Synergy between Zeolite Framework and Encapsulated Sulfur for

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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 Chem. Mater., Just Accepted Manuscript • DOI: 10.1021/acs.chemmater.8b02782 • Publication Date (Web): 02 Aug 2018 Downloaded from http://pubs.acs.org on August 9, 2018

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Chemistry of Materials

Eunhye Han,1 Young-Gu Kim,1 Hee-Man Yang,2 In-Ho Yoon,2 and Minkee Choi*,1 1

Department of Chemical and Biomolecular Engineering, Korea Advanced Institute of Science and Technology (KAIST), Daejeon 34141, Republic of Korea 2

Decommissioning Technology Research Division, Korea Atomic Energy Research Institute, 989-111 Daedukdaero, Yuseong, Daejeon 34057, Republic of Korea

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 ionexchanged Cs+ cations in zeolites, indicating significant Lewis acid-base interaction. According to the hard soft acid base (HSAB) theory, the enhanced Cs+ ion-exchange 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 sulfurmodified chabazite and mordenite showed highly promising Cs+ capture ability in the presence of various competing ions. nuclear sites, where immense heat and radioactivity are emitted.9-11 Conventional clays and zeolites are highly economic and commercially available, but they generally exhibit low Cs+ ion-exchange selectivities in the presence of various competing ions.8,12,13 More recently, new ionexchange 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 ion-exchange 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 low-cost ion-exchange materials with enhanced ion-exchange selectivity for Cs+ still remains a challenge.

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 longlived 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 sea water.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 137Cs+ capture because they are able to withstand the harsh conditions of

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In the present work, we demonstrate a facile strategy to remarkably increase the Cs+ ion-exchange selectivity of 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 (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 sulfur—a low-cost and abundant soft base—into 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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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 (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 pre-degassed 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+ ionexchange was carried out by stirring 1 g of Na+-form CHA and S-CHA samples in 2,000 mL of a 100 ppm Cs+ solution prepared by dissolving CsCl (99%, Tokyo Chemical Industry) in deionized water (18.3 MΩ·cm). Temperatureprogrammed 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 pre-degassing at 373 K for 5 h.

Scheme 1. Synthesis of Zeolites Encapsulating Sulfur for Enhanced Ion-Exchanged Selectivity towards Cs+

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, a 100 ppm 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 Cs+ solution, which was stirred at room temperature. Cs+ concentration of the supernatants collected after various equilibration times was analyzed by inductively coupled

EXPERIMENTAL SECTION Material Synthesis. NaA (Molecular Sieve 4A, Aldrich), NaX (Molecular Sieve 13X, Aldrich), and Namordenite (MOR) (CBV 10A, Zeolyst International) were purchased and used as received. Chabazite (CHA) was synthesized using a reported method,39 followed by ionexchange with 0.533 M NaNO3 solution three times at 328


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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 model:41 dqt/dt = k2(qe – qt)2

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 zeolites or 0.111 g Szeolite 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 = 2,000 mL gzeolite-1).

(1)

where k2 is the pseudo-second order rate constant (gzeolite mg-1min-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/k2qe2 + t/qe

RESULTS AND DISCUSSION

(2)

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.

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): q = (Ci – Ce)M/m

In Figure 1, the X-ray diffraction (XRD) patterns of CHA, S-CHA, 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 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 Szeolite composites (i.e., S-NaA, S-NaX, and S-MOR) showed XRD peaks only corresponding to the original zeolite structures (Figure S1), and no peak for sulfur was observed.

(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 = qmaxbCe/(1 + bCe)

(4)

where qmax is the maximum capacity of Cs+ (mg gzeolite-1), Ce is the equilibrium Cs+ concentration (ppm), and b is the Langmuir affinity 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 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 = 2,000 mL gzeolite-1. The Cs+ removal (%) and distribution coefficient (Kd, mL g-1) were calculated as follows: Cs+

removal (%) = (Ci – Ce)/Ci × 100

Kd (mL gzeolite

-1)

= V(Ci – Ce)/Cem

The pore structure analysis based on CO2 adsorption at 273 K revealed that the micropore volume of NaA, NaX, Table 1. Properties of Zeolites and S-Zeolites Containing 10 wt% Sulfur Sample NaA

(5)

S-NaA

(6)

NaX

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).

S-NaX CHA S-CHA

Cs+ ion-exchange experiments were also carried out using a simulated ground water solution, which was prepared by dissolving 1 ppm Cs+ in a background solution containing 125 ppm of Na+, 25 ppm of Ca2+, 10 ppm Mg2+, and 5 ppm of K+. This composition was similar to the ground water composition reported elsewhere.19 NaCl (99%, Samchun Chemicals), CaCl2·2H2O (99%, Junsei

MOR S-MOR aMicropore

Si/Al 1.0 1.2 2.0 6.5

Unit cell

Vmicroa

composition

(cm3 g-1)

Na12(AlO2)12(SiO2)12

0.243

S5.2Na12(AlO2)12(SiO2)12

0.019

Na86(AlO2)86(SiO2)106

0.308

S46.5Na86(AlO2)86(SiO2)106

0.148

Na8.9K2.1(AlO2)11(SiO2)22

0.219

S7.8Na8.9K2.1(AlO2)11(SiO2)22

0.092

Na6.4(AlO2)6.4(SiO2)41.6

0.219

S10.5Na6.4(AlO2)6.4(SiO2)41.6

0.098

volumes (Vmicro) were determined from the CO2 adsorption isotherms measured at 273 K using the DubininAstakhov method.



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

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Figure 2. HAADF-STEM image and EDS elemental mapping (Si is indicated as yellow and S is indicated as red) of (a) SNaA, (b) S-NaX, (c) S-CHA, and (d) S-MOR. The distributions of other elements (Na and Al) are also provided in Figure S2.

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 energy-dispersive 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 S-zeolite samples. No sulfur zoning at the exterior of zeolite crystallites was observed, indicating highly effective sulfur encapsulation in the 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 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 molS81),43 comparable to the heat of typical chemical reactions. Earlier crystallographic analysis of sulfur-loaded NaA indicated that two nonequivalent sulfur atoms alternated

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 12-membered 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.

Scheme 2. Interaction between Ionic Zeolite Surface and Highly Polarizable Sulfur


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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+ 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 sulfur peaks in S-CHA compared with those of elemental sulfur. This result confirms the strong polarization of elemental sulfur within zeolite micropores.

Table 2. Pseudo-Second Order Kinetics of the Cs+ IonExchange of Zeolites and S-Zeolite Samples qe (mg gzeolite-1)a

k (gzeolite mg-1 min-1)a

R2

qe (exp) (mg gzeolite-1)b

NaA

213

5.26 × 10-3

0.99

211

S-NaA

208

3.11 × 10-3

0.99

209

NaX

208

1.00 × 10-2

0.99

210

S-NaX

200

4.10 × 10-3

0.99

202

CHA

370

7.44 × 10-4

0.98

367

S-CHA

345

6.78 ×

10-4

0.99

345

MOR

270

1.05 × 10-2

0.99

267

S-MOR

250

8.42 × 10-3

0.99

250

aq e

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 pseudo-second order model. bqe (exp) is the experimentally determined qe value. bq e

Cs+ is Ion-Exchange Properties ofqe Sulfur-Loaded (exp) the experimentally determined value.

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. HAADFSTEM 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 XRD pattern of S-CHA did not change at all after treatment in water for 24 h, which confirmed the high structural stability of S-zeolite 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 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. 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.

Figure 3. S 2p XPS spectra of (a) elemental sulfur and (b) SCHA.

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 ioninduced dipole interaction between extra-framework cations (e.g., Na+ ions) and sulfur is the essential driving force for selective sulfur encapsulation in zeolite micropores.



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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 SCHA. (d) MOR and S-MOR. The solid lines indicate Langmuir fitting curves, the parameters of which are summarized in Table 3.

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-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 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).

Table 3. Langmuir Fitting Results of the Cs+ IonExchange Isotherms of Zeolites and S-Zeolite Samples

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) 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

qmax (mg gzeolite-1)

b (kg mg-1)

R2

NaA

289

0.0464

0.99

S-NaA

187

0.129

0.98

NaX

308

0.0577

0.99

S-NaX

231

0.254

0.97

CHA

428

1.78

0.99

S-CHA

346

6.16

0.99

MOR

239

2.25

0.99

S-MOR

196

9.51

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).


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Figure 5. Cs+ removal (%) and distribution coefficient (Kd) of 1 ppm 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 SCHA. (g), (h) MOR and S-MOR.



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 278–440%, after sulfur loading (Table 3). These results indicate that sulfur loading can generally increase the Cs+ selectivity of various aluminosilicate zeolites.

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the S-zeolite samples always exhibited larger Kd values and more complete Cs+ removal than the bare zeolites, whereas such behavior was more pronounced for Ca2+ than for Na+. This result shows that the encapsulated sulfur reduced the detrimental effects of the competing ions by increasing the selectivity to Cs+. This is consistent with the larger b values of S-zeolite composites compared with those of zeolites in the Langmuir fitting (Table 3), which indicated a higher affinity to Cs+. Conversely, there was little correlation between the maximum capacities of Cs+ in Langmuir fitting (qmax) and the Cs+ removal/distribution coefficients in the presence of competing ions. The results indicate that the Cs+ capture ability of ion-exchange materials in the presence of competing ions is more affected by the Cs+ selectivity than the maximum Cs+ capacity.

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 Szeolite 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 cationexchange 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 zeolites and the selectivity enhancement resulting from sulfur encapsulation.

The Cs+ ion-exchange behaviors of the samples in solutions containing Na+ and Ca2+ as competing cations were also investigated. The initial Cs+ concentration was fixed at 1 ppm, while the molar ratios of Na+/Cs+ and Ca2+/Cs+ were varied from 0 to 10000. As shown in Figure 5, the Cs+ removal (%) and distribution coefficients (Kd, mL g-1) of all samples monotonically decreased with increasing Na+/Cs+ and Ca2+/Cs+ ratios, indicating the detrimental effects of competing cations in Cs+ ion-exchange. Notably,

Origin of the Cs+ Selectivity Enhancement Resulting from Sulfur Encapsulation. Our ionexchange 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 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

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).

Figure 6. Cs+ distribution coefficients (Kd, bars) and removal (%, numbers above bars) of zeolite and S-zeolite samples in a simulated ground water solution containing 1 ppm Cs+, 125 ppm Na+, 25 ppm Ca2+, 10 ppm Mg2+, and 5 ppm K+.


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presence of substantial Lewis acid-base interaction between Cs+ and sulfur.

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).

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) 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 ionexchange selectivity toward Cs+ in the presence of various alkali and alkaline earth cations.

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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 ionexchange materials for Cs+ capture, but also for the removal of various toxic heavy metal cations that are even softer than Cs+.

Supporting Information. This material is available free of charge via the Internet at http://pubs.acs.org. Additional XRD patterns, HAADF-STEM images, SEM image, EDS elemental mappings, kinetics of Cs+ exchange, and tap densities of zeolites and S-zeolites

* E-mail: [email protected].

The authors declare no competing financial interest.



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