Article pubs.acs.org/cm
Cite This: Chem. Mater. 2018, 30, 5777−5785
Synergy between Zeolite Framework and Encapsulated Sulfur for Enhanced Ion-Exchange Selectivity to Radioactive Cesium Eunhye Han,† Young-Gu Kim,† Hee-Man Yang,‡ In-Ho Yoon,‡ and Minkee Choi*,† †
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Department of Chemical and Biomolecular Engineering, Korea Advanced Institute of Science and Technology (KAIST), Daejeon 34141, Republic of Korea ‡ Decommissioning Technology Research Division, Korea Atomic Energy Research Institute, 989-111 Daedukdaero, Yuseong, Daejeon 34057, Republic of Korea S Supporting Information *
ABSTRACT: To eliminate the radioisotope 137Cs+ from contaminated water, various inorganic ion-exchange materials have been developed. Many selective ion-exchange materials are relatively expensive and difficult to prepare, whereas conventional materials such as aluminosilicate zeolites lack ion-exchange selectivity in the presence of competing cations. Here, we report a simple but powerful strategy to significantly increase the Cs+ selectivity of conventional zeolites. We demonstrate that encapsulation of elemental sulfur in the micropores of zeolites (NaA, NaX, chabazite, and mordenite) via vacuum sublimation can remarkably increase the selectivity toward Cs+ in the presence of competing ions. It appears that the elemental sulfur does not provide additional adsorption sites for Cs+ ions but increases the ion-exchange selectivity toward Cs+ by providing additional interaction. Various analyses show that sulfur partially donates its electron to the ion-exchanged Cs+ cations in zeolites, indicating significant Lewis acid−base interaction. According to the hard soft acid base (HSAB) theory, the enhanced Cs+ ionexchange selectivity can be explained by the fact that sulfur, a soft Lewis base, interacts more strongly with Cs+, which is a softer Lewis acid than other alkali and alkaline earth metal cations. Because of the high intrinsic Cs+ selectivity of bare zeolites and selectivity enhancement resulting from sulfur encapsulation, the sulfur-modified chabazite and mordenite showed highly promising Cs+ capture ability in the presence of various competing ions.
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generally exhibit low Cs+ ion-exchange selectivities in the presence of various competing ions.8,12,13 More recently, new ion-exchange materials with improved Cs+ selectivities have been developed. Representative materials include silicotitanates,14−18 vanadosilicates,19 metal sulfides,11,20−31 and metal hexacyanoferrates.4,32−35 In particular, crystalline silicotitanate (CST) is a commercially developed ultrahighly selective ionexchange material that has been used for Hanford tank cleanup and also successfully used at Fukushima Daiichi.15 However, compared with conventional clays and zeolites, these new materials are relatively expensive and difficult to synthesize. Considering that these materials are not regenerated in general cleanup processes (they are directly stored as solid wastes), the cost of ion-exchange materials could be an important issue in practical applications. In this respect, the development of lowcost ion-exchange materials with enhanced ion-exchange selectivity for Cs+ still remains a challenge. In the present work, we demonstrate a facile strategy to remarkably increase the Cs+ ion-exchange selectivity of
INTRODUCTION Nuclear power generation has attracted much attention as a cost-effective and clean energy source that is free of greenhouse gas emissions, but it requires rigorous downstream treatment of radioisotopes produced during operation. 137Cs+ is one of the most hazardous and long-lived fission products (half-life of ∼30 years) that can produce high-energy gamma and beta particles.1,2 The high solubility of Cs+ ions in water allows them to spread rapidly into the environment through ground and seawater.3 Because contaminated water generally contains various competing ions (Na+, K+, Mg2+, and Ca2+) of much higher concentrations than that of 137Cs+, adsorbents must be able to remove Cs+ selectively. Low-concentration Cs+ in large volumes of contaminated water is generally concentrated into a small solid volume using ion-exchange materials with very high Cs+ selectivity.4−8 The solid wastes with concentrated 137Cs+ are then stabilized by mixing with cement or vitrification before disposal/storage.4−8 Inorganic ion-exchange materials have been widely studied for 137 Cs+ capture because they are able to withstand the harsh conditions of nuclear sites, where immense heat and radioactivity are emitted.9−11 Conventional clays and zeolites are highly economic and commercially available, but they © 2018 American Chemical Society
Received: July 2, 2018 Revised: August 1, 2018 Published: August 2, 2018 5777
DOI: 10.1021/acs.chemmater.8b02782 Chem. Mater. 2018, 30, 5777−5785
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Chemistry of Materials
(ramp: 1 K min−1) for 10 h. After the heat treatment, the sample was cooled slowly to room temperature and collected after cutting the sealed glass tube. The resultant materials were designated “S-zeolite”, where “zeolite” indicates the type of zeolite initially used for sulfur loading (NaA, NaX, CHA, or MOR). Material Characterization. Powder X-ray diffraction (XRD) patterns were recorded with a D2-PHASER (Bruker) operating at 30 kV and 10 mA with Cu Kα as an X-ray source. CO2 adsorption isotherms were measured with a Micromeritics ASAP 2050 at 273 K, and micropore volumes were estimated using the Dubinin−Astakhov method.40 Before the measurement, the zeolites were degassed under vacuum at 593 K. The S-zeolite composites were degassed at room temperature because the samples were already predegassed during the synthesis, and elevated temperatures may lead to the evaporation of elemental sulfur. High-angle annular dark-field scanning transmission electron microscopy (HAADF-STEM) images were acquired using a Talos F200X microscope (FEI) operating at 200 kV. Scanning electron microscope (SEM) images were acquired using a Magellan 400 (FEI). X-ray photoelectron spectroscopy (XPS) was carried out using a K-Alpha (Thermo Scientific) equipped with a microfocused monochromator X-ray source. Samples were mounted on adhesive copper tape, and the binding energies were calibrated with respect to the C 1s binding energy (284.8 eV). Before Cs 3d XPS analysis, full Cs+ ion-exchange was carried out by stirring 1 g of Na+-form CHA and S-CHA samples in 2000 mL of 100 ppm of Cs+ solution prepared by dissolving CsCl (99%, Tokyo Chemical Industry) in deionized water (18.3 MΩ·cm). Temperature-programmed desorption of sulfur from Na+- and Cs+-form S-CHA samples was carried out using a TGA N-1000 (Scinco) instrument at a temperature ramp of 2 K min−1 after predegassing at 373 K for 5 h. Cs+ Ion-Exchange Experiments. All Cs+ ion-exchange experiments were conducted in batch mode under magnetic stirring (400 rpm) at room temperature. To investigate the ion-exchange properties of Cs+ in the absence of competing cations, 100 ppm of Cs+ solution prepared by dissolving CsCl in deionized water was used. For kinetic analysis of Cs+ ion-exchange, 0.100 g of zeolites or S-zeolite samples was added to 400 mL of the 100 ppm of Cs+ solution, which was stirred at room temperature. Cs+ concentration of the supernatants collected after various equilibration times was analyzed by inductively coupled plasma mass spectroscopy (ICP-MS, Agilent ICP-MS 7700S) with 0.1 ppb level accuracy. The ion-exchange kinetics were analyzed using a pseudo-second-order model41
conventional aluminosilicate zeolites. Elemental sulfur was encapsulated within the micropores of various zeolites via vacuum sublimation to increase selectivity toward Cs + (Scheme 1). The strategy is based on the hard soft acid base Scheme 1. Synthesis of Zeolites Encapsulating Sulfur for Enhanced Ion-Exchanged Selectivity towards Cs+
(HSAB) theory, which is the generalized concept of Lewis acid−base theory, stating that a hard acid has a high affinity with a hard base, while a soft acid exhibits a high affinity with a soft base (“hard” species, in general, have small atomic radii, high effective nuclear charges, and low polarizability, whereas “soft” species possess the opposite characteristics).36−38 Because Cs+ is a relatively soft Lewis acid among the alkali and alkaline earth cations,38 we expected that introducing sulfura low-cost and abundant soft baseinto zeolites could increase the selectivity toward Cs+ in the presence of competing, relatively hard cations (e.g., Na+, Mg2+, and Ca2+). It is notable that the high Cs+ selectivities of recently developed metal sulfides11,20−31 and metal hexacyanoferrates4,32−35 can also be attributed to their soft frame anions, i.e., S2− and CN−, respectively.
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dqt /dt = k 2(qe − qt)2
(1)
where k2 is the pseudo-second-order rate constant (gzeolite mg−1 min−1); qe is the amount of Cs+ captured at equilibrium (mg gzeolite−1); and qt is the amount of Cs+ captured (mg gzeolite−1) at time “t” (min). The integrated form of the equation can be rearranged as follows: t /qt = 1/k 2qe 2 + t /qe
(2)
+
The Cs ion-exchange isotherms were collected using the 100 ppm Cs+ solution. Typically, 0.02−0.3 g of zeolite or S-zeolite samples was added to 200 mL of the solution, which was stirred at room temperature. After 3 h of equilibration, the Cs+ concentration in solution was analyzed using ICP-MS. The Cs+ uptake amount (q, mg gzeolite−1) was calculated by normalizing the mass of captured Cs+ by the mass of zeolite (not by the total mass of a composite)
EXPERIMENTAL SECTION
Material Synthesis. NaA (Molecular Sieve 4A, Aldrich), NaX (Molecular Sieve 13X, Aldrich), and Na-mordenite (MOR) (CBV 10A, Zeolyst International) were purchased and used as received. Chabazite (CHA) was synthesized using a reported method,39 followed by ion-exchange with 0.533 M NaNO3 solution three times at 328 K. Silicalite was also synthesized following a method reported in the literature.39 Prior to the encapsulation of elemental sulfur, all zeolites were degassed at 593 K (ramp: 2 K min−1) for 4 h in a plug-flow reactor under an Ar flow. After cooling to room temperature, 2 g of degassed zeolites was mixed with 0.222 g of elemental sulfur (10 wt % nominal sulfur loading). The mixtures were then introduced into a glass tube with an inner diameter of 3 cm. The glass tube was degassed under vacuum for 20 min and sealed using a torch under vacuum. The vacuum-sealed tube was heated at 593 K
q = (Ci − Ce)M /m
(3) +
where Ci is the initial Cs concentration (ppm); Ce is the equilibrium Cs+ concentration (ppm); M is the mass of the solution (kg); and m is the mass of the zeolite (gzeolite). Each isotherm was fitted using the Langmuir model q = qmax bCe/(1 + bCe)
(4) +
−1
where qmax is the maximum capacity of Cs (mg gzeolite ); Ce is the equilibrium Cs+ concentration (ppm); and b is the Langmuir affinity 5778
DOI: 10.1021/acs.chemmater.8b02782 Chem. Mater. 2018, 30, 5777−5785
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Chemistry of Materials constant (kg mg−1), which represents the affinity between Cs+ and solids. Cs+ ion-exchange experiments in the presence of competing Na+ and Ca2+ cations were conducted using solutions prepared by dissolving various concentrations of NaCl (99%, Samchun Chemicals) and CaCl2·2H2O (99%, Junsei Chemicals) in 1 ppm of Cs+ solution (molar ratios of Na+/Cs+ and Ca2+/Cs+ were controlled at 500, 1000, 2000, 4000, and 10000). For the ion-exchange experiments, 0.100 g of zeolites or 0.111 g of S-zeolite samples was added to 200 mL of the solutions, which were stirred at room temperature. The ratio between the volume of solution and the mass of zeolite was fixed at V/m = 2000 mL gzeolite−1. The Cs+ removal (%) and distribution coefficient (Kd, mL g−1) were calculated as follows
Cs+ removal (%) = (Ci − Ce)/Ci × 100
(5)
Kd (mL g zeolite−1) = V (Ci − Ce)/Cem
(6)
where Ci is the initial Cs+ concentration (ppm); Ce is the equilibrium Cs+ concentration (ppm); V is the volume of solution (mL); and m is the mass of the zeolite (gzeolite). Cs+ ion-exchange experiments were also carried out using a simulated groundwater solution, which was prepared by dissolving 1 ppm of Cs+ in a background solution containing 125 ppm of Na+, 25 ppm of Ca2+, 10 ppm of Mg2+, and 5 ppm of K+. This composition was similar to the groundwater composition reported elsewhere.19 NaCl (99%, Samchun Chemicals), CaCl2·2H2O (99%, Junsei Chemicals), MgCl2·6H2O (98%, Junsei Chemicals), KCl (99.5%, Junsei Chemicals), and CsCl (99%, Tokyo Chemical Industry) were used as cation sources. For the ion-exchange experiments, 0.100 g of zeolites or 0.111 g of S-zeolite samples was added to 200 mL of the solution, which was stirred at room temperature (the ratio between the volume of solution and the mass of the zeolite was fixed at V/m = 2000 mL gzeolite−1).
Figure 1. X-ray diffraction patterns of CHA, S-CHA, and a physical mixture of CHA and elemental sulfur (10 wt % sulfur). Arrows indicate the peaks corresponding to elemental sulfur.
corresponding to both elemental sulfur and CHA zeolite. This result indicates that vacuum sublimation can distribute elemental sulfur into a highly dispersed phase in the S-CHA sample, which is not detectable by XRD. Similar to the case of S-CHA, other S-zeolite composites (i.e., S-NaA, S-NaX, and SMOR) showed XRD peaks only corresponding to the original zeolite structures (Figure S1), and no peak for sulfur was observed. The pore structure analysis based on CO2 adsorption at 273 K revealed that the micropore volume of NaA, NaX, CHA, and MOR decreased significantly after the sulfur loading via vacuum sublimation (Table 1). Considering the small loading of sulfur (10 wt %), these significant decreases indicated selective sulfur encapsulation within zeolite micropores as a highly dispersed species. These results are also consistent with earlier XRD data. HAADF-STEM images and energydispersive X-ray spectroscopy (EDS) elemental mappings (Figure 2 and Figure S2) also confirmed that sulfur was uniformly distributed over the entire zeolite crystallites of all Szeolite samples. No sulfur zoning at the exterior of zeolite crystallites was observed, indicating highly effective sulfur encapsulation in the micropores of these zeolites. Under ambient conditions, elemental sulfur mainly exists as a cyclo-octasulfur form, S8, which has a molecular diameter of 6.9 Å. In large-pore zeolites such as NaX and MOR having 12membered pore apertures (pore diameter: 6.5−7.4 Å), the diffusion and encapsulation of elemental sulfur in micropores are not difficult to explain. However, it is quite surprising that elemental sulfur can also be selectively encapsulated within the small micropores of NaA and CHA having 8-membered pore apertures (3.8−4.0 Å). This could be explained by the final sublimation temperature (593 K) used in the present synthesis. It is known that an increasing fraction of sulfur can exist as S2 vapor as the temperature increases above 523 K.42 S2 has a molecular size of 3.7 Å and thus can readily diffuse into the small micropores of these zeolites. According to an earlier report by Barrer and Whiteman,43 sulfur adsorption in zeolites is fully reversible. However, the adsorption isotherms showed very steep sulfur uptake even at low sulfur pressures and elevated temperatures (530−600 K). These results showed that the nature of sulfur adsorption in zeolites is physical rather than chemical, even though its
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RESULTS AND DISCUSSION Encapsulation of Sulfur in Zeolite Micropores. Nominally 10 wt % elemental sulfur was loaded into NaA, NaX, chabazite (CHA), and mordenite (MOR) via vacuum sublimation at 593 K (see Experimental Section). The resultant composite materials were designated “S-zeolite”, where “zeolite” indicates the type of a zeolite initially used for sulfur hosting (NaA, NaX, CHA, and MOR). The basic properties of samples are summarized in Table 1. In Figure 1, the X-ray diffraction (XRD) patterns of CHA, SCHA, and a simple physical mixture of elemental sulfur and CHA are shown. The S-CHA sample showed only the peaks characteristic of the CHA structure, while no peak corresponding to elemental sulfur was observed. In contrast, the physical mixture of elemental sulfur and CHA showed XRD peaks Table 1. Properties of Zeolites and S-Zeolites Containing 10 wt % Sulfur sample
Si/Al
unit cell composition
Vmicroa (cm3 g−1)
NaA S-NaA NaX S-NaX CHA S-CHA MOR S-MOR
1.0
Na12(AlO2)12(SiO2)12 S5.2Na12(AlO2)12(SiO2)12 Na86(AlO2)86(SiO2)106 S46.5Na86(AlO2)86(SiO2)106 Na8.9K2.1(AlO2)11(SiO2)22 S7.8Na8.9K2.1(AlO2)11(SiO2)22 Na6.4(AlO2)6.4(SiO2)41.6 S10.5Na6.4(AlO2)6.4(SiO2)41.6
0.243 0.019 0.308 0.148 0.219 0.092 0.219 0.098
1.2 2.0 6.5
a
Micropore volumes (Vmicro) were determined from the CO2 adsorption isotherms measured at 273 K using the Dubinin−Astakhov method. 5779
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interacts with electron-rich sulfur species (δ−) via ion-induced dipole interaction and the electronegative O atoms (δ−) of a zeolite framework interact with electron-deficient sulfur species (δ+) via dipole−induced dipole interaction. As shown in Figure 3, the S 2p XPS spectra showed remarkable broadening of the
Figure 2. HAADF-STEM image and EDS elemental mapping (Si is indicated as yellow, and S is indicated as red) of (a) S-NaA, (b) SNaX, (c) S-CHA, and (d) S-MOR. The distributions of other elements (Na and Al) are also provided in Figure S2.
strength is unusually strong. The heats of sulfur adsorption on CaA and NaX zeolites were estimated to be 25−33 kcal molS−1 (200−264 kcal molS8−1),43 comparable to the heat of typical chemical reactions. Earlier crystallographic analysis of sulfurloaded NaA indicated that two nonequivalent sulfur atoms alternated to form the S8 ring in the zeolite micropores.44 The first group of sulfurs showed a S−Na distance of 2.80 Å, which corresponds to the sum of the van der Waals radius of S (1.85 Å) and the ionic radius of Na+ (0.95 Å). The second group of sulfurs showed a S−O distance of 3.21 Å, which is similar to the sum of the van der Waals radii of S (1.85 Å) and O (1.40 Å). These crystallographic results indicate the presence of intimate S−Na+ and S−O interactions. Considering the high polarizability of sulfur, the formation of an induced electric octupole in S8 is highly likely (Scheme 2),44 in which Na+
Figure 3. S 2p XPS spectra of (a) elemental sulfur and (b) S-CHA.
sulfur peaks in S-CHA compared with those of elemental sulfur. This result confirms the strong polarization of elemental sulfur within zeolite micropores. To further investigate the role of extra-framework metal cations (i.e., Na+) on sulfur encapsulation, we also tried to load sulfur into a purely siliceous zeolite, silicalite. As shown in the SEM image and EDS mapping (Figure S3), sulfur remained on the external surface of zeolite crystallites, even after the same vacuum sublimation procedure. The results clearly show that strong ion-induced dipole interaction between extra-framework cations (e.g., Na+ ions) and sulfur is the essential driving force for selective sulfur encapsulation in zeolite micropores. Cs+ Ion-Exchange Properties of Sulfur-Loaded Zeolites. The Cs+ ion-exchange properties of zeolites and S-zeolite samples were investigated at room temperature. It is noteworthy that, in the case of S-zeolite samples, the sulfur concentration of the solution obtained after each ion-exchange experiment was always undetectably small (our ICP-MS setup can detect sulfur levels of 10 ppb), indicating that the encapsulated sulfur in zeolite micropores is highly stable against leaching. HAADF-STEM images and EDS elemental mappings of the S-CHA sample after full Cs+ exchange (Figure S4) showed that all elements including Si, Al, S, and Cs were uniformly distributed over zeolite crystallites (no sulfur zoning). The result clearly shows that the S-zeolite composites are highly stable during the ion-exchange process in an aqueous environment. We also confirmed that the XRD pattern of S-CHA did not change at all after treatment in water
Scheme 2. Interaction between Ionic Zeolite Surface and Highly Polarizable Sulfur
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DOI: 10.1021/acs.chemmater.8b02782 Chem. Mater. 2018, 30, 5777−5785
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Chemistry of Materials for 24 h, which confirmed the high structural stability of Szeolite composites in water (Figure S5). As shown in Table 2, the zeolites and S-zeolite composites showed Cs+ ion-exchange kinetics that were well fitted with the
Nevertheless, all samples showed reasonable kinetics, and complete equilibrium could be reached within 3 h. Therefore, in the measurement of Cs+ ion-exchange isotherms afterward, we used an equilibration time of 3 h for each point. The Cs+ ion-exchange isotherms of the samples are shown in Figure 4. The isotherms were fitted with the Langmuir model (trend lines in Figure 4), and the fitting results are summarized in Table 3. It needs to be noted that, in the isotherms of S-
Table 2. Pseudo-Second-Order Kinetics of the Cs+ IonExchange of Zeolites and S-Zeolite Samples qe (mg gzeolite−1)a NaA S-NaA NaX S-NaX CHA S-CHA MOR S-MOR
213 208 208 200 370 345 270 250
k (gzeolite mg−1 min−1)a 5.26 3.11 1.00 4.10 7.44 6.78 1.05 8.42
× × × × × × × ×
10−3 10−3 10−2 10−3 10−4 10−4 10−2 10−3
R2
qe(exp) (mg gzeolite−1)b
0.99 0.99 0.99 0.99 0.98 0.99 0.99 0.99
211 209 210 202 367 345 267 250
Table 3. Langmuir Fitting Results of the Cs+ Ion-Exchange Isotherms of Zeolites and S-Zeolite Samplesa NaA S-NaA NaX S-NaX CHA S-CHA MOR S-MOR
qe is the amount of Cs+ captured at equilibrium (mg gzeolite−1), and k is the pseudo-second-order rate constant (gzeolite mg−1 min−1) obtained by fitting the Cs+ ion-exchange kinetics with a pseudosecond-order model. bqe (exp) is the experimentally determined qe value. a
qmax (mg gzeolite−1)
b (kg mg−1)
R2
289 187 308 231 428 346 239 196
0.0464 0.129 0.0577 0.254 1.78 6.16 2.25 9.51
0.99 0.98 0.99 0.97 0.99 0.99 0.99 0.99
Cs+ ion-exchange isotherms in Figure 4 were fitted with the Langmuir model, q = qmaxbCe/(1 + bCe), wherein qmax is the maximum capacity of Cs+ (mg gzeolite−1) and b is the Langmuir affinity constant (kg mg−1). a
pseudo-second-order kinetic model (kinetic data and fitting curves are provided in Figure S6 and Figure S7).41 The results showed that S-zeolite samples had somewhat decreased kinetic constants (k2 values in Table 2) compared with those of pristine zeolites. These results implied that the encapsulated sulfur retarded the Cs+ diffusion within zeolite micropores.
zeolite composites, the Cs+ uptake was normalized not by the total mass of the composite but by the mass of the zeolite in the composites (90 wt % with respect to the composite
Figure 4. Cs+ ion-exchange isotherms of zeolite and S-zeolite samples. (a) NaA and S-NaA. (b) NaX and S-NaX. (c) CHA and S-CHA. (d) MOR and S-MOR. The solid lines indicate Langmuir fitting curves, the parameters of which are summarized in Table 3. 5781
DOI: 10.1021/acs.chemmater.8b02782 Chem. Mater. 2018, 30, 5777−5785
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Figure 5. Cs+ removal (%) and distribution coefficient (Kd) of 1 ppm of Cs+ by zeolites and S-zeolite composites in the presence of Na+ (a, c, e, and g) and Ca2+ (b, d, f, and h) as competing cations. (a), (b) NaA and S-NaA. (c), (d) NaX and S-NaX. (e), (f) CHA and S-CHA. (g), (h) MOR and S-MOR.
weight) to obtain fundamental insight into the Cs+ capture mechanism. This calculation method is also practically meaningful because Cs+ uptake per a fixed solid “volume” is important in terms of column applications and subsequent solid waste disposal. Because sulfur is selectively encapsulated
within the micropores of zeolites, the packed solid volume per zeolite mass does not change after sulfur loading (Table S1). As shown in the isotherms (Figure 4), the bare CHA (Si/Al = 2.0) and MOR (Si/Al = 6.5) zeolites showed steeper Cs+ uptake than more Al-rich zeolites such as NaA (Si/Al = 1.0) 5782
DOI: 10.1021/acs.chemmater.8b02782 Chem. Mater. 2018, 30, 5777−5785
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Chemistry of Materials and NaX (Si/Al = 1.2). Consequently, the Langmuir affinity constants (b, Table 3) of CHA and MOR were an order of magnitude larger, indicating higher affinity to Cs+. This can be explained by the fact that too closely located ion-exchange sites (i.e., Al) within Al-rich zeolites cannot effectively exchange bulky Cs+ cations (ionic radius: 1.67 Å) because of steric repulsion. Compared with the corresponding bare zeolites, the S-zeolite composites always showed significantly steeper Cs+ uptakes at low Cs+ concentrations. Therefore, all zeolites showed an increase in Langmuir affinity constant (b), ranging from 278 to 440%, after sulfur loading (Table 3). These results indicate that sulfur loading can generally increase the Cs+ selectivity of various aluminosilicate zeolites. On the other hand, the maximum Cs+ uptake of zeolites at high equilibrium concentrations somewhat decreased after sulfur loading (Figure 4). As a result, the S-zeolite composites showed 18−35% decreases in the maximum capacity of Cs+ in Langmuir fitting (qmax, Table 3) compared with the bare zeolites. We confirmed that pure elemental sulfur did not show any detectable Cs+ uptake, indicating that sulfur cannot provide additional adsorption sites for Cs+ in addition to the original cation-exchange sites of zeolites. The decreased qmax after sulfur loading can be attributed to the partial blockage of zeolite micropores by the encapsulated sulfur, which can decrease the Cs+ accessibility. This is supported by the fact that the S-NaA sample, with almost zero residual pore volume (Table 1), showed the greatest loss of qmax (35%), whereas other S-zeolite samples, with higher residual pore volumes, showed less substantial loss of qmax (99.8%) because of the high intrinsic Cs+ selectivity of bare
Figure 6. Cs+ distribution coefficients (Kd, bars) and removal (%, numbers above bars) of zeolite and S-zeolite samples in a simulated groundwater solution containing 1 ppm of Cs+, 125 ppm of Na+, 25 ppm of Ca2+, 10 ppm of Mg2+, and 5 ppm of K+.
zeolites and the selectivity enhancement resulting from sulfur encapsulation. Origin of the Cs+ Selectivity Enhancement Resulting from Sulfur Encapsulation. Our ion-exchange experiments showed that the sulfur in zeolite micropores increased the Cs+ selectivity on the existing cation-exchange sites of zeolites by providing additional interaction with Cs+. To elucidate the nature of the interaction between Cs+ and sulfur, we investigated the Cs 3d XPS spectra of the CHA and S-CHA samples after full Cs+ ion-exchange (Figure 7a). The S-CHA
Figure 7. (a) Cs 3d XPS spectra of the CHA and S-CHA samples after full Cs+ exchange. (b) Thermogravimetric analysis (TGA) data of the S-CHA before (Na+-form) and after full Cs+ exchange (Cs+form).
sample showed a Cs 3d5/2 peak at significantly lower binding energy (−0.55 eV) than that of CHA, which implied that sulfur partially donated an electron to Cs+. This result supports the presence of substantial Lewis acid−base interaction between Cs+ and sulfur. We also carried out temperature-programmed desorption of sulfur from the Na+- and Cs+-form S-CHA samples under N2 flow (Figure 7b). For the full sulfur desorption, the Cs+-form S-CHA required substantially higher temperature (>627 K) 5783
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Chemistry of Materials than the Na+-form S-CHA (>593 K), confirming stronger Cs+−S interaction than Na+−S interaction. These results are consistent with the HSAB theory, in which sulfur is considered a soft base and thus is expected to exhibit higher affinity toward soft acids.36−38 Cs+ is chemically softer than other alkali and alkaline earth cations;38 thus, sulfur might preferentially increase the ion-exchange selectivity toward Cs+ in the presence of various alkali and alkaline earth cations.
(2) Yoshida, N.; Kanda, J. Tracking the Fukushima Radionuclides. Science 2012, 336, 1115−1116. (3) Avery, S. V. Caesium Accumulation by Microorganisms: Uptake Mechanisms, Cation Competition, Compartmentalization and Toxicity. J. Ind. Microbiol. 1995, 14, 76−84. (4) Haas, P. A. A. Review of Information on Ferrocyanide Solids for Removal of Cesium from Solutions. Sep. Sci. Technol. 1993, 28, 2479− 2506. (5) Todd, T. A.; Brewer, K. N.; Law, J. D.; Wood, D. J.; Garn, T. G.; Tillotson, R. D.; Tullock, P. A.; Wade, E. L. In Development and Application of Separation Technologies for the Treatment of Radioactive Wastes, Proceedings of the Waste Management Symposia; Tucson, AZ, March, 1997. (6) El-Kamash, A. M.; El-Naggar, M. R.; El-Dessouky, M. I. Immobilization of Cesium and Strontium Radionuclides in ZeoliteCement Blends. J. Hazard. Mater. 2006, 136, 310−316. (7) Andrews, M. K.; Feilinger, T. L.; Ferrara, D. M.; Harbour, J. R.; Herman, D. T. Vitrification of Cesium-Loaded Crystalline Silicotitanate (CST) in the Shielded Cells Melter; Technical Report for Westinghouse Savannah River Company Savannah River Technology Center, September, 1997. (8) Sinha, P. K.; Panicker, P. K.; Amalraj, R. V.; Krishnasamy, V. Treatment of Radioactive Liquid Waste Containing Caesium by Indigenously Available Synthetic Zeolites: A Comparative Study. Waste Manage. 1995, 15, 149−157. (9) Olatunji, M. A.; Khandaker, M. U.; Mahmud, H. N. M. E.; Amin, Y. M. Influence of Adsorption Parameters on Cesium Uptake from Aqueous Solutions- a brief review. RSC Adv. 2015, 5, 71658−71683. (10) Sylvester, P.; Clearfield, A. The Removal of Strontium and Cesium from Simulated Hanford Groundwater Using Inorganic Ion Exchange Materials. Solvent Extr. Ion Exch. 1998, 16, 1527−1539. (11) Manos, M. J.; Kanatzidis, M. G. Metal Sulfide Ion Exchangers: Superior Sorbents for the Capture of Toxic and Nuclear WasteRelated Metal Ions. Chem. Sci. 2016, 7, 4804−4824. (12) Mimura, H.; Kanno, T. Distribution and Fixation of Cesium and Strontium in Zeolite A and Chabazite. J. Nucl. Sci. Technol. 1985, 22, 284−291. (13) Borai, E. H.; Harjula, R.; malinen, L.; Paajanen, A. Efficient Removal of Cesium from Low-Level Radioactive Liquid Waste Using Natural and Impregnated Zeolite Minerals. J. Hazard. Mater. 2009, 172, 416−422. (14) Anthony, R. G.; Philip, C. V.; Dosch, R. G. Selective Adsorption and Ion Exchange of Metal Cations and Anions with Silico-Titanates and Layered Titanates. Waste Manage. 1993, 13, 503−512. (15) Nenoff, T. M.; Krumhansl, J. L. Cs+ Removal from Seawater by Commercially Available Molecular Sieves. Solvent Extr. Ion Exch. 2012, 30, 33−40. (16) Anthony, R. G.; Dosch, R. G.; Gu, D.; Philip, C. V. Use of Silicotitanates for Removing Cesium and Strontium from Defense Waste. Ind. Eng. Chem. Res. 1994, 33, 2702−2705. (17) Poojary, D. M.; Cahill, R. A.; Clearfield, A. Synthesis, Crystal Structures, and Ion-Exchange Properties of a Novel Porous Titanosilicate. Chem. Mater. 1994, 6, 2364−2368. (18) Celestian, A. J.; Kubicki, J. D.; Hanson, J.; Clearfield, A.; Parise, J. B. The Mechanism Responsible for Extraordinary Cs Ion Selectivity in Crystalline Silicotitanate. J. Am. Chem. Soc. 2008, 130, 11689− 11694. (19) Datta, S. J.; Moon, W. K.; Choi, D. Y.; Hwang, I. C.; Yoon, K. B. A Novel Vanadosilicate with Hexadeca-Coordinated Cs+ Ions as a Highly Effective Cs+ Remover. Angew. Chem., Int. Ed. 2014, 53, 7203−7208. (20) Manos, M. J.; Iyer, R. G.; Quarez, E.; Liao, J. H.; Kanatzidis, M. G. {Sn[Zn4Sn4S17]}6‑: A Robust Open Framework Based on MetalLinked Penta-Supertetrahedral [Zn4Sn4S17]10‑ Clusters with IonExchange Properties. Angew. Chem. 2005, 117, 3618−3621. (21) Manos, M. J.; Chrissafis, K.; Kanatzidis, M. G. Unique Pore Selectivity for Cs+ and Exceptionally High NH4+ Exchange Capacity
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CONCLUSION We demonstrated that the encapsulation of elemental sulfur in the micropores of zeolites via sublimation can significantly increase the ion-exchange selectivity toward Cs+ in the presence of various competing ions. It was confirmed that the encapsulated sulfur did not provide independent adsorption sites for Cs+ ions but rather increased the ion selectivity by providing additional Lewis acid−base interaction with Cs+. Various analyses showed that the elemental sulfur partially donated its electron to the ion-exchanged Cs+ cations in zeolites. According to the HSAB theory, the enhanced Cs+ ion selectivity can be explained by the fact that sulfur, one of the soft Lewis bases, provides electrons more efficiently to Cs+, which is a softer acid than other chemically harder alkali and alkaline earth metal cations. Because these composite materials were synthesized using commercially available zeolites and very cheap elemental sulfur, their syntheses are highly economic and scalable. We believe that the present strategy will be promising not only for the design of ion-exchange materials for Cs+ capture but also for the removal of various toxic heavy metal cations that are even softer than Cs+.
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ASSOCIATED CONTENT
* Supporting Information S
The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.chemmater.8b02782. Additional XRD patterns, HAADF-STEM images, SEM image, EDS elemental mappings, kinetics of Cs + exchange, and tap densities of zeolites and S-zeolites (PDF)
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AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. ORCID
Minkee Choi: 0000-0003-0827-2572 Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS This work was supported by the National Research Foundation of Korea grant funded by the Korean government (MSIP) (NRF-2017M2A8A501548). This work was also supported by Basic Science Research Program through the National Research Foundation of Korea (NRF2017R1A2B22002346).
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REFERENCES
(1) Delacroix, D.; Guerre, J. P.; Leblanc, P.; Hickman, C. Radionuclide and Radiation Protection Data Handbook 2002. Radiat. Prot. Dosim. 2002, 98, 1−168. 5784
DOI: 10.1021/acs.chemmater.8b02782 Chem. Mater. 2018, 30, 5777−5785
Article
Chemistry of Materials of the Chalcogenide Material K6Sn[Zn4Sn4S17]. J. Am. Chem. Soc. 2006, 128, 8875−8883. (22) Ding, N.; Kanatzidis, M. G. Permeable Layers with Large Windows in [(CH3CH2CH2)2NH2]5In5Sb6S19·1.45 H2O: High IonExchange Capacity, Size Discrimination, and Selectivity for Cs Ions. Chem. Mater. 2007, 19, 3867−3869. (23) Feng, M.-L.; Kong, D.-N.; Xie, Z.-L.; Huang, X.-Y. ThreeDimensional Chiral Microporous Germanium Antimony Sulfide with Ion-Exchange Properties. Angew. Chem., Int. Ed. 2008, 47, 8623− 8626. (24) Manos, M. J.; Kanatzidis, M. G. Highly Efficient and Rapid Cs+ Uptake by the Layered Metal Sulfide K2xMnxSn3‑xS6 (KMS-1). J. Am. Chem. Soc. 2009, 131, 6599−6607. (25) Ding, N.; Kanatzidis, M. G. Selective Incarceration of Caesium Ions by Venus Flytrap Action of a Flexible Framework Sulfide. Nat. Chem. 2010, 2, 187−191. (26) Mertz, J. L.; Fard, Z. H.; Malliakas, C. D.; Manos, M. J.; Kanatzidis, M. G. Selective Removal of Cs+, Sr2+, and Ni2+ by K2xMgxSn3−xS6 (x = 0.5−1) (KMS-2) Relevant to Nuclear Waste Remediation. Chem. Mater. 2013, 25, 2116−2127. (27) Qi, X.-H.; Du, K.-Z.; Feng, M.-L.; Li, J.-R.; Du, C.-F.; Zhang, B.; Huang, X.-Y. A Two-Dimensionally Microporous Thiostannate with Superior Cs+ and Sr2+ Ion-Exchange Property. J. Mater. Chem. A 2015, 3, 5665−5673. (28) Yang, H.; Luo, M.; Luo, L.; Wang, H.; Hu, D.; Lin, J.; Wang, X.; Wang, Y.; Wang, S.; Bu, X.; Feng, P.; Wu, T. Highly Selective and Rapid Uptake of Radionuclide Cesium Based on Robust Zeolitic Chalcogenide via Stepwise Ion-Exchange Strategy. Chem. Mater. 2016, 28, 8774−8780. (29) Kanatzidis, M. G.; Mertz, J. L.; Manos, E. Chalcogenide Compounds for the Remediation of Nuclear and Heavy Metal Wastes. U.S. Patent US 9056263 B2, June 16, 2015. (30) Sarma, D.; Malliakas, C. D.; Subrahmanyam, K. S.; Islam, S. M.; Kanatzidis, M. G. K2xSn4‑xS8‑x (x = 0.65−1): A New Metal Sulfide for Rapid and Selective Removal of Cs+, Sr2+ and UO22+ Ions. Chem. Sci. 2016, 7, 1121−1132. (31) Zhang, B.; Feng, M.-L.; Cui, H.-H.; Du, C.-F.; Qi, X.-H.; Shen, N.-N.; Huang, X.-Y. Syntheses, Crystal Structures, Ion-Exchange, and Photocatalytic Properties of Two Amine-Directed Ge−Sb−S Compounds. Inorg. Chem. 2015, 54, 8474−8481. (32) Lin, Y.; Fryxell, G. E.; Wu, H.; Engelhard, M. Selective Sorption of Cesium Using Self-Assembled Monolayers on Mesoporous Supports. Environ. Sci. Technol. 2001, 35, 3962−3966. (33) Sangvanich, T.; Sukwarotwat, V.; Wiacek, R. J.; Grudzien, R. M.; Fryxell, G. E.; Addleman, R. S.; Timchalk, C.; Yantasee, W. Selective Capture of Cesium and Thallium from Natural Waters and Simulated Wastes with Copper Ferrocyanide Functionalized Mesoporous Silica. J. Hazard. Mater. 2010, 182, 225−231. (34) Hu, M.; Furukawa, S.; Ohtani, R.; Sukegawa, H.; Nemoto, Y.; Reboul, J.; Kitagawa, S.; Yamauchi, Y. Synthesis of Prussian Blue Nanoparticles with a Hollow Interior by Controlled Chemical Etching. Angew. Chem., Int. Ed. 2012, 51, 984−988. (35) Torad, N. L.; Hu, M.; Imura, M.; Naito, M.; Yamauchi, Y. Large Cs Adsorption Capability of Nanostructured Prussian Blue Particles with High Accessible Surface Areas. J. Mater. Chem. 2012, 22, 18261−18267. (36) Pearson, R. G. Hard and Soft Acids and Bases. J. Am. Chem. Soc. 1963, 85, 3533−3539. (37) Pearson, R. G. Acids and Bases. Science 1966, 151, 172−177. (38) Pearson, R. G. Absolute Electronegativity and Hardness: Application to Inorganic Chemistry. Inorg. Chem. 1988, 27, 734−740. (39) Robson, H.; Lillerud, K. P. In Verified Syntheses of Zeolitic Materials, 2nd revised ed.; International Zeolite Association: Amsterdam, Netherlands, 2001; pp 126, 201. (40) Hutson, N. D.; Yang, R. T. Theoretical Basis for the DubininRadushkevitch (D-R) Adsorption Isotherm Equation. Adsorption 1997, 3, 189−195. (41) Ho, Y. S.; McKay, G. Pseudo-Second Order Model for Sorption Processes. Process Biochem. 1999, 34, 451−465.
(42) Meyer, B. Elemental Sulfur. Chem. Rev. 1976, 76, 367−388. (43) Barrer, R. M.; Whiteman, J. L. Intracrystalline Sorption of Sulphur and Phosphorus by Some Porous Crystals. J. Chem. Soc. A 1967, 13−18. (44) Seff, K. The Crystal Structure of a Sulfur Sorption Complex of Zeolite 4A. J. Phys. Chem. 1972, 76, 2601−2602.
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DOI: 10.1021/acs.chemmater.8b02782 Chem. Mater. 2018, 30, 5777−5785
