Effect of silicate on U(VI) sorption to γ-Al2O3: Batch and EXAFS studies

Chemical Engineering Journal 269 (2015) 371–378

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Effect of silicate on U(VI) sorption to c-Al2O3: Batch and EXAFS studies Huiyang Mei a,b, Xiaoli Tan a,c,⇑, Shujun Yu a,b, Xuemei Ren a, Changlun Chen a, Xiangke Wang c,d,e,⇑ a

Institute of Plasma Physics, Chinese Academy of Sciences, P.O. Box 1126, Hefei 230031, PR China University of Science and Technology of China, Hefei 230026, PR China c School for Radiological and Interdisciplinary Sciences (RAD-X), Soochow University, 215123 Suzhou, PR China d Collaborative Innovation Center of Radiation Medicine of Jiangsu Higher Education Institutions, PR China e Faculty of Engineering, King Abdulaziz University, Jeddah 21589, Saudi Arabia b

h i g h l i g h t s  Effects of U(VI) and silicate on their interaction with

c-Al2O3 is investigated.

 Investigate the microstructure of the ternary complexes by EXAFS spectroscopy.  Obtain the sorption mechanism of U(VI) and silicate on

a r t i c l e

i n f o

Article history: Received 25 December 2014 Received in revised form 26 January 2015 Accepted 29 January 2015 Available online 7 February 2015 Keywords: U(VI) Silicate c-Al2O3 Sorption EXAFS

c-Al2O3.

a b s t r a c t The effect of soluble silicate on the sorption of U(VI) to c-Al2O3 was investigated by batch experiments and extended X-ray absorption fine structure (EXAFS) method. The presence of silicate enhanced the sorption of U(VI) on c-Al2O3 surface and the sorption was attributed to inner-sphere surface complexation. The structure of the adsorbed U(VI) and silicate on c-Al2O3 was investigated in the analysis of EXAFS spectra. The fitting of the experimental EXAFS data was obtained by including two uranium coordination shells with 2 axial (Oax) and 5 equatorial (Oeq) oxygen atoms at 1.79 ± 0.02 and 2.43 ± 0.02 Å, respectively, and the third coordination shells with Al atom at 3.35 Å. Silicate contributed to the formation of ternary inner-sphere surface complexes, acting as ‘‘bridge’’ between U(VI) and c-Al2O3 and enhanced the sorption of U(VI). The observations suggested that the interactions between U(VI) and silicate were important in controlling U(VI) retention. Ó 2015 Elsevier B.V. All rights reserved.

1. Introduction Migration of radionuclides in natural aqueous system deserves ongoing attention in environmental research. Knowledge of the sorption of radionuclides in the environment is vital for predicting the risk caused by long-term storage of nuclear waste and also for assessing the migration of radionuclides in the environment near population centers. Natural uranium is mined for nuclear fuel for reactors and exists in the ecosphere as a mineral as well as a waste of the nuclear industry. Uranium is a toxic and representative radioactive element which is usually found in the environment in the hexavalent form as an oxo-cation (UO2+ 2 ). Studies of the

⇑ Corresponding authors at: Institute of Plasma Physics, Chinese Academy of Sciences, P.O. Box 1126, Hefei 230031, PR China. Tel.: +86 551 65592788; fax: +86 551 65591310 (X. Tan). School for Radiological and Interdisciplinary Sciences (RADX), Soochow University, 215123 Suzhou, PR China. Tel.: +86 551 65592788; fax: +86 551 65591310 (X. Wang). E-mail addresses: [email protected] (X. Tan), [email protected] (X. Wang). http://dx.doi.org/10.1016/j.cej.2015.01.121 1385-8947/Ó 2015 Elsevier B.V. All rights reserved.

sorption of U(VI) ions onto mineral/water interface provide information on their geochemical behavior in peculiar environments where radionuclides’ migration is a concern, such as waste repository sites and mines [1–5]. Sorption of U(VI) ion onto mineral surfaces has been extensively studied since this process has a significant effect on transport properties [1–12]. Various experimental techniques, such as batch experiments, extended X-ray absorption fine-structure spectroscopy (EXAFS) [6,7], timeresolved laser-induced fluorescence (TRLFS) spectroscopy [7–9], and density functional theory (DFT), etc. [9–12], are used to explore the sorption of U(VI) on the surfaces of (hydr)oxides and clay minerals. Most researches were mainly focused on the investigation of a single solute adsorbed on minerals [1,6–12]. However, natural environments are usually multi-component or multiphase systems, where the effect of a solute on the sorption of the contaminant on minerals could not be simply ignored [13,14]. The hazardous cations and anions have been proved to affect the surface reactivity of the minerals as well as their sorption property in

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the subsurface, and their migration, transformation and bio-availability [15]. A few researchers have evaluated the potential influences of anions (phosphate, carbonate, and sulfate) [16–19], competitive cations [20,21] and humic substances [19,22,23] on the sorption of the U(VI) ions. The enhancement of the sorption by the formation of cation–anion-surface ternary complexes have been proposed in the sorption of U(VI) on minerals. In the presence of phosphate, U(VI)-phosphate precipitation was found for the interaction of high concentration of U(VI) with phosphate on oxides [16,18]. While the presence of calcium ions decreased the sorption of U(VI) due to the formation of highly soluble and stable uranyl–carbonate and calcium–uranyl–carbonate-complexe [20,21]. Logue et al. [22] and Krˇepelová et al. [23] have also found that surface coated organic complexing agents significantly influenced the sorption of U(VI) by altering the natural chemistry and physics of the minerals. Soluble silicates come into the natural system due to the dissolution of clays and silica (present as an impurity in many minerals), or the degradation of glass canisters. As one of the major species found in environment, the interactions of silicates with U(VI) have great importance in predictively modeling the geochemical behavior of nuclear wastes in underground repository [24]. However, as far as we know, only a few reports were related to the solution complexation of U(VI) with silicate [5,25,26]. The U(VI)–silicate complexation complicates the investigation of U(VI) sorption. However, studies of silicates sorption properties are rather scarce though their presence on adsorbent surfaces can modify surface charge and site availability, which could influence whole surface reactivity of the adsorbents [27–29]. Indeed, the presence of the soluble silicates may have significant consequences on the sorption of radionuclides [30,31]. They can compete with the objective elements and thus involve a decrease of sorption capacities of studied adsorbent, and also can modify the surface affinity for the neutral and ionic species by forming chemical bonds with surface active groups or by modifying the electrical double layer. It was reported that the dissolved silicates could increase significantly the retention of Cs(I) onto the surface of magnetite [30]. Whereas precipitation, controlled largely by sediment Si release, became increasingly important at longer time and higher sorbate amount [31]. As mentioned above, an insight into the effect of silicate on radionuclide retention behavior is of importance to understand the physicochemical behavior of radionuclides in natural environments [30,31], and few was focused on the effect of silicate on the retention of U(VI) at mineral/water interface. In a comprehensive study on the mechanism and structure of U(VI) sorption on natural Si-/Al- and Fe-rich gels, Allard et al. proposed a two step U-uptake process [32]. The initial complexation of U(VI) by Si or Al, followed by trapping of these complexes within hydrous ferric oxides during precipitation. Ulrich et al. suggested that the influence of dissolved silicate on U(VI)–ferrihydrite surface complexation in Fe(III)-rich acidic mine water could not be detectable in slightly acidic conditions using EXAFS [33]. It was undeniably that the uranyl–silicate complexation could possible play a role in the migration of U(VI) in natural environment. However, understanding of how adsorbed silicate affects U(VI) sorption is currently unclear and needs further study. In addition, the priorities of U(VI) ions and silicate interaction with minerals received little attention when they coexisted in the system. In this work, we aimed to explain the mechanism of U(VI) transport in subsurface media influenced by soluble silicate. c-Al2O3 was selected as Si-free adsorbent in order to simplify the sorption system. Batch and spectroscopic techniques were adopted to investigate the interactions of U(VI) and silicate on the c-Al2O3/ water interface. EXAFS spectroscopy was utilized to investigate the microstructure of U(VI)–silicate complexes binding to sorption sites and thus to evaluate the sorption mechanism at molecular scale.

2. Experimental 2.1. Materials and methods The c-Al2O3 powder (Aluminum Oxide C, purity of 99.6%, specific surface area of 100 m2/g, primary particle size of 20 nm and site density of 1 nm2) was used directly without further purification [14,34]. c-Al2O3 has an pHpzc value of 9.8, which accords with the published values [35]. Analytical-grade uranyl nitrate hexahydrate (UO2(NO3)26H2O) was dissolved in Milli-Q water to prepare uranyl stock solution, which was further diluted to the needed concentration. Silicate solution was prepared with analytical grade Na2SiO3H2O. The purchased chemicals were of analytical purity and used without further purification. Milli-Q water was used to prepare the solutions in the experiments. The experiments of U(VI) and/or silicate adsorption on c-Al2O3 were conducted in polyethylene tubes by employing batch technique under ambient conditions. Prior to initializing U(VI) sorption, the c-Al2O3 solids were hydrated and NaNO3, Na2SiO3 solution stock solution were added in polyethylene tubes to attain the desired concentrations of individual components. The suspensions were adjusted to acidic pH value before U(VI) addition. Negligible amount of 0.01 or 0.1 mol/L HNO3 or NaOH were added in each tube to achieve the desired pH values of the suspensions. After the solutions were shaken for 24 h to attain equilibrium, the liquid and solid phases were separated by centrifugation at 9000 rpm for 30 min. The centrifugation was controlled at a constant temperature. U(VI) contents in the supernatant were analyzed through the formation of uranyl–chlorophosphonazo(III) complex at wavelength 670 nm, and silicate contents were analyzed by inductively coupled plasma-atomic emission spectroscopy (ICP-AES). The amounts of U(VI) and silicate adsorbed on c-Al2O3 were figured from the difference between the initial and equilibrium concentration (C0 and Ce, respectively, mmol/L) (sorption percentage (%) = (C0  Ce)/C0  100%, and Cs = (C0  Ce)/m  V, where Cs(mmol/g) denotes the concentration of ions adhered to the solid phase, V(L) denotes the volume of the suspension, and m(g) denotes the mass of the adsorbents). For spectroscopic analysis, sorption experiments were conducted with 1.0 g/L c-Al2O3, 0.01 mol/L NaNO3 and 0.5 mmol/L U(VI) at pH 5.0 ± 0.1. The solid was separated by filtration and subsequently washed with 0.01 mol/L NaNO3 solution to remove non-adsorbed U(VI), and then the solid phases were dried under vacuum and ambient temperature. Only part of free water on solid phases was removed in this treatment, and this treatment did not result in any surface species modification [14].

2.2. EXAFS analysis For EXAFS analysis, samples were examined in situ by centrifuging the suspensions and the wet sorbent pastes were used for EXAFS measurements. Uranium L3-edge X-ray absorption spectra were recorded at Shanghai Synchrotron Radiation Facility (SSRF, China) in fluorescence modes. The electron storage ring energy was 3.5 GeV and the maximum storage beam current was 220 mA. The energy of X-ray was detuned by using a fixed-exit double-crystal Si (1 1 1) monochromator. Ionization chambers with N2 atmosphere were used to collect the Ni K-edge spectra in transmission mode at room temperature. A multi-element pixel high purity Ge solid-state detector was used to collect the fluorescence signal. IFEFFIT and their graphical interfaces ATHENA and ARTEMIS were used for background subtraction and fitting [36]. The input parameter to ATHENA that determined the maximum frequency of the background, Rbkg, was set to 1.1 Å [37]. Radial structure functions (RSFs) were obtained by Fourier transformed k3-weighted

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v(k) functions between 3.5 and 12.2 Å1 with a Kaizer–Bessel win-

3. Results and discussion 3.1. Surface structure and the influence of silicate on pHpzc value of hydrated c-Al2O3 It has been shown that surface hydroxyl groups are easily formed on clean Al-terminated c-Al2O3 (1 0 0) and (1 1 0) surfaces (Fig. 1) by theoretical and experimental studies [12,39,40], and the hydrated surfaces are examined with OH coverage of 8.8 OH/ nm2 for the (1 0 0) surface and 8.9–11.8 OH/nm2 for the (1 1 0) surface [39,40]. The surfaces of c-Al2O3 include a mixture of singly, doubly and triply coordinated oxygen atoms, and these surface oxygen atoms need to be stabilized by the protonation and hydrogen bonding to water molecules. The reactivity of c-Al2O3 depends on the terminating layers exposed by each surface and directly relates to the degree of distortion of the surface in the presence of the water. It was found that silicate and/or U(VI) interacted with the c-Al2O3 surface via their OH group. In natural aquatic environments, the presence of silicate may influence the U(VI) uptake by modifying the c-Al2O3 surface or complex with U(VI). Zeta potential (ZP) is a function of the surface coverage by charged species at a given pH, which is theoretically determined by the activity of the species in solution. The ZP values of c-Al2O3 before and after silicate sorption are measured and shown in Fig. 2. The results agree with the reported the point of zero charge (pHpzc) value of 9.8 [34,35]. As predicted, the pHpzc value decreases to 8.5 when silicate anions are added. It can be explained that the presence of silicate makes hydrated c-Al2O3 surface more negatively. Chemical interaction rather than pure electrostatic interaction could be concluded from the changes of both ZP and pHpzc at c-Al2O3/water interface [35]. Such considerations to predict the influence of silicate on U(VI) to c-Al2O3 are evaluated later. 3.2. The effect of silicate on the retention of U(VI) as a function of pH pH is one of the most important factors affecting the retention of actinide ions, since the solution pH usually has great influence

Al2O3

60

Al2O3+silicate 40

zeta potential (mV)

dow function and a dk value of 0.5 Å1. Ab initio U–Oax, U–Oeq, and U–Si scattering parameters were calculated for a self consistent potential based on the crystal structure of soddyite, (UO2)2SiO42H2O [38], using FEFF7. The U–Al scattering amplitudes and phases derived from soddyite can also be obtained by replaced Si for Al atoms [7]. The solution of UO2+ 2 in the presence of silicate was additionally analyzed with the U–Si path. With regard to the samples uranyl adsorbed onto c-Al2O3 in the presence/absence of silicate, the equivalent U–Al path was used. All fitting operations were performed to Fourier transform (FT) spectra in R-space between 0.9 and 4.0 Å. Values for the energy shift, DE0, and coordination numbers, N, of Oax and Oeq were determined during an initial fit and kept fixed during further refinements bond distances, R, and Debye–Waller factors, r2, were allowed to adjust freely.

20

0 2

-20

4

6

8

10

12

pH

-40

Fig. 2. The zeta potentials of c-Al2O3 before and after silicate sorption versus pH. m/V = 1 g(c-Al2O3)/L, I = 0.01 mol/L NaNO3, C(silicate)initial = 0.5 mmol/L.

on the surface properties of adsorbents and the hydrolysis, complexation, and the precipitation of radionuclides [19,31]. The retention of U(VI) on c-Al2O3 in the absence and presence of silicate as a function of pH is shown in Fig. 3A. The pH-dependence of U(VI) sorption on c-Al2O3 is similar to that observed for U(VI) sorption on silica, iron hydroxide, and clay minerals in open systems [18,32,33]. In the presence of silicate, U(VI) retention increases up to 10–20% higher than that without silicate in all pH range. The increase of removal percentage of U(VI) sorption suggesting that the silicate has a positive effect on U(VI) retention. At pH < 6.0, the positively charged U(VI) species in solution may exchange with H from the surface hydroxyl groups of c-Al2O3, and the increase in free hydroxyl concentration is the cause of the steep rise of U(VI) sorption. Thus low retention of U(VI) on cAl2O3 occurs at pH < 6.0 in the absence of silicate owing to the strong electrostatic repulsion between the positive charged cAl2O3 and U(VI) ions [31]. In the presence of silicate, the percentage of U(VI) retention increases in pH range of 4.0–6.0. The addition of silicate has a similar effect as that of adding OH, which can decrease the electrostatic potential near the surface. Meanwhile, silicate can form complex with U(VI) as UO2SiO(OH) 3 and then reduce the electropositivity of U(VI) [25,41]. All these factors lead to the increased U(VI) loading near the surface, and hence result in the increase of U(VI) sorption on c-Al2O3 surface. In pH range of 6.0–9.8, U(VI) exists mainly in the neutral U(VI) species and the soluble U(VI)–carbonate complexes. Since these species are neutral and anionic, electrostatic force between U(VI) and cAl2O3 can be negligible. While due to the significant sorption of silicate (Fig. 3B), U(VI) sorption increases by the formation of ternary UO2–SiO4–Al2O3 complexes. However, we could not exclude the possibility of the formation of Al2O3–UO2–SiO4 or U(VI)–silica colloids with reaction continuing [34]. At pH > 9.8, the formation of dissolved U(VI) carbonate complexes cause the decrease in U(VI) sorption. However, the adsorbed silicate can facilely form complexes with U(VI) and then promote U(VI) on c-Al2O3 surface.

Fig. 1. Side views of hydrated (1 0 0) (A) and (1 1 0) (B and C) terminations of c-Al2O3. Red balls, oxygen atoms. Purple balls, aluminum atoms. White balls, hydrogen atoms. (For interpretation of the references to colour in this figure legend, the reader is referred to the web version of this article.)

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80

60

40

20

no silicate 0.5mmol/L silicate

0

4

5

6

7

8

9

10

11

3.3. Kinetics of U(VI) sorption 12

pH 70

(B)

50

40

30

20

no uranyl -2

6.72x10 mmol/L uranyl

10

0 3

5

7

9

11

pH Fig. 3. Influence of pH on the mutual effects of U(VI) (A) and silicate (B) on their interaction with c-Al2O3. C(U(VI))initial = 6.72  102 mmol/L, m/V = 1 g(c-Al2O3)/L, I = 0.01 mol/L NaNO3, C(silicate)initial = 0.5 mmol/L.

The results implies that the complexation ability of U(VI) with silicate is stronger and silicates may act as a ‘‘bridge’’ between the surface of c-Al2O3 and U(VI). The greater affinity of U(VI) for silanol surface sites than for magnetite surface sites may explain such a behavior. Similar positive effects of silicate on the retention of Cs(I) on magnetite were also observed by Marmier and Fromage [30]. These results suggest that interactions between U(VI) and silicate are significant in controlling U(VI) retention. Our experimental data indicate strong silicate sorption for those data points where U(VI) sorption is enhanced by silicate (Fig. 3B). The curve shows a regular increase of sorption from pH 3.0 to 9.8, and then decreases with further increasing pH up to 12.0. This is coincident with the trend of silicate sorption on goethite and ferrihydrite [24,27]. The effect of pH can be explained in terms of pHpzc of c-Al2O3 and the distribution of silicate species. The surface charge of c-Al2O3 is positive at pH < pHpzc (i.e., 9.8) and in turn negative at pH > pHpzc. From the silicate species distribution diagram, the silicate exists as neutral H4SiO4 and negatively-charged H2SiO2 4 species at pH < 9.8 [24]. The increased sorption of silicate with increasing pH values is ascribed to the increase of H2SiO2 4 proportion. At pH > 9.8, the sorption of silicate is not favored with increasing pH. This is due to the retention of OH on the surface of c-Al2O3, which leads to the formation of a new negative charged layer and decreases the affinity of c-Al2O3 surfaces towards silicate. Meanwhile, OH competes strongly with silicate for active sites,

Retained U(VI) concentration,qt(mmol/g)

0.07

(A) 0.06

0.05

0.04

0.03

0.02

pH 5.0 pH 7.0 pH 9.0

0.01

solid: Al2O3/UO22+ open: Al2O3/silicate/UO22+

0.00 -2

0

2

4

6

8

10

12

14

16

Time(hours) 600

pH 5.0 pH 7.0 pH 9.0

(B) 500

-1

Removal percentage (%)

60

To obtain a realistic view of the U(VI) sorption on c-Al2O3, kinetic experiments have been performed with/without silicate. The results at pH 5.0, 7.0 and 9.0 are illustrated in Fig. 4, from which it clearly shows that the sorption of U(VI) increases as the presence of silicate. The sorption of U(VI) on c-Al2O3 is rapid during the first 3 h and then maintains the high level with increasing contact time. In the following experiments, the reaction time is fixed to 24 h to ensure the sorption equilibrium of U(VI) on c-Al2O3. The fast sorption velocity implies that strong chemical sorption rather than physical sorption contributes to the sorption of U(VI)

t/qt(h•g•mmol )

Removal percentage (%)

resulting in the reduction of sorption capacity. The U(VI) can promote silicate sorption in the pH window where there is an appreciable concentration of silicate in solution and also an appreciable amount of c-Al2O3-adsorbed U(VI). A plausible explanation for the increased sorption of silicate is the formation of a ternary surface complex containing both U(VI) and silicate, which increases the sorption of both U(VI) and silicate [35]. These results together suggest that a ternary surface complex formed under these conditions and should be included in the modeling in order to correctly predict U(VI) sorption in the presence of silicate. The structure of the ternary surface complex is investigated in the analysis of spectroscopic data.

(A)

100

400

300

200

100

solid: Al2O3/UO22+ open: Al2O3/silicate/UO22+

0 0

2

4

6

8

10

12

14

16

Time(hours) Fig. 4. U(VI) retained by c-Al2O3 in the absence and the presence of silicate as a function of contact time. C(U(VI))initial = 6.72  102 mmol/L, m/V = 1 g(c-Al2O3)/L, I = 0.01 mol/L NaNO3, C(silicate)initial = 0.5 mmol/L, pH = 5.0, 7.0 and 9.0 ± 0.1, respectively.

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t 1 1 ¼ þ t qt kq2e qe

ð1Þ

where k (g/(mmol h)) is the pseudo-second-order rate constant of sorption, qt (mmol/g of dry weight) is the amount of U(VI) adsorbed on the surface of c-Al2O3 at time t (h), and qe (mmol/g of dry weight) is the equilibrium sorption capacity. Linear plot feature of t/qt vs. t is achieved and shown in Fig. 4. The k value calculated from the slope and intercept is shown in Table 1. The equilibrium sorption capacity of U(VI) on c-Al2O3 is smaller than that of U(VI) sorption in the presence of silicate suggests the enhancement of silicate. The correlation coefficient of the pseudo-second-order rate equation for the linear plots is very close to 1, which suggests that kinetic sorption can be described by a pseudo-second-order model well. This result shows that the sorption of U(VI) in the presence/ absence of silicate might be controlled by similar mechanisms as inner-sphere complexation. 3.4. Sorption isotherms of U(VI) The sorption isotherms of U(VI) on c-Al2O3 in the presence and absence of silicate are shown in Fig. 5. One can see that the sorption isotherms of U(VI) in the presence of silicate are much higher than that in the absence of silicate. The significant positive effect of silicate on U(VI) sorption can be explained by the metal complexation with the surface of silicate, and suggests the surface-binding of U(VI) via silicate bridge between the surface and radionuclide U(VI). In order to gain a better understanding of the mechanism and to quantify the sorption data, Langmuir (qe = abCe/(1 + bCe)) and Freundlich (qe = KfC1/n e ) isotherm models are conducted to simulate the sorption of U(VI) on c-Al2O3. Herein, a is the maximum sorption capacity; b is the Langmuir sorption constant; and Kf and 1/n are the Freundlich constants [43]. The sorption of U(VI) on c-Al2O3 is simulated well by Langmuir model and Freundlich model. The maximum sorption capacity of U(VI) sorption on silicate adsorbed c-Al2O3 are higher than that of U(VI) sorption on bare c-Al2O3, which also indicates that silicate have stronger affinity for U(VI) sorption than bare c-Al2O3 [43] (Table 2). 3.5. EXAFS analysis EXAFS is very important to understand the interaction mechanism and microstructure of radionuclides at the solid/water interface [44–47]. Fig. 6 shows the raw k3-weighted EXAFS data and their corresponding radial structure functions (RSFs) for UO2+ 2 in aqueous solution in the absence/presence of silicate. The theoretical curve fits are also shown, and the structural results

Table 1 Kinetic parameters of U(VI) sorption on c-Al2O3 at various pH with/without silicate concentrations. System pH5.0Al2O3/UO2+ 2 pH5.0Al2O3/silicate/UO2+ 2 pH7.0Al2O3/UO2+ 2 pH7.0Al2O3/silicate/UO2+ 2 pH9.0Al2O3/UO2+ 2 pH9.0Al2O3/silicate/UO2+ 2

qe (mmol/g) 2

2.71  10 3.38  102 5.60  102 6.03  102 5.55  102 6.01  102

k

R2

574.6 350.2 649.2 1333.8 203.2 517.8

0.997 0.998 0.999 0.999 0.999 0.999

0.08

0.07

0.06

qe(mmol/g)

on c-Al2O3 [14,34]. The presence of silicate at three pH values increases the rate of U(VI) sorption. The total quantity of U(VI) adsorbed at pH 5.0, 7.0 and 9.0 agree with the data in Fig. 3. In order to investigate the sorption mechanism, kinetic models are generally used to test the experimental data. Pseudo-secondorder equation is applied to simulate the kinetic sorption data. The pseudo-second-order rate equation is expressed as [42]:

0.05

0.04

0.03

0.02

pH5.0Al2O3/UO22+ pH5.0Al2O3/silicate/UO22+

0.01 0.00

0.02

0.04

0.06

0.08

0.10

0.12

Ce(mmol/L) Fig. 5. Sorption of U(VI) on c-Al2O3 in the absence and the presence of silicate. pH = 5.0 ± 0.1, m/V = 1 g(c-Al2O3)/L, I = 0.01 mol/L NaNO3, C(silicate)initial = 0.5 mmol/L. Solid line: Langmuir model, dash line: Freundlich model.

are summarized in Table 3. The peaks are shifted to lower R values as a result of the phase shifts related to the absorber scatterer interactions (0.2–0.5 Å). In the absence of silicate, the UO2+ 2 spectrum shows axial (Uax–O) and the equatorial (Ueq–O) shells with the best theoretical fit of 2 Oax at 1.76 Å and 5 Oeq at 2.41 Å. This is in agreement with the structural results obtained previously for the fully hydrated uranyl ion [6,7,10–12]. An inspection of the UO2+ 2 /silicate data shows that the structural transformation is taking place in the axial region, and the dioxo peak is compromised by Si–O correlation that has been adequately removed. Two axial oxygen atoms, Oax, centered at a U–O distance of 1.77 Å from the apices of the bipyramid. In the equatorial plane, there are a total of five oxygen atoms. The spectrum of the UO2+ 2 /silicate sample shows a different feature with a peak appearing at a higher R value than the Oeq peak for the hydrated uranyl ion. For UO2+ 2 /silicate as reference, the peak at 3 Å on the uncorrected spectrum attributing to the multiple scattering effects of the U–Oax shell provided a poor fit to the data. The peak feature at 3 Å attributes to the backscatter from the nearest Si neighbor. And curve fit to this spectrum indicates the structural changes could be fit primarily by change in appearance of the U–Si bond. The peak at 3.15 Å can be ascribed to a uranyl–silicate mononuclear, bidentate U–Si interaction as shown in Fig. 6. The data and fit ranges, along with the goodness-of-fit parameters for this fit, are illustrated in Table 3. The curve fits support the formation of complex UO2OSi(OH)+3, which is analogous to the known structure of uranyl orthosilicate, or soddyite, (UO2)2SiO42H2O solid [48]. Two sorption samples show distinct evidence for the existence of uranium in the uranyl form. The large peaks around 1.79 Å were identified as the backscatter of the two (axial) oxygen atoms bonded linearly with the central uranium atom and were observed in many previous studies. This has been confirmed by curve fits, giving 2 oxygen atoms at 1.79 and 1.80 Å for each sample. The second shell is composed of 5 equatorial oxygen atoms at distances 2.43–2.44 Å. The U–Oeq distances are close to those of U atoms coordinated by 5 equatorial oxygen atoms of aqueous species or of outer-sphere surface complexes onto clays [1,7]. The U–Oeq distances are also close to an average of bond lengths reported for 5–6 oxygen atoms split in two equatorial subshells of adsorbed U(VI) as a bidentate inner-sphere surface complex onto c-Al2O3 [1]. The absence of equatorial splitting for our sorption samples with/without silicate may indicate the outer-sphere surface complex and/or multiple uranyl coordination environments.

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Table 2 Parameter of sorption models for U(VI) on c-Al2O3. System

Langmuir model a (mmol tg

pH5.0Al2O3/UO2+ 2 pH5.0Al2O3/silicate/UO2+ 2

0.138 0.247

1

)

Freundlich model b (L mmol

1

)

11.99 10.98

2

R

Kf (mmol g1)

n

R2

0.970 0.971

0.106 0.349

1.57 1.54

0.978 0.967

Fig. 6. Raw k3-weighted v(k) spectra of reference sample and sorption samples (A) and their corresponding radial structure functions (RSFs) (B) (magnitude, symbols; fitted data, solid lines).

Table 3 Structure parameters derived from EXAFS analysis of reference samples and sorption samples at U LIII-edge. Sample

Shell

R(Å)a

Nb

r2(Å2)c

Rfd

UO2+ 2 (aq)

U–Oax U–Oeq

1.79 2.44

2 5

0.0002 0.0076

0.026

UO2+ 2 /silicate(aq)

U–Oax U–Oeq U–Si

1.77 2.43 3.16

2 5 0.33

0.0039 0.0045 0.0015

0.034

Al2O3/UO2+ 2

U–Oax U–Oeq U–Al

1.79 2.43 3.35

2 5 0.120

0.0037 0.0087 0.0063

0.013

Al2O3/silicate/UO2+ 2

U–Oax U–Oeq U–Al

1.80 2.43 3.35

2 5 0.306

0.0031 0.0092 0.0036

0.023

2+ 2+ UO2+ 2 (aq) and UO2 /silicate(aq) are named as reference samples, whereas the other samples of c-Al2O3 with adsorbed UO2 are named as sorption sample. a Interatomic distance. b Number of neighbor atoms. c Debye–Waller factor. d The residual factor.

For the Al2O3/UO2+ 2 sorption sample, a better fit was achieved by including a U–Al shell. The procedure gives U–Al bond distances of 3.35 Å and coordination numbers of 0.12. The presence of detectable U–Al interactions verifies that inner-sphere surface complexation is a main mechanism of sorption of uranyl at acidic pH [23]. The U–Al bond length of 3.35 Å matches well with the U–Al distance of the model structure of Hennig et al. with the uranyl unit (RU–Oeq = 2.44 Å) in bidentate, edge-sharing connection with [AlO6] octahedral (1.85 < RAl–O < 1.97 Å) [49]. The UO2+ 2 cation has a strong affinity to c-Al2O3 and may form an inner-sphere,

mononuclear, edge-sharing complex with alumina octahedral (Fig. 7A) as was proposed from U LIII-edge EXAFS spectroscopy. Silicate ions are also expected to influence the surface complexation of U(VI), since SiO44 ligands are known to influence the surface of c-Al2O3 and form complex with UO2+ 2 . The influence of silicate can be judged by comparing the U LIII-edge k3-weighted EXAFS oscillations and the corresponded RSFs of samples with silicate illustrated in Fig. 6. It is evident that both samples show comparable EXAFS oscillations and FTs without any significant differences. The obtained U–Oax and U–Oeq distances (see Table 3)

H. Mei et al. / Chemical Engineering Journal 269 (2015) 371–378

377

Fig. 7. Different sorption complexes to hydrated or silicate coated c-Al2O3 surface. U(VI) via complexation with two oxygens bound to the Al atom (A); two oxygens bound to the separate Si atom (B); oxygens bound to the one Si atom and one Al atom (C).

are also very similar. The RSF peaks at R + DR 3.0 Å occur in both spectra, irrespective of the presence of silicate. This is evidence for either Al or Si atoms in the range of 3–4 Å. Other authors have suggested that the peak feature at 3.0 Å on RSF spectrum for U(VI) onto Al-silicates is the superposition of the multiple scattering (MS) effects of the U–Oax shell and of the interactions between U(VI) and Al–silicates [23]. The U–Si and U–Al interactions are indistinguishable by EXAFS spectroscopy. It is impossible to distinguish between Si and Al for their similar scattering amplitude and phase functions [23]. These single and multiple scattering paths appear in the same place of the FTs; disturbance between them makes the accurate interpretation of this spectrum difficult, especially considering the limited data range. The best approach is to be a fit of the shell with Si/Al including fixed Debye–waller factors together with the constrained parameters of the U–Oax MS contribution. This procedure gives distances of around 3.35 Å and the coordination number of 0.306. The differences in coordination number are likely owing to the differences in presence of silicate for the different complexes. Compared with Al2O3/UO2+ 2 sample, the increased NU–Al can be assigned to the contribution of the silicate, which implies a sorption to the surface through complexation with two oxygen atoms bonded to the Al atom (Fig 7A); and/or two oxygen atoms bonded to the separate Si atom (Fig 7B); and/or oxygen atoms bonded to the one Si atom and one Al atom (Fig 7C); a bidentate complexation to hydrated or silicate coated c-Al2O3 surface. In the sorption samples where bidentate, mononuclear complexes have been observed by noting the presence of Al atom at 3.35 Å. It is not proved that all the U(VI) ions adsorbed were bounded in this bidentate geometry, actually, multiple adsorption geometries can form simultaneously. In the presence of silicate, the surface adsorbed silicate supports the formation of ternary U(VI)– silicate surface complexes. Fitting the peak by a U–Si shell at 3.15 Å is in line with a bidentate linkage of silicate to the equatorial U(VI) oxygen atoms. As mentioned in this study, it is hard to distinguish Al and Si at distances 3.0 Å using EXAFS spectroscopy. Without the way to resolve the Al or Si shells or to quantify the relative coordination numbers of Al or Si atoms at different distances, the accurate influence of silicate on the U(VI) on c-Al2O3 cannot be determined. It is known that U(VI) may present in multiple geometries but cannot be observed due to the limitations of EXAFS spectroscopy mentioned above. Determination of the types U(VI) sorption geometries formed on c-Al2O3 and obtained the potential influences of coexist ions may help constrain the possible sorption complexes forming on real environmental substrates.

4. Conclusion The results show that the presence of silicate promotes the sorption of U(VI) on c-Al2O3. The EXAFS analyses suggest that a

ternary surface complex is formed under these conditions and should be included in the modeling in order to correctly predict U(VI) sorption in the presence of silicate. This is readily explained the ability of silicate, a naturally occurring and ubiquitous oxyanion, to: (1) provide potential sorption sites and thereby increase the total quantity of U(VI) adsorbed, and (2) cooperative sorption with U(VI). These two processes may ultimately result in a decrease in the mobility and potential bioavailability of U(VI). The influence of silicate on U(VI) retention in natural systems may help us to better understand the mobility and movement of radionuclides in natural environments and to determine their bioavailability. Acknowledgments Financial supports from National Natural Science Foundation of China (21377132, 41273134, 21307135, 21225730, 91326202), Chinese National Fusion Project for ITER (No. 2013GB110004), the Jiangsu Provincial Key Laboratory of Radiation Medicine and Protection and the Priority Academic Program Development of Jiangsu Higher Education Institutions are acknowledged. References [1] E.R. Sylwester, E.A. Hudson, P.G. Allen, The structure of uranium (VI) sorption complexes on silica, alumina, and montmorillonite, Geochim. Cosmochim. Acta 64 (2000) 2431–2438. [2] D.E. Latta, B. Mishra, R.E. Cook, K.M. Kemner, M.I. Boyanov, Stable U(IV) complexes from at high-affinity mineral surface sites, Environ. Sci. Technol. 48 (2014) 1683–1691. [3] D.M. Giaquinta, L. Soderholm, S.E. Yuchs, S.R. Wasserman, The speciation of uranium in a smectite clay: evidence for catalysed uranyl reduction, Radiochim. Acta 76 (1997) 113–121. [4] D.E. Giammar, J.G. Hering, Time scales for sorption–desorption and surface precipitation of uranyl on goethite, Environ. Sci. Technol. 35 (2001) 3332– 3337. [5] V. Wheaton, D. Majumdar, K. Balasubramanian, L. Chauffe, P.G. Allen, A comparative theoretical study of uranyl silicate complexes, Chem. Phys. Lett. 371 (2003) 349–359. [6] J.G. Catalano, T.P. Trainor, P.J. Eng, G.A. Waychunas, G.E. Brown Jr, CTR diffraction and grazing-incidence EXAFS study of U(VI) adsorption onto aAl2O3 and a-FeO3 (102)surface, Geochim. Cosmochim. Acta 69 (2005) 3555– 3572. [7] A. Froideval, C. Gaillard, R. Barillon, I. Rossini, J.L. Hazemann, Uranyl sorption species at low coverage on Al-hydroxide: TRLFS and XAFS studies, Geochim. Cosmochim. Acta 70 (2006) 5270–5284. [8] H.S. Chang, G.V. Korshin, Z. Wang, J.M. Zachara, Adsorption of uranyl on gibbsite: a time resolved laser-induced fluorescence spectroscopy study, Environ. Sci. Technol. 40 (2006) 1244–1249. [9] N. Baumann, V. Brendler, T. Arnold, G. Geipel, G. Bernhard, Uranyl sorption onto gibbsite studied by time-resolved laser-induced fluorescence spectroscopy (TRLFS), J. Colloid Interface Sci. 290 (2005) 318–324. [10] T. Hattori, T. Saito, K. Ishida, A.C. Scheinost, T. Tsuneda, S. Nagasaki, S. Tanaka, The structure of monomeric and dimeric uranyl adsorption complexes on gibbsite: a combined DFT and EXAFS study, Geochim. Cosmochim. Acta 73 (2009) 5975–5988. [11] L.V. Moskaleva, V.A. Nasluzov, N. Rösch, Modeling adsorption of the uranyl dication on the hydroxylated alpha-Al2O3 (0 0 0 1) surface in an aqueous medium. Density functional study, Langmuir 22 (2006) 2141–2145.

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