Article pubs.acs.org/est
Impact of Interactions between Natural Organic Matter and Metal Oxides on the Desorption Kinetics of Uranium from Heterogeneous Colloidal Suspensions Yu Yang,†,§,* James E. Saiers,† and Mark O. Barnett‡ †
School of Forestry and Environmental Studies, Yale University, 195 Prospect Street, New Haven, Connecticut, 06511, United States Department of Civil Engineering, 238 Harbert Engineering Center, Auburn University, Auburn, Alabama, 36849, United States
‡
S Supporting Information *
ABSTRACT: Colloids play an important role in governing the transport of radionuclides in geologic environments. As naturally occurring colloidal suspensions are compositionally heterogeneous, the subsurface fate of radionuclides may be sensitive to interactions among different kinds of colloids. Therefore, we investigated the adsorption equilibrium and desorption kinetics of uranium (U(VI)) in experiments conducted with compositionally homogeneous suspensions of colloidal SiO2, ZnO, hydrous ferric oxide (HFO) or humic acids (HAs) as well as heterogeneous suspensions consisting of a colloidal metal oxide and HA. We found that interactions between HAs and ZnO or HFO greatly inhibited the sorption of U onto colloids in the heterogeneous suspensions. HA−ZnO interactions enhanced the desorption of U from the heterogeneous colloidal suspensions, while the association between HA and SiO2 or HFO inhibited U desorption. Molecular-level characterizations reveal that HFO interacted with HAs by electrostatic interactions, association with aliphatic/aromatic carbon and inner-sphere complexation with carboxyl functional groups, while SiO2 and ZnO mainly associated with HAs by weak interactions (e.g., van der Waals interactions). The present findings indicate that interactions between HA and metal-oxide colloids can substantially influence the desorption of U(VI) from these particles, thereby potentially affecting the mobility of this radionuclide in groundwater.
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INTRODUCTION
mobilization of organic phases in groundwater environments and leaching from soils.5 COM and mineral colloids do not exist separately but interact in ways that depend on their chemical compositions as well as the pH, ionic composition, and ionic strength of groundwater. For instance, COM binds with metal oxides mainly through electrostatic and ligand-exchange interactions, for which the phenolic and carboxyl moieties of COM play a critical role.10,11 Interactions among COM and inorganic colloids can greatly influence their sorption of radionuclides. Bouby et al.12 found that the addition of humic acids (HAs) promotes the release of Th(IV) and Eu(IV) from clay colloids, even for those colloids which have been aged for three years. Although these studies reveal that colloid−colloid interactions are important in governing radionuclide adsorption and desorption reactions, most published research has focused on radionuclide sorption to compositionally homogeneous suspensions of either inorganic or organic colloids.5 As natural groundwater suspensions are inherently heterogeneous with
Nuclear weapons testing and radionuclear materials storage has led to radionuclide contamination of groundwaters at sites around the world.1 Aqueous colloids are critical agents of radionuclide transport within geologic environments, and colloid-facilitated transport of radionuclides has received considerable attention due to the health risks associated with human exposure to these toxic constituents in groundwater.1,2 A comprehensive understanding of the sorptive associations between colloids and radionuclides is important for predicting radionuclide transport and optimizing remediation strategies for contaminated sites.3,4 The composition of groundwater colloids is variable and governed by the parent material (i.e., source rock or sediments) and the tendency of dissimilar colloids to resist deposition and remain stable in aqueous suspension.5,6 Novikov et al.7 found that colloids composed of amorphous iron hydroxide, rancieite, and hematite were responsible for the long-range transport of plutonium (Pu) and uranium (U) beneath the Mayak Production Facility in Russia. In addition to inorganic particles, colloidal forms of organic matter (COM) are present in groundwater and serve as vectors of metals and radionuclide transport.5,8,9 Groundwater COM is generated by in situ © 2013 American Chemical Society
Received: Revised: Accepted: Published: 2661
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respect to chemical composition,13 continued advances in our understanding of radionuclide transport in real subsurface environments rely on analysis of additional observations of radionuclide desorption from colloidal suspensions composed of mixtures of inorganic and organic particles. To improve descriptions of the facilitated transport of U by natural colloid suspensions, we seek to improve understanding of how associations between COM and metal-oxide colloids affect the adsorption and desorption of U. In this study, we (1) measured the equilibrium sorption of U on compositionally homogeneous suspensions of COM or metal oxides colloids as well as heterogeneous suspensions composed of mixtures of COM and metal-oxide colloids; (2) analyzed the desorption kinetics of U from the homogeneous and heterogeneous suspensions; and (3) characterized the molecular-level interactions between COM and metal-oxide colloids and evaluated the relationships between these interactions and U desorption kinetics.
of SiO2, ZnO, HFO, HA1 or HA2 and heterogeneous suspensions consisting of either HA1 or HA2 and one of the colloidal metal oxides (Supporting Informatio (SI), Figure S1). The use of dialysis bags (1000 Da cutoff, VWR, Radnor, PA, United States) enabled detection of the sorptive associations of U with portions of the colloidal pool that were too small to be separated from the aqueous phase by conventional membrane filtration. In preparation for the adsorption experiments, a background solution (0.01 M NaNO3, pH = 7.8) was bubbled with air for at least 2 d to reach equilibrium with the atmosphere. Four milliliters of a colloidal suspension (pH = 7.8) were transferred into a dialysis bag (1000 Da cutoff) that was immersed in 100 mL of the 0.01 M NaNO3 solution that contained 1 mg/L U. The concentrations of colloids inside the dialysis bag were specified to ensure that colloids removed sufficient quantities of U from solution so that adsorption could be reliably quantified, and, for both the single-colloid and mixed-colloid experiments, the colloid concentrations equaled 5, 5, 400, 40, and 20 mg/L for HA1, HA2, SiO2, ZnO, and HFO, respectively. In experiments with the heterogeneous suspensions, HA1 or HA2 was mixed separately with SiO2, ZnO, or HFO for 12 h prior to addition to the dialysis bag to achieve interaction equilibrium between the COM and metal oxides.20 Sampling began soon after the colloid-containing dialysis bag was immersed in the U-NaNO3 solution. Samples from both inside and outside the dialysis bag were collected at intervals over a period of 400 h by pipets and analyzed for U concentrations. U(VI) was measured by a kinetic phosphorescence analyzer (KPA) based on the phosphorescence of U(VI) after mixing with Uraplex (Chemchek, Richland, WA, U.S.).15 The detection limit for U was 0.001 mg/L. U Desorption Experiments. The continuously stirring flow system (CSFS) developed in our previous study15 was applied to analyze the kinetics of U desorption from singlecolloid and mixed-colloid suspensions. Predetermined amounts of colloid stock suspension and U stock solution were mixed with the NaNO3 background solution. The colloid concentrations were the same as those used in the adsorption experiments and the U concentration equaled 4 mg/L, which is within the range of concentrations reported for some actinidecontaminated areas, such as the Hanford site in Washington state.18 The U-colloid mixtures were pre-equilibrated for 24 h, which, according to published studies, is a period sufficient to achieve sorption equilibrium.4,19 The equilibrated U-colloid mixture was transferred by pipets into a dialysis bag (1000 Da cutoff, VWR, Radnor, PA, United States) that, in turn, was immersed in a 100 mL of a colloid-free background solution (0.01 M NaNO3, pH = 7.8). To maintain near-zero concentrations of U in the background solution outside the dialysis bag, the solution was continuously refreshed by a peristaltic pump at a rate of 1 mL/min. At time intervals ranging from 60 to 1440 min, 0.1 mL of solution from inside the dialysis bag was collected and analyzed for U concentrations with the KPA. Model Analysis. Quantifying Adsorption. We analyzed the results of the adsorption experiments largely for the purposes of quantifying the adsorption of U to the colloids under equilibrium conditions. A simple empirical model was used to compute temporal changes in U concentrations inside and outside the dialysis bag for times leading up to the establishment of steady-state, and the modeled concentrations
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METHODS AND MATERIALS Materials. Humic Acids. The HAs were progressively extracted from a peat soil seven times with 0.1 M Na4P2O7 and then six times with 0.1 M NaOH at a 10:1 (v:w) solutionto-solid ratio, for which the details have been described in previous studies.14−16 The extracted HAs were precipitated by adding concentrated HCl to acidify the solution and demineralized by rinsing five times with HCl/HF followed by rinsing five times with DI water.14−16 The first extraction (HA1) and last extraction (HA2) were used in this study as model humic substances given that they have relatively large differences in their chemical structures as determined in our previous studies.14−16 Stock mixtures of aqueous-phase HAs were made by adding 0.1 g of the deashed HAs to 2 mL of 0.5 M NaOH and filtering the resulting suspensions through 2.7 μm membranes to collect particles with size smaller than 2.7 μm (Millipore, Billerica, MA, U.S.). The colloidal fraction of the HAs were collected through dialysis of the aqueous-phase HAs using a Spectra/Por dialysis bag (1000 Da cutoff, VWR, Radnor, PA, U.S.).15 Concentrations of colloidal HAs (with size between 1000 Da and 2.7 μm) were measured by a total organic carbon (TOC) analyzer (TOC-V, Shimadzu, Japan). These HAs have been characterized by nuclear magnetic resonance (NMR), X-ray photoelectron spectroscopy (XPS) and synchrotron-based techniques.14−16 Metal-Oxide Colloids. Colloidal suspensions of SiO2 (47−49 wt %, pH 8.5−9.0, diameter 20−30 nm) (Nissan Chemical America Corporation, Houston, TX, US) and ZnO (50 wt %, pH 7, diameter 0.90 for all the experiments (SI, Table S1). The estimates of log Kd (mL/g) were 4.5 ± 0.54, 5.0 ± 0.14, and 6.1 ± 0.85 for U on SiO2, ZnO and HFO, respectively, indicating that HFO had a higher sorption affinity for U compared to SiO2 and ZnO. The log Kd for U on HFO was much higher than the reported value (∼3.5 at pH of 8) for colloidal magnetite,19 probably because the HFO in this experiment was more amorphous than synthesized magnetite. Previous studies showed that ferrihydrite associates with U through inner-sphere complexation of the type (FeO2)UO2.30 Such inner-sphere complexation between U and iron oxide, compared to weaker interactions between U and other colloids (as discussed below), 2664
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Figure 2. (A−C) Measurements (symbols) and model simulations (dashed lines computed with eq 5) of U desorption. The residual fraction of total U expresses the ratio of total U concentration inside the colloid-containing dialysis bag at time t to the initial total U concentration inside the dialysis bag. (D−F) Calculations of colloid-bound residual fraction: (i) dotted lines are derived from the experimental data and computed by using the bestfit values of Ccoll0 and k1 in eq 6 and (ii) solid lines are computed with eq 7, which are based on the assumption that HA-colloid interactions do not affect U desorption kinetics.
may be responsible for the higher sorption of U on HFO under the experimental conditions tested in this study. The log Kd (mL/g) values for the adsorption of U on HA1 and HA2 were 5.8 ± 1.1 and 5.7 ± 1.0, respectively, which are greater than corresponding values for SiO2 and ZnO, but similar to HFO. The addition of HAs to ZnO and HFO decreased U concentrations inside the dialysis bags at all sample times (Figure 1) relative to the experiments with metal colloids only, indicating that the presence of HAs within the ZnO and HFO suspensions reduced the fraction of colloid-bound U. U concentrations inside dialysis bag with SiO2 were increased by addition of HA1 but decreased by HA2 (Figure 1). Calculations of Ceqin‑pred made by application of eq 4 were much higher than corresponding measured values (Ceqin), except for the case of SiO2−HA1 (SI, Figure S2). The ratio of Ceqin to Ceqin‑pred can be taken as an index of the impact of interactions between metal oxides and HAs on the equilibrium sorption of U. Calculated ratios were the lowest for HFO, with the values of 0.37 and 0.25 for HFO-HA1 and HFO-HA2, respectively. These low ratios for HFO-HA indicate that interactions
between HFO and HAs led to substantially lower U sorption. Gu et al.20,31 have demonstrated that iron oxide can strongly bind HAs through ligand exchange of carboxyl/hydroxyl functional groups on HAs and the surfaces of iron oxides. These interactions likely made some of the moieties on both HFO and HAs unavailable to U, and consequently decreased the U sorption. For all the metal oxides, the ratios for mixtures with HA2 were lower than those for mixtures with HA1, which is possibly due to more extensive interactions between metal oxides and HA2. Our previous studies showed that HA2 has more hydrophobic functional groups compared to HA1, and the hydrophobic carbon may interact with greater portions of the metal-oxide surfaces through general van der Waals and hydrophobic interactions.14−16 Uranium Desorption from Colloids. The total U residual fraction (i.e., proportion of U remaining inside the dialysis bag) decreased more slowly in the experiments with HFO than in experiments with SiO2 or ZnO (Figure 2A−C), partially because of the higher initial quantity of U bound to HFO than either SiO2 or ZnO. Best-fit values of the desorption rate 2665
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constant (k1) obtained by fitting eq 5 to these data equal 7.7, 9.3, and 5.7 × 10−4 min−1 for SiO2, ZnO, and HFO, respectively (SI, Table S2). The rate constants for U desorption from HA1 and HA2 (determined previously15) equal 1.7 and 5.1 × 10−4 min−1, which are lower than those for the metal colloids and indicative of stronger interactions between HAs and U. Using values of k1 and Ccoll0 estimated by fitting eq 5 to the time-series data on total U concentrations (Ctot), the residual fraction of colloid-bound U (Ccoll/Ccoll0) can be computed from eq 6. Consideration of the calculations made with eq 6 show that HFO-associated U desorbed more slowly than U bound to ZnO and SiO2, such that 44% of U associated with HFO remained inside the dialysis bags after 1,440 min, while 26−33% of U bound by SiO2 and ZnO remained within dialysis bags after this time (Figure 2D−F, dotted lines). The addition of HAs altered the desorption kinetics of U, and the time-series data on total U residual fraction were closely reproduced by the desorption model (Figure 2A−C). The best-fit values of the desorption rate constant (i.e., k1 in eq 5) depended on HA composition and varied from 2.5 to 3.0, 4.6−7.0, and 3.1−3.7 × 10−4 min−1 for SiO2, ZnO, and HFO, respectively (SI, Table S2). Using these estimates of k1 in eq 6 shows that residual fraction of colloid-bound U after 1440 min of desorption exceeded 0.25 in all cases and was greater in the experiments with HA than in the experiments without HA (Figure 2D−F, dotted lines). We compared these time-series computations of colloid-bound residual fraction to those calculated under the assumption that interactions between metal oxide colloids and HA exerted negligible effects on U desorption (Figure 2D−F, solid lines). This comparison suggests that interactions between HA and metal-oxide colloids decreased the U desorption rate for SiO 2̊ HA2 and HFO̊ HA2, but increased the desorption rate for ZnO− HA1. The differences for the other treatments are comparatively small. We also calculated the average desorption rate (%/min) of U from colloids by dividing the total desorption fraction by the duration of the experiment. The results suggest that interactions between metal oxides and HAs decreased the average desorption rate of U by 35% and 33% for SiO2̊ HA2 and HFO̊ HA2, respectively, but increased the value for ZnOHA1 by 1.4 times (SI, Figure S4). A possible explanation for this behavior is that HA2 occupied sorption sites on SiO2 and HFO that would otherwise readily release bound U into solution, while the association between ZnO and HA1 made some slower-desorption sites unavailable for binding U. Interactions between HAs and Metal-Oxide Colloids and Their Impact on U Desorption. The average diameters for colloidal SiO2, ZnO and HFO were 19 ± 6, 32 ± 15 and 19 ± 11 nm, and size of colloidal HA1 and HA2 were much larger than metal oxides with the diameter of 105 ± 25 and 109 ± 29 nm, respectively (SI, Figure S5). There were no significant correlations between the size of colloids and the log Kd of U adsorption on colloids or the rate constants for U desorption from colloids. These results suggest that U adsorption and desorption was insensitive to colloid size, at least within our experiments. HAs can interact with colloidal metal oxides through electrostatic interactions, for which the surface charge is a key governing factor. The surface charge differed substantially among three metal oxides, with the average Z-potential values of −12.7 ± 0.7, −2.8 ± 6.8, and 31.1 ± 4.9 mV for SiO2, ZnO, and HFO, respectively (Figure 3). The Z-potential values indicate the experimental pH value was above the point of zero
Figure 3. Zeta potentials of single-colloid (homogeneous) suspensions of SiO2, ZnO, HFO, HA1, and HA2 and heterogeneous suspensions comprising HAs and metal oxides, measured under the same conditions as used in adsorption and desorption experiments.
charge (pHzpc) for SiO2 and ZnO, but lower than that of HFO, which is consistent with previous studies on the pHzpc for metal oxides.32 HAs had a negative surface charge with Z-potentials of −10.5 ± 0.5 and −28.5 ± 0.4 mV for HA1 and HA2, respectively. The opposite charges of HAs and HFO made their electrostatic attractive interactions the strongest among the colloidal mixtures tested. Previous calculations have demonstrated that the major aqueous U species under our experimental conditions were UO2CO3, UO2(CO3)22−, and UO2(CO3)34−, owing to the strong complexation between uranyl and carbonate.33 As a result, electrostatic attractive interactions may also play an important role in the adsorption of U by HFO, and this adsorption could be inhibited by HAs due to shielding of the positive surface charge on HFO through HA-HFO interactions. This proposed mechanism may account for the observed reduction in the amount of U adsorbed by HFO in experiments conducted in the presence of HA (Figure 1). Published studies have shown that HAs can affect the microbial reduction of U(VI);34 however, our previous experiments demonstrate that HAs cannot reduce U(VI) under the abiotic conditions tested in this study.15 ATR-FTIR spectra can provide additional information on the nature of sorptive associations between HA and metal colloids.11,24−26 There were limited signals detected in aqueous HAs, probably due to its low concentration of 5 mg/L (Figure 4). Only one peak at 2930 cm−1 corresponding to aliphatic carbon was detected in the spectra for HA2, but this peak was not detected in the spectra for HA1. This difference between HA1 and HA2 is consistent with the higher fraction of aliphatic carbon in HA2 compared to HA1, measured by 1H NMR, 13C NMR, and synchrotron-based analysis.14−16 A remarkable change in the ATR-FTIR spectra of HAs after interactions with metal colloids was a new peak emerging around 1440 cm−1 for HA-HFO, which has been attributed to the innersphere complexation of carboxyl and metals.24−26 This peak is more significant in HA1-HFO than HA2-HFO due to the higher fraction of carboxyl groups in HA1 as reported in previous studies.14−16 The interactions between Fe in HFO and carboxyl moieties in HAs were further supported by the EPR spectra analysis (SI, Figure S6). There were two fingerprint peaks at 1541 and 2859 G for iron in HFO colloids, and these peaks became sharper and their position shifted to 1536 and 3009 G for the heterogeneous mixture of HFO and HA1. The shift of the peak with higher G is significant, indicating the chemical micro2666
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Figure 4. Fourier Transform Infrared Spectroscopy in Attenuated Total Reflection Mode (ATR-FTIR) spectra of metal oxides and their mixture with HAs, under the same concentration and conditions as used in the adsorption and desorption experiments. The resolutions of spectra were relatively low, primarily because we used low concentrations of colloids, consistent with those used in our sorption and desorption experiments.
fractions of U bound with other functional groups, and decreased the desorption rate of U from the heterogeneous colloids (Figure 2). Our previous studies showed that HA2 has more alkyl carbons compared with HA1.14−16 This condition is consistent with the more significant impact of interactions between SiO2 and HA2 on U desorption relative to HA1. ZnO associated with both aromatic and aliphatic carbons in HA. Aromatic carbon in HA can interact with U through electron donor−receptor interactions, which are stronger than associations between U and aliphatic carbon, and aromatic carbon-bound U desorbed relatively slowly.15,35 The occupation of aromatic domains in HAs by associations between HAs and ZnO can decrease the fraction of U bound by aromatic domains and increase the desorption rate of U from ZnO-HA colloidal particles. The impact of ZnO-HA1 interaction on U desorption was more substantial than ZnO-HA2, consistent with the higher fraction of aromatic carbon in HA1.14,15 HFO can associate with HAs through electrostatic interactions, association with aliphatic or aromatic carbon, and specific inner-sphere complexation interactions with carboxyl functional groups. Nonspecific electrostatic interactions between HFO and HA only decreased the sorption of U on colloids, while not changing the desorption rate of colloid-bound U. Occupation of carboxyl functional groups on the HAs by their interactions with HFO may increase the desorption rate of U, as carboxyl functional group-bound U desorbed relatively slowly. The decreased desorption rate of U due to interactions between HA2 and HFO (Figure 2) was most probably due to the interactions between aliphatic carbon and HFO. Such effects were more predominant for HA2 compared to HA1, which is in accord with the higher fractions of aliphatic carbon in HA2 compared to HA1.14,15
environment of unpaired electrons of Fe(III) had been changed substantially. This change in unpaired electrons of Fe(III) suggests the unpaired electrons in HFO associated with HA1 through inner-sphere complexation, which is consistent with the ATR-FTIR analysis. In addition, the ITC analysis implies that the interaction between HFO and HA1 is an exothermic reaction (SI, Figure S6), which can facilitate the association between HFO and COM in the natural environment. The stronger interactions between HFO and carboxyl functional groups, together with the more substantial electrostatic interactions, are responsible for the greater impacts of interactions between HFO and HAs on the sorption of U compared to SiO2 or ZnO. Other significant changes in the ATR-FTIR spectra induced by the addition of HAs are the development of new peaks at 1620 cm−1 for HAs-ZnO and HAs-HFO. These peaks likely reflect the interactions between aromatic functional groups and metals through electron donor-receptor interactions and ligand exchange between metal oxides and phenolic functional groups.11,35 The peak at 2930 cm−1 diminished after complexation with metal oxides, reflecting the association between aliphatic carbon and metals. These metal oxides may associate with aliphatic carbon through van der Waals force or hydrogen bonds with the O-alkyl carbon. The ATR-FTIR spectra revealed that SiO2 interacted with aliphatic carbon in HAs through weaker interactions, possibly van der Waals force and hydrogen bonding. The aliphatic carbon can associate with a great amount of U, and the desorption of aliphatic carbon-sorbed U is relatively fast compared to U bound with other moieties.16 Therefore, the interactions between SiO2 and HAs made aliphatic sorption sites in HAs partially unavailable to U, led to the higher 2667
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(3) Miller, A. W.; Rodriguez, D. R.; Honeyman, B. D. Upscaling sorption/desorption processes in reactive transport models to describe metal/radionuclide transport: A critical review. Environ. Sci. Technol. 2010, 44, 7796−8007. (4) Missana, T.; Garcia-Gutierrez, M.; Alonso, U. Kinetics and irreversibility of cesium and uranium sorption onto bentonite colloids in a deep granitic environment. Appl. Clay Sci. 2004, 26, 137−150. (5) Sen, T. K.; Khilar, K. C. Review on subsurface colloids and colloid-associated contaminant transport in saturated porous media. Adv. Colloid Interfac. 2006, 119, 71−96. (6) de Jonge, L. W.; Kjaergaard, C.; Moldrup, P. Colloids and colloidfacilitated transport of contaminants in soils: An introduction. Vadose Zone J. 2004, 3, 321−325. (7) Novikov, A. P.; Kalmykov, S. N.; Utsunomiya, S.; Ewing, R. C.; Horreard, F.; Merkulov, A.; Clark, S. B.; Tkachev, V. V.; Myasoedov, B. F. Colloid transport of plutonium in the far-field of the Mayak production association, Russia. Science 2006, 314, 638−641. (8) Sharma, P.; Rolle, M.; Kocar, B.; Fendorf, S.; Kappler, A. Influence of natural organic matter on as transport and retention. Environ. Sci. Technol. 2011, 45, 546−553. (9) Geckeis, H.; Manh, T. N.; Bouby, M.; Kim, J. I. Aquatic colloids relevant to radionuclide migration: Characterization by size fractionation and ICP-mass spectrometric detection. Colloids Surf., A 2003, 217, 101−108. (10) Yoon, T. H.; Johnson, S. B.; Brown, G. E. Adsorption of suwannee river fulvic acid on aluminum oxyhydroxide surfaces: An in situ atr-ftir study. Langmuir 2004, 20, 5655−5658. (11) Yang, K.; Lin, D. H.; Xing, B. S. Interactions of humic acid with nanosized inorganic oxides. Langmuir 2009, 25, 3571−3576. (12) Bouby, M.; Geckeis, H.; Luetzenkirchen, J.; Mihai, S.; Schafer, T. Interaction of bentonite colloids with Cs, Eu, Th and U in presence of humic acid: A flow field-flow fractionation study. Geochim. Cosmochim. Acta 2011, 75, 3866−3880. (13) James, F. R.; David, J. C. Ronald, B. Particle-size and element distribution of soil colloids. Soil Sci. Soc. Am. J. 2005, 69, 1173−1184. (14) Yang, Y.; Shu, L.; Wang, X. L.; Xing, B. S.; Tao, S. Impact of demineralizing humic acid and humin on organic matter structural properties and sorption mechanisms of phenanthrene. Environ. Sci. Technol. 2011, 45, 3996−4002. (15) Yang, Y.; Saiers, J. E.; Xu, N.; Minasian, S. G.; Tyliszczak, T.; Kozimor, S. A.; Shuh, D. K.; Barnett, M. O. 2012. Impact of natural organic matter on uranium transport through saturated geologic materials: From molecular to column scale. Environ. Sci. Technol. 2012, 46, 5931−5938. (16) Wang, X. L.; Guo, X. Y.; Yang, Y.; Tao, S.; Xing, B. S. Sorption mechanisms of phenanthrene, lindane and atrazine with various humic acid fractions from a single soil sample. Environ. Sci. Technol. 2011, 45, 2124−2130. (17) Kalmykova, Y.; Rauch, S.; Stromvall, A. M.; Morrison, G.; Stolpe, B.; Hassellov, M. Colloid-facilitated metal transport in peat filters. Water Environ. Res. 2010, 82, 506−511. (18) Zachara, J. M.; Brown, C.; Christensen, J.; Davis, J. A.; Dresel, E.; Liu, C. X.; Kelly, S.; Mckinley, J.; Serne, J.; Um, W. A site-wide perspective on uranium geochemistry at the Hanford site. 2007. US Department of Energy Publications. Paper 286. (http:// digitalcommons.unl.edu/usdoepub/286). (19) Missana, T.; García-Gutiérrez, M.; Fernńdez, V. Uranium (VI) sorption on colloidal magnetite under anoxic environment: experimental study and surface complexation modeling. Geochim. Cosmochim. Ac. 2003, 67, 2543−2550. (20) Gu, B. H.; Schmitt, J.; Chen, Z.; Liang, L. Y.; Mccarthy, J. F. Adsorption and desorption of different organic-matter fractions on iron-oxide. Geochim. Cosmochim. Acta 1995, 59, 219−229. (21) Saleh, N. B.; Pfefferle, L. D.; Elimelech, M. Aggregation kinetics of multiwalled carbon nanotubes in aquatic systems: Measurements and environmental implications. Environ. Sci. Technol. 2008, 42, 7963− 7969. (22) Saleh, N. B.; Pfefferle, L. D.; Elimelech, M. Influence of biomacromolecules and humic acid on the aggregation kinetics of
Environmental Implications. Natural organic matter is ubiquitous in groundwaters and may occur at concentrations that exceed 20 mg/L.36,37 This common groundwater constituent plays an important role in governing the reduction, sorption, and transport of uranium and other actinides at contaminated sites, such as the Nevada Test Site, the Chernobyl Site and in the vicinities of uranium-mill tailing sites.36−38 HAs are also important for the mobilization of actinides from contaminated soils and recently HAs have been applied to attenuate U mobility in waste plumes.39,40 Similar to HAs, inorganic colloids can bind U and influence its mobility in groundwaters and soil waters. Our findings show that HAs interact with metal-oxide colloids and demonstrate that these interactions affect U adsorption and desorption in ways that depend on the compositions of the organic matter and mineral colloids. When taken in context with published studies, our results imply that inferences on processes that govern the fate and transport of U in subsurface environments, such as oxidation−reduction reactions and sorptive partitioning between aqueous and solid phases, will rely on thorough characterization of the compositional variability of the suspended particulate matter. Furthermore, faithful description of colloid-facilitated transport of U, and perhaps other radionuclides, in real subsurface environments, will require extending theoretical frameworks in order to account for the influences of HA-mineral colloid interactions on the adsorption and desorption of radionuclides from mobile particulate phases.
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ASSOCIATED CONTENT
S Supporting Information *
Parameters for simulating the sorption and desorption kinetics of U (Tables S1, S2); Scheme for experimental setup (Figure S1); other complementary information for U sorption and desorption (Figures S2−S6). This material is available free of charge via the Internet at http://pubs.acs.org.
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AUTHOR INFORMATION
Corresponding Author
*Tel: 203-9012423. E-mail:
[email protected]. Present Address
§ Department of Chemical and Environmental Engineering, Yale University, 9 Hillhouse Avenue, New Haven, Connecticut, 06511.
Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS This research was supported by the Office of Science (BER), U.S. Department of Energy Grant DE-FG02-08ER6463. We are grateful to the helpful comments of four reviewers and the associate editor, which led to substantial improvement of this manuscript. We thank Dr. Brudvig and Kari Brumback Young for help with EPR analysis.
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REFERENCES
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dx.doi.org/10.1021/es304013r | Environ. Sci. Technol. 2013, 47, 2661−2669
