Low concentration of Fe(II) to enhance the precipitation of U(VI) under neutral oxygen-rich conditions

Journal Pre-proofs Low concentration of Fe(II) to enhance the precipitation of U(VI) under neutral oxygen-rich conditions Yanpei Xie, Qi Fang, Mi Li, Sainan Wang, Yingfeng Luo, Xiaoyan Wu, Junwen Lv, Wenfa Tan, Hongqiang Wang, Kaixuan Tan PII: DOI: Reference:

S0048-9697(19)34819-3 https://doi.org/10.1016/j.scitotenv.2019.134827 STOTEN 134827

To appear in:

Science of the Total Environment

Received Date: Revised Date: Accepted Date:

27 July 2019 19 September 2019 3 October 2019

Please cite this article as: Y. Xie, Q. Fang, M. Li, S. Wang, Y. Luo, X. Wu, J. Lv, W. Tan, H. Wang, K. Tan, Low concentration of Fe(II) to enhance the precipitation of U(VI) under neutral oxygen-rich conditions, Science of the Total Environment (2019), doi: https://doi.org/10.1016/j.scitotenv.2019.134827

This is a PDF file of an article that has undergone enhancements after acceptance, such as the addition of a cover page and metadata, and formatting for readability, but it is not yet the definitive version of record. This version will undergo additional copyediting, typesetting and review before it is published in its final form, but we are providing this version to give early visibility of the article. Please note that, during the production process, errors may be discovered which could affect the content, and all legal disclaimers that apply to the journal pertain.

© 2019 Elsevier B.V. All rights reserved.

Low concentration of Fe(II) to enhance the precipitation of U(VI) under neutral oxygen-rich conditions Yanpei Xie a, Qi Fang a, b, *, Mi Li a, b, Sainan Wang a, Yingfeng Luo a, Xiaoyan Wu a, b,

Junwen Lv a, b, Wenfa Tan a, b, Hongqiang Wang a, b, Kaixuan Tan a

a

School of Resource Environment and Safety Engineering, University of South

China, Hengyang 421001, PR China b

Hengyang Key Laboratory of Soil Pollution Control and Remediation, University of

South China, Hengyang 421001, PR China

* Corresponding author. E-mail address: [email protected]

Abstract: Immobilization of U(VI) by naturally ubiquitous ferrous ions (Fe(II)) has been considered as an efficient and ecofriendly method to retard the migration of aqueous U(VI) at many nuclear sites and surface environments. In this study, we conducted Fe-U coprecipitation experiments to investigate the mechanism and stability of uranium (U) precipitation induced by a small quantity of Fe(II) under oxygen-rich conditions. The experimental results suggest that the sedimentation rates of U(VI) by Fe(II) under neutral oxygen-rich conditions are more than 96%, which is about 36% higher than those without Fe(II) and 16% higher than those under oxygen-free

1

conditions. The Fe-U coprecipitates were observed to remain stable under slightly acidic to neutral and oxygen-rich conditions. While Fe(II) primarily settles down as low-crystalline iron oxide hydroxide. U(VI) mainly precipitates as three forms: 16– 20% of U forms uranyl hydroxide and metaschoepite, which is absorbed on the surface of the solids; 52–56% of U is absorbed as discrete uranyl phases at the internal pores of iron oxide hydroxide; and 27–29% of U is probably incorporated into the FeO(OH) structure as U(V) and U(VI). The U(V) generated via one-electron reduction is somewhat resistant to the oxidation of O2 and the acid dissolution. In addition, nearly 70% of U and only about 15% of Fe could be extracted in 24 h by a hydrochloric acid solution with the H+ concentration ([H+]) of 0.01 M, revealing that U(VI) immobilization by a low concentration of Fe(II) combined with O2 has potential applications in the separation and recycling of aqueous uranium. Key words: U(VI); Fe(II); oxygen; coprecipitation; mechanism; stability

2

1. Introduction Uranium is a radioactive nuclide with high chemical toxicity, which is widely distributed in the soil and groundwater around uranium mining, processing, and disposing sites (Wang et al., 2019; Wu et al., 2019). The migration and release of U in these sites pose serious threats to both environment and human health, and hence are a cause of extensive public concern (Huang et al., 2018; Liu et al., 2019). The solubility and toxicity of U in natural environment depend largely on its valence states: oxidized U(VI) with high solubility and mobility in groundwater is highly toxic, whereas, reduction of U(Ⅵ) to U(IV) by the bacteria and abiotic reductants can achieve a dramatic reduction of dissolved U due to phenomena such as precipitation as uraninite (UO2) or other uranium-bearing solid phases (e.g., U3O8/β–U3O7/U4O9) (Chakraborty et al., 2010; Chen et al., 2017); adsorption of the U(IV) species (Latta et al., 2014); and formation of non-crystalline U(IV) phases (Bi et al., 2016). Reduction of soluble U(VI) to immobile U(IV) is commonly considered as is a desirable strategy for the retardation and long-term storage of nuclear uranium in soil and groundwater. However, the reduced U(IV) species can be readily oxidized and are thus re-released due to the extended reaction periods or the changing redox conditions (Crane et al., 2011; Cerrato et al., 2013; Chen et al., 2017). Therefore, it is of significance to find an effective method to achieve the retention and long-term stabilization of aqueous U(VI) under varying environments. Recent studies have increasingly focused on the sequestration of aqueous U(VI) by iron-based materials (e.g., zero-valent iron, ferrihydrite, goethite, magnetite, green 3

rust, hematite and Fe(II)/Fe(III) ions, et.al.). There are four primary mechanisms of uranium immobilization by iron-based materials. Firstly, the adsorption of U(VI) by iron compounds and iron minerals (Guo et al., 2009; Zhao et al., 2012; Wei et al., 2014), which has been reported to be a fast but ultimately reversible method for the short-term immobilization of aqueous U (Massey et al., 2014); Secondly, reductive precipitation of U(VI) as less soluble substances by zero-valent iron and Fe(II)-bearing minerals (e.g., magnetite, pyrite and pyrrhotite) (Skomurski et al., 2011; Yang et al., 2014; Li et al., 2015; Liu et al., 2017). Fe(II) is suggested to be the main electron donor for the reduction of U(VI) by these reductive iron materials (Bae et al., 2015; Tsarev et al., 2016; Tsarev et al., 2017); Thirdly, the precipitation of U(VI) as stable uranyl minerals (e.g., (meta)schoepite, uranophane, liebigite, and (meta)torbernite), which is reported to be an efficient method for the retention of aqueous U(VI) under aerobic environments (Burns et al., 2005; Massey et al., 2014). Taylor and Duff et al. have verified the precipitation of (meta)schoepite phases during the reaction between U(VI) and Fe(II)/Fe(III) (Taylor et al., 2015; Duff et al., 2002); Lastly, the incorporation of U(V)/U(VI) into host iron (oxyhydr)oxide minerals, for which several possible embedding pathways have been suggested in previous studies: for instance, the coprecipitation process of U(VI) and Fe(II)/Fe(III) (Ilton et al., 2012; Roberts et al., 2017; McBriarty et al., 2018), during the Fe(II)-aroused transformation of ferrihydrite to goethite and magnetite (Massey et al., 2014; Marshall et al., 2015), and through the dissolution of structural Fe(II) in magnetite (Ilton et al., 2010; Yuan et al., 2015a; Yuan et al., 2015b). Thus, Fe(II) apparently plays a vital role in the 4

immobilization of U(VI) through mechanism such as reduction, precipitation, and structural incorporation. The recent studies on the coprecipitation of U(VI) and Fe(II)/Fe(III) can be divided into two groups based on the reaction conditions: Firstly, under oxygen-free conditions, a study on the coprecipitation of U(VI) and Fe(II)/Fe(III) revealed the appearance and stabilization of U(V) via incorporation into the structure of magnetite and green rust (Roberts et al., 2017), while the reduction of U(VI) by sole Fe(II) mainly generated stable U(IV) solids (Du et al., 2011; Taylor et al., 2015). Secondly, Doornbusch and Mei et al. showed that under aerated conditions, U(VI) or U(V) could be sequestered into goethite or lepidocrocite through the coprecipitation of Fe(II) and U(VI) (Doornbusch et al., 2015; Mei et al., 2018). Furthermore, Ivan and Huber et al. (Huber et al., 2012; Pidchenko et al., 2017) have demonstrated that the adsorbed U(VI) and the U(V) that is structurally incorporated into magnetite could remain stable for over 550 and 142 days under anoxic and ambient conditions, respectively. However, former investigations on the retention of aqueous U(VI) by Fe(II) had mainly focused on the oxygen-free or aerated conditions, while the stability of the immobilized U(V) under more complicated conditions such as oxygen-rich conditions remains unexplored. Thus, it is essential to understand the mechanism and stability of Fe-U coprecipitation under oxygen-rich conditions, which is easy to implement in practical applications. During the coprecipitation of U(VI) and Fe(II), the phase of the Fe precipitates and the retention mechanism of U(VI) are of great concern to the relative content of 5

Fe(II) to U(VI) (Doornbusch et al. 2015, Mei et al. 2018). The recent experiments on the coprecipitation of Fe(II)/Fe(III) and U(VI) were conducted with relatively high Fe/U molar ratios ranging from 4.76 to 1000 (Du et al., 2011; Doornbusch et al., 2015; Taylor et al., 2015; Roberts et al., 2017; Mei et al., 2018). However, the effect of a higher Fe/U molar ratio on improving the sedimentation efficiency of U remains undetermined. Therefore, an investigation of the mechanism and stability of Fe-U coprecipitation induced by a low concentration of Fe(II) is significant, owing to prominent advantages such as less production of solid residue and less negative impact on the environment. Therefore, in this study, we conducted experiments on the coprecipitation of Fe(II) and U(VI) with a low Fe-U molar ratio under oxygen-rich conditions to identify the mechanism and stability of U precipitation induced by low concentration of Fe(II). The following are the objectives of this research: (i) to understand the redox mechanism of U(VI) and Fe(II) and ascertain the precipitation mechanism of U through an investigation of the combined forms of sedimentary U and Fe; and (ii) to assess the stability of stationary U under more complex environmental conditions by measuring the re-release content of U under oxygen-rich conditions at different pH values. 2. Material and methods 2.1. Materials All chemicals used in this study are of analytical grade. The stock solutions of FeCl2, NaOH and HCl were prepared by dissolving and diluting chemicals in doubly 6

deionized (DDI) water purged with N2 (99.99%). The U(VI) stock solution was prepared by dissolving U3O8 (of >99% purity) in concentrated hydrochloric acid, hydrogen peroxide and nitric acid solutions, and then diluted to the desired concentration. 2.2. Coprecipitation experiment N2-purged DDI and U(VI) stock solutions were compounded in Erlenmeyer flasks to achieve an initial concentration of 50 mg/L. The Fe(II) stock solutions were then added to the U(VI) solutions to obtain variable Fe/U molar ratios of 0.43, 0.85, 2.13, 4.26 and 8.52, with a total solution volume of 400 mL. The pH value of each solution was adjusted to 7.0 (±0.1) by adding 1.0 M NaOH solution. Then the Erlenmeyer flasks were placed to a vibrator operating at 150 rpm, at 25 ℃. High-purity oxygen (99.9%) was continuously injected into the suspensions during the whole experimental process to ensure that there is enough dissolved oxygen in the solutions suspensions. Coprecipitation was allowed for 24 h, after which, the supernatant samples were separated by 0.22 μm filters to measure the concentrations of Fe and U in the solutions: the concentration of U(VI) in the supernatants was determined by Arsenazo III spectrophotometry at a wavelength of 652 nm while the concentrations of Fe(II) and Fe(III) were determined by the phenanthroline spectrophotometric method. The precipitation rate of U after interaction with Fe(II) was calculated by using Eq. (1) : P = (C0-Ce)/C0*100%

(1)

where P is the U(VI) sedimentation rate (%), C0 is the initial concentration of 7

U(VI) (mg/L), and Ce is the final U(VI) concentration in the filtrate after interaction (mg/L). Solids labeled as FeU+O2, which is short for coprecipitates of Fe and U at the Fe/U molar ratio of 2 under oxygen-rich neutral conditions, were aged for another 24 h followed by centrifugation at 9000 rpm for 20 min. Then, the precipitates were washed three times with N2-purged DDI and dried at 60 ℃ for 8 h for later use, such as for the acid dissolution experiments and characterization. Similarly, the samples marked as Fe+O2, which is short for precipitates formed by 25 mg/L Fe(II) under oxygen-rich conditions at the pH value of 7.0 (±0.1), were obtained for characterization to identify the difference between the FeU+O2 and Fe+O2 samples. 2.3. Dissolution experiment The modified BCR sequential extraction procedure was adapted to identify the occurrence of uranium in the co-precipitates. The extraction procedure has been discussed in the supporting information (SI). Acid dissolution experiments were carried out for 24 h at 25 ℃ at a shaking speed of 150 rpm to examine the stability of the samples (FeU+O2) as a function of the [H+] of hydrochloric acid solution; the measurements were recorded at 10-6, 10-5, 10-4, 10-3, 0.01, 0.1, 0.79, 0.98, 2.0 3.0, 4.0 and 5.0 M. The series of dissolution times were set as 0.25, 0.5, 1, 2, 4, 6, 8, 12, 24, 36 and 48 h to investigate the effect of time on the release of sedimentary U at a [H+] value of 0.01 M. Afterward, the supernatant samples were separated by 0.22 μm filters to measure the concentrations of Fe and U. The remainder solids after acid dissolution at 0.01 M [H+] for 0.5 h were dried at 60 8

℃ for 8 h for the subsequent characterization. 2.4. Characterization The crystal phase and structure of the samples before and after acid dissolution were identified using X-ray diffraction (XRD, Bruker, D8 Advance) with a Cu Kα1 radiation and sweep range of 5−95° at a step size of 0.01°. The working voltage and the current of XRD measurements were 40 kV and 40 mA, respectively. The Fourier transform infrared (FTIR) spectra of the precipitates were measured over a wave number range of 400−4000 cm−1. In addition, scanning electron microscopy (SEM) combined with energy dispersive spectroscopy (EDS) was conducted to determine the morphology and composition of the precipitates. The working voltages of SEM (SUPRATM55) and EDS (Oxford Aztec X-Max 80) were 5 kV and 18 kV, respectively. High resolution transmission electron microscopy (HRTEM) was carried out with an JEM2100F electron microscope operated at 200 kV acceleration voltage. X-ray photoelectron spectroscopy (XPS) was performed using a Thermo Scientific ESCALAB 250Xi to determine the oxidation states of the main elements in the Fe-U coprecipitates; an Al Kα X-ray source was used for the XPS analysis. The C 1s peak (284.8 eV) was employed as a reference to calibrate the binding energies and the “CasaXPS” software was used to process the obtained XPS data. 3. Results and discussion 3.1. Effect of dissolved oxygen and Fe/U molar ratio on the sedimentation rate of U(VI) We first compared the sedimentation rates of U(VI) by Fe(II) with/without oxygen under neutral conditions. As shown in Fig. 1, the precipitation efficiency of 9

U(VI) with oxygen is much higher than that without oxygen. Moreover, the sedimentation rates of U(VI) at Fe/U molar ratios of 0, 0.85 and 8.52 are about 60%, 94%, and 96%, respectively, under oxygen-rich conditions and about 16%, 79%, and 82%, respectively, under oxygen-free conditions. Thus, we observe that the sedimentation rates of aqueous U(VI) with and without Fe(II) under oxygen-rich conditions are nearly 16% and 44% higher than those under oxygen-free conditions, respectively. Hence, the oxygen-rich system seems to be considerably more favorable for U(VI) precipitation. Fig. 2 shows the sedimentation rates of U(VI) at different Fe/U molar ratios. Under acidic conditions, we observe that a relatively low Fe/U molar ratio is more favorable for the precipitation of U(VI) compared with a higher one. The U(VI) sedimentation rate increases rapidly with Fe/U molar ratio and reaches its peak around Fe/U molar ratio of 2, beyond which it slightly decreases. Meanwhile, under neutral conditions, the U(VI) sedimentation rates with the presence of Fe(II) are all about 96% and are less affected by Fe/U molar ratios. 3.2. Occurrence states of U in the coprecipitates The modified BCR data of the solids are shown in Fig. S1. The adsorbed state of U accounts for approximately 23%, while about 46% is in the oxidized state, which indicates that U is partially precipitated as U(VI). The reduction state of U in the precipitates is about 21%, indicating that some of U(VI) in the solution might be reduced by Fe(II) under oxygen-enriched conditions, while Mei et al. (Mei et al., 2018) also demonstrated the production of U(V) during the coprecipitation of U(VI) 10

and Fe(II) under ambient conditions. The remaining U, which is about 10%, corresponds to the residual state. 3.3. Solubility of Fe-U coprecipitates The acid dissolution experiments were performed by examining the congruency of Fe and U dissolution to better understand the structure and solubility of the Fe-U coprecipitates. Fig. 3 shows the cumulative dissolution of Fe and U for 24 h as a function of [H+] in the solutions. We observe that there is hardly any dissolution of the Fe-U coprecipitates under near neutral or slightly acidic conditions at [H+] values ranging from 10-6 to 10-4 M. At [H+] values in the range of 0.001-0.1 M, we observe that the dissolution of U occurs earlier than that of Fe, followed by a congruent release of both when [H+] is in the range of 1-4 M. Furthermore, we observe that U begins to release from the co-precipitates at a [H+] value of 0.001 M while the release of Fe is observed at a [H+] value of 0.01 M. About 14±2% and 72±1% of the total U is observed to dissolve at [H+] values of 0.001 and 0.01 M, respectively, while only about 15±1% of the total Fe is observed to dissolve at a [H+] value of 0.01 M. Therefore, it is essential to determine the critical point at which Fe begins to release and identify the corresponding amount of the dissolved U(VI). We selected the [H+] value of 0.01 M to trace the variation of the accumulated solubility of Fe and U with time so that we can identify the critical point at which the release of Fe begins. As shown in Fig. 4, about 16–20% of U and a small amount of Fe dissolves during the first 0.5 h, indicating that this partial U is not related to the Fe precipitates and is likely to be absorbed on the surface of the precipitates. 11

Subsequently, about 52–56% of the U is gradually released and stabilized after 24 h accompanied by the gradual release of 18–19% of iron, suggesting the congruence of the partial U and Fe precipitation. Finally, the relationship between the release of Fe and U was analyzed. As shown in Fig. 5, the release process of U from the Fe-U coprecipitates can be divided into three stages: The first stage refers to a separate release of about 16–20% of sedimentary U before the dissolution of the Fe precipitates. The second stage involves a rapid release of about 52-56% of U along with a small quantity of Fe (18-19%) dissolution. The last stage is related to the release of the residual 27–29% of U and the main body of the Fe precipitates. Apparently, the last two parts of about 80–84% of U, corresponding to the U existing in the samples after acid dissolution, are considered to be distributed within the iron (oxyhydr)oxides, probably entrapped by the iron (oxyhydr)oxides particles or incorporated into their structure. 3.4. Structure of Fe-U tcoprecipitates XRD studies were conducted to compare the crystalline structure of the precipitates before and after acid dissolution (Fig. 6). We observe that the dominant characteristic

peaks

of

the

precipitates

correspond

to

metaschoepite

[(UO2)8O2(OH)12(H2O)10], this is consistent with the results of Duff and Taylor et al. (Taylor et al., 2015; Duff et al., 2002), according to which U(VI) is precipitated as a discrete crystalline uranyl phase at a relatively high initial concentration. The diffraction peaks centered at 19.98°, 23.8° and 27.8° are ascribed to uranyl hydroxide UO2(OH)2, while the peaks at 2θ = 11.9° and 26.8° are attributed to the (110) and 12

(130) crystal planes of iron oxide hydroxide [FeO(OH)], respectively. We also observe that the precipitates aged for a longer time exhibited sharper characteristic diffraction peaks, indicating the higher crystalline degree of the solids. Furthermore, we observe that while the significant characteristic peaks of metaschoepite and uranyl hydroxide were missing in the XRD pattern of the samples after acid dissolution, some minor peaks of the species remained. These results indicate the dissolution of partial crystalline uranium oxides after acid dissolution at 0.01 M of [H+] for 0.5 h, and the presence of uranyl compounds in samples after acid dissolution. In addition, the peak of FeU+O2 samples at 2θ≈11.9° is much sharper than that of samples after acid dissolution，which might result from the partial overlap of the peaks of iron oxide hydroxide and crystalline metaschoepite at 2θ = 11.9°. The result further reveal the low crystallinity of the Fe precipitates. While some previous studies have revealed that Fe mainly precipitates as well crystallized lepidocrocite during the coprecipitation of Fe(II) and U(VI) under neutral aerobic conditions at room temperature (Mei et al., 2018). The low crystallinity of the iron oxide hydroxide obtained in this study could be attributed to the low initial concentration of Fe(II). The absence of the iron oxide hydroxide diffraction peaks in the Fe+O2 samples indicates that the presence of U(VI) might facilitate the crystallization of the Fe precipitates. The FT-IR spectra of the samples before and after acid dissolution are shown in Fig. 7. The bands centered around 3384–3429 cm-1 and 1628 cm-1 are assigned to the stretching and deforming vibrations of the OH- groups and H2O molecules (Liu et al., 2007; Ristić et al., 2007), indicating the presence of small amounts of structural 13

hydroxyl groups and adsorbed H2O molecules on the solids(Supattarasakda et al., 2013). The relatively weak peaks at 914 cm-1 of the FeU+O2 samples before and after acid dissolution are caused by the stretching vibrations of the O=U=O moiety (Ding et al., 2018; Zhao et al., 2018; Li et al., 2019; Zhang et al., 2019a). Besides, the slight shift in the O=U=O peaks compared to the reported wavenumbers of 930 cm-1 might be ascribed to the de-protonation of a U(VI)-coordinated water molecule at neutral conditions and the formation of a neutral surface species, such as metaschoepite or uranyl hydroxide (Hadjittofi et al., 2014). The presence of the O=U=O group in the samples after acid dissolution also demonstrates the distribution of the uranyl components within the iron oxide hydroxide particles. The strong peaks around 470 and 530 cm-1 can be attributed to the symmetric Fe-O bonds vibrations (Jing et al., 2004; Supattarasakda et al., 2013; Lu et al., 2018). The peaks at 794 and 1034 cm-1 corresponding to the typical goethite and lepidocrocite bands indicate the presence of the FeO(OH) structure in the precipitates, which is also demonstrated by the XRD spectra (Jesús et al., 2000; Doornbusch et al., 2015; Mei et al., 2018). The broadening of the 1034 cm-1 band for FeO(OH) is probably due to the presence of carbonates as amorphous iron(III)-hydroxide is very susceptible to CO2 (Ristić et al., 2007). The Tyndall effect experiment shows that an obvious light path appears in the suspension using the red-light source as the incident light, which elucidates the formation of iron oxide hydroxide colloid. 3.5. Morphology and composition of Fe-U coprecipitates The morphologies and composition of the precipitates (FeU+O2) before and after 14

acid dissolution are identified by high resolution TEM. The FeU+O2 samples exhibit a dense main structure with many thin films on the edge (Fig 8a), and the thin films disappear after acid dissolution (Fig 8c). Combined with the result of acid dissolution experiments and XRD analysis, the shin films might correspond to the crystalline metaschoepite and uranyl hydroxide, adsorbing on the main structure of the iron oxide hydroxide. The well-defined rings of the SAED pattern in Fig 8a also indicate the polycrystallinity of the Fe-U coprecipitates (FeU-O2) (Mei et al., 2018). Fig 8b shows the clear lattice fringes with the d-spacing of 0.29 and 0.19 nm, which approximately correspond to the (101) and (110) lattice planes of the iron oxide hydroxide. The needle-like structures are observed in Fig 8c-d, which might due to the formation of a small amount of goethite (Qian et al., 2014). Moreover, it is observed that no obvious lattice fringes appear in the Fig 8d, which might result from the partly broken lattice structure during the acid dissolution (Fu et al., 2019; Geng et al., 2020). The SEM images also show that the FeU+O2 sediments present a rough surface with a very platy structure interspersed on the surfaces of the spherical aggregates (Fig. S2). In the case of the samples after acid dissolution, the platy structure disappears after acid dissolution, accompanied by the appearance of some deep gullies on the surface of the precipitates. This result is consistent with the HRTEM analysis. Moreover, the decline of the relative content of U in the samples after acid dissolution in EDS analysis (Fig. 8g-h) also testifies the release of U after the acid dissolution. 3.6. Chemical states of Uranium in the precipitates XPS is a standard method to characterize the oxidation states of U. The overall 15

XPS spectra of the samples before and after acid dissolution (Fig. 9a) show that Fe, O, and U with clear peaks are the main elements in the co-precipitates. The single O 1s peak of the samples after acid dissolution could be decomposed into three peaks at about 530.43, 531.88 and 533.15 eV, which correspond to O2-, OH- and adsorbed H2O, respectively (Liu et al., 2014). Fig. 9c and d show the high-resolution Fe 2p spectra of the samples before and after acid dissolution. The Fe 2p3/2 peaks of the two samples at ~711.29 and 711.43 eV are narrower and stronger than their Fe 2p1/2 peaks at ~725.07 and 725.03 eV, respectively. The Fe 2p3/2 peaks at ~710.9 and 712.5 eV correspond to the binding energy of Fe(III) including hematite and iron oxide hydroxide (Li et al., 2007; Wen et al., 2016; Jin et al., 2018; Zhang et al., 2019b). The distinct Fe 2p3/2 satellites centered at ~8.01–8.05 eV above the primary peaks are also a crucial characteristic of Fe(III) (Yamashita et al., 2008). For the crystal structure of iron oxide hydroxides, three quarters of the Fe(III) cations of the γ- compounds are octahedral and one quarter tetrahedral, while all Fe(III) cations of the α-compounds are octahedral (Grosvenor et al., 2004; Mei et al., 2018). In this study, the Fe 2p3/2 spectra are well fitted by two peaks corresponding to octahedral and tetrahedral Fe(III), which indicate that the Fe precipitates are likely to be γ- compounds. Moreover, the great comparability of the two high-resolution Fe 2p spectra also indicate that the Fe precipitates did not dissolve after the acid dissolution at 0.01 M of [H+] for 0.5 h. The high-resolution U 4f spectra of the two precipitates are shown in Fig. 10e and f. Both the primary U 4f7/2 (382.1–382.2 eV) and U 4f5/2 (392.9–393.1 eV), with 16

an energy separation of ~10.8 eV, are decomposed into two peaks, indicating two chemical states of U in the solids. Ilton et al. (Ilton et al., 2011; Mei et al., 2018) reported that the binding energy separation between the primary spin-orbit split peaks of U(IV), U(V), and U(VI) and their associated satellite peaks are ∼6.0−7.0, ∼7.8−8.5, and ∼4.0 eV, respectively. The U 4f7/2 satellites of the two precipitates at ~8.5 eV above their corresponding primary peaks demonstrate the presence of U(V) in the precipitates before and after acid dissolution. Similarly, the satellites of the two precipitates at ~3.2 eV above their corresponding primary peaks prove the presence of U(VI) in the samples. Moreover, the higher relative contents of U(V) within samples after acid dissolution also indicates the release of U(VI) during acid dissolution and the strong stability of U(V) under acidic conditions. 3.7. Mechanism of uranium precipitation by low concentration of Fe(II) under neutral oxygen-rich conditions Figure 11 shows the mechanism of Fe-U coprecipitation with the presence of a low concentration of Fe(II) under neutral oxygen-rich conditions. During the precipitation of Fe and U, the reduction of U(VI) to U(V) is suggested to occur primarily through one-electron transfer from Fe(II) to U(VI) as shown in equation (2). The result is consistent with former investigations, which reported that U(V) is stabilized into the crystal structure of goethite or magnetite through one electron transfer from soluble or structural Fe(II) to U(VI) (Massey et al., 2014; Doornbusch et al., 2015; Yuan et al., 2015a; Yuan et al., 2015b). Then, the direct precipitation of the oxidized Fe(III) to Fe(OH)3 occurs instantly as shown in equation (3), accompanied 17

by the precipitation of a few unoxidized Fe(II) as shown in equation (4) and the subsequent oxidation of Fe(OH)2 by oxygen as shown in equation (5). The subsequent transformation of Fe(OH)3 into a stable FeO(OH) structure is confirmed by XRD, FTIR, and HRTEM analyses. Zhang et al. also revealed the analogical precipitation process of Fe(II) (Zhang et al., 2019a). Furthermore, the Tyndall effect experiment demonstrated the presence of iron oxide hydroxide colloform, which flocculates and settles quickly when the shaking stops. The XPS analysis reveals that the iron oxide hydroxide include both octahedral and tetrahedral Fe(III) cations, which correspond to the γ- compounds. Thus, the structure of γ- FeO(OH) is likely to be formed in this study. U(VI)O22+ + Fe2+→ U(V)O2+ + Fe3+

(2)

Fe3+ + 3OH- ↔ Fe(OH)3

(3)

Fe2+ + 2OH- ↔ Fe(OH)2

(4)

4Fe(OH)2 + O2 + 2H2O ↔ 4Fe(OH)3

(5)

The XPS and BCR analyses demonstrate that the reduced U(V) is recalcitrant to extensive acid dissolution and oxidation of oxygen. Indeed, the structurally incorporated U(V) was reportedly stable in magnetite for up to 550 days under anoxic conditions and 142 days under ambient conditions (Huber et al., 2012; Pidchenko et al., 2017). Moreover, U(V) was also detected in hematite (but not U(IV)) at an Eh value of +0.21 V (pH=7.1~7.3) (Ilton et al., 2012), indicating that the U(V) incorporated into the iron oxide hydroxides may be stable over extended time periods and

varying

environments.

However, 18

the

disproportionation

of

UO2+

(2U(V)→U(IV)+U(VI)) on the surface of particulate magnetite was also highlighted in some recent literatures (Yuan et al., 2015b). Thus, the partial U(V), which remains stable under oxygen-rich conditions and is recalcitrant to the acid-dissolving, is suggested to be incorporated into the iron oxide hydroxide structure during the transformation of Fe(OH)3 into FeO(OH). Meanwhile, the incorporation of U(VI) (with ionic radii of 0.73-0.81 Å) into the structure of the iron oxide hydroxide is also suggested to be achieved through the substitution of trans-dioxo uranyl-like configurations for structural Fe(III) (with ionic radii of 0.645 Å) (Ilton et al., 2012; Marshall et al., 2014; McBriarty et al., 2018; Duff et al., 2002). Therefore, it is likely that the partial U(VI) is also incorporated into the iron oxide hydroxide structure. The XRD spectra show the formation of metaschoepite and uranyl oxides during the coprecipitation of U(VI) and Fe(II), and some of whose distinct characteristic peaks disappear after acid dissolution at 0.01 M of [H+] for 0.5 h. According to the acid dissolution experiments, U probably exists in three forms within the precipitates: firstly, 16-20% of the precipitated U forms the discrete crystalline uranyl phase, i.e., metaschoepite and uranyl oxides, which is adsorbed on the surface of the precipitates. Secondly, 52–56% of U might be adsorbed on the internal channel and clearance of iron oxide hydroxides as the discrete uranyl compounds, which is entrapped by the iron oxides. Ilton et al. reported that up to 30% of the total U is adsorbed in the pores within the hematite structure with uranyl components during the coprecipitation of U(VI) and Fe(III) (Ilton et al., 2012). Finally, 27–29% of U(V) and U(VI) is considered to be incorporated into the FeO(OH) structure during the oxidation, 19

deprotonation and precipitation of Fe(II). Thus, these results show that the presence of a low concentration of Fe(II) under neutral oxygen-rich conditions leads to the reduction of U(VI) to U(V) and significantly facilitates the immobilization of aqueous U(VI). Meanwhile, O2 plays a significant role in the enhanced Fe-U coprecipitation. Moreover, the solids exhibit strong stability under oxygen-rich or weak-acid conditions. 4. Conclusions The mechanism of the immobilization of U(VI) induced by a small quantity of Fe(II) under neutral oxygen-rich conditions and the stability of stationary U were investigated by an acid dissolution experiment and characterization analysis of the precipitates. The main conclusions are as follows: (1) The sedimentation rate of U(VI) induced by a low concentration of Fe(II) under oxygen-rich conditions is approximately 96% and about 16% higher than that under oxygen-free conditions, and the optimal Fe/U molar ratio is about 2. (2) Fe mainly precipitates as low-crystalline iron oxide hydroxide FeO(OH). The Fe-U coprecipitates remain stable under neutral, slightly acidic and oxygen-rich conditions. (3) The U(V) generated via one-electron reduction is somewhat recalcitrant to the oxidation of O2 and acid dissolution, suggesting its incorporation into the iron oxide hydroxide structure. (4) The immobilization of U(VI) by a low concentration of Fe(II) under neutral oxygen-rich conditions refers to three primary mechanisms namely surface 20

adsorption, encapsulation and incorporation into Fe (oxyhydr)oxides. First, 16– 20% of U forms metaschoepite and uranyl hydroxide, which is adsorbed on the surface of solids. Then 52-56% of U is adsorbed at the internal pores of the iron oxide hydroxide as discrete uranyl phases. Finally, 27–29% of the residual uranium is likely to be incorporated into the iron oxide hydroxide structure as U(V) and U(VI). (5) Nearly 70% of the U could be extracted by a hydrochloric acid solution with 0.01 M [H+] for 24 h; this reveals that U immobilization by low concentration of Fe(II) combined with oxygen has potential applications in uranium separation and recycling. In general, the coprecipitation of U(VI) and small amounts of Fe(II) under oxygen-rich neutral conditions is a promising method for U-contained wastewater treatment, which offers the advantages such as enhanced aqueous U(VI) sequestration efficiency, fewer production of solid residue, lower costs as well as the reduction of negative impact on the environment. Acknowledgments This work was supported by the National Natural Science Foundation of China (11705082, 51874180 and51704169), Natural Science Foundation of Hunan Province (2018JJ3434), the double first class construct program of USC (2017SYL05) and the graduate research innovation project of USC (193YXC005). References Bae, S., Hanna, K., 2015. Reactivity of nanoscale zero-valent iron in unbuffered 21

systems: effect of pH and Fe(II) dissolution. Environ. Sci. Technol. 49(17), 10536-10543. Bi, Y., Stylo, M., Bernier-Latmani, R., Hayes, K.F., 2016. Rapid mobilization of noncrystalline U(IV) coupled with FeS oxidation. Environ. Sci. Technol. 50(3), 1403-1411. Burns, P.C., Hawthorne F.C., 2005. U6+ minerals and inorganic compounds: insights into an expanded structural hierarchy of crystal structures. Can. Mineral. 43, 1839–1894. Cerrato, J.M., Ashner, M.N., Alessi, D.S., Lezama-Pacheco, J.S., Bernier-latmani, R., Bargar, J.R., Giammar, D.E., 2013. Relative reactivity of biogenic and chemogenic uraninite and biogenic noncrystalline U(IV). Environ. Sci. Technol. 47(17), 9756-9763. Chakraborty, S., Favre, F., Banerjee, D., Scheinost, A.C., Mullet, M., Ehrhardt, J.J. , Brendle, J., Vidal, L., Charlet, L., 2010. U(VI) sorption and reduction by Fe(II) sorbed on montmorillonite. Environ. Sci. Technol. 44(10), 3779-3785. Chen, A., Shang, C., Shao, J., Zhang, J., Huang, H., 2017. The application of iron-based technologies in uranium remediation: A review. Sci. Total Environ. 575, 1291-1306. Crane, R.A., Dickinson, M., Popescu, I.C., Scott, T.B., 2011. Magnetite and zero-valent iron nanoparticles for the remediation of uranium contaminated environmental water. Water Res. 45(9), 2931-2942. Ding, L., Tan, W.F., Xie, S.B., Mumford, K., Lv, J.W., Wang, H.Q., Fang, Q., Zhang, X.W., Wu, X.Y., M. Li. 2018. Uranium adsorption and subsequent re-oxidation under aerobic conditions by Leifsonia sp. - Coated biochar as green trapping agent. Environ. Pollut. 242(Pt A), 778-787. 22

Doornbusch, B., Bunney, K., Gan, B.K., Jones, F., Gräfe, M., 2015. Iron oxide formation from FeCl2 solutions in the presence of uranyl (UO22+) cations and carbonate rich media. Geochim. Cosmochim. Acta 158, 22-47. Du, X., Boonchayaanant, B., Wu, W.M., Fendorf, S., Bargar, J., Criddle, C.S., 2011. Reduction of uranium(VI) by soluble iron(II) conforms with thermodynamic predictions. Environ. Sci. Technol. 45(11), 4718-4725. Duff, M.C., Coughlin, J.U., Hunter, D.B., 2002. Uranium co-precipitation with iron oxide minerals. Geochim. Cosmochim. Acta 66(20), 3533-3547. Fu, Y., Qin, L., Huang, D., Zeng, G.M., Lai, C., Li, B.S., He, J.F., Yi, H., Zhang, M.M., Cheng, M., Wen X.F., 2019. Chitosan functionalized activated coke for Au nanoparticles anchoring: Green synthesis and catalytic activities in hydrogenation of nitrophenols and azo dyes. Appl. Catal., B 255, 117740. Geng, H.H., Wang, F., Yan, C.C., Tian, Z.J., Chen, H.L., Zhou, B.H., Yuan, R.F., Yao, J., 2020. Leaching behavior of metals from iron tailings under varying pH and low-molecular-weight organic acids. J. Hazard. Mater. 383, 121136. Grosvenor, A.P., Kobe, B.A., Biesinger, M.C., McIntyre, N.S., 2004. Investigation of multiplet splitting of Fe 2p XPS spectra and bonding in iron compounds. Surf. Interface Anal. 36(12), 1564-1574. Guo, Z., Li, Y., Wu, W., 2009. Sorption of U(VI) on goethite: effects of pH, ionic strength, phosphate, carbonate and fulvic acid. Appl. Radiat. Isot. 67(6), 996-1000. Hadjittofi, L., Pashalidis, I., 2014. Uranium sorption from aqueous solutions by activated biochar fibres investigated by FTIR spectroscopy and batch experiments. J. Radioanal. Nucl. Chem. 304(2), 897-904. Huang, S., Pang, H., Li, L., Jiang, S., Wen, T., Zhuang, L., Hu, B., Wang X., 2018. 23

Unexpected ultrafast and high adsorption of U(VI) and Eu(III) from solution using porous Al2O3 microspheres derived from MIL-53. Chem. Eng. J. 353, 157-166. Huber, F., Schild, D., Vitova, T., Rothe, J., Kirsch, R., Schäfer, T., 2012. U(VI) removal kinetics in presence of synthetic magnetite nanoparticles. Geochim. Cosmochim. Acta 96, 154-173. Ilton, E.S., Bagus, P.S., 2011. XPS determination of uranium oxidation states. Surf. Interface Anal. 43(13), 1549-1560. Ilton, E.S., Boily, J.F., Buck, E.C., Skomurski, F.N., Rosso, K.M., Cahill, C.L. ， Bargar J.R., Felmy A.R., 2010. Influence of dynamical conditions on the reduction of U(VI) at the magnetite-solution interface. Environ. Sci. Technol. 44(1), 170-176. Ilton, E.S., Pacheco, J.S., Bargar, J.R., Shi, Z., Liu, J., Kovarik, L., Engelhard, M.H., Felmy, A.R., 2012. Reduction of U(VI) incorporated in the structure of hematite. Environ. Sci. Technol. 46(17), 9428-9436. Jesús, C., Vidal, B., José, T., 2000. Effect of phosphate on the formation of nanophase lepidocrocite from Fe(II) sulfate. Clays Clay Miner. 48(5). Jin, Q., Zhang, S., Wen, T., Wang, J., Gu, P., Zhao, G., Wang, X., Chen, Z., Hayat, T., Wang, X., 2018. Simultaneous adsorption and oxidative degradation of Bisphenol A by zero-valent iron/iron carbide nanoparticles encapsulated in N-doped carbon matrix. Environ. Pollut. 243, 218-227. Jing Z.H., Wu S.H., 2004. Synthesis and characterization of monodisperse hematite nanoparticles modified by surfactants via hydrothermal approach. Mater. Lett. 58(27-28), 3637-3640. Latta, D.E., Mishra, B., Cook, R.E., Kemner, K.M., Boyanov, M.I., 2014. Stable 24

U(IV) complexes form at high-affinity mineral surface sites. Environ. Sci. Technol. 48(3), 1683-1691. Li, L., Huang S., Wen T., Ma R., Yin L., Li J., Chen Z., Hayat T., Hu B., Wang X., 2019. Fabrication of carboxyl and amino functionalized carbonaceous microspheres and their enhanced adsorption behaviors of U(VI). J. Colloid Interface Sci. 543, 225-236. Li, X.Q., Zhang, W.X., 2007. Sequestration of metal cations with zerovalent iron nanoparticles-A study with high resolution X-ray photoelectron spectroscopy (HR-XPS). J. Phys. Chem. C 111(19), 6939-6946. Li, Z.J., Wang, L., Yuan, L.Y., Xiao, C.L., Mei, L., Zheng, L.R., Zhang, J., Yang, J.H., Zhao, Y.L., Zhu, Z.T., Chai, Z.F., Shi, W.Q., 2015. Efficient removal of uranium from aqueous solution by zero-valent iron nanoparticle and its graphene composite. J. Hazard. Mater. 290, 26-33. Liu, A., Liu, J., Pan, B., Zhang, W.X., 2014. Formation of lepidocrocite (γ-FeOOH) from oxidation of nanoscale zero-valent iron (nZVI) in oxygenated water. RSC Adv. 4(101), 57377-57382. Liu, H., Wei, Y., Li, P., Zhang, Y., Sun, Y., 2007. Catalytic synthesis of nanosized hematite particles in solution. Mater. Chem. Phys. 102(1), 1-6. Liu, H., Zhu, Y., Xu, B., Li, P., Sun, Y., Chen, T., 2017. Mechanical investigation of U(VI) on pyrrhotite by batch, EXAFS and modeling techniques. J. Hazard. Mater. 322(Pt B), 488-498. Liu, X., Ma, R., Wang, X., Ma, Y., Yang, Y., Zhuang, L., Zhang, S., Jehan, R., Chen, J., Wang, X., 2019. Graphene oxide-based materials for efficient removal of heavy metal ions from aqueous solution: A review. Environ. Pollut. 252(Pt A), 62-73. 25

Lu, B.Q., Li, M., Zhang, X.W., Huang, C.M., Wu X.Y., Fang Q., 2018. Immobilization of uranium into magnetite from aqueous solution by electrodepositing approach. J. Hazard. Mater. 343, 255-265. Marshall T.A., Morris K., Law, G.T.W., Mosselmans J.F.W., Bost, P., Roberts, H., Shaw, S., 2015. Uranium fate during crystallization of magnetite from ferrihydrite in conditions relevant to the disposal of radioactive waste. Mineral. Mag. 79(6), 1265-1274. Marshall, T.A., Morris, K., Law, G.T.W., Livens, F.R., Mosselmans, J.F.W., Bots, P., Shaw, S., 2014. Incorporation of Uranium into Hematite during Crystallization from Ferrihydrite. Environ. Sci. Technol. 48(7), 3724-3731. Massey, M.S., Lezama-Pacheco, J.S., Jones, M.E., Ilton, E.S., Cerrato, J.M., Bargar, J.R., Fendorf, S., 2014. Competing retention pathways of uranium upon reaction with Fe(II). Geochim. Cosmochim. Acta 142, 166-185. McBriarty, M.E., Kerisit, S., Bylaska, E.J., Shaw, S., Morris, K., Ilton, E.S., 2018. Iron vacancies accommodate uranyl incorporation into hematite. Environ. Sci. Technol. 52(11), 6282-6290. Mei, H., Tan, X., Tan, L., Meng, Y., Chen, C., Fang, M., Wang, X., 2018. Retention of U(VI) by the formation of Fe precipitates from oxidation of Fe(II). ACS Earth Space Chem. 2(10), 968-976. Pidchenko, I., Kvashnina, K.O., Yokosawa, T., Finck, N., Bahl, S., Schild, D., Polly, R., Bohnert, E., Rossberg, A., Gottlicher, J., Dardenne, K., Rothe, J., Schafer, T., Geckeis, H., Vitova, T., 2017. Uranium redox transformations after U(VI) coprecipitation with magnetite nanoparticles. Environ. Sci. Technol. 51(4), 2217-2225. Qian, L.P., M. M.H., Cheng, D.H., 2014. Adsorption and desorption of uranium on 26

nano goethite and nano alumina. J. Radioanal. Nucl. Chem. 303(1), 161-170. Ristić, M., Grave, E.D., Musić, S., Popović, S., Orehovec, Z., 2007. Transformation of low crystalline ferrihydrite to α-Fe2O3 in the solid state. J. Mol. Struct. 834-836, 454-460. Roberts, H.E., Morris, K., Law, G. T.W., Mosselmans, J. F.W., Bots, P., Kvashnina, K., Shaw, S., 2017. Uranium(V) incorporation mechanisms and stability in Fe(II)/Fe(III) (oxyhydr)oxides. Environ. Sci. Technol. Lett. 4(10), 421-426. Skomurski, F.N., Ilton, E.S., Engelhard, M.H., Arey, B.W., Rosso, K.M., 2011. Heterogeneous reduction of U6+ by structural Fe2+ from theory and experiment. Geochim. Cosmochim. Acta 75(22), 7277-7290. Supattarasakda, K., Petcharoen, K., Permpool, T., Sirivat, A., Lerdwijitjarud, W., 2013. Control of hematite nanoparticle size and shape by the chemical precipitation method. Powder Technol. 249, 353-359. Taylor, S.D., Marcano, M.C., Rosso, K.M., Becker, U., 2015. An experimental and ab initio study on the abiotic reduction of uranyl by ferrous iron. Geochim. Cosmochim. Acta 156, 154-172. Tsarev, S., Collins, R.N., Fahy, A., Waite T.D., 2016. Reduced uranium phases produced from anaerobic reaction with nanoscale zerovalent iron. Environ. Sci. Technol. 50(5), 2595-2601. Tsarev, S., Collins R.N., Ilton, E.S., Fahy, A., Waite, T.D., 2017. The short-term reduction of uranium by nanoscale zero-valent iron (nZVI): role of oxide shell, reduction mechanism and the formation of U(V)-carbonate phases. Environ. Sci.: Nano 4(6), 1304-1313. Wang, X., Chen, L., Wang, L., Fan, Q., Pan, D., Li, J., Chi, F., Xie, Y., Yu, S., Xiao, C., Luo, F., Wang, J., Wang, X., Chen, C., Wu, W., Shi, W., Wang, S., Wang, 27

X., 2019. Synthesis of novel nanomaterials and their application in efficient removal of radionuclides. Sci. China: Chem. 62(8), 933-967. Wei, J.B., Zhang, X.F., Liu, Q., Li, Z.S.,. Liu, L.H., Wang, J., 2014. Magnetic separation of uranium by CoFe2O4 hollow spheres. Chem. Eng. J. 241, 228-234. Wen, T., Wang, X., Wang, J., Chen, Z., Li, J., Hu, J., Hayat, T., Alsaedi, A., Grambow, B., Wang, X., 2016. A strategically designed porous magnetic N-doped Fe/Fe3C@C matrix and its highly efficient uranium(VI) remediation. Inorg. Chem. Front. 3(10), 1227-1235. Wu, X., Huang, Q., Mao, Y., Wang, X., Wang, Y., Hu, Q., Wang, H., Wang, X., 2019. Sensors for determination of uranium: A review. TrAC, Trends Anal. Chem. 118, 89-111. Yamashita, T., Hayes P., 2008. Analysis of XPS spectra of Fe2+ and Fe3+ ions in oxide materials. Appl. Surf. Sci. 254(8), 2441-2449. Yang, Z., Kang, M., Ma, B., Xie, J., Chen, F., Charlet, L., Liu, C., 2014. Inhibition of U(VI) reduction by synthetic and natural pyrite. Environ. Sci. Technol. 48(18), 10716-10724. Yuan, K., Ilton, E.S., Antonio, M.R., Li, Z., Cook, P.J., Becker, U., 2015a. Electrochemical and Spectroscopic Evidence on the One-Electron Reduction of U(VI) to U(V) on Magnetite. Environ. Sci. Technol. 49(10), 6206-6213. Yuan, K., Renock, D., Ewing, R.C., Becker, U., 2015b. Uranium reduction on magnetite: Probing for pentavalent uranium using electrochemical methods. Geochim. Cosmochim. Acta 156, 194-206. Zhang, Q., Zhao, D., Ding, Y., Chen, Y., Li, F., Alsaedi, A., Hayat, T., Chen, C., 2019a. Synthesis of Fe–Ni/graphene oxide composite and its highly efficient removal of uranium(VI) from aqueous solution. J. Cleaner Prod. 230, 1305-1315. 28

Zhang, Q., Zhao, D., Feng, S., Wang, Y., Jin, J., Alsaedi, A., Hayat, T., Chen, C., 2019b. Synthesis of nanoscale zero-valent iron loaded chitosan for synergistically enhanced removal of U(VI) based on adsorption and reduction. J. Colloid Interface Sci. 552, 735-743. Zhao, D., Wang, X., Yang, S., Guo, Z., Sheng, G., 2012. Impact of water quality parameters on the sorption of U(VI) onto hematite. J. Environ. Radioact. 103(1), 20-29. Zhao, D., Zhu, H., Wu, C., Feng, S., Alsaedi, A., Hayat, T., Chen, C., 2018. Facile synthesis of magnetic Fe3O4/graphene composites for enhanced U(VI) sorption. Appl. Surf. Sci. 444.

Figure captions: Fig. 1. Sedimentation rate of U(VI) versus reaction time under oxygen-rich/ oxygen-free conditions for high and low Fe/U molar ratios. Fig. 2. Sedimentation rate of U(VI) versus different Fe/U molar ratios. Fig. 3. Cumulative dissolution of U and Fe from the Fe-U coprecipitations as a function of [H+]. Fig. 4. Accumulated dissolution of U and Fe for different dissolving times at 0.01 mol/L of [H+]. Fig. 5. Uranium release with iron from the Fe-U coprecipitations. Fig. 6. XRD patterns of the Fe+O2 and FeU+O2 samples at different durations and samples after acid dissolution. Fig. 7. FT-IR spectra of the Fe+O2 and FeU+O2 samples and samples after acid 29

dissolution. Fig. 8. HRTEM images and EDS spectra of the solids before and after acid dissolution. (a-b, FeU+O2; c-d, FeU+O2 after acid dissolution) Fig. 9. XPS survey spectra of the samples before and after acid dissolution. Fig. 10. Schematic diagram of Fe-U coprecipitation under neutral oxygen-rich conditions. Highlights: 

The precipitation rate of U(Ⅵ) by Fe(Ⅱ) under O2-rich conditions is up to 96%.



16–20% of U is adsorbed on the surface of the solids as uranyl compounds.



52–56% of U is absorbed at the internal pores of Fe precipitates as uranyl phases.



27-29% of U is likely to be incorporated into the FeO(OH) structure.



U(Ⅵ) partially precipitates as U(V) that remains stable under O2-rich conditions.

30

Fig. 1. Sedimentation rate of U(VI) versus reaction time under oxygen-rich/ oxygen-free conditions for high and low Fe/U molar ratios.

Fig. 2. Sedimentation rate of U(VI) versus different Fe/U molar ratios.

Fig. 3. Cumulative dissolution of U and Fe from the Fe-U coprecipitations as a function of [H+].

31

Fig. 4. Accumulated dissolution of U and Fe for different dissolving times at 0.01 mol/L of [H+].

Fig. 5. Uranium release with iron from the Fe-U coprecipitations.

32

Fig. 6. XRD patterns of the Fe+O2 and FeU+O2 samples at different durations and samples after acid dissolution.

Fig. 7. FT-IR spectra of the Fe+O2 and FeU+O2 samples and samples after acid dissolution.

33

Fig. 8. HRTEM images and EDS spectra of the solids before and after acid dissolution. (a-b, FeU+O2; c-d, FeU+O2 after acid dissolution) (b)

34

(a)

(d)

(c)

Fig. 9. XPS survey spectra of the samples before and after acid dissolution.

35

Fig. 10. Schematic diagram of Fe-U coprecipitation under neutral oxygen-rich conditions.

36