Comparative Investigation of Fe2O3 and Fe1âxS Nanostructures for

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Comparative Investigation of Fe2O3 and Fe1xS Nanostructures for Uranium Decontamination Ran Ma, Ling Yin, Lei Li, Sai Zhang, Tao Wen, Chenlu Zhang, Xiangxue Wang, Zhongshan Chen, Tasawar Hayat, and Xiangke Wang ACS Appl. Nano Mater., Just Accepted Manuscript • DOI: 10.1021/acsanm.8b01059 • Publication Date (Web): 03 Oct 2018 Downloaded from http://pubs.acs.org on October 6, 2018

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ACS Applied Nano Materials

Comparative Investigation of Fe2O3 and Fe1-xS Nanostructures for Uranium Decontamination Ran Ma,† Ling Yin,† Lei Li,† Sai Zhang,† Tao Wen,†,* Chenlu Zhang,† Xiangxue Wang,† Zhongshan Chen,† Tasawar Hayat,‡ Xiangke Wang†,‡* †

College of Environment Science and Engineering, North China Electric Power

University, Beijing 102206, P.R. China ‡

NAAM Research Group, Faculty of Science, King Abdulaziz University, Jeddah,

21589, Saudi Arabia *: Corresponding author. Fax (Tel): +86-10-61772890; Email: [email protected] (T. Wen); [email protected] (X. Wang).

1

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ABSTRACT: Nowadays, more and more researchers pay close attention to the decontamination of radioactive uranium because of its high toxicity and long half-life in the environment. Although many strategies for the development of highly effective materials have been reported, the narrow pH operating range and high salinity seriously limit their widespread application in radioactive waste water treatment. With the purpose of mitigating these drawbacks, herein we synthesized highly uniform iron sulfide (Fe1-xS) through heating vulcanization of iron oxide (Fe2O3) microcubes. A comparative study of U(VI) removal on Fe1-xS and Fe2O3 was systemically investigated under a series of environmental conditions. Batch experiments illustrated that the U(VI) sorption activities on Fe2O3 and Fe1-xS nanostructures were different under various pH conditions, and U(VI) sorption on Fe1-xS exhibited strong tolerance to ionic strength compared to that on Fe2O3 sample. The sorption of U(VI) on Fe2O3 and Fe1-xS could reach equilibrium rapidly within a few minutes, which presented as a spontaneous endothermic process.. The competitive Ca2+ ions presented in relative high concentrations could strongly affect the distribution coefficients (Kd) of U(VI) on Fe2O3 and Fe1-xS owing to the forming of ternary Ca-U-carbonate complexes, whereas the Kd values were scarcely influenced by Mg2+ concentrations. Furthermore, the application of Fe1-xS in different wastewater treatment presented superior sorption performance than Fe2O3. Combining the FTIR and XPS analysis, the enrichment of U(VI) on Fe1-xS was mainly attributed to the formation of UO22+···S2- bonding. Besides, we compared the sorption capacity of Fe1-xS with other systems. The studies further revealed that Fe1-xS could be served as a promising sulfide-based scavenger in 2


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the effective removal of U(VI) from various wastewater samples. KEYWORDS: Fe2O3; Fe1-xS; U(VI); Sorption; Covalent bonds

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INTRODUCTION As one of the most common and clean fission fuels, uranium has become popular with the rapid development of nuclear technology and adhibition of nuclear energy.1 However, the utilization efficiency of uranium fuel is extremely low owing to the limitation of technology, resulting in the generation of considerable amounts of nuclear waste (approximately 2300 t) per year.2 In addition, the abundant isotopes of uranium have long half-lives (U235, 7.13×108 years; U238, 4.51×109 years).3 In general, uranium is regarded as the last heavy meal element in natural environment, which causes detrimental effects on human health and aquatic life, on the other hand, uranium and its byproducts can also produce various radioactive particles,4 such as the short range of alpha particles, high energy beta and gamma particles, as a result of disrupting DNA sequences, resulting cell death, giving rise to genetic variation and even causing cancers.5 Therefore, developing progressive technologies and materials for highly-performance uranium removal in aqueous solutions remains a challenge. Nowadays,

various

technologies

have

been

applied

to

dispose

uranium-contaminated water, such as ion exchange,6 photocatalytic reduction,7 electrochemical redox process,8 coagulation method9 and so on. Because these technologies have some disadvantages, they cannot be widely used in actual waste water treatment due to the harsh conditions requirement for ion exchange, lower utilization of solar energy for photocatalytic reduction, higher cost for electrochemical redox as well as longer reaction time for coagulation method. Compared with the aforementioned technologies, sorption technique is considered to be one of the most 4


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efficient methods owing to its cost-effective, simple to operate and wide feasibility. Therefore, multiple sorbents have been conducted to capture U(VI) from wastewater. Some traditional sorbents including active carbon,10 diatomite1,11 bentonite,12 fly ash13 and so on possess outstanding sorption capacity owing to their large surface area and rich porosity, and other nanomaterials like nano-Fe3O4 14 and graphene oxides15 also exhibit a remarkable sorption performance for their oxygen-containing functional groups. Unfortunately, the O2- ligands could be easily protonated (≡SOH + H+ → ≡SOH2+) under relatively lower pH conditions. And these protonated materials are inefficient for uranium scavenging on account of electrostatic repulsion. In recent years, some metal sulfides have been manufactured as efficient U(VI) sorbents to overcome the limitation of O2- ligands, such as layered metal sulfides KMS-1, polysulfide/layered double hydroxide composites Sx-LDH, FeS2.16, 17 Based on Lewis acid sense, uranyl ions are regarded as hard cations whereas S2- ions are soft ligands. Nevertheless, U(VI) is regarded as a softer Lewis acid center, in which the strong bonding compounds of S2- with UO22+ could attenuate U(VI) efficiently.18-20 However, the synthesis these metal sulfides is complicated and time-consuming, so simple preparation methods need to be developed. Motivated by the above considerations, we herein presented a large scale and facile synthesis of Fe1-xS microcubes by sulfurizing Fe2O3. Both of the two materials were performed to decrease U(VI) concentration in aqueous solution. The structures of as-prepared Fe2O3 and Fe1-xS were analyzed by the typical scanning electron microscopy (SEM), transmission electron microscopy (TEM) and TEM element 5



mapping, Brunauer–Emmett–Teller (BET) analysis, X-ray diffraction (XRD), and magnetic properties. To explore the functional groups on the surface and understand the valence structures of Fe2O3 and Fe1-xS, X-ray photoelectron spectroscopy (XPS) and Fourier transform infrared (FTIR) analysis were carried out, respectively. The comparative sorption behaviors of U(VI) on sulfide- and oxygen-based materials were systematically investigated under desired experimental conditions. The integrated features of excellent sorption capacity, high salt tolerance and easy magnetic separation demonstrated the outstanding sorption performance of U(VI) on Fe1-xS compared with that of U(VI) on Fe2O3. Furthermore, the materials were adopted to attenuate U(VI) concentrations from various wastewater system. EXPERIMENTAL SECTION Synthesis of Fe2O3 Microcubes Precursor. The used reagents were purcharsed in analytical grade without further purification. Based on a previous study,21 Fe2O3 microcubes were fabricated via a facile precipitation method. Typically, the distilling flask containing 50 mL of FeCl3 (2 M) was dissolved in a preheated oil bath until the temperature increased to 75 oC. Then, 50 mL of NaOH (5.4 M) was poured into the above suspension under vigorous stirring. The resulting solution was further stirred for 5 min and refluxed at 100 oC for 4 d, followed by withdrawing via centrifugation and rinsing with deionized water and ethanol for several times until the liquid supernatant was near neutral. The final products were obtained after fully dring the materials under vacuum at 60 oC overnight. 6


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Synthesis of Fe1-xS Microcubes. Firstly, the Fe2O3 precursors were mixed homogeneously with thioacetamide (TAA, with weight ratio of 1:3). After that, the resulting mixture was transfered into a sealed alumina boat. Finally, the product was obtained by directly annealing the mixture at 500 oC for 3 h under high purity Ar atmosphere (99.999%) with a heating rate of 1 oC/min. The synthesis of Fe2O3 and Fe1-xS were schematically illustrated in Scheme 1. Material Characterization. The SEM images were observed by JEOL JSM-6330F (Japan). TEM images, TEM element mapping images and energy dispersive X-ray spectroscopy (EDS) were performed on a Hitachi-7650 microscope (Japan). The XRD patterns were achieved on a Rigaku D/max-2500 apparatus with Cu-Kα radiation with the scan rate

of

2θ

=

0.05

o

/s

(λ

=

1.54178

Å). According

to

the

BET

(Barrett–Emmett–Teller) method, the N2 adsorption-desorption isotherms were measured by Micromeritics 3Flex system. Magnetic measurements were recorded on an MPMS-XL SQUID magnetometer. The FTIR spectra were obtained by Bruker Tensor 27 FTIR spectrophotometer, and XPS survey was collected by ESCALAB Mark II spectrometer. The Zeta-potential values were obtained by ZETASIEZER 3000 HSA. Batch Sorption Experiments. The U(VI) sorption performances on Fe2O3 and Fe1-xS were carried out by batch technique under different environmental conditions. In a typical procedure, 0.1 M 7



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NaNO3, 0.3 g L-1 Fe2O3 or Fe1-xS and 10 mg L-1 U(VI) solution were added into the 10 mL polyethylene tubes to obtain desired concentrations, and then adjusted the pH values by adding different concentrations of NaOH and HNO3 solutions. The sorption isotherms and thermodynamic experiments at different temperatures were conducted under increasing initial U(VI) concentrations (from 1 mg L-1 to 60 mg L-1 for Fe1-xS and from 1 mg L-1 to 30 mg L-1 for Fe2O3). Besides, sorption kinetics of U(VI) (10 mg L-1) on Fe2O3 and Fe1-xS were measured to evaluate the effect of contact time (0 ~ 240 min). The sorption effect of coexisting ions (i.e. Ca2+ and Mg2+) and different water samples were also investigated. The abovementioned suspensions were shaken for 240 min to attain sorption equilibration, and then the materials and liquid phase were segregated by high-speed centrifugation at 8000 rpm. The remaining U(VI) concentrations were determined by inductively coupled plasma mass spectrometry (ICP-MS). The sorption percentage (%), sorption capacity (qe, mg g-1) and distribution coefficient (Kd, mL g-1) were defined by the following equations:

Sorption(%) =

(C0 − Ce ) × 100% C0

(1)

qe (mg • g −1 ) =

V × (C0 − Ce ) m

(2)

K d (mL • g −1 ) =

C 0 − Ce V × Ce m

(3)

where C0 (mg L-1) and Ce (mg L-1) were the initial U(VI) concentration and equilibrium concentration after reaction. m (mg) and V (mL) were the quality of sorbents and the total volume of the solution, respectively. RESULTS AND DISCUSSION 8


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Characterization of Fe2O3 and Fe1-xS. The SEM and TEM images of Fe2O3 and Fe1-xS samples were schematically described in Figure 1. As shown in Figure 1a and 1b, the morphology of Fe2O3 at different magnifications was perfect microcubes with an average size dimension ~ 500 nm, while the Fe1-xS inherited cubic structures with a lesser extent of aggregation, and the surface of Fe1-xS became rough after sulfurization (Figure 1c and 1d). Besides, the size of Fe1-xS was almost invariable after react with TAA under 500 oC anneal. To gain the insight into the interior structures of Fe2O3 and Fe1-xS, both products were further examined by TEM, EDS and the corresponding element mapping images. The TEM image of Fe2O3 presented the single cube (Figure 1e). After reacting with TAA at 500 oC (Figure 1g), the Fe2O3 microcubes were converted into Fe1-xS with interior pores, revealing that the sulfurizing process resulted in the collapse of interior structure. The EDS analysis determined that the Fe/O and Fe/S weight ratios were ~2.3 and ~1.6 for Fe2O3 and Fe1-xS samples, respectively (Figure 1f and 1h). Notably, the x value was calculated as 0.067, in accordance with pyrrhotite (0 ≤ x ≤ 0.125). In addition, the corresponding TEM element mapping images (Figure 1i and 1j) diagrammed the homogenous distribution of Fe and O elements in Fe2O3 and Fe and S elements on Fe1-xS microstructures, indicating the successful sulfuration of iron oxide all over the whole sample. The crystal phases of as-prepared Fe2O3 and Fe1-xS samples were analyzed by XRD patterns (Figure 2a). The diffraction peaks of Fe2O3 at 2θ = 24.1°, 33.2°, 9



35.6°, 40.9°, 49.5°, 54.1°, 62.5° and 64.0° were accorded with the (012), (104), (110), (113), (024), (116), (214) and (300) facets of hematite (JCPDS NO.33-0664). After sulfurization, the characteristic peaks of Fe1-xS at 2θ = 29.8°, 33.8°, 43.7° and 53.1° corresponded to the (200), (2011), (2022) and (220) facets of pyrrhotite (JCPDS NO.29-0726). The peaks of Fe1-xS were in good agreement with pyrrhotite phase without distinct other phases of Fe2O3, demonstrating that the Fe1-xS product was perfectly synthesized in this work. Interestingly, the peaks of Fe1-xS still existed after U(VI) adsorption (Figure S1a). While some peaks of Fe2O3 appeared after the Fe1-xS the reaction with U(VI), indicating that partial Fe(II) was oxidized and formed the corrossion byproduct. However, the peaks of Fe1-xS were still remained as dominated signal, and the phenomenon could be interpreted as the surface oxidization of Fe1-xS. The presence of O2 in the system should be eliminated beforehand to acquire high perfomance of Fe1-xS towards U(VI). In order to further understand the structural properties of the iron sulfide phase, the Rietveld refinement of XRD data was conducted by using FullPROF program (Figure S1b). The corresponding lattice parameters a and c of Fe1-xS were 6.899 Å and 63.211 Å, respectively. Interestingly, the refined values fit well with JCPDS card data of pyrrhotite (a = 6.896 Å and c = 63.222 Å), which was consistent with the previous study. 21 According to the magnetization curves (Figure 2b), the sulfuration process significantly enhanced the magnetic property of Fe1-xS (Ms = 2.1 emu g-1), which was much higher than that of Fe2O3 (Ms = 0.4 emu g-1). Interestingly, the enhanced 10


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magnetism of Fe1-xS made it easier for separation from aqueous solution (inset of Figure 2b). According to the IUPAC classification, the N2 adsorption-desorption isotherms of Fe1-xS measured at 77.35 K were type IV characteristic with a H3 hysteresis loop for mesoporous architecture (Figure 2c). For comparison, the isotherm curve of Fe2O3 demonstrated type III without distinct hysteresis loop, indicating that the interaction between sorbent and adsorbate was very weak.22, 23 Notably, the surface area of Fe1-xS (12.67 m2 g-1) was much lower than that of Fe2O3 (26.35 m2 g-1). However, the surface area of the as-prepared sorbent is one of the possible factors, but not a crucial factor to afford the high-performance elimination capability. In addition, the pore size of Fe1-xS was broadly distributed over the range of 12-80 nm and exhibited a hierarchically porous structure with a meso/macro multimodal pore size distribution, while the pore size distribution of Fe2O3 showed a narrow pore diameter at 58 nm, which signified that such hierarchically porous Fe1-xS could provide more accessible active sites than macroporous Fe2O3 to capture U(VI) from contaminated water.24 FTIR spectra of Fe2O3 (Figure 3a) and Fe1-xS (Figure 3b) before and after U(VI) sorption were investigated to understand its structural characterization. The FTIR spectra of Fe2O3 showed that the strong and wide peak at 3430 cm-1 was assigned to the –OH groups associated with the hydrogen bonding and adsorbed water molecules (Figure 3a). The peak at 1628 cm-1 was considered as H-O-H stretching vibration, which was ascribed to the bending modes of H2O. Obviously, the intensity of Fe-O peaks at 582 and 480 cm-1 were significantly enhanced after U(VI) sorption, which 11



indicated that sorption was mainly dominated by surface complexation (Fe-O···UO22+).25, 26 However, the absence of these peaks in the FTIR spectra of Fe1-xS further demonstrated the successful sulfuration of Fe2O3 (Figure 3b). In addition, one can see that the antisymmetric vibration of UO22+ at ~ 908 cm-1 for Fe2O3 and ~ 902 cm-1 for Fe1-xS appeared after U(VI) sorption, indicating the succssful sorption of U(VI) on the surface of Fe2O3 and Fe1-xS. The XPS survey (Figure 3c) further proved that U(VI) was adsorbed on the surface of Fe2O3 and Fe1-xS. From the spectra of U 4f (Figure 3d), the binding energy values at U 4f7/2 (381.6 eV, 52.4%) and U 4f5/2 (392.4 eV, 44.3%) for Fe2O3 and at 382.01 eV (59.1%) and 385.20 eV (37.1%) for Fe1-xS were splited owing to the unpaired electrons spin in the atomic shell, which were attributed to the hexavalent state of UO22+ onto the sorbents rather than the redox activity during the interaction with sulfide groups.27-31 Besides, the peak at 385.2 eV was index to the satellite signal.32 As shown in Figure 3e, the peak of O 1s in Fe2O3 before sorption was fitted with two peaks at 529.85 eV and 531.45 eV, corresponding to O2- anions (74.1%) in metal oxide lattice and surface OH- groups (25.9%), respectively.28 Interestingly, the peak at lower binding energy slightly shifted to 529.92 eV (73.3%) and the higher one shifted to 531.62 eV (26.7%) after U(VI) sorption. This phenomenon can be attributed to a new chemical bond (Fe-O-U(VI)), in which the coordination interaction could trigger O electron cloud shift towards U 6d.33 For comparison, the S 2p of Fe1-xS can be divided into four peaks (Figure 3f): i.e., the first at 161.4 eV corresponded to S2- bounded with metals (19.4%), the second at 12


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162.6 eV was attributed to S22- (11.9%), the third at 164 eV was ascribed to Sn2(46.2%), and the fouth at 168.5 eV was considered as SO42- (22.5%) due to the oxidization of Fe1-xS in air.34-36 Interestingly, all the four peaks shifted to higher binding energies after U(VI) sorption, and the relative ratios were changed to 7.6%, 25.3%, 22.7% and 44.4%, respectively, suggesting the formation of the polysulfide and different sulfur species after U(VI) sorption on Fe1-xS.37-39

Kinetic Studies of U(VI) Sorption on Fe2O3 and Fe1-xS. The kinetic data of U(VI) onto Fe2O3 (Figure S2) diagramed that the sorption of U(VI) on Fe2O3 reached equilibrium rapidly in a few minutes. In contrast, the removal process of U(VI) on Fe1-xS took place in two diverse stages: the fast sorption stage within the first reaction time of 10 min implied that the interaction mechanism between U(VI) and Fe1-xS was dominated by surface complexion or chemisorption instead of physical sorption,40 and the relative slow stage indicated that U(VI) ions gradually penetrated across Fe1-xS surface layer and then entered into interior layers to react with the inner functional groups. The experimental data were fitted by two typical kinetic models, which were shown as: Pseudo-first-order model:

ln(qe − qt ) = ln qe − k1t

(4)

Pseudo-second-order model:

1 t t = + 2 qt k 2 qe qe

(5)

13



where qt (mg g-1) and qe (mg g-1) were the amounts of target ions adsorbed on sorbents at time t (min) and after equilibrium, respectively. k1 (min-1) was the pseudo-first-order model constant, which was obtained by plotting ln (qe - qt) versus t. And k2 (g mg-1 min-1) represented the constant of the pseudo-second-order model. It was worth noting that qe and k2 can be calculated by the plot of t/qt as a function of t. The fitting parameters of the two models on Fe2O3 and Fe1-xS were summarized in Table 1. The higher correlation coefficients (R2) indicated that the kinetic sorption was well fitted by the pseudo-second-order model better than pseudo-first-order one. Besides, the sorption performances of Fe1-xS were much higher than that of Fe2O3, indicating that Fe1-xS was a promising material for the treatment of wastewater in the future.

Effect of Solution pH and Ionic Strength. Considering the various pH values of polluted water and real wastewater, it was of great significance to study the removal of U(VI) on Fe2O3 and Fe1-xS as a function of the solution pH. From Figure 4a, the sorption percentages of U(VI) on Fe2O3 increased slightly in the pH range of 2 ~ 4 (Process I), then increased significantly with pH increasing from 4 to 8 (Process II), and finally decreased weakly at pH > 8 (Process III). While the U(VI) sorption on Fe1-xS as a function of pH (Figure 4b) showed that the sorption efficiencies of U(VI) on Fe1-xS increased significantly in the pH range 2 ~ 8 and presented a high plateau at pH > 8. This phenomenon can be explained by the surface charge characteristics of Fe2O3 and Fe1-xS and U(VI) species under different pH conditions. As shown in Figure 4c, the surface charge of Fe2O3 (pHpzc= 8.81) and 14


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Fe1-xS (pHpzc= 7.04) were positively charged at pH < pHpzc and negatively charged at pH > pHpzc. Meanwhile, one can see from Figure 4d that U(VI) species were mainly cationic forms [i.e., UO22+, (UO2)3(OH)5+] over the pH range of 2 ~ 8 and anionic forms [i.e., (UO2)3(OH)7-, UO2(OH)3-] at pH of 8 ~ 11. The strong electrostatic repulsion between positively charged U(VI) species and the protonated sorbent surface leading to the relatively low adsorption efficiency at pH < pHpzc, and the decrease of zeta potential (pH from 4 to 8) enhanced the interaction between positive charged Fe2O3 and these cations, and thereby increased the sorption efficiency to some extent. Notably, it was found that Fe1-xS showed better sorption performance than Fe2O3 in acidic conditions. This can be attributed to the low affinity of S2- ligands for hard H+, and the strong binding energy between UO22+ and S2- sharply enhanced the removal efficiency of Fe1-xS. At lower pH conditions (pH < pHpzc), the target UO22+ was hard to be adsorbed on Fe2O3 with positive groups (≡SOH2+) on the surface due to electrostatic repulsion. In terms of the hard-soft acid-base theory, the S2- contained soft basic sites in Fe1-xS, displaying a relative high affinity for soft U(VI) ions. As a result, the protonation reaction of Fe1-xS sample was hardly occurred under low pH conditions than that of Fe2O3, which implied that the majority of oxygen-containing adsorbents were inefficient for capturing UO22+. Additionally, the removal process of U(VI) on Fe1-xS was scarcely influenced by ionic strength, whereas the U(VI) sorption percentages on Fe2O3 decreased with the increasing NaNO3 concentration at pH > 4. The increscent Na+ concentration could compete with the active sites of Fe2O3. According to the abovementioned results, one 15



can conclude that the dominate mechanisms of U(VI) adsorption onto Fe2O3 and Fe1-xS were outer-sphere surface complexation and inner-sphere surface complexation, respectively.41,42

Effect of Sorbent Content and Coexisted Ions The effects of sorbent content on U(VI) sorption by Fe2O3 and Fe1-xS were studied to assess the cost-effective in wastewater treatment (Figure S3a and S3b). The sorption efficiency of U(VI) on Fe1-xS increased rapidly from ~3.2 % to ~87% as the sorbent dosage increased from ~0.01 to 0.5 g L-1, and then remained high-level with the increase of Fe1-xS content, while the sorption efficiency of U(VI) on Fe2O3 increased gradually from ~4.7% to ~37% with an increase of solid dosage from ~0.01 to 1 g L-1. And the rapidly increasing tendency was mainly attributed to the exposing of more accessible active sites for U(VI) capture. Additionally, the values of distribution coefficient (Kd) were almost independent of sorbent dosage, which was very consistent with the physicochemical properties of Kd. The comparison of U(VI) sorption on both samples further demonstrated that Fe1-xS could be considered as a suitable and economical choice for U(VI) elimination from wastewater. The influences of coexisted cations on U(VI) uptake as a function of Ca2+ or Mg2+ concentrations were presented in Figure S3c and S3d. From the relationship between Kd and Ca/U molar ratios, one can obviously see that Kd values (from ~499 to 0 mL g-1 for Fe2O3 and from ~6340 to ~293 mL g-1 for Fe1-xS) decreased remarkably with the increase of Ca/U molar ratios from 10 to 300, in which the Kd values as a function of Mg/U molar ratios decreased slightly (from 732 to 397 mL g-1 for Fe2O3

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and from ~2749 to 2494 mL g-1 for Fe1-xS). In addition, the sorption efficiencies of Fe2O3 decreased (from 26% to 16%) in the presence of 10 ~ 300 ppm of Mg2+, while the sorption efficiency was remained substantially unchangeable for Fe1-xS (from 44% to 42%) with the increase of Mg/U molar ratio. The high sorption efficiency even in high Mg2+ concentrations can be attributed to the affinity of Fe1-xS for UO22+ arising from the soft-soft acid-base UO22+···S2- bonding interaction, which was much stronger than the hard-soft Mg2+···S2- interaction, whereas soft-hard acid-base UO22+···O2bonding interaction was weaker than hard-hard Mg2+···O2-. Besides, O2- anions have high affinity to proton ions, resulting in the inefficient capture of Fe2O3 for UO22+.19 It was worth noting that the high Ca2+ concentration could strongly compete with UO22+. In general, the reduction of sorption efficiency is mainly due to the forming of metal-U(VI)-carbonate species with dissolved CO2 in aqueous solutions (e.g., Ca2UO2(CO3)3), which could hinder U(VI) sorption on Fe2O3 and Fe1-xS. Furthermore, the decreasing degree of U(VI) sorption in the presence of Ca2+ was greater than that of U(VI) sorption coexisted with Mg2+ owing to the formation of Mg-U(VI)-CO32species (MgUO2(CO3)32-) under alkaline conditions at 298 K and PCO2 = 0.0314% atm,1,43 which had basically no distinct influence of Mg2+ on U(VI) sorption.

Sorption Isotherms and Thermodynamic Studies. The sorption isotherms of U(VI) on Fe2O3 and Fe1-xS at 298 K, 313 K and 328 K were studied to evaluate the sorption performance, and the relationship between qe and Ce at various temperatures were elucidated in Figure 5a-5c. It can be seen that the sorption capacities of U(VI) on Fe2O3 and Fe1-xS increased gradually with the increase

17



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of equilibrium concentrations until the maximum sorption capacities were reached. Comparing the U(VI) sorption capacities on Fe2O3 and Fe1-xS, one can see that the Fe1-xS showed a better scavenging efficiency and higher affinity to U(VI) than Fe2O3 in different temperatures. In order to better understand the sorption mechanism, Langmuir and Freundlich models were applied to simulate the experimental data, which were expressed as: Langmuir model:

Ce C 1 = + e qe b ⋅ qmax q max

Freundlich model:

Ln q e = Ln K +

1 Ln C e n

(6) (7)

where Ce (mg L-1) was the U(VI) concentration at equilibrium. qe (mg g-1) and qmax (mg g-1) were the amount of U(VI) adsorbed per unit weight of sorbent and the maximum sorption capacity, respectively. b (L mg-1) was noted as a constant of Langmuir model. n and K (mg g-1) represented the sorption intensity and sorption capacity. The Langmuir isotherm model was usually applied to monolayer and homogeneous sorption process, whereas the Freundlich model was appropriate to a multilayer and heterogeneous sorption.44 The corresponding sorption isotherms of two models were shown in Figure S4a–S4d, and the relevant parameters were listed in Table 2. The higher correlation coefficients of the Langmuir model than those of Freundlich model indicated that the sorption process were fitted by Langmuir isotherm model better, and the binding affinity was uniform for U(VI) sorption on Fe2O3 and Fe1-xS. In addition, the maximum sorption capacities of Fe2O3 and Fe1-xS increased with temperature increasing. According to the parameters of Langmuir model in Table 2, the sorption capacities of U(VI) on Fe1-xS were calculated to be 18


Page 19 of 33 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


21.23, 32.27 and 73.15 mg g-1 at 298, 313 and 328 K, respectively, which were significantly higher than those of U(VI) on Fe2O3, demonstrating that Fe1-xS exhibited better sorption abilities than Fe2O3.45 The enhancement of sorption capacity could be assigned to the strong interaction between U(VI) and S2-. Moreover, the individual Fe1-xS microcube was composed of numerous nanoparticles, which can facilitate the target ions diffusion from the solution to these nanoparticles. Moreover, we compared the sorption performance of Fe2O3 and Fe1-xS on U(VI) with that of other reported materials under different experimental conditions listed in Table S1. The thermodynamic data of Gibbs free energy change (∆G0) was determined by the following formula:

∆G 0 = − RT ln K 0

(8)

where R (8.3145 J mol-1·K-1) represented ideal gas constant, T (K) was Kelvin temperature. ln K0 as a sorption equilibrium constant can be obtained by the intercept of plotting ln Kd versus Ce. The standard enthalpy change (∆H0) and entropy change (∆S0) were obtained by plotting ∆G0 as a function of T (Figure 5d), which was described as:

∆G 0 = ∆H 0 − ∆S 0 T

(9)

As tabulated in Table 3, the positive ∆H0 value (16.65 kJ·mol-1) on Fe2O3 manifested the endothermic sorption. Similar result was obtained for Fe1-xS and the enthalpy became more positive (48.10 kJ·mol-1), indicating that the energy consumption of U(VI) dehydration process before sorption exceed the exothermicity of U(VI) attached on the surface of Fe1-xS. The positive ∆S0 values suggested that the 19



U(VI) sorption on Fe2O3 and Fe1-xS were a chaotic and spontaneous process. The negative ∆G0 values further illustrated that the sorption was spontaneous under the experimental conditions.44 On the other hand, the ∆G0 values got more negative with the temperature increasing, indicating that the sorption process was enhanced at higher temperature.46

Uranium Sorption on Fe2O3 and Fe1-xS in Different Water Systems. In order to evaluate the advantages and practicability of sulfide modified materials, U(VI) sorption over Fe2O3 and Fe1-xS in different water systems (synthetic water samples and real water samples) was further studied. The synthetic water samples containing 10 ppm U(VI) were conducted to simulate contaminated water by adding a certain foreign ions into Milli-Q water based on the previous study.19 And the sorption experiment was firstly performed in Milli-Q water without adding other ions, which served as blank control group. As shown in Figure 6, the U(VI) sorption percentage in Milli-Q water was calculated to be 39.3% for Fe1-xS, which was almost 9.65 times higher than that for Fe2O3 (4.07%). The concentration of coexisting ions (Ca2+, Ma2+, K+ and Na+) in synthetic contaminated potable water (Synthetic C-P water), synthetic contaminated seawater (Synthetic C-seawater) and synthetic original seawater (Synthetic O-seawater) were listed in Table S2. One can see that U(VI) sorption increased in the order of Synthetic C-P water > Synthetic C-seawater > Synthetic O-seawater (the same tendency for Fe2O3 and Fe1-xS). The Synthetic O-seawater had considerable amount of electrolyte ions compared with the concentration of target ions. Interestingly, the U(VI) sorption efficiency over Fe1-xS in Synthetic O-Seawater is the

20


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highest, which can be attributed to interaction of foreign ions with sorbent as well as sorbate. Further studies were performed in tap water and real seawater (obtained from Liu Heng Island, Zhejiang province of China) by adding 10 ppm U(VI) into tap water and real seawater in which the sorption capacities of 13.48 and 8.55 mg g-1 were achieved on Fe1-xS. From the relevant parameters listed in Table S2, one can conclude that Fe1-xS exhibited excellent U(VI) removal ability even in many simulated synthetic water samples and real seawater.

CONCLUSION To sum up, the Fe1-xS and Fe2O3 precursors were prepared by a high-yield facial method, and characterized by a sequence of techniques to better understand their morphologies and microstructures. Batch experiments were carried out to evaluate the sorption properties of U(VI) on Fe2O3 and Fe1-xS. As a result, the Fe1-xS showed a relative better sorption performance towards U(VI) than Fe2O3. The sorption of U(VI) was strongly dependent on solution pH. Interestingly, Fe1-xS was more resistant to ionic strength, while Fe2O3 was independent of ionic strength at acidic pH and strongly affected at higher pH values. The pseudo-second-order model could preferably describe the kinetics data, indicating that the U(VI) sorption process on Fe1-xS and Fe2O3 was the rate-controlling mechanism. In addition, as the temperature increased, the U(VI) adsorption performances on Fe2O3 and Fe1-xS increased, indicating that the adsorption was spontaneous. Satisfactorily, Fe1-xS could efficiently eliminate U(VI) ions in real contaminated water samples, suggesting that metal sulfide exhibited a better sorption performance than metal oxide. This work may provide a

21



facile and general approach to synthesize magnetic sulfide-based nanomaterials for nuclear wastewater and real wastewater treatment.

AUTHOR INFORMATION Corresponding Author * E-mail: [email protected] (T. Wen); [email protected] (X.K. Wang). Phone: +86-10-61772890. Fax: +86-10-61772890.

Notes The authors declare no competing financial interest.

ACKNOWLEDGEMENTS This work is supported by Science Challenge Project (TZ2016004), NSFC (21607042, 21707033, 21577032), and Fundamental Research Funds for the Central Universities (2018ZD11, 2017MS045).

ASSOCIATED CONTENT Supporting Information The Supporting Information is available free of charge on the ACS Publication website: XRD patterns of Fe1-xS before and after U(VI) sorption; Rietveld refined XRD pattern of Fe1-xS.; Kinetic curves of Fe2O3 and Fe1-xS; Plot of U(VI) sorption on Fe2O3 and Fe1-xS versus solid contents at I = 0.01 M NaNO3; Linear fitting of Langmuir model and Freundlich model on Fe2O3 and Fe1-xS; Comparison of the maximum sorption capacities of various sorbents for U(VI) in Table S1; Uranium sorption on Fe2O3 and Fe1-xS in different water samples in Table S2.

REFERENCES: 22


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(1) Wen, T.; Wang, X.; Wang, J.; Chen, Z.; Li, J.; Hu, J.; Hayat, T.; Alsaedi, A.; Grambow, B.; Wang, X., A Strategically Designed Porous Magnetic N-Doped Fe/Fe3C@C Matrix and Its Highly Efficient Uranium(VI) Remediation. Inorg. Chem. Front. 2016, 3, 1227-1235. (2) Nuclear Waste: Amounts and On-Site Storage; Nuclear Energy Institute: Washington, DC, 2011; http://www.nei.org/resources and stats /nuclear; statistics /nuclear waste amounts and onsite storage/(accessed December 20, 2012). (3) Urey, H. C., The Cosmic Abundances of Potassium, Uranium, and Thorium and the Heat Balances of the Earth, the Moon, and Mars. P. Natl. Acad. Sci. USA. 1955, 41, 127-144. (4) 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. (5) M. Legrand.; C. Elie.; J. Stefani.; N.Flores.;, C. Culeux a.; O. Delissen.; C. Ibanez.; P. Lestaevel .; P. Eriksson.; C. Dinocourt., Cell Proliferation and Cell Death are Disturbed During Prenatal and Postnatal Brain Development after Uranium Exposure. NeuroToxicology 2016, 52, 34-45. (6) Ladeira, A. C. Q.; Morais, C. A., Uranium Recovery from Industrial Effluent by Ion Exchange—Column Experiments. Miner. Eng. 2005, 18 (13-14), 1337-1340. (7) Ke, L.; Li, P.; Wu, X.; Jiang, S.; Luo, M.; Liu, Y.; Le, Z.; Sun, C.; Song, S., Graphene-Like Sulfur-Doped g-C3N4 for Photocatalytic Reduction Elimination of UO22+ under Visible Light. Appl. Catal. B- Environ. 2017, 205, 319-326. (8) Yuan, K.; Renock, D.; Ewing, R. C.; Becker, U., Uranium Reduction on Magnetite: Probing for Pentavalent Uranium Using Electrochemical Methods. Geochim. Cosmochim. Ac. 2015, 156, 194-206. (9) Kim, K.; Baek, Y.; Lee, K.; Chung, D.; Moon J., Treatment of Radioactive Waste Seawater by Coagulation – Flocculation Method Using Ferric Hydroxide and Poly Acrylamide. J. Nucl. Sci. Technol. 2016, 439-450. (10) Abbasi, W. A.; Streat, M., Adsorption of Uranium from Aqueous Solutions Using Activatedcarbon. Sep. Sci. Technol. 1978, 29, 1217-1230. (11) Sprynskyy, M.; Kovalchuk, I.; Buszewski, B., The Separation of Uranium Ions by Natural and Modified Diatomite from Aqueous Solution. J. Hazard. Mater. 2010, 181, 700-707. (12) Aytas, S.; Yurtlu, M.; Donat, R., Adsorption Characteristic of U(VI) Ion onto Thermally Activated Bentonite. J. Hazard. Mater. 2009, 172, 667-674. (13) Xie, X.; Zhou, S. K., Preparation of Modified Fly Ash and Its Adsorption Properties for U(VI) Uranium. Appl. Mech. Mater. 2017, 873, 107-111. (14) Ding, C.; Cheng, W.; Sun, Y.; Wang, X., Novel Fungus-Fe3O4 Bio-Nanocomposites as High Performance Adsorbents for the Removal of Radionuclides. J. Hazard. Mater. 2015, 295, 127-137. (15) Sun, Y.; Yang, S.; Chen, Y.; Ding, C.; Cheng, W.; Wang, X., Adsorption and Desorption of U(VI) on Functionalized Graphene Oxides: A Combined Experimental and Theoretical Study. Environ. Sci. Technol. 2015, 49, 4255-4262. (16) Eglizaud, By N.; Miserque, F.; Simoni, E.; Schlegel, M.; Descostes, M., 23



Uranium(VI) Interaction with Pyrite (FeS2) Chemical and Spectroscopic Studies. Radiochim. Acta 2006, 94, 651–656. (17) Scott, T. B.; Riba Tort, O.; Allen, G. C., Aqueous Uptake of Uranium onto Pyrite Surfaces; Reactivity of Fresh versus Weathered Material. Geochim. Cosmochim. Ac. 2007, 71, 5044-5053. (18) Manos, M. J.; Kanatzidis, M. G., Layered Metal Sulfides Capture Uranium from Seawater. J. Chem. Soc. 2012, 134, 16441-16446. (19) Ma, S.; Huang, L.; Ma, L.; Shim, Y.; Islam, S. M.; Wang, P.; Zhao, L.; Wang, S.; Sun, G.; Yang, X.; Kanatzidis, M. G., Efficient Uranium Capture by Polysulfide/Layered Double Hydroxide Composites. J. Chem. Soc. 2015, 137, 3670-3677. (20) Gu, P.; Zhang, S.; Li, X.; Wang, X.; Wen, T.; Jehan, R.; Alsaedi, A.; Hayat, T.; Wang, X., Recent Advances in Layered Double Hydroxide-Based Nanomaterials for the Removal of Radionuclides from Aqueous Solution. Environ. Pollut. 2018, 240, 493-505. (21) Li, L.; Peng, S.; Bucher, N.; Chen, H.; Shen, N.; Nagasubramanian, A.; Eldho, E.; Hartung, S.; Ramakrishna, S.; Srinivasan, M., Large-scale Synthesis of Highly Uniform Fe1−xS Nanostructures as a High-Rate Anode for Sodium Ion Batteries. Nano Energy 2017, 37, 81-89. (22) Sing, K.S., Reporting Physisorption Data for Gas/Solid Systems with Special Reference to the Determination of Surface Area and Porosity (Recommendations 1984). Pure and applied chemistry, 1985. 57, 603-619. (23) Naderi, M., Chapter Fourteen - Surface Area: Brunauer – Emmett – Teller (BET). Elsevier. 2015, 585-608. (24) Wen, T.; Wu, X.; Tan, X.; Wang, X.; Xu, A., One-Pot Synthesis of Water-Swellable Mg–Al Layered Double Hydroxides and Graphene Oxide Nanocomposites for Efficient Removal of As(V) from Aqueous Solutions. ACS Appl. Mater. Inter. 2013, 5, 3304-3311. (25) Wen, T.; Wang, J.; Li, X.; Huang, S.; Chen, Z.; Wang, S.; Hayat, T.; Alsaedi, A.; Wang, X., Production of a Generic Magnetic Fe3O4 Nanoparticles Decorated Tea Waste Composites for Highly Efficient Sorption of Cu(II) and Zn(II). J. Environ. Chem. Eng. 2017, 5, 3656-3666. (26) Liu, Z.; Chen, J.; Wu, Y.; Li, Y.; Zhao, J.; Na, P., Synthesis of Magnetic Orderly Mesoporous α-Fe2O3 Nanocluster Derived from MIL-100(Fe) for Rapid and Efficient Arsenic(III,V) Removal. J. Hazard. Mater. 2018, 343, 304-314. (27) Descostes, M.; Schlegel, M. L.; Eglizaud, N.; Descamps, F.; Miserque, F.; Simoni, E., Uptake of Uranium and Trace Elements in Pyrite (FeS2) Suspensions. Geochim. Cosmochim. Ac. 2010, 74, 1551-1562. (28) Riba, O.; Scott, T. B.; Vala Ragnarsdottir, K.; Allen, G. C., Reaction Mechanism of Uranyl in the Presence of Zero-Valent Iron Nanoparticles. Geochim. Cosmochim. Ac. 2008, 72, 4047-4057. (29) Zhang, R.; Chen, C.; Li, J.; Wang, X., Investigation of Interaction Between U(VI) and Carbonaceous Nanofibers by Batch Experiments and Modeling Study. J. Colloid Interf. Sci. 2015, 460, 237-246. 24


Page 24 of 33

Page 25 of 33 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(30) MUIR, A. R. PRATT, I. J.; NESBITT, H. W., X-Ray Photoelectron Spectroscopy of Some Uranium Oxide Phases. Geochim. Cosmochim. Ac. 1994, 827-841. (31) Huber, F.; Schild, D.; Vitova, T.; Rothe, J.; Kirsch, R.; Schäfer, T., U(VI) Removal Kinetics in Presence of Synthetic Magnetite Nanoparticles. Geochim. Cosmochim. Ac. 2012, 96, 154-173. (32) Song, S.; Huang, S.; Zhang, R.; Chen, Z.; Wen, T.; Wang, S.; Hayat, T.; Alsaedi, A.; Wang, X., Simultaneous Removal of U(VI) and Humic Acid on Defective TiO2−x Investigated by Batch and Spectroscopy Techniques. Chem. Eng. J. 2017, 325, 576-587. (33) Ding, C.; Cheng, W.; Nie, X.; Yi, F., Synergistic Mechanism of U(VI) Sequestration by Magnetite Graphene Oxide. Chem. Eng. J. 2017, 324, 113-121. (34) El Kamel, M.; Galtayries, A.; Vermaut, P.; Albinet, B.; Foulonneau, G.; Roumeau, X.; Roncin, B.; Marcus, P., Sulfidation Kinetics of Industrial Steels in a Refinery Crude Oil at 300 °C: Reactivity at the Nanometer Scale. Surf. Interface Anal. 2010, 42, 605-609. (35) Liao, F.; Światowska, J.; Maurice, V.; Seyeux, A.; Klein, L. H.; Zanna, S.; Marcus, P., Electrochemical Lithiation and Passivation Mechanisms of Iron Monosulfide Thin Film as Negative Electrode Material for Lithium-Ion Batteries Studied by Surface Analytical Techniques. Appl. Surf. Sci. 2013, 283, 888-899. (36) Pratt, A. R.; Muir, I. J.; Nesbitt, H. W., X-ray Photoelectron and Auger Electron Spectroscopic Studies of Pyrrhotite and Mechanism of Air Oxidation. Geochim. Cosmochim. Ac. 1994, 58, 827-841. (37) Wersin, P.; Hochella, M. F.; Persson, P.; Redden, G.; Leckie, J. O.; Harris, D. W., Interaction Between Aqueous Uranium (VI) and Sulfide Minerals: Spectroscopic Evidence for Sorption and Reduction. Geochim. Cosmochim. Ac. 1994, 58, 2829-2843. (38) Shao, D.; Wang, X.; Li, J.; Huang, Y.; Ren, X.; Hou, G.; Wang, X., Reductive Immobilization of Uranium by PAAM–FeS/Fe3O4 Magnetic Composites. Environmental Science: Water Res. Technol. 2015, 1, 169-176. (39) Hua, B.; Deng, B., Reductive Immobilization of Uranium (VI) by Amorphous Iron Sulfide. Environ. Sci. Technol. 2008, 42, 8703-8708. (40) Li, J.; Wang, X.; Zhao, G.; Chen, C.; Chai, Z.; Alsaedi, A.; Hayat, T.; Wang, X., Metal–Organic Framework-Based Materials: Superior Adsorbents for the Capture of Toxic and Radioactive Metal Ions. Chem. Soc. Rev. 2018, 47, 2322-2356. (41) Zhang, C.; Li, X.; Chen, Z. Wen, T.; Huang, S.; Hayat, T.; Alsaedi, A.; Wang, X., Synthesis of Ordered Mesoporous Carbonaceous Materials and Its Highly Efficient Capture of Uranium from Solutions. Sci. Chi. Chem. 2017, 61, 281-293. (42) Pang, H.; Huang, S.; Wu, Y.; Yang, D.; Wang, X.; Yu, S.; Chen, Z.; Alsaedi, A.; Hayat, T.; Wang, X., Efficient Elimination of U(VI) by Polyethyleneimine-Decorated Fly Ash. Inorg. Chem. Front. 2018. Doi:10.1039/C8QI00253C. (43) Dong, W.; Brooks, S. C., Determination of the Formation Constants of Ternary Complexes of Uranyl and Carbonate with Alkaline Earth Metals (Mg2+, Ca2+, Sr2+, and Ba2+) Using Anion Exchange Method. Environ. Sci. Technol. 2006, 40, 25



Page 26 of 33

4689-4695. (44) Huang, S.; Song, S.; Zhang, R.; Wen, T.; Wang, X.; Yu, S.; Song, W.; Hayat, T.; Alsaedi, A.; Wang, X., Construction of Layered Double Hydroxides/Hollow Carbon Microsphere Composites and Its Applications for Mutual Removal of Pb (II) and Humic Acid from Aqueous Solutions. ACS Sustain. Chem. Eng. 2017, 5, 11268-11279. (45) Gu, P.; Xing, J.; Wen, T.; Zhang, R.; Wang, J.; Zhao G. Hayat, T.; Ai, Y.; Lin, Z.; Wang, X., Experimental and Theoretical Calculation Investigation on Efficient Pb(II) Adsorption to Etched Ti3AlC2 Nanofibers and Nanosheets. Environ. Sci.: Nano, 2018, 5, 946-955. (46) Zhao, G.; Li, J.; Ren, X.; Chen, C.; Wang, X., Few-Layered Graphene Oxide Nanosheets as Superior Sorbents for Heavy Metal Ion Pollution Management. Environ. Sci. Technol. 2011, 45, 10454-10462.

Scheme 1. Schematic illustration for the fabrication of Fe2O3 and Fe1-xS microcubes.

Precipitation method

FeCl3

+

TAA, 500 oC Anneal

100 oC, Reflux NaOH

Fe1-xS

Fe2O3

According to the synthetic route of Fe2O3 and vulcanization process of Fe1-xS, we fabricated these cliparts by using ChemDraw, ChemOffice and 3Ds MAX.

(b)

(a)

2 µm

(c)

(d)

2 µm

200 nm

200 nm

Fe Kα1

(f)

(e)

(h)

(g) Wt % 70 30

Fe O

O Kα1

O Kα1

Fe Kα1

Wt % 62 38

Fe S S Kα1

Fe Lα1_2

Cu Kα1

Cu Kα1

200 nm 0

(i)

1

Fe Kα1

200 nm

Cu Lα1_2

2

3

Fe

O

250 nm

250 nm

4

5

6

7

8

9

10

(j)

250 nm

Cu Kα1

0

Fe

250 nm

1

2

3

4

5

6

7

8

250 nm

250 nm

Figure 1. Characterization of the as-synthesized Fe2O3 and Fe1-xS: different 26


9

S

10

Page 27 of 33

magnification SEM images of Fe2O3 microcubes (a, b) and Fe1-xS (c, d). TEM images of Fe2O3 (e) and Fe1-xS (g). EDS of Fe2O3 (f) and Fe1-xS (h). And the corresponding

(220)

(2022)

2 -1

Magnetization (emu⋅ g )

(200)

(a)

(2011)

TEM element mapping images of Fe2O3 (i) and Fe1-xS (j).

(214) (300)

(116)

(024)

(110) (113)

(104)

(012)

Intensity (a.u.)

Fe1-xS JCPDS Card NO.29-0726

Fe2O3 JCPDS Card NO.33-0664

(b)

Fe2O3

1

-1

Ms = 0.4 emu g

0 -1

-1

Ms = 2.1 emu g

-2 20

30

45

60

70

(c)

-1

35

Fe2O3 Adsorption

30

Fe2O3 Desorption

25

Fe1-xS Adsorption

20

Fe1-xS Desorption

Fe1-xS -10

0.06

3

40

40 50 2θ (degree)

dV/dlog(w) Pore Volume (cm g )

3

Quantity Adsorbed (cm /g STP)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


15 10 5 0

-5 0 5 Magnetic Field (KOe)

10 Fe2O3

(d)

Fe1-xS

0.05 0.04

d = 12∼80 nm

0.03 0.02

d = 58 nm

0.01 0.00

0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 1.1 Relative Pressure (P/Po)

10

100

Pore Diameter (nm)

Figure 2. Characterization of the as-prepared Fe2O3 and Fe1-xS. XRD patterns (a), Magnetization curves (b), N2 adsorption-desorption isotherms (c) and the corresponding pore size distributions (d).

27



(b) Intensity (a.u.)

Intensity (a.u.)

(a) Fe2O3 before sorption 3400

480

908

1628

Fe1-xS before sorption

Fe1-xS after sorption

Fe2O3 after sorption

902 582

3500

3000

2500 2000 1500 -1 Wavenumber (cm )

(c)

3500

500

(d) U 4f

Fe 2p O 1s

2500 2000 1500 -1 Wavenumber (cm )

1000

Fe1-xS

U 4f7/2

U 4f5/2

500

S 2p

Before soption

After sorption

3000

Intensity (a.u.)

Intensity (a.u.)

1000

Fe 2p

After sorption

Fe1-xS

U 4f

Fe2O3

U 4f7/2

U 4f5/2

Before soption Fe2O3

1000

800 600 400 Binding Energy (eV)

200

0

396

After Sorption

(e) O 1s

Before Sorption 2-

O (74.1%) -

OH (25.9%)

392

388 384 Binding Energy (eV)

380

Before Sorption 2-

2-

Sn

2-

SO4

2-

S2

S

Fe1-xS

Fe2O3

534

533

532

531

530

529

376

After Sorption

(f) S 2p Intensity (a.u.)

1200

Intensity (a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 28 of 33

528

172

Binding Energy (eV)

170

168

166

164

162

160

Binding Energy (eV)

Figure 3. FTIR spectra of Fe2O3 (a) and Fe1-xS (b) before and after U(VI) sorption. (c) XPS survey spectra of Fe2O3 and Fe1-xS before and after U(VI) sorption. (d) U 4f XPS spectra of Fe2O3 and Fe1-xS after U(VI) sorption.(e) O 1s XPS spectra of Fe2O3 and (f) S 2p XPS spectra of Fe1-xS before and after U(VI) sorption.

28


Page 29 of 33

100

(a)

(b)

100

Process ΙΙΙ

80 Fe2O3

60

Sorption (%)

Sorption (%)

80 Process ΙΙ

40 0.1 M NaNO3

20 0 3

4

5

6

7

8

9

10

40 0.1 M NaNO3 0.01 M NaNO3

0.001 M NaNO3

2

Fe1-xS

60

20

0.01 M NaNO3

Process Ι

0.001 M NaNO3

0 2

11

3

4

5

6

(c)

9

10

11

12

60

-

100

Fe1-xS

UO2(OH)3

2+

UO2

(UO2)3(OH)5

80

40

pH pzc

20

=8 .8

1

0 pHpzc=7.04

-20

8

(d)

Fe2O3

Species (%)

80

7 pH

pH

Zeta Potential (mV)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


+ -

(UO2)3(OH)7

60 +

(UO2)4(OH)7

40 2+

(UO2)2(OH)2

20

-40

+

UO2OH 0

-60 2

3

4

5

6

7

8

9

10

2

11

4

6

8

10

pH

pH

Figure 4. Sorption of U(VI) on (a) Fe2O3 and (b) Fe1-xS as a function of pH at various ionic strength concentrations. Conditions: T = 298 K, m/V = 0.3 g·L-1, and C0U = 10.0 ppm. (c) Zeta potentials of Fe2O3 and Fe1-xS and (d) relative distribution of U(VI) species as a function of solution pH.

29



35

20

(a) T = 298 K

-1

10

T = 313 K

25 qe (mg g )

-1

qe (mg g )

(b)

30

15

20 15 10

5

Fe2O3

0

Fe1-xS

0

0

5

10

15

20 25 30 -1 Ce (mg L )

35

40

45

0

80

5

10

15

20 25 30 -1 Ce (mg L )

35

40

45

10.5

(c)

70

10.0 T = 328K

60

Fe2O3

(d)

Fe1-xS

9.5

50

9.0 0

40

ln K

-1

Fe2O3

5

Fe1-xS

qe (mg g )

30 20

8.5

Fe2O3

8.0

Fe1-xS

7.5

10

7.0

0 6.5

0

5

10

15 -1 20 Ce (mg L )

25

30

0.0030

0.0031

0.0032 -1 1/T (K )

0.0033

0.0034

Figure 5. Sorption capacities of U(VI) on Fe2O3 and Fe1-xS at 298K (a), 313 K (b) and 328 K (c). (d) LnK0 was plotted vs. 1/T for U(VI) sorption on Fe2O3 and Fe1-xS. Conditions: pH = 5.0 ± 0.05, m/V = 0.3 g·L-1, and I = 0.01 M NaNO3.

16

Fe2O3

14

Fe1-xS

12 10 -1

qe (mg g )

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 30 of 33

8 6 4 2 0 Tap Milli-Q Synthetic Synthetic Synthetic Real Water C-P Water C-Seawater O-Seawater Water Seawater

Figure 6. Sorption performances of U(VI) on Fe2O3 and Fe1-xS in Milli-Q water, synthetic contaminated potable water, synthetic contaminated seawater, synthetic 30


Page 31 of 33 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


original seawater, tap water and real seawater.

31



Page 32 of 33

Table 1 Kinetic modeling of U(VI) sorption on Fe2O3 and Fe1-xS. Pseudo-second-order model

Pseudo-first-order model Sorbents k1 (min-1)

qe (mg∙g-1)

R2

k2 (g∙mg-1·min-1)

qe (mg·g-1)

R2

Fe2O3

0.67

3.67

0.92

2.61×10-2

4.35

0.987

Fe1-xS

0.25

11.17

0.93

1.91×10-2

12.38

0.999

Table 2 Parameters of isotherm models on Fe2O3 and Fe1-xS. Langmuir Samples

b (L·mg-1)

Fe2O3

Fe1-xS

Freundlich

T qmax(mg·g-1)

R2

K

R2

n

298K

0.21

5.70

0.950

1.45

0.67

0.864

313K

0.20

8.58

0.975

0.39

0.52

0.938

328K

0.16

12.08

0.993

0.29

0.57

0.949

298K

0.35

21.23

0.980

1.77×10-3

0.30

0.994

313K

1.53

32.27

0.997

2.17×10-8

0.16

0.961

328K

0.84

73.15

0.989

4.86×10-5

0.33

0.981

Table 3 Thermodynamic parameters of U(VI) sorption on Fe2O3 and Fe1-xS. Sorbent

Fe2O3

Fe1-xS

T

∆G0

∆H0

∆S0

∆G0

∆H0

∆S0

(K)

(kJ·mol-1)

(kJ·mol-1)

(J·mol-1·K-1)

(kJ·mol-1)

(kJ·mol-1)

(J·mol-1·K-1)

298

-16.54

313

-18.75

48.10

0.23

328

-19.90

-20.91 16.65

0.11

-24.75 -27.87

32


Page 33 of 33

TOC 500 oC, Sulfuration

Fe2O3 16

Fe1-xS

Fe2O3

14

Fe1-xS

12 10 -1

qe (mg g )

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


UO22+

8

UO22+

6 4 2 0 Tap Milli-Q Synthetic Synthetic Real Synthetic Water C-P Water C-Seawater O-Seawater Water Seawater

Various water treatments

33


Comparative Investigation of Fe2O3 and Fe1âxS Nanostructures for

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