Competitive Adsorption of Cd(II), Cr(VI), and Pb(II) onto

Article pubs.acs.org/est

Competitive Adsorption of Cd(II), Cr(VI), and Pb(II) onto Nanomaghemite: A Spectroscopic and Modeling Approach Michael Komárek,*,† Carla M. Koretsky,‡ Krishna J. Stephen,‡ Daniel S. Alessi,§ and Vladislav Chrastný†,∥ †

Department of Environmental Geosciences, Faculty of Environmental Sciences, Czech University of Life Sciences Prague, Kamýcká 129, Prague 6 − Suchdol, 165 21, Czech Republic ‡ Department of Geosciences, Western Michigan University, 1187 Rood Hall, Kalamazoo, Michigan 49008, United States § Department of Earth and Atmospheric Sciences, University of Alberta, 1-26 Earth Sciences Building, Edmonton, Alberta, T6G 2E3, Canada ∥ Department of Geochemistry, Czech Geological Survey, Geologická 6, 152 00, Prague 5, Czech Republic S Supporting Information *

ABSTRACT: A combined modeling and spectroscopic approach is used to describe Cd(II), Cr(VI), and Pb(II) adsorption onto nanomaghemite and nanomaghemite coated quartz. A pseudo-second order kinetic model fitted the adsorption data well. The sorption capacity of nanomaghemite was evaluated using a Langmuir isotherm model, and a diffuse double layer surface complexation model (DLM) was developed to describe metal adsorption. Adsorption mechanisms were assessed using X-ray photoelectron spectroscopy and X-ray absorption spectroscopy. Pb(II) adsorption occurs mainly via formation of inner-sphere complexes, whereas Cr(VI) likely adsorbs mainly as outer-sphere complexes and Cd(II) as a mixture of inner- and outer-sphere complexes. The simple DLM describes well the pH-dependence of single adsorption edges. However, it fails to adequately capture metal adsorption behavior over broad ranges of ionic strength or metal-loading on the sorbents. For systems with equimolar concentrations of Pb(II), Cd(II), and Cr(VI). Pb(II) adsorption was reasonably well predicted by the DLM, but predictions were poorer for Cr(VI) and Cd(II). This study demonstrates that a simple DLM can describe well the adsorption of the studied metals in mixed sorbate−sorbent systems, but only under narrow ranges of ionic strength or metal loading. The results also highlight the sorption potential of nanomaghemite for metals in complex systems.

1. INTRODUCTION

Due to the dependence of both the mobility and bioavailability of metal(loid)s on their chemical speciation, many environmental problems are best addressed using detailed process-level knowledge of metal(loid) speciation. Numerous spectroscopic techniques can provide information regarding the actual speciation of metal(loid)s, including extended X-ray absorption fine structure (EXAFS), X-ray absorption near edge structure (XANES), and X-ray photoelectron spectroscopy (XPS), among others.11−13 For some metals, such as Cr, isotope analyses can provide additional information about speciation, because oxidation and reduction reactions can be quantified using isotope fractionation.14,15 Isotope fractionation has also been observed due to metal coordination changes when Cd and Zn adsorb on Mn and Fe oxides to form inner-sphere complexes, especially at low ionic strengths.16,17

Adsorption of metals and metalloids onto secondary oxides in soils and sediments significantly influences their speciation and subsequent mobility and bioavailability. Due to their high reactivity and large specific surface areas, which can measure in the tens to hundreds of m2/g, nanosized oxides are important scavengers of soil contaminants.1 Maghemite is a common weathering product in soils of temperate, tropical, and subtropical climatic regions, usually forming during the oxidation of magnetite. Although not as abundant as goethite and hematite in soils, prior work demonstrates that natural maghemite is an important sorbent for metals and metalloids.2−9 Synthetic maghemite is thus a promising material for the removal of inorganic contaminants from aqueous solutions and soils as it is readily available, inexpensive and can be easily separated and recovered because it is magnetic. Additionally, nanomagnetite and subsequently nanomaghemite are possible oxidation products of nanozerovalent iron, an emerging amendment used for the remediation of contaminated groundwaters and soils.10 © XXXX American Chemical Society

Received: June 25, 2015 Revised: September 22, 2015 Accepted: October 12, 2015

A

DOI: 10.1021/acs.est.5b03063 Environ. Sci. Technol. XXXX, XXX, XXX−XXX

Article

Environmental Science & Technology Many empirical and mechanistic models have been developed to simulate metal(loid) speciation, including adsorption and precipitation.18−21 Mechanistic models, such as surface complexation models (SCMs), have been used to describe the binding of metal(loid)s onto various soil constituents, including solid and dissolved organic matter, clay minerals, bacteria and metal (hydr)oxides. SCMs are based on thermodynamics (mass law equations) and are analogous to speciation models used to describe complexation reactions in aqueous solutions. “Intrinsic” equilibrium or stability constants used in SCMs are less system-dependent than parameters derived using simpler semiempirical or empirical models, and thus they should be more transferable to other well-defined systems.19 Semiempirical isotherm and kinetic models can, however, be used to derive system-dependent information, such as adsorption capacities and reaction rates. Despite attempts to model metal(loid) adsorption onto natural and synthetic Fe oxides,22,23 data and models for systems with nanosized materials remain scarce. The aim of this study is to measure and model the competitive adsorption of Cd(II), Cr(VI), and Pb(II) onto nanosized maghemite and maghemite-coated quartz. The use of these complex adsorption systems allows the robustness of a SCM developed using the component additivity approach to be assessed, specifically, to determine whether relatively simple models can accurately capture interactions between multiple sorbents and sorbates. The metals were deliberately chosen to represent common contaminants with diverse geochemical behavior. Spectroscopic data (XPS and XAS) and isotopic fractionation data (Cd, Cr, and Pb) were used to gain insights into adsorption mechanisms. This, together with bulk adsorption and kinetic data, was used to constrain both semiempirical and thermodynamically based SCMs describing competitive metal(loid) adsorption over a wide range of solution conditions.

Figure 1. XRD analyses (a, nanomaghemite; b quartz) and TEM (c, d) images of the nanomaghemite coatings on quartz, m, nanomaghemite; q, quartz.

Table 1. Characteristics of Nano-Maghemite and Quartz, Reaction Stoichiometries and Stability Constants Used in DLM Calculationsa nanomaghemite specific surface area: 44.6 m2/g site density: 3.79 μmol/m2

log K

≡FeOH + H+ ⇌ ≡FeOH2+ ≡FeOH ⇌ ≡FeO− + H+ ≡FeOH + Cd2+ ⇌ ≡FeOCd+ + H+

5.71 −7.74 −2.28 ± 0.62 ≡FeOH + Pb2+ ⇌ ≡FeOPb+ + H+ 3.32 ± 1.25 ≡FeOH + H+ + CrO42‑ ⇌ ≡FeCrO4− + 12.4 ± 0.75 H2O quartz

2. MATERIALS AND METHODS 2.1. Materials. Synthetic nanomaghemite (γ-Fe2O3; < 50 nm) was obtained from Sigma Aldrich; natural crystalline quartz with an average particle size of 1.7 μm (Min-U-Sil 5) was purchased from U.S. Silica. XRD and TEM analyses were used to verify the phases (Figure 1). Qualitative analysis of the XRD patterns was performed using PANalytical X’Pert HighScore software (version 1.0d) and the ICDD PDF-2 database. Transmission electron microscopy images were obtained using a JEOL electron microscope model JEM-1230. Prior to adsorption experiments, the quartz was heated to 500 °C, washed with 4 M HNO3, rinsed with deionized water, centrifuged, and dried at room temperature to remove possible impurities. Nanomaghemite coatings (Figure 1) on the quartz particles were prepared according to Schwertmann and Cornell.24 XRF (Delta, Olympus) analyses showed that the mixed particles contained 4.42 wt.% of the nanomaghemite coating. Specific surface areas for the nanomaghemite and quartz were determined at atmospheric pressure using a Quanta-Chrome Nova Surface and Pore Analyzer model 2200e. Three replicates of 2 g samples of each solid were degassed at ∼ 80 °C for 24 h and analyzed using 11-point N2 BET (Table 1). All chemicals used in the experiments were of analytical grade or better. Solutions of Cd(II), Cr(VI), and Pb(II) were prepared from single-metal ICP standards (Analytica, Czech Republic). 2.2. Potentiometric Titrations. Acid−base titration data were used to optimize surface protonation models and to

specific surface area: 9.4 m2/g

23.9 (16.1, 39.2) 35.5 (23.5, 60.0) 10.5 (6.88, 18.0)

26

site density: 16.6 μmol/m

2 21

≡SiOH ≡SiOH ≡SiOH ≡SiOH

V(Y) (V(Y)min, V(Y)max)

log K

+ H+ ⇌ ≡SiOH2+ ⇌ ≡SiO− + H+ + Cd2+ ⇌ ≡SiOCd+ + H+ + Pb2+ ⇌ ≡SiOPb+ + H+

−1.1 21 −8.1 21 −5.3 42 −1.74 26

a

The log K values for metal complexation reactions are presented as averages ± SD based on best-fit values derived for individual adsorption edges at each metal loading and ionic strength. Average goodness-of-fit parameters of Cd(II), Cr(VI) and Pb(II) surface complexation models obtained in this study are presented as (V(Y)) at 95% confidence intervals.

determine protonation/deprotonation constants of nanomaghemite. Acid (1 M HNO3) and base (1 M NaOH) were used for the titration experiments. The titrations were performed using 40 g/L nanomaghemite suspensions in background electrolyte (0.01 M NaNO3) using a CO2-free chamber with N2 as the inert gas. The titration procedure was fully automated using a TitroLine Alpha Plus Titrator (Schott, Germany). The equilibrium condition was set to a pH drift lower than 0.60 mV/h.23 Protonation/deprotonation constants and site densities were obtained using a diffuse double layer model (DLM) and the ProtoFit 2.1 software.25 B


Article

Environmental Science & Technology 2.3. Adsorption Experiments. 2.3.1. Adsorption Kinetics. To evaluate the time required to reach adsorption equilibrium and to compare adsorption rates with prior studies, kinetic batch adsorption experiments were performed for all the studied metals with nanomaghemite (2 g/L) at three different pH values (3, 4.5, 6 for Cr(VI) and Pb(II); 6, 7, 8 for Cd(II), as there was no Cd adsorption at pH values below 6) in open atmosphere and using 0.01 M NaNO3 as the background electrolyte. The pH was manually set and maintained using HNO3 and NaOH solutions. The initial metal concentrations were 0.1 mM. Samples were shaken in a reciprocal shaker for 1−180 min and subsequently filtered using 0.2 μm syringe nylon filters. The possible presence of the maghemite nanoparticles in the filtrates was monitored by analyzing them for Fe, which was below detection limits in all supernatant solutions after filtration (0.2-μm) (data not shown), probably due to aggregation of the nanomaghemite particles. Therefore, 0.2 μm filters were subsequently used for all adsorption experiments. Metal concentrations were analyzed in the filtrates using ICP-OES (720 Series, Agilent Technologies). The difference between initial and final values provided the amounts sorbed per gram of sorbent. Each experiment was conducted in triplicate and results are given as averages and standard deviations. 2.3.2. Adsorption Isotherms. Adsorption isotherms were measured in triplicate in open atmosphere using a procedure similar to that for the kinetic experiment, using single- and multimetal (Cd(II)+Pb(II)) solutions, 2 g/L nanomaghemite and 0.01 M NaNO3 as the background electrolyte. The pH values of the single-metal solutions were set to 3, 4.5, and 6 for Cr(VI) or Pb(II), and 6, 7, and 8 for Cd(II). For solutions containing Cd(II) with Pb(II), the pH was set to 3, 4.5, and 6. Cr(VI) was not added to the multimetal solution as it resulted in the precipitation of PbCrO4 and CdCrO4, especially at higher metal concentrations. The adsorption time was set to 3 h based on the results from the kinetic experiments, which ensured that equilibrium was achieved. 2.3.3. Adsorption Edges. Adsorption edges were measured according to Reich et al.26 using batch experiments under atmospheric conditions at different metal concentrations (10−6 and 10−5 M) and ionic strengths (0.001, 0.01, 0.1 M NaNO3). The investigated systems included nanomaghemite (0.2 and 2 g/L) and nanomaghemite-coated quartz (2 g/L). The log K values for protonation, deprotonation and metal (Cd and Pb) adsorption onto quartz were taken from previous studies (Table 1). The metals were added either alone or in combination to evaluate competitive effects. The solids were added to batch reactors and left to equilibrate on a rotary shaker for 24 h at room temperature. The slurry was then titrated using HNO3/NaOH and 10 mLaliquots were removed at ∼ 0.5 pH intervals and left to equilibrate for an additional 24 h on a rotating shaker. The pH of each aliquot was then measured, the samples centrifuged and filtered using 0.2 μm syringe nylon filters and the supernatants analyzed for metals using ICP-OES or for Cr(VI) on select samples using the colorimetric diphenylcarbazide method.27 2.4. Adsorption Modeling. 2.4.1. Adsorption Kinetics and Isotherms. Data from the kinetic experiments were fitted to pseudo-first-order and pseudo-second-order equations. Langmuir and Freundlich isotherms were initially developed to model gas adsorption, but are now commonly used to describe aqueous metal sorption. Because they are empirical models, the fitted parameter values do not provide information

about sorption mechanisms. Nevertheless, Langmuir isotherms can give information on sorption capacities.28,29 Isotherm parameters were extracted from the measured data using nonlinear least-squares regression.30,31 2.4.2. Surface Complexation Modeling. Adsorption edges for single metal nanomaghemite-only systems were modeled using FITEQL 4.0 with the default thermodynamic data for aqueous species from Visual MINTEQ.32 Stability constants were optimized for each individual edge using a DLM with a given reaction stoichiometry (Tables 1 and S1). Davies equation activity corrections were used for the aqueous species, and the concentration of CO 2 (aq) was set assuming equilibrium with an atmospheric concentration of 380 ppm. Only adsorption data on each edge (i.e., at > 2% and < 98% adsorption) was used to optimize stability constants. Visual MINTEQ 3.133 was used to calculate surface complexation model predictions for systems with mixed metals and sorbents using stability constants obtained from the singlemetal adsorption edges. The goodness-of-fit V(Y) of the edges obtained from FITEQL to the measured data was assessed according to Heinrich et al.34 and Reich et al.26 and is generally considered to be reasonable if it is between 1 and 20. 2.5. XPS and XAS Analyses. X-ray photoelectron spectroscopy (XPS) analyses were performed to evaluate the elemental composition and chemical oxidation states of surface and near-surface metal species using an Omicron Nanotechnology ESCAProbeP spectrometer. The X-ray source was monochromated at 1486.7 eV and spectra were measured stepwise with a binding energy step of 0.05 eV. The NIST Xray photoelectron spectroscopy database was used for the identification of the spectra. X-ray absorption spectroscopy (XAS) at the Cr and Fe Kedges was conducted at beamline HXMA of the Canadian Light Source (CLS). Nanomaghemite samples were diluted at a 25:1 mass ratio with boron nitride, and pressed into pellets that were subsequently covered with Kapton tape. A double-crystal Si(111) monochromator, detuned 50−60% to reject higher harmonic frequencies, was used to select energies. For continuous energy monitoring and calibration, a Cr or Fe foil and a pin diode detector were placed after the I0 ion chamber and before the sample position along the beam. Fluorescence from the sample was collected using a 32-element Ge solid state detector. Spectra of iron(II) chloride and synthetic goethite, and potassium dichromate and synthetic chromium(III) hydroxide, were used as references for X-ray absorption near edge structure (XANES) linear combination fitting (LCF) of sample valence states. Background subtraction, splining, and fitting of spectra was conducted using the Athena analysis package.35 2.6. Isotope Analyses. Cadmium, Cr and Pb isotope analyses were performed to investigate possible fractionation of Cd, Cr, and Pb during adsorption. Before analysis, Cd, Cr and Pb were separated from matrix elements using anion exchange chromatography.36−39 All Cr, Cd, and Pb isotope measurements were performed on a double-focusing multicollector inductively coupled plasma mass spectrometer (MC ICP-MS, Neptune, Thermo Scientific, Germany) equipped with nine Faraday detectors.

3. RESULTS AND DISCUSSION 3.1. Adsorption Experiments. 3.1.1. Adsorption Kinetics. Adsorption equilibrium is generally reached after 15−30 min for Cr(VI), 60 min for Cd(II) and 45−60 min for Pb(II) C


Article

Environmental Science & Technology (Figure 2). Because the experiments were performed in the presence of atmospheric CO2, otavite (CdCO3) at pH 8 and

Figure 2. Adsorption kinetics for 0.1 mM of metals on 2 g/L nanomaghemite in 0.01 M NaNO3, qt, adsorbed concentration of metal at time t.

hydrocerrusite (Pb3(CO3)2(OH)2) at pH 7 and 8 are supersaturated and may precipitate according to speciation modeling using VisualMINTEQ. All three metals follow a pseudo-second order kinetics model (Table S2), in agreement with prior studies of metal adsorption onto oxides, especially maghemite.2,40,41 Few prior studies report kinetic parameters for metal adsorption onto (nano)maghemite, but the pseudosecond order kinetic model parameters for Cr(VI) are in accordance with the study of Jiang et al.2 3.1.2. Adsorption Isotherms. Adsorption of all three metals on the nanomaghemite follows expected trends over the studied pH range, that is, adsorption of Cd(II) and Pb(II) increase with increasing pH, and Cr(VI) adsorption decreases. There is no significant influence of Cd(II) or Pb(II) competition on maximum sorption of the other metal. This is likely because the isotherms with mixed metals were measured at low pH. Under these conditions there is significant adsorption of Pb(II), but a lower affinity of Cd(II) for the nanomaghemite surface is observed. Precipitation of otavite (CdCO3) and hydrocerrusite (Pb3(CO3)2(OH)2) may occur at pH ≥ 6 based on saturation indices calculated using VisualMINTEQ. Therefore the data measured for this pH range were not included in the isotherm models. Both the Langmuir and the Freundlich model fit the adsorption isotherms adequately, with the former producing a slightly better fit to the data (Table S3). The fitted Langmuir parameters demonstrate that Pb(II) is the most efficiently adsorbed of the metals, as expected. The adsorption maxima are comparable to those reported in other studies of metal(loid) adsorption onto maghemite, including nanomaghemite.2,6 Although it is not possible to identify adsorption mechanisms from isotherm data alone, parameters derived from the isotherms, for example, adsorption maxima, can be compared with other studies. 3.1.3. Adsorption Mechanisms. Wide scan XPS spectra for experiments with maximum metal adsorption (Figure S1, Table S3) were used to obtain insight into Cd(II), Cr(VI), and Pb(II) adsorption onto the nanomaghemite. The XPS spectra were compared to the NIST X-ray photoelectron spectroscopy database, which was used to identify binding energies of the metals. No change in Fe(III) valence due to metal adsorption was observed, in agreement with prior studies.6,41 Reduction of Cr(VI) to Cr(III) has been reported to occur during Cr(VI) adsorption on goethite,43 biosorbents and mixed maghemitemagnetite nanoparticles, with Fe(II) in the magnetite serving as an electron donor.13,15,41,44 However, the XPS spectra demonstrate that redox reactions did not occur during adsorption of metals on nanomaghemite in this study.

Figure 3. Langmuir adsorption isotherms for Cd(II), Cr(VI), and Pb(II) on 2 g/L nanomaghemite in 0.01 M NaNO3 as background electrolyte, data for which otavite and hydrocerrusite are calculated to be supersaturated were excluded from the model; a, c, e: single metal solutions; b, d: competition isotherms with equimolar Cd(II) and Pb(II); ce: equilibrium metal ion concentration in the solution after adsorption, qe: adsorbed metal concentration at equilibrium (n = 3).

XPS data identified Cr(VI) adsorbed to the nanomaghemite,45 possibly as an electrostatically bound outer-sphere complex, which is also supported by the observation that adsorption of negatively charged Cr(VI) species only occurs on a positively charged surface (i.e., at pH < pHzpc) and by the strong ionic-strength dependence of Cr(VI) adsorption on the nanomaghemite (see below). In contrast, the XPS spectra suggest that Cd(II) and Pb(II) may be present as inner-sphere complexes as indicated by the binding energies corresponding to Cd-O (∼405 eV) and Pb-O (∼138 eV) bonds. This interpretation is also supported by 100% adsorption of Pb(II) below the pHzpc (Figures 4 and 5). Previous XPS studies41 also demonstrated that a significant portion of Cd(II) is complexed with oxygen atoms at the nanomaghemite surface, indicating the presence of inner-sphere complexes. However, Cd(II) adsorbs at pH values around and above the pHzpc and shows some dependence on ionic strength (Figures 4 and 5), suggesting that a mixture of outer- and inner-sphere complexes may form. Speciation calculations using VisualMINTEQ indicate supersaturation of otavite (CdCO3) at pH > 7.4 in the presence of atmospheric CO2. However, because most of the Cd(II) adsorbs to the nanomaghemite surface at lower pH, precipitation likely does not occur. Isotope fractionation has been previously observed during adsorption of Zn onto hematite and goethite17 and Cd onto birnessite.16 The Cd isotopic fractionation is thought to be related to distortion of inner-sphere surface complexes compared to the symmetrical octahedral configuration of Cd(H 2 O) 6 2+ in solution, which results in preferential scavenging of lighter isotopes onto oxide surfaces at low ionic strength.16 No Cd fractionation was observed in this study (Table S4), perhaps due to the higher ionic strengths used in this study, or to a combination of outer- and inner-sphere Cd complex formation on the nanomaghemite. In the case of Pb, D


Article

Environmental Science & Technology

Figure 4. Modeled adsorption edges for Cd(II), Cr(VI), Pb(II) and nanomaghemite (0.2 and 2 g/L) from solutions containing single metals (a−c) and equimolar concentrations of all metals (d−f). Symbols represent experimental data; solid and dashed lines represent respective modeled edges of the metals at 10−6 M and 10−5 M, respectively. Vertical lines indicate otavite (CdCO3) and hydrocerrusite (Pb3(CO3)2(OH)2) supersaturation for aqueous metal concentrations of 10−6 M (solid line) and 10−5 M (dashed line) in the absence of adsorption.

Figure 5. Modeled adsorption edges for Cd(II), Cr(VI), Pb(II), and nanomaghemite coated quartz (2 g/L total solid: 0.088 g/L nanomaghemite, 1.912 g/L quartz) from solutions containing single metals (a−c) and all the metals at equimolar concentrations (d−f). Symbols represent experimental data; solid and dashed lines represent respective modeled edges of the metals at 10−6 M and 10−5 M, respectively. Vertical lines indicate otavite (CdCO3) and hydrocerrusite (Pb3(CO3)2(OH)2) supersaturation for aqueous metal concentrations of 10−6 M (solid line) and 10−5 M (dashed line) in the absence of adsorption.

no fractionation is expected. The fractionation of Cr isotopes during sorption is mainly associated with Cr(VI) reduction to

Cr(III). In agreement with the XPS spectra, which indicate that redox reactions do not occur during adsorption of Cr(VI) on E


Article


ionic strength experiments, possibly because Cr(VI) forms outer-sphere complexes that are not adequately described by the DLM. Aqueous carbonate complexes CdCO3(aq) and PbCO3(aq) become significant at pH values above the pH of ∼100% metal adsorption (8.0 and 7.0, respectively).53,54 Additionally, the formation of aqueous ternary complexes, such as Pb(OH)CO3− is only significant at pH values above 8.0, which is well above the pH of 100% adsorption.53 Thus, metal−carbonate surface complexes were not included in the modeling. Carbonate might adsorb on the nanomaghemite surface at lower pH as observed for example on goethite and thus compete for surface sites with the metals especially at high pCO2.55,56 However, without carbonate adsorption data, it is difficult to justify adding this reaction, and thus another adjustable parameter, to the model. In the absence of metal adsorption, otavite and hydrocerrusite are supersaturated at high pH (>7.4 for 10−5 M Cd and > 8.0 for 10−6 M Cd; > 6.8 for 10−5 M Pb(II) and > 7.3 for 10−6 M Pb(II)). Hydrocerrusite is supersaturated well above the pH at which 100% Pb(II) adsorption occurs, and thus is unlikely to influence Pb(II) sorption. While otavite supersaturation in the absence of adsorption occurs somewhat below the predicted pH of 100% adsorption for some of the edges, precipitation likely affects few, if any, of the measured data, because adsorption of Cd(II) at lower pH will reduce the aqueous Cd(II) concentration, resulting in supersaturation of otavite at higher pH values (Figures 4a−d). The poor performance of the DLM in describing the ionicstrength dependent sorption of Cr(VI) on nanomaghemite in the absence of Cd(II) or Pb(II) is even more apparent in experiments with equimolar concentrations of Cd(II), Cr(VI), and Pb(II). In these experiments, the DLM only provides reasonable agreement with Cr(VI) adsorption data for systems with lower concentrations of the metals and at lower ionic strengths (Figure 4e). At the higher loading and highest ionic strength, the DLM dramatically overpredicts adsorption of Cr(VI) on the nanomaghemite in the presence of Cd(II) and Pb(II). As for the Cr(VI)-only system, this may be due to the lack of distinction between inner- and outer-sphere complexes in the diffuse layer model. In contrast, the DLM provides reasonable predictions for Pb(II) and Cd(II) adsorption in the mixed sorbate systems that are of similar accuracy to those in the absence of competing ions. Pb(II) adsorption on nanomaghemite in the presence of Cr(VI) and Cd(II) is well described as a function of pH, metal concentration, and ionic strength. Cd(II) adsorption is also reasonably well predicted in experiments with equimolar levels of the metals at lower loadings. At higher loadings, sorption of Cd(II) reaches 100% at a pH near 8, but the DLM underpredicts Cd(II) adsoprtion in the presence of Cr(VI) and Pb(II) at this pH. Thermodynamically based surface complexation models are generally assumed to have a significant advantage over more empirical adsorption models because they should be capable of correctly predicting adsorption in more complex systems than those used to derive the adsorption stability constants. Specifically, the component additivity model assumes that stability constants derived for ion adsorption on single sorbents can be combined to correctly predict adsorption in the presence of multiple sorbents.57 This approach typically assumes that no significant sorbent-sorbent interactions (e.g., site blocking) occur. DLM parameters derived in this study using pure nanomaghemite systems were used to predict

the nanomaghemite, no fractionation of Cr isotopes was observed (Table S4). Chromium is commonly found in the hexavalent and trivalent states in nature. Therefore, XAS analyses were used to investigate the valence state of Cr bound to nanomaghemite. Linear combination fitting (LCF) of the first Cr K-edge XAS scan indicates that more than 86% of total Cr adsorbed to the nanomaghemite is hexavalent (Figure S2ac; Table S6), as expected. Maghemite, an Fe(III) mineral, should exhibit little or no capacity to reduce Cr(VI). However, subsequent scans of the same sample demonstrated reduction of Cr(VI) and a resulting increase in the fraction of Cr(III) observed in the XANES LCFs (Figure S2ac; Table S6). After nine scans of 48.9 min each, 73% of Cr(VI) was reduced to Cr(III) in the beam. Fitting a first-order kinetics model (t1/2 = 224 min) to the Cr(VI) reduction data (Figure S3) demonstrates that all Cr on the nanomaghemite was initially hexavalent. No beam-induced reduction of Cr(VI) in the K2Cr2O7 standard was observed. Thus, while nanomaghemite cannot itself reduce Cr(VI), it may allow generation of radical species at the water-nanomaghemite interface46 that can reduce Cr(VI) and other redox-active metals. Fe K-edge XANES are nearly identical for nanomaghemite with and without adsorbed Cr(VI), with no change in bulk Fe valence state observed over time in either sample, precluding reduction of a significant fraction of the total Fe due to beam damage (Figure S2b). 3.1.4. Surface Complexation Modeling. The protonation and deprotonation log K values and the site density for the nanomaghemite surface modeled in ProtoFit 2.1 using a DLM with a single amphoteric site (Table 1) are comparable to those from prior studies.23,47,48 The calculated pHzpc values for the nanomaghemite and quartz were 6.72 and 2.09, respectively. The adsorption edges of Cd(II), Cr(VI), and Pb(II) on nanomaghemite were modeled using a single monodentate surface complex and excluding carbonate adsorption. Although some XAS studies suggest that Pb binds to hydrous ferric oxide via inner-sphere bidentate complexes, bidentate complexes resulted in poorer model fits compared to monodentate ones.26,49,50 The log K values extracted from individual adsorption experiments with different ionic strengths and loadings were averaged to obtain a single log K describing each metal-solid system (Figure 4a−c; Table 1). The DLM average stability constants describe the pHdependent Cd(II), Cr(VI), and Pb(II) adsorption on purenanomaghemite reasonably well. However, as observed in prior studies with many other sorbents and sorbates,26,51 a simple DLM does not correctly capture the ionic strength or loading dependence, even for single metal systems (Figure 4a−c). The standard deviation of the log K values obtained for individual adsorption edges (various metal concentrations and ionic strengths) was the highest for Pb(II), resulting in especially poor representation of the observed ionic strength and loading dependence. For studies in which broad ranges of ionic strength and/or metal loading are considered, it has frequently been found that the DLM does not produce invariant stability constants.26,51,52 We suggest that the DLM should only be extrapolated with care because of this issue, and that it may be better to use ionic-strength dependent stability constants when applying a DLM.51 Cd(II) and Pb(II) adsorption was underpredicted for the lowest ionic strengths (0.001 M) and metal concentrations (10−6 M) and somewhat overpredicted for higher metal concentrations (10−5 M) and ionic strengths (0.1 M). Cr(VI) adsorption was overpredicted for the lowest F


Article


Cr(VI) outer-sphere complexes, which are suggested both by the spectroscopic data and the observed electrostatic effects (ionic strength dependence), and which are not adequately described by a simple DLM. This study suggests that a simple component additivity approach can provide useful descriptions of complex systems involving Cd(II) and Pb(II), but that the results are only as good as those for the end member systems used to constrain the model (i.e., synthetic nanomaghemite and natural quartz).

Cd(II), Pb(II) and Cr(VI) adsorption, both singly and in systems with metal mixtures, on nanomaghemite coated quartz. DLM parameters for Pb(II) and Cd(II) on quartz were taken from prior studies (Table 1).21,26,42 Cr(VI) adsorption on pure quartz was negligible for the conditions used in this study (data not shown), therefore, Cr(VI) adsorption on quartz was excluded from the component additivity model. As for pure nanomaghemite, the pH-dependent adsorption of Cd(II) and Pb(II) is reasonably well described on nanomaghemite coated quartz both in the absence and presence of competing metals, but the dependence on ionic strength and metal loading is not (Figures 5a−f). In the absence of competing metals, Cd(II) adsorption is somewhat underpredicted at the lower metal concentration (10−6 M) and lowest ionic strength (0.001 M) and overpredicted at the higher metal concentration (10−5 M) and at highest ionic strengths (0.01−0.1 M). In the presence of competing metals, Cd(II) on nanomaghemite coated quartz tends to be underpredicted, especially at higher loadings. Pb(II) adsorption on nanomaghemite coated quartz is overpredicted for the low metal loading and slightly underpredicted for the high loading (Figure 5). The component additivity model dramatically overpredicts adsorption of Cr(VI) on nanomaghemite coated quartz under most conditions at pH < 7. This may be due to the charge repulsion between the negatively charged aqueous chromate and the negatively charged quartz surface, which has a pHzpc of ∼2. As for the pure nanomaghemite system, the strong ionic strength dependence of Cr(VI) adsorption is not captured by the DLM (especially at the lowest ionic strength), which cannot differentiate between inner- and outer-sphere surface complexes. The DLM was most successful at predicting Cr(VI) adsorption onto pure nanomaghemite at a solid concentration of 2 g/L. Lowering the nanomaghemite concentrations to 0.2 g/L for multimetal systems with pure nanomaghemite and to 0.088 g/L for nanomaghemite coatings on quartz resulted in much poorer fits to the data, demonstrating that the Cr(VI) DLM cannot be readily extrapolated over large ranges in solid concentration. The SCM was also used to calculate metal adsorption for the measured isotherms, but generally resulted in poor fits to the measured data (data not shown). This may be because the isotherms were measured at significantly higher metal concentrations than those for which the SCM was parametrized. This could reflect the difficulty described above with extrapolating the DLM over broad ranges of solution conditions, or could indicate that the mechanisms of sorption are not identical at the high and low metal loadings. This study demonstrates that nanomaghemite may be an important scavenger of Cr(VI), Pb(II), and Cd(II) from aqueous solutions and thus could be a useful sorbent for water and soil remediation. Adsorption of Pb(II) occurs mainly through the formation of inner-sphere complexes, while Cd(II) was likely adsorbed as a mixture of inner- and outer-sphere complexes. A simple DLM can readily reproduce pHdependent adsorption edges at a single ionic strength or loading conditions, but does not produce good fits to Pb(II) or Cd(II) adsorption data over broad ranges of ionic strength or metal loading. The component additivity model captures some of the metal competition effects and interactions of nanomaghemite with quartz particles reasonably well for Pb(II) and Cd(II). Cr(VI) adsorption was more poorly described by the DLM in both simple and more complex systems as compared to Cd(II) and Pb(II). This may be due to the formation of

■

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.est.5b03063. Details about potentiometric titrations, adsorption modeling, including used equations, kinetic/isotherm parameters, XPS, XAS spectra and isotope data are available as Supporting Information (PDF)

■

AUTHOR INFORMATION

Corresponding Author

*Phone: +420224383857; fax: +4202343837; e-mail: [email protected]. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS M.K. is thankful for the support from the Fulbright Commission, from the Czech Science Foundation (project 15-07117S) and from the Internal Grant Agency of the Czech University of Life Sciences Prague (project 20154202). K.S. was supported by the Western Michigan University Alliances Graduate Education and the Professoriate program. D.A. received funding from the Natural Sciences and Engineering Research Council of Canada Discovery program. Jeremy Fein, ́ Jennifer E. S. Szymanowski, Hana Šillerová, Sylva Č ihalová , Petr Sajdl, Marie Králová, Vojtěch Ettler, Amy Troy, Thomas J. Reich, and Md. Samrat Alam are thanked for their assistance during experiments and analyses. We thank four reviewers for their valuable comments and suggestions.

■

REFERENCES

(1) Waychunas, G. A.; Kim, C. S.; Banfield, J. F. Nanoparticulate iron oxide minerals in soils and sediments: Unique properties and contaminant scavenging mechanisms. J. Nanopart. Res. 2005, 7 (4− 5), 409−433. (2) Jiang, W.; Pelaez, M.; Dionysiou, D. D.; Entezari, M. H.; Tsoutsou, D.; O’Shea, K. Chromium(VI) removal by maghemite nanoparticles. Chem. Eng. J. 2013, 222, 527−533. (3) Tuutijärvi, T.; Lu, J.; Sillanpäa,̈ M.; Chen, G. As(V) adsorption on maghemite nanoparticles. J. Hazard. Mater. 2009, 166 (2−3), 1415− 1420. (4) Chowdhury, S. R.; Yanful, E. K. Arsenic and chromium removal by mixed magnetite-maghemite nanoparticles and the effect of phosphate on removal. J. Environ. Manage. 2010, 91 (11), 2238−2247. (5) Roy, A.; Bhattacharya, J. Removal of Cu(II), Zn(II) and Pb(II) from water using microwave-assisted synthesized maghemite nanotubes. Chem. Eng. J. 2012, 211−212, 493−500. (6) Hu, J.; Chen, G.; Lo, I. M. C. Removal and recovery of Cr(VI) from wastewater by maghemite nanoparticles. Water Res. 2005, 39 (18), 4528−4536. (7) Komarek, M.; Vanek, A.; Ettler, V. Chemical stabilization of metals and arsenic in contaminated soils using oxides - A review. Environ. Pollut. 2013, 172, 9−22. G


Article

Environmental Science & Technology (8) Michalkova, Z.; Komarek, M.; Sillerova, H.; Della Puppa, L.; Joussein, E.; Bordas, F.; Vanek, A.; Vanek, O.; Ettler, V. Evaluating the potential of three Fe- and Mn-(nano)oxides for the stabilization of Cd, Cu and Pb in contaminated soils. J. Environ. Manage. 2014, 146, 226− 234. (9) Taylor, R. M.; Schwertmann, U. Maghemite in soils and its origin. I. Properties and observations on soil maghemites. Clay Miner. 1974, 10 (4), 289−298. (10) Filip, J.; Karlický, F.; Marušaḱ , Z.; Lazar, P.; Č erník, M.; Otyepka, M.; Zbořil, R. Anaerobic reaction of nanoscale zerovalent iron with water: mechanism and kinetics. J. Phys. Chem. C 2014, 118 (25), 13817−13825. (11) Morin, G.; Ona-Nguema, G.; Wang, Y.; Menguy, N.; Juillot, F.; Proux, O.; Guyot, F.; Calas, G.; Brown, G. E., Jr. Extended X-ray absorption fine structure analysis of arsenite and arsenate adsorption on maghemite. Environ. Sci. Technol. 2008, 42 (7), 2361−2366. (12) Sparks, D. L. Elucidating the fundamental chemistry of soils: past and recent achievements and future frontiers. Geoderma 2001, 100 (3−4), 303−319. (13) Chowdhury, S. R.; Yanful, E. K.; Pratt, A. R. Chemical states in XPS and Raman analysis during removal of Cr(VI) from contaminated water by mixed maghemite-magnetite nanoparticles. J. Hazard. Mater. 2012, 235−236, 246−256. (14) Døssing, L. N.; Dideriksen, K.; Stipp, S. L. S.; Frei, R. Reduction of hexavalent chromium by ferrous iron: A process of chromium isotope fractionation and its relevance to natural environments. Chem. Geol. 2011, 285 (1−4), 157−166. (15) Sillerova, H.; Chrastny, V.; Cadkova, E.; Komarek, M. Isotope fractionation and spectroscopic analysis as an evidence of Cr(VI) reduction during biosorption. Chemosphere 2014, 95, 402−407. (16) Wasylenki, L. E.; Swihart, J. W.; Romaniello, S. J. Cadmium isotope fractionation during adsorption to Mn oxyhydroxide at low and high ionic strength. Geochim. Cosmochim. Acta 2014, 140, 212− 226. (17) Pokrovsky, O. S.; Viers, J.; Freydier, R. Zinc stable isotope fractionation during its adsorption on oxides and hydroxides. J. Colloid Interface Sci. 2005, 291 (1), 192−200. (18) Goldberg, S. Use of surface complexation models in soil chemical systems. Adv. Agron. 1992, 47, 233−329. (19) Koretsky, C. The significance of surface complexation reactions in hydrologic systems: a geochemist’s perspective. J. Hydrol. 2000, 230 (3−4), 127−171. (20) Bradl, H. B. Adsorption of heavy metal ions on soils and soils constituents. J. Colloid Interface Sci. 2004, 277 (1), 1−18. (21) Sverjensky, D. A.; Sahai, N. Theoretical prediction of single-site surface-protonation equilibrium constants for oxides and silicates in water. Geochim. Cosmochim. Acta 1996, 60 (20), 3773−3797. (22) Dzombak, D. A.; Morel, F. M. M. Surface Complexation Modeling. Hydrous Ferric Oxide; Wiley-Interscience: New York, 1990; p 416. (23) Jolsterå, R.; Gunneriusson, L.; Holmgren, A. Surface complexation modeling of Fe3O4-H+ and Mg(II) sorption onto maghemite and magnetite. J. Colloid Interface Sci. 2012, 386 (1), 260−267. (24) Schwertmann, U.; Cornell, R. M. Iron Oxides in the Laboratory: Preparation and Characterization; Wiley-VCH Verlag GmgH: Weinheim, Germany, 2000; p 188. (25) Turner, B. F.; Fein, J. B. Protofit: A program for determining surface protonation constants from titration data. Comput. Geosci. 2006, 32 (9), 1344−1356. (26) Reich, T. J.; Das, S.; Koretsky, C. M.; Lund, T. J.; Landry, C. J. Surface complexation modeling of Pb(II) adsorption on mixtures of hydrous ferric oxide, quartz and kaolinite. Chem. Geol. 2010, 275 (3− 4), 262−271. (27) Clesceri, L. S.; Greenberg, A. W.; Eaton, A. D., Standard Methods for the Examination of Water and Wastewater, 20th ed.; American Public Health Association: Washington DC, USA, 1992; p 1325. (28) Hinz, C. Description of sorption data with isotherm equations. Geoderma 2001, 99 (3−4), 225−243.

(29) Limousin, G.; Gaudet, J. P.; Charlet, L.; Szenknect, S.; Barthès, V.; Krimissa, M. Sorption isotherms: A review on physical bases, modeling and measurement. Appl. Geochem. 2007, 22 (2), 249−275. (30) Bolster, C. H.; Hornberger, G. M. On the use of linearized Langmuir equations. Soil Sci. Soc. Am. J. 2007, 71 (6), 1796−1806. (31) Bolster, C. H. Revisiting a statistical shortcoming when fitting the Langmuir model to sorption data. J. Environ. Qual. 2008, 37 (5), 1986−1992. (32) Herbelin, A. L.; Westall, J. FITEQLA Computer Program for Determination of Chemical Equilibrium Constants from Experimental Data, Dept of Chemistry Rep 99−01, Oregon State University, Corvallis, OR, USA: 1999. (33) Gustafsson, J. P. Visual MINTEQ. Version 3.1.; Division of Land and Water Resources, Royal Institute of Technology: Stockholm, Sweden: 2013. (34) Heinrich, H. T. M.; Bremer, P. J.; McQuillan, A. J.; Daughney, C. J. Modelling of the acid−base properties of two thermophilic bacteria at different growth times. Geochim. Cosmochim. Acta 2008, 72 (17), 4185−4200. (35) Ravel, B.; Newville, M. Athena, artemis, hephaestus: data analysis for X-ray absorption spectroscopy using IFEFFIT. J. Synchrotron Radiat. 2005, 12, 537−541. (36) Gao, B.; Liu, Y.; Sun, K.; Liang, X.; Peng, P.; Sheng, G.; Fu, J. Precise determination of cadmium and lead isotopic compositions in river sediments. Anal. Chim. Acta 2008, 612 (1), 114−120. (37) Wombacher, F.; Rehkämper, M.; Mezger, K.; Münker, C. Stable isotope compositions of cadmium in geological materials and meteorites determined by multiple-collector ICPMS. Geochim. Cosmochim. Acta 2003, 67 (23), 4639−4654. (38) Chrastny, V.; Rohovec, J.; Cadkova, E.; Pasava, J.; Farkas, J.; Novak, M. A new method for low-temperature decomposition of chromites and dichromium trioxide using bromic acid evaluated by chromium isotope measurements. Geostand. Geoanal. Res. 2014, 38 (1), 103−110. (39) Bullen, T. D.; Wang, Y. Chromium stable isotopes as a new tool for forensic hydrology at sites contaminated by anthropogenic chromium. In Proceedings of the 12th International Symposium on Water-Rock Interaction, Kunming, China, 31 July to 5 August, 2007; Taylor and Francis, CRC Press, London, UK: Kunming, China, 2007. (40) Della Puppa, L.; Komarek, M.; Bordas, F.; Bollinger, J. C.; Joussein, E. Adsorption of copper, cadmium, lead and zinc onto a synthetic manganese oxide. J. Colloid Interface Sci. 2013, 399, 99−106. (41) Chowdhury, S. R.; Yanful, E. K. Kinetics of cadmium(II) uptake by mixed maghemite-magnetite nanoparticles. J. Environ. Manage. 2013, 129 (C), 642−651. (42) Schaller, M. S.; Koretsky, C. M.; Lund, T. J.; Landry, C. J. Surface complexation modeling of Cd(II) adsorption on mixtures of hydrous ferric oxide, quartz and kaolinite. J. Colloid Interface Sci. 2009, 339 (2), 302−309. (43) Abdel-Samad, H.; Watson, P. R. An XPS study of the adsorption of chromate on goethite (a-FeOOH). Appl. Surf. Sci. 1997, 108, 371− 377. (44) Hua, M.; Zhang, S.; Pan, B.; Zhang, W.; Lv, L.; Zhang, Q. Heavy metal removal from water/wastewater by nanosized metal oxides: a review. J. Hazard. Mater. 2012, 211−212, 317−31. (45) Biesinger, M. C.; Payne, B. P.; Grosvenor, A. P.; Lau, L. W. M.; Gerson, A. R.; Smart, R. S. C. Resolving surface chemical states in XPS analysis of first row transition metals, oxides and hydroxides: Cr, Mn, Fe, Co and Ni. Appl. Surf. Sci. 2011, 257 (7), 2717−2730. (46) Chirita, M.; Grozescu, I. Fe2O3 − Nanoparticles, Physical properties and their photochemical and photoelectrochemical applications. Chem. Bull. ″POLITEHNICA″ Univ. (Timişoara) 2009, 54 (1), 1−8. (47) Jolsterå, R.; Gunneriusson, L.; Forsling, W. Adsorption and surface complex modeling of silicates on maghemite in aqueous suspensions. J. Colloid Interface Sci. 2010, 342 (2), 493−498. (48) Stumm, W.; Morgan, J. J., Aquatic Chemistry. Chemical Equlibria and Rates in Natural Waters. 3rd ed.; John Wiley & Sons, Inc.: Hoboken, NJ, 1996. H


Article

Environmental Science & Technology (49) Xu, Y.; Boonfueng, T.; Axe, L.; Maeng, S.; Tyson, T. Surface complexation of Pb(II) on amorphous iron oxide and manganese oxide: spectroscopic and time studies. J. Colloid Interface Sci. 2006, 299 (1), 28−40. (50) Trivedi, P.; Dyer, J. A.; Sparks, D. L. Lead sorption onto ferrihydrite. 1. A macroscopic and spectroscopic assessment. Environ. Sci. Technol. 2003, 37 (5), 908−914. (51) Reich, T. J.; Koretsky, C. M. Adsorption of Cr(VI) on γ-alumina in the presence and absence of CO2: Comparison of three surface complexation models. Geochim. Cosmochim. Acta 2011, 75 (22), 7006−7017. (52) Landry, C. J.; Koretsky, C. M.; Lund, T. J.; Schaller, M.; Das, S. Surface complexation modeling of Co(II) adsorption on mixtures of hydrous ferric oxide, quartz and kaolinite. Geochim. Cosmochim. Acta 2009, 73 (13), 3723−3737. (53) Powell, K. J.; Brown, P. L.; Byrne, R. H.; Gajda, T.; Hefter, G.; Leuz, A.-K.; Sjö berg, S.; Wanner, H. Chemical speciation of environmentally significant metals with inorganic ligands. Part 3: The Pb2+ + OH−, Cl−, CO32−, SO42−, and PO43− systems (IUPAC Technical Report). Pure Appl. Chem. 2009, 81 (12), 2425−2476. (54) Powell, K. J.; Brown, P. L.; Byrne, R. H.; Gajda, T.; Hefter, G.; Leuz, A.-K.; Sjö berg, S.; Wanner, H. Chemical speciation of environmentally significant metals with inorganic ligands. Part 4: The Cd2+ + OH−, Cl−, CO32−, SO42−, and PO43− systems (IUPAC Technical Report). Pure Appl. Chem. 2011, 83 (5), 1163−1214. (55) Van Geen, A.; Robertson, A. P.; Leckie, J. O. Complexation of carbonate species at the goethite surface: Implications for adsorption of metal ions in natural waters. Geochim. Cosmochim. Acta 1994, 58 (9), 2073−2086. (56) Villalobos, M.; Perez-Gallegos, A. Goethite surface reactivity: a macroscopic investigation unifying proton, chromate, carbonate and lead(II) adsorption. J. Colloid Interface Sci. 2008, 326, 307−323. (57) Davis, J. A.; Coston, J. A.; Kent, D. B.; Fuller, C. C. Application of the surface complexation concept to complex mineral assemblages. Environ. Sci. Technol. 1998, 32 (19), 2820−2828.

I


Competitive Adsorption of Cd(II), Cr(VI), and Pb(II) onto

Recommend Documents