ATR-FTIR and Flow Microcalorimetry Studies on the Initial Binding

Article pubs.acs.org/JPCA

ATR-FTIR and Flow Microcalorimetry Studies on the Initial Binding Kinetics of Arsenicals at the Organic−Hematite Interface Arthur Situm,† Mohammad A. Rahman,† Nicholas Allen,‡,§ Nadine Kabengi,‡,§ and Hind A. Al-Abadleh*,† †

Department of Chemistry and Biochemistry, Wilfrid Laurier University, Waterloo, ON N2L 3C5, Canada Department of Geoscience and §Department of Chemistry, Georgia State University, Atlanta, Georgia 30303, United States

‡

S Supporting Information *

ABSTRACT: The environmental fate of arsenic compounds depends on their surface interactions with geosorbents that include clays, metal oxides, and natural organic matter (NOM). While a number of batch studies reported that NOM inhibits the uptake of arsenicals, it remains unclear how different classes of organic functional groups affect their binding mechanisms. We report herein the adsorption kinetics of arsenate and dimethylarsinic acid (DMA) with hematite nanoparticles pre-exposed to three types of low molecular weight organics: citrate, oxalate, and pyrocatechol as representatives to the majority of reactive organic functional groups in NOM. These studies were conducted using attenuated total internal reflection Fourier transform infrared spectroscopy (ATR-FTIR) and flow microcalorimetry at pH 7 with an emphasis on the role that electrolytes (KCl, NaCl, and KBr) play in the adsorption process. Results show that (1) negatively charged carboxylate versus hydrophobic phenyl groups influence amounts and initial rates of arsenicals adsorption on hematite nanoparticles to varying degrees depending on the type of complexes they form, (2) the type of electrolytes affects initial adsorption rate of DMA to a greater extent than arsenate when oxalate is present on the surface, and (3) the extent of organics retention by hematite nanoparticles is influenced by the type of the desorbing agent.

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ternary As−Fe−OM complexes/colloids.16,20,22−26 Kinetic batch measurements showed a reduction in the adsorption rates for As(V) and As(III) on hematite-,16 goethite-,21 and ferrihydrite-NOM27 systems accompanied by displacement of sorbed NOM in significant quantities. Direct evidence for the formation of ternary As−Fe−OM complexes/colloids was provided using attenuated total internal reflectance Fourier transform infrared spectroscopy (ATR-FTIR),16,24 Mössbauer spectroscopy,24,28 and extended X-ray absorption fine structure spectroscopy (EXAFS).29 Earlier batch experiments by Ritter et al.,22 Bauer and Blodau30 showed linear associations between the amount of complexed arsenate and the Fe(III) content of the NOM in the dissolved and colloidal phases. The stability of these complexes were found to depend on the pH, relative size of NOM, and ratio of Fe to OM.26,31 Silva et al.29 probed the type of functional groups responsible for iron binding in NOM and found that esterification of carboxyl groups reduced iron loading by 84%. In addition to the above spectroscopy studies, calorimetric measurements of adsorption to metal oxides proved to be powerful in elucidating binding mechanisms. When available,

INTRODUCTION Natural organic matter (NOM) is an important component of soil and plays a role in many types of chemical and biological reactions. NOM is secreted by plant roots and microbes1−6 or introduced by the decaying biological matter and can serve as agricultural supplements.4,6 It enhances the bioavailability of nutrients, such as iron and phosphorus, acts as food source for microscopic organisms, and aids in water retention, further enhancing soil fertility.6 Once introduced, NOM retention in soils is heavily correlated with iron and aluminum oxide content.2,5 As a result, they change the surface properties of the metal oxide/water interface such as surface charge and functional groups, hydrophobicity, interfacial water structure, adhesion strength, and dissolution rate of the metal oxide.5,7−14 Such molecular-level changes have direct consequences on the mobility and bioavailability of pollutants such as arsenic (As) compounds emitted from natural and anthropogenic sources. NOM also influences the efficiency of iron-based materials used in As removal technologies.15 NOM influences As mobility via three main mechanisms that include reductive dissolution of iron minerals, or direct redox reactions leading to cycling between arsenite (As(III)) and arsenate (As(V)),16,17 competition for binding sites at metal oxide/mineral surfaces,18−21 and formation of colloidal or dissolved metal bridging complexes with As, referred to as © 2017 American Chemical Society

Received: April 11, 2017 Revised: July 7, 2017 Published: July 10, 2017 5569

DOI: 10.1021/acs.jpca.7b03426 J. Phys. Chem. A 2017, 121, 5569−5579

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The Journal of Physical Chemistry A

electrolyte type on the kinetics. The term “no added electrolyte” is used herein to refer to experiments conducted using Millipore water only, where the source of cations and anions would be mostly from HCl and NaOH used to adjust the pH. The concentration of these ions is at least 10× lower than that obtained by adding KCl, NaCl, or KBr to get 10 mM concentration. Powder of organic compounds used herein (as received without further purification) include oxalic acid dihydrate (C2H2O4·2H2O, 99%+, ACS reagent, Sigma-Aldrich), citric acid monohydrate (C6H8O7·H2O, 99%+, ACS reagent, Alfa Aesar), catechol (99%+, Sigma-Aldrich), sodium hydrogen arsenate heptahydrate (Na2HAsO4·7H2O, ACS reagent, J. T. Baker), and DMA (sodium cacodylate trihydrate, C2H6AsO2Na·3H2O, 98%, Sigma-Aldrich). Caution: The aforementioned arsenical compounds are highly toxic via inhalation and skin contact and are carcinogens. While As(III) may be more toxic, it is also more mobile and hence its adsorption onto metal oxide surfaces is limited.38 Table S1 shows the structure and pKa values of these compounds. The pKa values are key to understanding the type of surface complexes that these compounds form, and they aid in interpreting spectral data and customizing surface complexation models for extracting binding thermodynamics.36 All the solutions were prepared freshly before the start of each experiment and were covered with parafilm until used. The experiments were conducted at concentrations of organics and arsenicals (listed in the Supporting Information) that are 8−10× lower than the detection limit of the ZnSe flow cell for aqueous phase concentrations. Hence, the recorded spectra only show features due to surface species. Hematite nanoparticles (α-Fe2O3, 98+%, US Research Nanomaterials, Inc., batch US3180) were characterized for these experiments to obtain BET surface area (N2 gas, 54 ± 5 m2/g, Aveka Inc.), isoelectric point (IEP) (zeta potential titration, 8.45, Aveka Inc.), particle shape and size (TEM Philips C12 at 120 keV, spherical, 39 ± 11 nm average diameter, Canadian Centre for Electron Microscopy, McMaster University). The thickness of the hematite film deposited on the ZnSe ATR crystal was determined to be 7.5 ± 1.4 μm (see reference36 for details). The film was stable under the 80 min aqueous flow conditions described in the following section as judged by the intensity of the single beam spectra below 1000 cm−1 where lattice Fe−O vibrations exist. Additional details on film preparation are in the Supporting Information. ATR-FTIR Spectroscopy Experiments. ATR-FTIR spectra were collected as a function of time on a freshly prepared hematite film using a HATRPlus accessory (Pike Technologies) installed in a Nicolet 8700 FTIR spectrometer (Thermo Instruments) equipped with Mercury−Cadmium-Telluride (MCT-A) detector. The ATR flow cell used in the experiments contains a 60° ZnSe crystal IRE (80 × 10 × 4 mm, 100 μL). Aqueous solutions were flowed at a rate of 2 mL min−1 across the ATR flow cell using Tygon tubes (0.8 mm i.d., Masterflex) and a compact pump (Masterflex L/S). The combination of this flow rate and the mass used in our experiments (6 mg) were shown previously to minimize diffusion contribution to the initial adsorption rate.39 Faster flow rate results in detaching the hematite film before the end of the experiment. Singlebeam ATR-FTIR spectra were collected at 8 cm−1 resolution throughout the experiments. More details are provided in the Supporting Information. Flow Microcalorimetric Measurements. These measurements were conducted using oxalate and pyrocatechol followed

these thermodynamic investigations are obtained through either Calvet-type differential microcalorimetry (CDM) or isothermal titration calorimetry (ITC). For example, Gu et al.32 used CDM to study the adsorption of NOM on hematite and reported exothermic heats ca. 125−320 kJ per mole of carboxyl/hydroxyl functional groups. Lin et al.33 used ITC and reported a strongly exothermic interaction (ΔH = −16.4 ± 0.9 kJ mol−1) for the binding of catechol to zinc oxide quantum dots. Similarly, Benoit et al.34 reported the adsorption of aromatic acids, monochlorophenols and aliphatic acids on aluminum oxides by titration calorimetry, and reported exothermic adsorption enthalpies for all organic ligands studied. The temperature dependence of the adsorption of the two organics 1,10-phenanthroline and 2,2′-bipyridyl on hematite was studied by Ferreiro and de Bussetti,35 who reported the process to be exothermic. Careful examination of the aforementioned studies highlights several areas that warrant further investigation at the molecular level pertaining to the mechanism of arsenicals binding to NOM complexed to iron (oxyhydr)oxides. The relatively long time scales in hours employed in the above measurements complicates the interpretation given that a number of processes could be at play: adsorption/desorption, redox, and precipitation. Also, given the structural complexity of NOM, it is impossible to establish structure−reactivity relationships that will shed light into how different functional groups react with As compounds. Moreover, few studies investigated the binding of organoarsenicals to NOM compared to metal(oxyhydr)oxides. Methylated and aromatic organoarsenicals have As−O bonds, substituted methyl and phenyl functional groups, and hence can interact with NOM through additional pathways compared to arsenate. The main objective of our studies herein is to establish structure−reactivity relationships from the initial adsorption kinetics and heat release upon arsenicals binding to organics on the surface of hematite nanoparticles. These measurements are obtained using ATR-FTIR spectroscopy and flow microcalorimetry because these techniques allow for real time and in situ data collection in the flow mode. This is in contrast to the ex situ techniques mentioned above that used wet paste samples for analysis. Also, achieving this objective demands using model compounds representative of the majority of reactive functional groups in NOM that will modify the charge and hydrophobicity of the hematite/water interface, namely, oxalate, citrate, and pyrocatechol. The results presented herein built on our earlier work to understand the type and thermodynamics of surface complexes that arsenicals and the above low molecular weight organics form on hematite nanoparticles.36,37 Also, in order to investigate the effect of organic substituents on arsenate binding to the organichematite/water interface, experiments were conducted with dimethylarsinic acid (DMA).

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EXPERIMENTAL SECTION Chemicals. Solutions were made by dissolving the desired organic compound in 0.01 M KCl (ACS grade, 99.0%+, EMD) solution prepared fresh in 18 MΩ-cm Millipore water and adjusted to pH 7 by using dilute and concentrated HCl (6 N, Ricca Chemical Company) and NaOH (GR ACS, 99.0−100%, EMD) solutions. The concentration of KCl varied for ionicstrength-dependent studies. Experiments were also conducted using NaCl (99%, ACS reagent, BDH) and KBr (anhydrous, 99%+, ACS reagent, Sigma-Aldrich) to probe the effect of 5570


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Figure 1. Representative ATR-FTIR spectra collected as a function of time (1, 2, 5, 10, 20, 30 min) upon flowing 1 mM arsenate (a−c) or 1 mM DMA in (d−f) on the hematite nanoparticles with adsorbed citrate (CA), oxalate (OA), and pyrocatechol (PC), respectively. Solutions were prepared in 10 mM KCl and adjusted to pH 7. The bar to the left represents the absorbance scale for each set of spectra. Spectra were obtained by referencing to the spectrum collected after 80 min of flowing the respective organic compound (see example in Figure S1).

Figure 2. Kinetic curves for the desorption of organics by either 1 mM arsenate (iAs) (a−c) or DMA (d−f) at pH 7 and 10 mM KCl. The concurrent adsorption kinetic curves for the arsenicals are also shown. Error bars represent standard deviation from averaging 3−4 experiments, each on a freshly prepared film.

at pH 7 (5.2(2) × 1012 and 4.1(2) × 1012 molecules cm−2 for arsenate and DMA, respectively). These values were calculated using equations and parameters in the Supporting Information of reference.37 Spectra in Figure 1 were obtained by referencing to the spectrum collected after 80 min flow of 1 mM solution of the respective organic solution at pH 7. Figure S1 shows the reference spectra of adsorbed citrate, oxalate, and pyrocatechol prior to flowing the arsenicals. The spectral range between 1000 and 750 cm−1 contains positive absorbance features assigned to the stretching vibrations of As−O bonds, v(As−O), in adsorbed arsenate and DMA. In Sabur et al.,37 the adsorption of arsenate and DMA on clean hematite was studied in detail and is used to guide the interpretation of spectra reported herein. Spectral features shown in Figure 1 are attributed to the formation of mostly inner-sphere bidentate binuclear arsenate complexes and a mix of monodentate and outer-sphere DMA complexes. The data in Figure 1 show the frequency of the most intense feature of adsorbed DMA at 830 cm−1 instead of 837 cm−1 reported earlier on clean hematite.41 This difference

by arsenate adsorption. The surface charge was also probed before and after exposure of hematite nanoparticles to the organics. The flow adsorption microcalorimetry (FAMC) setup used in this study was custom-designed and constructed inhouse. A description of the instrumentation and operation of the FAMC setup has been detailed previously.37,40 A brief description outlining the basic principles of obtaining data and experimental procedures relevant to this study is provided in the Supporting Information.

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RESULTS AND DISCUSSION Adsorption Kinetics of Arsenicals on Organics− Hematite Nanoparticles Interface. Figure 1 shows initial kinetic ATR-FTIR spectra upon flowing 1 mM arsenate or DMA solution on hematite nanoparticles with adsorbed citrate, oxalate, and pyrocatechol at pH 7 and 10 mM KCl. This solution concentration was shown to result in about 90% of the saturation surface coverage on “clean” hematite (i.e., no adsorbed organics) from adsorption isotherm measurements 5571


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outer-sphere compelexes and pyrocatechol forms a mix of protonated mono- and bidentate binuclear complexes.36 In last section, additional evidence is provided from flow microcalorimetry studies in support of this interpretation. Figures S2 and S3 show the effect of electrolyte type on the adsorption kinetics of arsenicals on hematite nanoparticles with and without adsorbed organics at pH 7. The slope of the leastsquares linear fits in Figures S2 and S3 are listed in Table S3 and represents kads,i for arsenate and DMA adsorption kinetics on different surfaces. In these figures, robs were obtained from data collected during initial times of surface interactions (0−2 min). Values of robs from longer times (2−10 min, empty markers, Figure S2a,c) are shown for comparison. It is clear that values from initial times are more sensitive to concentrationdependence than longer times, which are usually analyzed in kinetic bulk batch measurements. For concentration-dependent arsenate and DMA adsorption kinetics, hematite surfaces preexposed to oxalate were only studied using 10 mM NaCl and KBr (Figure S3a,b) and compared to similar experiments on clean surfaces, no oxalate. This is because at pH 7, oxalate introduces carboxylate groups and forms mostly double negatively charged outer-sphere complexes. These complexes are more sensitive to changes in ionic strength and electrolyte type than inner-sphere complexes. In addition, concentrationdependent arsenate and DMA adsorption kinetic experiments were conducted on hematite surfaces pre-exposed to pyrocatechol at pH 7 and 10 mM KCl for comparison with rates from clean surfaces. Because pyrocatechol introduces uncharged inner-sphere complexes to the surface that are insensitive to the concentration and type of the electrolyte,36 using 10 mM NaCl or KBr for arsenate and DMA adsorption kinetic experiments were inferred to provide no new physical insight relative to KCl. As for hematite surfaces pre-exposed to citrate, adsorption kinetic experiments of arsenate and DMA conducted at pH 7 and 10 mM KCl resulted in robs,i values at 0.19(0.01) and 0.17(0.01) min−1, respectively. When these values are compared to the data in Figure S2a,c, they fall within the uncertainty of initial adsorption rates on clean surfaces. Hence, we inferred that the presence of surface citrate does not influence arsenicals binding kinetics to the same degree as oxalate and pyrocatechol, and was not studied any further. The absolute values of kads,i are about 10× higher for arsenate than DMA irrespective of the surface and type of electrolyte, which is in line with the spectral interpretation that arsenate forms mostly inner-sphere bidentate binuclear complexes and DMA forms mostly outer-sphere complexes. This is further supported by the lower values of the y-intercept (i.e., kdes) for arsenate than DMA. The reasoning behind the link between kads and the type of surface complexes stems from the relationship between kinetics and thermodynamics of bond formation for each type. As we reported earlier from thermodynamic measurements,37 the binding affinity constant (Keq) from the triple layer surface complexation model for arsenate on hematite nanoparticles is nearly 6 orders of magnitude higher than that of DMA at 25 °C (logKeq = 12.93 ± 0.06 compared to 7.30 ± 0.04, respectively). These numbers reflect the higher binding strength of covalently bonded innersphere complexes of arsenate relative to the H-bonded outersphere complexes of DMA. Such differences in binding strength are reflected in the As−Fe interatomic distances and adsorption Gibbs free energy from DFT calculations on two iron clusters.45 For arsenate bidentate binuclear complexes, the values are 3.30 and 3.26 Å, and ΔG°ads = −75.4 kJ mol−1,

is not attributed as a shift since the resolution of the FTIR instrument is set at 8 cm−1 for both cases. The former value is close to that observed for deprotonated aqueous phase DMA under basic conditions even in D 2O (see Supporting Information in reference 42). This clearly suggests that the adsorption of DMA on hematite nanoparticles with adsorbed organics favors the formation of outer-sphere complexes. The adsorption reaction continued for 80 min, and as detailed below, baseline-corrected ATR absorbances [A(ṽ)] at 875 and 830 cm−1 shown in Figure 1 were used to generate arsenate and DMA adsorption kinetic curves (right axes, Figure 2), respectively. Calculations of the surface coverage of arsenicals reached are shown in Table S2 in the Supporting Information relative to the clean surface. To confirm reaching a plateau in Figure 2 for arsenate and DMA, experiments at longer time frames were conducted and show very little change in absorbance values beyond 80 min. On average and within the uncertainty of the measurements, the numbers show that the presence of adsorbed citrate and pyrocatechol reduces amount of surface arsenate by 17 ± 4 and 9 ± 2% relative to the clean surface, whereas surface oxalate does not significantly change the amount of surface arsenate. On the other hand, the presence of surface organics (citrate, oxalate, and pyrocatechol) resulted in an overall increase in the amount of surface DMA by 14 ± 3, 159 ± 37, and 41 ± 7%, respectivley. Similar findings were observed using 10 mM NaCl and KBr, where the surface coverage of DMA in the presence of surface oxalate is found to be higher by 14 ± 3 and 40 ± 7%, respectively. These trends are explained in the next section in light of the role that cations play in attracting negatively charged arsenicals to organics on the surface of hematite nanoparticles. The spectral range between 1800 and 1000 cm−1 contains negative absorbance features assigned to the symmetric and asymmetric stretching vibration of −CO2 functional groups, v(CO2), and v(CO2Fe) in citrate and oxalate, and ring vibrations in pyrocactechol. These modes are coupled with bending vibrations of uncomplexed −CO2H and −CH2 functional groups, δ(CO2H), δ(CH2), and δ(OH). We refer the reader to reference 36 for detailed interpretation of these features in relation to the type of surface complexes that these organics form as summarized in the introduction above. Desorption kinetic curves for each organic compound due to flowing arsenate and DMA solutions were obtained by subtracting the negative absorbances at 1574, 1658, and 1258 cm−1 shown in Figure 1 from the initial absorbances prior to flowing the arsenicals. These curves are shown in the left axes of Figure 2. The kinetic curves in Figure 2 contain a wealth of information on the initial rates of arsenicals adsorption, initial desorption rates of adsorbed organics, and their residual amounts on the surface after arsenical coverage reached the plateau region. To extract adsorption kinetic rates for arsenicals from the data shown in Figure 2, we used the first-order Langmuir adsorption kinetic model that we employed in our earlier work on the adsorption kinetics of arsenicals and phosphate on clean iron (oxyhydr)oxide surfaces.39,43,44 In this model, the observed adsorption rate, robs, is related to the pseudo first-order adsorption and desorption rate constants: robs,i = kads,i[As(aq)] + kdes, where i stands for initial. Figures S2 and S3 in the Supporting Information show plots for robs versus the bulk aqueous concentration of arsenate (upper panel) and DMA (lower panel) for three hematite surfaces: clean, with adsorbed oxalate, and pyrocatechol. These two surface organics represent two extremes given that at pH 7, oxalate forms mostly 5572


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with adsorbed oxalate relative to a clean surface suggests that arsenate ligand exchange rate is faster with outer-sphere oxalate than with covalently bonded hydroxyl and protonated sites on the surface. The addition of electrolytes at 10 mM concentration leads to layering of anions and cations near the hematite surface.48 Interfacial layering of electrolyte ions at the iron (oxyhdyr)oxide surface were shown to be important from modeling48 and experimental studies.46,47,49,50 On a hematite surface with adsorbed outer-sphere oxalate, enhanced electrostatic attraction layering will lead to concentrating more positive charge at the interface than a clean surface. In the case of a clean surface and within the uncertainty of the measurement, adding either KCl or NaCl had the same effect on slightly enhancing adsorption rate constants of arsenate. One plausible explanation to this finding is through charge density: the ionic radii of K+(aq) and Na+(aq) ions in aqueous solutions (0.141 ± 0.008 and 0.097 ± 0.006 nm, respectively)51 is smaller than of Cl−(aq) ions (0.180 ± 0.008 nm),51 and hence, charge density is higher for the cations. This higher positive charge density is sufficient to screen the negative charge density from Cl−(aq) ions, hence rendering the surface “positive”. However, adding KBr resulted in a decrease in the initial adsorption rate constant relative to no added electrolyte. This observation could be explained by the ineffectiveness of K+(aq) ions in fully screening the negative charges from Br−(aq) ions, which are highly polarizable (α = 5.90 Å3 compared to 0.78 Å3 for K+(aq)).52 Hence, unfavorable electrostatic repulsion between incoming arsenate and a “negatively” charged surface would explain the reduction in the adsorption rate. Data in Table S3 also show that adsorption of arsenate on hematite nanoparticles with adsorbed pyrocatechol is slower even in the presence of KCl as an electrolyte. Pyrocatechol not only forms a mix of protonated monodentate and bidentate binuclear complexes36 with hematite at pH 7, but also introduces a benzene ring rendering the interface hydrophobic. Hence, the reduction in the rate of arsenate adsorption could be explained by the reduction in the number of surface sites favored by arsenate due to adsorbed pyrocatechol, and also the significant reduction in positive surface charge that muted electrostatic attraction. Given the higher thermodynamic favorability of inner-sphere formation of uncharged pyrocatechol surface complexes relative to the formation of negatively charged outer-sphere oxalate complexes, as stated above, different electrolyte types will have little effect on arsenate binding for the former type of complexes. Since NOM has a large density of hydrophobic benzene rings, our findings reported herein suggest that it is these organic functional groups that are responsible for reducing the rate of arsenate adsorption as summarized in the introduction. To further support the above interpretations, in the last section, we provide results from flow microcalorimetric measurements of arsenate adsorption on hematite nanoparticles with adsorbed oxalate and pyrocatechol. The data in Figure 3b show higher rate constants in the presence of electrolytes for the adsorption of DMA on hematite surfaces with adsorbed oxalate than on clean surfaces. This enhancement is more pronounced than that observed for arsenate in Figure 3a. There is also a clear decrease in the adsorption rate constants when Na+(aq) is used instead of K+(aq), and when Br−(aq) replaces Cl−(aq). Scheme 1b shows a pictorial representation of the molecular-level picture inferred from these trends. As stated above, DMA forms a mix of inner-

respectively. For DMA outer-sphere complexes, the values are 4.81 and 4.72 Å, and ΔG°ads = −39.9 kJ mol−1, respectively. Furthermore, because Keq is related to adsorption and desorption rate constants through Keq = kads/kdes, it follows that values of kads will be higher for strongly bonded innersphere bidentate binuclear (with lower kdes, Table S3) compared to outer-sphere complexes that have lower kads and higher kdes. Regardless of the electrolyte type, the presence of surface oxalate on hematite nanoparticles increased values of kads,i for arsenate adsorption by 1.5−2.5x relative to clean hematite (Table S3). This increase is higher than that observed for the control experiments with no added electrolytes (ca. 1.3). A similar trend was observed for DMA adsorption on hematite nanoparticles with surface oxalate in the presence of any of the electrolytes. Figure 3 highlights differences among the electro-

Figure 3. Effect of electrolytes type on the kinetics of arsenicals adsorption on hematite nanoparticles with and without adsorbed organics at pH 7. Raw data used to make these plots are listed in Table S3. OA(ads) = adsorbed oxalate.

lytes in initial adsorption rate constants of arsenicals on different surfaces relative to the control values with no added electrolytes. The cations K+/Na+ and anions Cl−/Br− differ in polarizability, ionic radii, surface loading, and complexation constants on hematite particles.46,47 Bar graphs in Figure 3 were generated by referencing values of kads,i for a given a surface type to those with no added electrolytes. The data in Figure 3a show slightly higher rate constants in the presence of NaCl and KBr relative to KCl for arsenate adsorption on hematite nanoparticles with adsorbed oxalate. Scheme 1a shows a pictorial representation of the molecularlevel picture inferred from these trends. The surface is positively charged at pH 7 with no added electrolytes, below the IEP of hematite. Arsenate adsorption is driven by electrostatics and ligand exchange with FeOH and  FeOH+2 sites. Under these conditions, adsorbed oxalate forms mostly negatively charged outer-sphere complexes. With no added electrolytes in solution, the increase in kads,i for surfaces 5573


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Scheme 1. Inferred Molecular-Level Pictures That Explain the Observed Trends in the Adsorption Kinetics of (a) Arsenate; and (b) DMA on Hematite Nanoparticles in the Absence and Presence of Surface Oxalate as a Function of Different Electrolytes

introduces hydrophobic benzene groups that point away from the surface, and hence can interact with the two methyl groups of DMA instead of the AsO bonds. This explains why the band assigned to v(AsO) in Figure 1f appears narrower than those in Figure 1d and 1e and resembles more the band observed for aqueous phase DMA.42 This finding highlights the effect that organic substituents on arsenate have on their binding mechanism to hematite surfaces with adsorbed organics. To summarize this section: (1) the type of functional groups of adsorbed organics (negatively charged carboxylate versus hydrophobic phenyl) influence amounts and rates of arsenicals adsorption on hematite nanoparticles, (2) organic substituents on arsenate as in DMA shift surface speciation to more outersphere and do interact with hydrophobic groups when organics are present on the surface, and (3) the type of electrolytes affects rate of DMA adsorption to a greater extent than arsenate when organics are present on the surface. Concurrent Desorption of Adsorbed Organics by Arsenicals and Electrolyte Anions in Solution. Figure 2 shows kinetic curves for the desorption of adsorbed organics due to flowing 1 mM arsenate or DMA at pH 7 and 10 mM KCl. For comparison, similar studies were conducted as a function of KCl concentration in the absence of arsenicals to better quantify their effectiveness as desorbing agents (DesA) for low molecular weight organics. The highest KCl concentration used was 1 M as an extreme case in order to

monodentate and outer-sphere complexes on clean hematite, and mostly outer-sphere complexes on surfaces with adsorbed oxalate. The extent of outer-sphere complexation of oxyanions is sensitive to electrolytes type and concentration that changes the ionic strength.53 Since the concentration of all electrolytes studied herein was 10 mM (10× concentration of DMA), the trends in Figure 3b can be explained by a reduction in the rate of ligand exchange between DMA and smaller cations and bigger anions relative to K+(aq) and Cl−(aq), respectively. The relatively high charge-to-size ratio of Na+(aq) ions will lead to stronger interactions with the surface (with and without adsorbed oxalate)48 than K+(aq), hence more resistance to exchange with incoming DMA molecules. This effect is more pronounced with oxyanions that form outer-sphere complexes than inner-sphere bidentate complexes as in the case of arsenate. The case of Br−(aq) ions can be explained in the same manner as arsenate: the ineffectiveness of K+(aq) ions in fully screening the negative charges from Br−(aq) ions renders the surface “negative” and hence electrostatically unfavorable for the adsorption of DMA. For DMA adsorption on hematite particles with adsorbed pyrocatechol, the data in Table S3 show a 1.7× increase in kads,i relative to a clean surface with 10 mM KCl. This is opposite to the reduction observed for the case of arsenate (Table S3), and suggests that the methyl groups on DMA play a role in the adsorption process. As mentioned above, pyrocatechol 5574


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strongly bonded inner-sphere bidentate binuclear arsenate complexes on hematite nanoparticles. The initial desorption rate of outer-sphere oxalate show the largest dependency on the concentration and type of DesA at pH 7. The inset in Figure 4 shows the observed k′des,i as a function of [Cl−(aq)] clearly displaying nonfirst order kinetics. As stated in the caption, the least-square fit to this data is obtained with n = 0.13, suggesting a complex desorption mechanism that might have contributions from electrostatics, ligand exchange, or changes to interfacial water structure. Decoupling these different processes could be possible via theoretical calculations of the dynamics at the hematite/water interface. Using DMA as a DesA further increases the observed desorption rate of oxalate to 5× that with little Cl−(aq), which reflects a more efficient exchange with outer-sphere oxalate. By far, arsenate at 1 mM is the most efficient DesA in exchanging with oxalate, nearly 25× faster than 10 mM Cl−(aq). This increase in the desorption rate of weakly bonded oxalate is more pronounced than with citrate and pyrocatechol desorption confirming earlier interpretation from in situ infrared spectral data36 that the latter two organics form stronger inner-sphere complexes with hematite nanoparticles than outer-sphere oxalate. Table S5 quantifies the surface coverage of organics before and after desorption by Cl−(aq) and arsenicals over the 80 min time frame shown in Figure 2. Within the uncertainty of the measurements, nearly 70% of adsorbed citrate and pyrocatechol are removed by arsenate compared to 100% of oxalate. The flow of DMA results in removing nearly 40% of adsorbed citrate and pyrocatechol compared to 86% of oxalate. This data suggests that depending on the type of organic surface complexes, hematite nanoparticles can retain certain functional groups to some extent even in the presence of strong DesA like arsenate. The retention of these organics explains the trends observed in the surface coverage of arsenicals discussed above. Residual strongly bonded citrate and pyrocatechol block sites favored by arsenate leading to a slight reduction in its surface coverage relative to a clean surface. On the other hand, the increase in the amount of adsorbed DMA relative to the clean surface in the presence of residual citrate and oxalate suggests that these organics provide additional binding sites unavailable on the clean surface. At first glance, this finding is counterintuitive given that these organics and DMA are negatively charged. The surface coverage estimates in Table S2 were obtained from experiments conducted using 10 mM KCl. Similar increase in DMA surface coverage was observed in experiments conducted using 10 mM NaCl and KBr as summarized above. As a result, it is very likely that the cations, K+(aq) and Na+(aq) play a role in attracting and stabilizing DMA outer-sphere complexes. The extent to which K+(aq) and Na+(aq) form ion pairs with DMA can be inferred from their relative concentration: the cations exist at 10× the concentration of DMA, hence a very small portion would be involved in ion pair formation with DMA. Since the largest effect was observed with adsorbed oxalate, control experiments were conducted where no salt (i.e., cations) was added to the aqueous background solutions and those of oxalate and DMA prior to flowing over clean hematite nanoparticles for 80 min. Under these conditions, DMA desorbed 71% of surface oxalate, and the residual organics resulted in a 40% reduction in DMA surface coverage relative to the clean surface. To summarize this section, the extent of organics retention by hematite nanoparticles is influenced by the type of the desorbing agent.

compare the desorption kinetics using 1 mM arsenicals in 0.01 M KCl. As mentioned in the Experimental Section, the desorption of surface arsenicals was carried out for 80 min, but only the first 5 min were used to examine the effect of DesA concentration on the initial desorption rates of adsorbed organics. Longer desorption times were found to be insensitive to concentration as with adsorption rates described above. To extract values for the initial desorption rate constant, k′des,i, experimental data were plotted in the linear form of the Langmuir desorption model, ln[A(t)/Ainitial] = −k′des,i ·t, where A(t) is the absorbance at a given ṽ and Ainitial is the absorbance before the desorption starts, k′des,i = kdes,i·[DesA(aq)]n, where n is the order of desorption. Figure S4 shows the linear form of the desorption experimental data in Figure 2 per the Langmuir model. The slopes obtained from the linear least-squares fits (i.e., k′des,i) are listed in Table S4. The absolute values of k′des,i clearly show that arsenate is the most effective desorbing agent with initial desorption rates that are ca. 2× higher for desorbing surface citrate and pyrocatechol, and 5× higher for desorbing oxalate than DMA and Cl−(aq). To further highlight how the type and concentration of DesA affect initial desorption rates of the three types of surface organics studied herein, for each surface type, values of k′des,i in Table S4 were normalized to those from flowing [Cl−(aq)] = 10−3 mM at pH 7. The results are shown in Figure 4, which can

Figure 4. Dependency of the normalized initial desorption rate parameter, k′des,i, on the type and concentration of desorbing agents at pH 7. Values of k′des,i were normalized to those at 10−3 mM Cl−(aq), which were obtained from linear least-squares fits shown in Figure S4 with slopes listed in Table S4. The inset shows the dependency of the observed k′des,i on [Cl−(aq)] for adsorbed oxalate, and best fit parameters are n = 0.13 and kdes,i = 0.04 min−1 mM−0.13.

be explained by the nature of surface complexes that each adsorbed organic molecule form as summarized in the introduction. Both citrate and pyrocatechol form mostly inner-sphere complexes and hence display the least dependency of k′des,i to increasing concentration of weak DesA like Cl−(aq) and DMA. A 2−3× increase in k′des,i is observed when arsenate is used as a desorbing agent because the exchange with adsorbed citrate and pyrocatechol lead to the formation of 5575


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Table 1. Mass-Normalized Net Total Energies (mJ g−1) Released for the Adsorption of Organics on Hematite Nanoparticles and the Adsorption of Arsenate on Three Types of Surfaces at pH 7 and 5

The favorability of arsenate to form strongly bonded innersphere complexes drive its ability to effectively remove the majority of adsorbed organics. This is in contrast to weaker desorbing agents that mostly form outer-sphere complexes with exchangeable organics leading to the retention of strongly bonded organic complexes. Insights into the Organics−Hematite Interface and Arsenate Adsorption from Flow Microcalorimetric Measurements. Figure 5a shows the raw data (calorimetric

adsorption of organics on clean hematite oxalate

adsorption of arsenate

pyrocatechol

pH 7

pH 5

pH 7

pH 5

surface type

pH 7

pH 5

−195

−2288

−1320

−874

clean hematite pyrocatechol/ hematite oxalate/hematite

−2213 −1823

−3460 −3790

−1277

−2094

The results in Table 1 at pH 7 indicate a significantly more energetic adsorption reaction for pyrocatechol relative to oxalate, further supporting the spectral interpretation of the type of surface complexes each organic compound forms as described above: pyrocatechol forms mostly bidentate binuclear versus outer-sphere surface complexes of oxalate.36 Also the values in Table 1 correlate with binding strengths of oxalate and pyrocatechol surface complexes determined by surface complexation modeling.37 On the other hand, the values at pH 5 suggest that pyrocatechol does not adsorb strongly at that pH. This finding suggests increasing dominance of protonated monodentate complexes of pyrocatechol.36 This result is further supported by the little change in energy following cycles of chloride and nitrate flow as explained below. As for oxalate, the higher energies for adsorption at pH 5 relative to pH 7 are consistent with the formation of predominately bidentate mononuclear complexes under acidic conditions.36 This finding is further supported by the change in the energy following cycles of chloride and nitrate flow over oxalate/ hematite particles to 79% at pH 5 compared to 59% at pH 7 (see below). Guided by the kinetic results shown in Figure 4 for organics desorption due to flowing arsenicals, the effect of surface organics on arsenate adsorption energies relative to clean hematite nanoparticle surfaces was investigated at pH 7 and 5. Calorimetric experiments with DMA were not performed because the highest desorption rates were observed using arsenate (Figure 4). We hypothesized that the calorimetric response among the different surfaces will be clearly differentiated with arsenate than other desorbing agents. Heat signals generated by arsenate adsorption onto the hematite nanoparticle samples are shown in Figure 5c,d. The reaction of arsenate with the three surfaces: clean, and pre-exposed to pyrocatechol and oxalate was strongly exothermic. Table 1 lists the total heats generated per arsenate adsorption event (i.e., exothermic adsorption which also includes the simultaneous endothermic desorption reaction of surface organics) normalized to the mass of hematite nanoparticles. The energy released at pH 5 was found to be higher than that at pH 7. Moreover, Table S6 in the Supporting Information lists total energies of the calorimetric peaks associated with the Cl− displacement of exchangeable NO3− (Cl/N) pre- and post- the organics treatment at pH 7 and 5. A brief discussion on the interpretation of these energies is also provided in the Supporting Information. The Cl/N exchange energies preexposure to organics are lower at pH 7 than 5, which is due to the surface being more positively charged at pH 5, ca. four pH units below the IEP. Following the adsorption of pyrocatechol and oxalate at pH 7, the positive surface charge from FeOH+2

Figure 5. Calorimetric response obtained for the adsorption of pyrocatechol and oxalate on clean hematite nanoparticles surface at (a) pH 7 and (b) pH 5. Also shown is the calorimetric response following the adsorption of arsenate on clean surface and in the presence of adsorbed pyrocatechol and oxalate at (c) pH 7 and (d) pH 5. An increase in voltage resulting in a positive peak corresponds to a release of energy and hence an exothermic reaction. This data is not normalized to the amount of hematite nanoparticles used in each experiment. See text for normalized energies.

signal versus time) for flowing the organic solutions over clean hematite nanoparticles. In these studies, pyrocatechol and oxalate were chosen for the same reason listed above for the adsorption kinetics experiments at pH 7: these two organics represent the two extremes in forming inner- and outer-sphere complexes and introducing a benzene ring versus a negatively charged adsorbate, respectively. Calorimetric studies were also conducted at pH 5 because our earlier pH envelope studies36 showed enhancement in oxalate uptake to hematite nanoparticles and increased concentration of bidentate mononuclear complexes. Hence, we hypothesized that heat released at pH 5 due to oxalate adsorption would be larger than at pH 7, and hence, these findings prove that oxalate forms mostly outersphere at the latter pH. The flow rate of each organic compound was uninterrupted until the thermal signal returned to the initial baseline, which is taken to indicate that the reaction has come to completion. In the case of pyrocatechol adsorption, the signal took an additional 10 min to return to baseline relative to oxalate adsorption, indicating a slightly slower adsorption process for pyrocatechol. As described in the Supporting Information, the area under the peaks in Figure 5 was used to calculate the energy released from these exothermic adsorption reactions. Table 1 lists values for mass-normalized total energies of organics adsorption on hematite nanoparticles at pH 7 and 5. 5576


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found to be 90 and 85%, respectively. When these numbers are compared to the %energies of Cl/N exchange post-organics adsorption, it becomes clear that arsenate adsorption removed most of adsorbed pyrocatechol to the degree that the % Cl/N energy exchange is close to that calculated following arsenate adsorption on clean particles. This observation is consistent with the trend in energy values reported in Table 1. However, arsenate adsorption on oxalate/hematite caused additional change in the % Cl/N energy exchange to 85% from 79% following the organics adsorption. This 6% difference is smaller than the 18% calculated at pH 7, which is explained by less desorption of strongly bonded surface oxalate by arsenate at pH 5. Further investigations and mass-balance measurements using other techniques are warranted.

sites had decreased 49 and 59%, respectively, as Cl/N cycles have revealed (Table S6, Supporting Information). These measurements were obtained immediately following the end of the organics adsorption, i.e., before any complexes of organics have had the significant time and opportunity to be desorbed and flushed by repeated NaCl and NaNO3 flow. Successive multiple Cl/N cycles were also performed, where it was observed that the reduction in the Cl/N energy post-oxalate adsorption is partially reversed (i.e., ∼30% of the reduction is reversed). This recovery in energy increase and thus in positive surface charge probably resulted from the loss of some outersphere and/or weakly bonded oxalate complexes with the flow of NaNO3 and NaCl solutions. In contrast, the same experiments failed to measure any recovery in the positive surface charge following pyrocatechol adsorption, suggesting that aqueous chloride and nitrate are unable to desorb/ exchange with the strongly bonded pyrocatechol. The relatively lower %energy exchange for pyrocatechol (49 ± 10% compared to 59 ± 10% for oxalate) further supports the interpretation that pyrocatechol is likely forming more irreversible and bulky bidentate complexes, especially bidentate binuclear, that are lower in number than oxalate complexes. This will give rise to lower number of available FeOH+2 sites available for Cl/N exchange relative to the clean surface and hence smaller exchange energy post-exposure to pyrocatechol. Benjamin54 modeled bidentate versus monodentate adsorption on pristine surfaces to account for the effect of molecular size on the number of available binding sites. In our case, the pyrocatechol is larger in size (as reflected by its molecular weight and structure) resulting in a higher steric crowding on the surface relative to oxalate. A similar analysis of the exchange energies of Cl/N was done at pH 5 (Table S6, Supporting Information). The surface charge decreased by ca. 79% following the adsorption of oxalate, and there was no change detected following the adsorption of pyrocatechol. These observations suggest formation of strongly bonded oxalate complexes, more specifically bidentate mononuclear,36 driven by the increase in the concentration of positive surface charge and singly protonated oxalate, HOA−(aq), species under acidic conditions. As for pyrocatechol, the results suggest increased formation of weaker complexes at pH 5 than pH 7, likely protonated monodentate.36 In addition, arsenate adsorption on the above surfaces with organics further reduced the positive surface charge evident by the reductions in the energies of Cl/N exchange post-arsenate by an additional 33 and 18% for pyrocatechol and oxalate, respectively (Table S6, Supporting Information). These results reveal that at pH 7, arsenate not only adsorbs on unoccupied sites and hence further reduces the number of sites available for Cl/N exchange, but desorbs surface organics and occupies the sites previously held by the organics. For example, in the case of arsenate adsorption on oxalate/hematite particles, arsenate desorbed weakly bonded oxalate, occupied these sites, and also adsorbed onto an additional 18% of charge-bearing sites. These calorimetric observations strongly suggest that arsenate desorbed more surface oxalate than pyrocatechol, hence consuming more energy in the desorption process leading to lower values of the net heat released. This conclusion further supports the differences in the kinetic curves shown in Figure 2b,c, where arsenate clearly desorbed more surface oxalate than pyrocatechol. At pH 5, the %energies of Cl/N exchange post-arsenate adsorption on pyrocatechol/and oxalate/hematite particles was

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CONCLUSIONS AND SIGNIFICANCE The adsorption kinetics and heat release of arsenate and DMA with hematite nanoparticles pre-exposed to three types of low molecular weight organics: citrate, oxalate, and pyrocatechol were conducted using ATR-FTIR and flow microcalorimetry. The results highlight that differences in the structure of arsenicals and surface organics, and the type of electrolyte present in solution affect the reactivity of hematite nanoparticles. More specifically, negatively charged carboxylate versus hydrophobic phenyl groups influence amounts and rates of arsenicals adsorption on hematite nanoparticles, the type of electrolytes affects adsorption rate of DMA to a greater extent than arsenate when oxalate is present on the surface, and the extent of organics retention by hematite nanoparticles is influenced by the type of the desorbing agent. Given the ubiquity of organic−metal oxide interfaces in environmental systems, these results are significant because, unlike studies involving large molecular weight NOM, low molecular weight organics can either have an inhibiting or a stimulating effect on arsenicals binding to iron (oxyhydr)oxides. Our studies also highlight that in situ flow mode and timedependent spectroscopic and microcalorimetric experiments can amplify trends and molecular-level processes otherwise missed in bulk batch adsorption experiments. Kinetic and thermodynamic models need to be developed that incorporate molecular-level processes involved in the retention and release of organics and arsenicals in particular in the presence of iron (oxyhydr)oxides.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpca.7b03426. Detailed experimental procedures; figures and tables showing data analysis (PDF)

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Phone: (519)884-0710, ext.2873. Fax: (519)746-0677. ORCID

Nadine Kabengi: 0000-0002-1901-0992 Hind A. Al-Abadleh: 0000-0002-9425-0646 Notes

The authors declare no competing financial interest. 5577


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(19) Meier, M.; Namjesnik-Dejanovic, K.; Maurice, P. A.; Chin, Y. P.; Aiken, G. R. Fractionation of Aquatic Natural Organic Matter Upon Sorption to Goethite and Kaolinite. Chem. Geol. 1999, 157 (3-4), 275−284. (20) Mladenov, N.; Zheng, Y.; Simone, B.; Bilinski, T. M.; McKnight, D. M.; Nemergut, D.; Radloff, K. A.; Rahman, M. M.; Ahmed, K. M. Dissolved Organic Matter Quality in a Shallow Aquifer of Bangladesh: Implications for Arsenic Mobility. Environ. Sci. Technol. 2015, 49, 10815−10824. (21) Weng, L.; Van Riemsdijk, W. H.; Hiemstra, T. Effects of Fulvic and Humic Acids on Arsenate Adsorption to Goethite: Experiments and Modeling. Environ. Sci. Technol. 2009, 43, 7198−7204. (22) Ritter, K.; Aiken, G. R.; Ranville, J. F.; Bauer, M.; Macalady, D. L. Evidence for the Aquatic Binding of Arsenate by Natural Organic Matter−Suspended Fe(III). Environ. Sci. Technol. 2006, 40, 5380− 5387. (23) Mikutta, C.; Kretzschmar, R. Spectroscopic Evidence for Ternary Complex Formation between Arsenate and Ferric Iron Complexes of Humic Substances. Environ. Sci. Technol. 2011, 45, 9550−9557. (24) Sharma, P.; Ofner, J.; Kappler, A. Formation of Binary and Ternary Colloids and Dissolved Complexes of Organic Matter, Fe and As. Environ. Sci. Technol. 2010, 44, 4479−4485. (25) Wang, S. L.; Mulligan, C. N. Enhanced Mobilization of Arsenic and Heavy Metals from Mine Tailings by Humic Acid. Chemosphere 2009, 74, 274−279. (26) Liu, G.; Fernandez, A.; Cai, Y. Complexation of Arsenite with Humic Acid in the Presence of Ferric Iron. Environ. Sci. Technol. 2011, 45, 3210−3216. (27) Mikutta, R.; Lorenz, D.; Guggenberger, G.; Haumaier, L.; Freund, A. Properties and Reactivity of Fe-Organic Matter Associations Formed by Coprecipitation Versus Adsorption: Clues from Arsenate Batch Adsorption. Geochim. Cosmochim. Acta 2014, 144, 258−276. (28) ThomasArrigo, L. K.; Mikutta, C.; Byrne, J.; Barmettler, K.; Kappler, A.; Kretzschmar, R. Iron and Arsenic Speciation and Distribution in Organic Flocs from Streambeds of an Arsenic-Enriched Peatland. Environ. Sci. Technol. 2014, 48, 13218−13228. (29) Silva, G. C.; Vasconcelos, I. F.; de Carvalho, R. P.; Dantas, M. S.; Ciminelli, V. S. Molecular Modeling of Iron and Arsenic Interactions with Carboxy Groups in Natural Biomass. Environ. Chem. 2009, 6, 350−356. (30) Bauer, M.; Blodau, C. Arsenic Distribution in the Dissolved, Colloidal and Particulate Size Fraction of Experimental Solutions Rich in Dissolved Organic Matter and Ferric Iron. Geochim. Cosmochim. Acta 2009, 73, 529−542. (31) Sundman, A.; Karlsson, T.; Sjoberg, S.; Persson, P. Complexation and Precipitation Reactions in the Ternary As(V)−Fe(III)−OM (Organic Matter) System. Geochim. Cosmochim. Acta 2014, 145, 297− 314. (32) Gu, B.; Schmitt, J.; Chen, Z.; Liang, L.; McCarthy, J. F. Adsorption and Desorption of Natural Organic Matter on Iron Oxide: Mechanisms and Models. Environ. Sci. Technol. 1994, 28, 38−46. (33) Lin, W.; Walter, J.; Burger, A.; Maid, H.; Hirsch, A.; Peukert, W.; Segets, D. A General Approach to Study the Thermodynamics of Ligand Adsorption to Colloidal Surfaces Demonstrated by Means of Catechols Binding to Zinc Oxide Quantum Dots. Chem. Mater. 2015, 27, 358−369. (34) Benoit, P.; Hering, J. G.; Stumm, W. Comparative Study of the Adsorption of Organic Ligands on Aluminum Oxide by Titration Calorimetry. Appl. Geochem. 1993, 8, 127−139. (35) Ferreiro, E. A.; de Bussetti, S. G. Thermodynamic Parameters of Adsorption of 1,10-Phenanthroline and 2,2′-Bipyridyl on Hematite, Kaolinite and Montmorillonites. Colloids Surf., A 2007, 301, 117−128. (36) Situm, A.; Rahman, M. A.; Goldberg, S.; Al-Abadleh, H. A. Spectral Characterization and Surface Complexation Modeling of Organics on Hematite Nanoparticles: Role of Electrolytes in the Binding Mechanism. Environ. Sci.: Nano 2016, 3, 910−926.

ACKNOWLEDGMENTS H.A.A. acknowledge partial funding from Laurier, NSERC and Early Researcher Award from Ontario’s Ministry of Research and Innovation. N.K. acknowledge the support of the U.S. Department of Energy (DOE) Office of Science Early Career Research Program Grant number 211267.

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