the arsenate system, as outlined in the following section. Water Content of Iron(III) (Hydr)oxides. Although phosphate was incorporated in larger quantities than arsenate, the size of the incorporated oxyanions must also be considered to explain why larger primary particle sizes were observed for the arsenate system. The ionic radius of arsenate as H2AsO4− is

Figure 4. Results from (A) TGA and (B) HRXRD, showing the effects of oxyanions on iron(III) (hydr)oxides. TGA results show that the low arsenate system contained the most water. Broadening of HRXRD peaks in the systems containing oxyanions indicates that the incorporation of arsenate and phosphate resulted in decreased precipitate crystallinity.

2.20 Å, slightly larger than that of phosphate, H2PO4−, which is 2.03 Å.50 This difference may contribute slightly to the larger particle sizes in the case of arsenate or phosphate that forms direct inner-sphere bonds with Fe-octahedra. In addition, outersphere complexes may form concurrently between the arsenate or phosphate and the iron(III) (hydr)oxide surface.51 If so, then the oxyanions will remain solvated by water in aqueous environments. The hydrated ionic radius is then used to describe the radius of both the ion in solution and tightly bound water.52 The hydrated ionic radii of H2AsO4− and F

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H2PO4− are 5.9 Å53 and 3.02 Å,54 respectively. If they remain inside the particles at the early nucleation stage, then these outer-sphere complexes may provide increased water content for precipitates in the oxyanion-containing systems. Moreover, the incorporation of arsenate and phosphate anions in the iron(III) (hydr)oxide precipitates will hinder crystallization, as evidenced by HRXRD results, which can also lead to higher water content.55 Although the volume of water in these precipitates will not be included in calculated invariant values, incorporated water will impact Rg values. Therefore, the significant difference in solvated oxyanions radius, water inclusion, and lesser extent of phase crystallization could contribute to larger particle sizes for the arsenate system, despite less arsenate incorporation than in the phosphate system. Recent literature has reported that smaller, nanostable iron oxyhydroxide phases, such as ferrihydrite, can contain excess water.56 To test this hypothesis for our system, we investigated water incorporation into precipitates using TGA on homogeneously formed particles. For all homogeneous nanoparticle slurries, there was significant mass loss between 107 and 440 °C (Figure 4A). The Fe(III) only system contained 14.9% water by mass, while the phosphate and arsenate systems contained 17.1% and 21.6%, respectively. Results for heterogeneous precipitates were corroborated by investigation of water content using AFM on freshly precipitated and oven-dried samples (Section S2 and Figure S3 in the Supporting Information). Together, these investigations indicate that the poorly crystalline iron(III) (hydr)oxide precipitates that formed in the presence of oxyanions contained more water, and that this effect was strongest for the arsenate system. The greater degree of hydration for oxyanion-containing nanoparticles can in turn help to explain the observed stability of larger and less crystalline precipitates in arsenate and phosphate systems, because more hydrated particles will have a lower surface energy and higher thermodynamic stability.57

provided by this study can be used to develop more rigorous reactive transport models of contaminant fate and transport in natural and engineered aquatic systems, such as arseniccontaminated groundwater aquifers and acid mine drainage sites. Moreover, many additional factors can influence the nucleation and growth rates of iron(III) (hydr)oxides in the presence of arsenate and phosphate anions, and these must be considered further. One factor is the oxyanion-induced change in interfacial energy of iron(III) (hydr)oxides, which can affect their nucleation and growth by influencing the balance between substrate−solution, precipitate−solution, and precipitate−substrate interfacial energies. Second, sorption of arsenate and phosphate oxyanions can also affect the formation of iron(III) (hydr)oxides by passivation of the surface or by catalytic promotion of nucleation on the iron(III) (hydr)oxide surface. Another important future consideration is the presence of arsenite (As(III)) and Fe(II) in addition to Fe(III) and arsenate (As(V)). In natural surface and groundwater, both oxidation states can be present simultaneously and have significant interactions, including redox transformations65 and electron/atom exchange.66 We also need to better understand mineral reactivity changes in order to accurately predict the fate and transport of these nanoparticles and associated contaminants. For example, the incorporation of contaminants, such as arsenic, into iron(III) (hydr)oxide can greatly impact nanoparticle reaction rates and pathways.64 Incorporation of these impurities adds defects to the nanoparticles which can impact their photochemical behavior by changing their band gap structures, as well as affecting the phase stability of precipitates. Continuing to advance our knowledge of this system will allow us to better model contaminant interactions with iron(III) (hydr)oxides, as well as improve analytical techniques to observe nanoscale interfacial reactions in environmentally relevant systems.

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ENVIRONMENTAL IMPLICATIONS Previously, much of the research related to arsenate and phosphate interactions has focused on oxyanion adsorption onto preformed or more crystallized iron(III) (hydr)oxides as a means of contaminant remediation. Numerous studies have shown the formation of inner-sphere bidentate bridging structures between arsenate or phosphate tetrahedra and iron octahedra on the iron(III) (hydr)oxide surfaces.25,58−63 There have also been extensive studies on the competitive effects of phosphate and arsenate on sorption by iron(III) (hydr)oxides. However, no previous studies have accomplished in situ, timeresolved observation of iron(III) (hydr)oxide nucleation and growth on quartz substrates in the presence of these oxyanions. Our new results indicate that arsenate and phosphate can significantly impact the nucleation and growth of iron(III) (hydr)oxides and the extent of water molecules in their structure. In particular, arsenate presence will promote growth. The sizes of these particles can greatly impact their electronic structure. When particles are small enough, the band gap increases compared to larger particles, affecting their redox potential and allowing these nanoparticles to act as semiconductors in environmental systems.64 The larger band gap facilitates redox reactions or photoredox reactions which would not be possible for bulk minerals. Furthermore, new quantitative information such as the heterogeneous particle sizes and arsenic and phosphate incorporation percentages

ASSOCIATED CONTENT

S Supporting Information *

The experimental setup, GISAXS fitting procedures, and additional testing. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We are grateful for support received from Washington University’s Faculty Start-up Grant and National Science Foundation (EAR-1424927). CWN acknowledges the generous support of the Mr. and Mrs. Spencer T. Olin Fellowship. We wish to thank the Environmental NanoChemistry Group members, Dr. Soenke Seifert at APS, and Dr. Christopher Kim for valuable discussion. Use of the Advanced Photon Source (Sector 11-BM for HRXRD and Sector 12 ID-B for GISAXS) at Argonne National Laboratory was supported by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences, under Contract No. DE-AC0206CH11357. G

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