Fluoride adsorption on manganese carbonate: Ion-exchange based on the surface carbonate-like groups and hydroxyl groups

Accepted Manuscript Fluoride adsorption on manganese carbonate: Ion-exchange based on the surface carbonate-like groups and hydroxyl groups Yong-Xing Zhang, Yong Jia PII: DOI: Reference:

S0021-9797(17)31121-9 https://doi.org/10.1016/j.jcis.2017.09.090 YJCIS 22842

To appear in:

Journal of Colloid and Interface Science

Received Date: Revised Date: Accepted Date:

2 July 2017 27 August 2017 23 September 2017

Please cite this article as: Y-X. Zhang, Y. Jia, Fluoride adsorption on manganese carbonate: Ion-exchange based on the surface carbonate-like groups and hydroxyl groups, Journal of Colloid and Interface Science (2017), doi: https:// doi.org/10.1016/j.jcis.2017.09.090

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Fluoride adsorption on manganese carbonate: Ion-exchange based on the surface carbonate-like groups and hydroxyl groups Yong-Xing Zhanga, Yong Jiab,

a

School of Physics and Electronic Information, Huaibei Normal University, Huaibei 235000, PR China b

School of Pharmacy, Anhui University of Chinese Medicine, Hefei 230012, PR China



Corresponding author. E-mail: [email protected] (Y. Jia). 1

Abstract Manganese carbonate (MnCO3) nanowires and microcubes were controlled synthesized by a simple ethylene glycol (EG) mediated solution method. The volume ratios of EG and water has a decisive impact on the morphology of MnCO3 products. With a decreased water volume, MnCO3 products were transformed from microcubes to nanowires. The obtained MnCO3

nanowires were characterized by X-ray

diffraction, field emission scanning electron microscopy, transmission electron microscopy, and nitrogen adsorption–desorption isotherm. The adsorption properties of the MnCO3 products towards fluoride were investigated. The adsorption capacities of the nanowires were higher than that of the microtubes. According to the Langmuir model, the maximum adsorption capacity was 46.80 mg g-1 at pH 7.0. The adsorption capacity was 11.58 mg g-1 when the equilibrium fluoride concentration was just below the WHO guideline of 1.5 mg L-1. The kinetic data were well fitted to pseudo-secondorder model. The fluoride removal was attributed to the ion-exchange based on the surface hydroxyl groups and carbonate-like groups, which was revealed by Fourier transform infrared absorption spectroscopy and X-ray photoelectron spectroscopy. Keywords: Manganese carbonate; Fluoride adsorption; Surface carbonate-like groups; Surface hydroxyl groups; Ion-exchange

2

1. Introduction Nowadays, the synthesis of micro and nanocrystals with controllable size and shapes has attracted great attentions owing to the deep understanding of the shape evolution of crystals, and their structural controlled performance [1, 2]. Because of the important applications in the fields of biomedical [3], supercapacitor[4], environment[5], catalyst[6], lithium ion batteries[7], and

sacrificial template[8],

manganese carbonate (MnCO3) materials have attracted attentions. In addition, MnCO3 can also be utilized to prepare composite materials[9, 10]. Moreover, MnCO3 is a very importantly precursor for the synthesize of manganese oxide[11, 12]. Thus, the synthesis of MnCO3 with a controlled structure is primary important. Up to today, many shapes of MnCO3 have been successfully synthesized, including nanowires [1315], nanorods [15, 16], microtubes [12, 14, 17, 18], microspheres [12, 19], ellipsoidalshaped [18], lotus shaped [20], steep rhombohedra crystals [21], dumbbell-like [22], peanut-like [23], and MnCO3 nanoplates [24]. Most of the above reported synthesis methods can prepare only one type of MnCO3 products. By using a complex detached device, Lei et al. [14] reported the preparation of MnCO3 microcubes in pure water. With the presence of cetyltrimethylammonium bromide, n-hexane and n-hexanol, MnCO3 nanowires were obtained. Thus, it is still a challenge to synthesis MnCO3 products with a controlled structure using a simple method. In addition, fluoride occurring in natural ground water is a serious worldwide threat and has attracted great attentions. Adsorption has proved to be an effective strategy for fluoride removal from water owing to its easy operation, low cost, and treatment stability [25]. For the reported inorganic adsorbents, the most frequently proposed fluoride removal mechanism is an ion-exchange process between the surface hydroxyl group and fluoride ions. Obviously, this type of ion-exchange is suffered 3

from high pH values based on the Le Chatelier’s principle. Therefore, the investigation of new types of surface ion-exchange groups will bring us an opportunity to reduce this influence. Recently, besides of the surface hydroxyl groups, the ion-exchange between the surface carbonate groups and fluoride was proposed [26]. Furthermore, previous works have demonstrated the fluoride removal performance and mechanism of calcite, Ca-Zn(OH)2CO3 and aluminum carbonate [27-29]. These results suggested that both of the adsorbed surface carbonate-like groups and the lattice carbonates close to the surface could participate in the ionexchange between fluoride[26-29]. Herein, we demonstrated a simple EG assisted hydrothermal method for the controlled synthesis of MnCO3 products. The morphology of the MnCO3 products, from nanowires to microtubes, were well controlled by the volume of EG and water. Furthermore, the fluoride removal performance of the MnCO3 was studied. The maximum adsorption capacity of the MnCO3 nanowires at neutral pH was 46.80 mg g-1. The adsorption capacity was 11.58 mg g-1 when the equilibrium fluoride concentration was just below the WHO guideline of 1.5 mg L-1 [30]. Furthermore, the fluoride removal was an ion-exchange process based on the surface hydroxyl groups and the carbonate-like groups, which was revealed by FTIR and XPS. 2. Materials and methods 2.1. Preparation of MnCO3 Manganese chloride tetrahydrate, EG, ethanol, and urea were purchased from the Shanghai Chemical Reagents Company, and were of analytical grade. These reagents were directly used without further purification. For the synthesis of MnCO3, a certain amount of manganese chloride tetrahydrate was dissolved in the mixture solution of deionized water and EG at room temperature. The concentration of the Mn2+ was 0.1 4

mol L-1. Then, a certain amount of urea was added in the above solution, and its concentration was 0.4 mol L-1. After stirring for 10 min, the above solution was transferred into a conical flask with a stopper, and then heated at 100 oC for 10 h. Then, allowed to cool to room temperature naturally. The resulting precipitates were separated by centrifugation, and washed with deionized water and alcohol several times. The obtained precipitates were dried in an oven at 60 oC. 2.2. Characterization The structures of the obtained products were characterized by X-ray diffraction (X’Pert ProMPD, Cu–Kα radiation, wavelength 1.54056 Å), field emission scanning electron microscopy (FEI Sirion 200 FEG, operated at 10 kV), transmission electron microscopy (JEOL-2010, operated at 200 kV), Fourier transform infrared spectroscopy (Nicolet Analytical Instruments, NEXUS-870), and X-ray photoelectron spectroscopy (VG ESCALAB MKII spectrometer, Mg KR X-ray source, 1253.6 eV, 120 W) analyses. Nitrogen adsorption–desorption measurement was performed using a Micromeritics ASAP 2020 M+C instrument with a degassing temperature of 80 °C, and using Barrett–Emmett–Teller (BET) calculations for the surface area, and Barret– Joyner–Halender (BJH) calculations for the pore size distribution. 2.3. Adsorption experiments The adsorption experiments were performed according to our previous works [31]. The stock solution of fluoride was prepared with deionized water using NaF, and stored in a polypropylene bottle under a dark and cool conditions. The fluoride working solutions were obtained by appropriately diluting the stock solution with deionized water. The adsorption isotherm was studied by varying initial fluoride concentration from 5 to 200 mg L -1 (5, 10, 20, 50, 75, 100, 150 and 200 mg L-1). The experiments were performed in 15 mL polypropylene flasks containing 10 mL of 5

fluoride working solution and 10 mg of adsorbent. The flasks were shaken at 150 rpm in a shaker at 25 ºC. The solution pH value was adjusted every 4 h by using 0.1 mol L-1 HCl or/and NaOH to around pH 7.0 ± 0.1. After 24 h, the adsorbent was separated by centrifugation, and the residual fluoride concentration was measured by the fluoride ion selective electrode (PF-202-CF) by use of total ionic strength adjustment buffer solution to maintain pH 5.0 and to eliminate the interference effect of complexing ions [31]. All of the experiments were conducted three times and the averages of the results were used for data analysis. The amount of the adsorbed fluoride (qe in mg g-1) was calculated by Eq. (1):

qe 

(C0  Ce )V m

(1)

where C0 and Ce (mg L-1) are the initial and final equilibrium concentrations of fluoride, V (mL) is the volume of the solution and m (mg) is mass of the adsorbent. In addition, the residual adsorbent was centrifuged and soaked into a 0.05 mol L-1 of NaOH solution for 24 h. After filtration, the obtained adsorbent was rinsed with deionized water until the filtrate was neutral. The regenerated adsorbent was also dried at 60 oC. With the initial fluoride concentration of 200 mg L-1, the above adsorption–desorption experiments were continued performed for five times, and the corresponding adsorption capacities were also calculated according to Eq. (1). Kinetics study experiments were carried out by adding 300 mg of adsorbent into a 500 mL of polypropylene vessel, containing 300 mL of fluoride solution with a initial concentration of 5 mg L-1. After adjusting the pH to 7.0 ± 0.1, the suspension was put on a magnetic stirrer with a stirring speed of 2000 rpm. About 4 mL of the suspension was sampled using a pipette after stirring 3, 5, 10, 20, 30, 40, 50, 60, 120, 180, 240 and 300 min. Then, the adsorbent was separated, and the residual fluoride 6

concentrations were measured. The influence of solution pH on the fluoride removal rate was studied with the initial fluoride concentration of 200.0 mg L-1, and with 1.0 g L-1 of adsorbent. The solution pH was also adjusted every 4 h. 3. Results and discussion 3.1. Characterization It is well known that urea will hydrolyze in water at high temperature to form carbonate ions according to the following reactions [18, 32]: CO(NH2)2 → NH3 + HNCO →NH4+ + NCO−

(2)

NCO− + OH− + H2O → NH3 + CO32−

(3)

Then, the dissolved Mn2+ ions will react with the generated CO32− anions. According to these simple reactions, the obtained product synthesized with a 75 : 25 volume ratios of EG and water was characterized by XRD. From Fig. 1, all the diffraction peaks can be indexed to the a single phase of hexagonal MnCO3 (JCPDS, NO. 441472), which also suggested that the MnCO3 product was well-crystallized. The product was further characterized by SEM and TEM, and the results are shown in Fig. 2. From Fig. 2a and b, large amounts of nanowires with tens of micrometers in length were observed. The diameter of the nanowires was less than 100 nm, which was confirmed by TEM images shown in Fig. 2c. Furthermore, from the highmagnification TEM image shown in Fig. 2d, the obtained nanowires were composed of several thin ones. The diameter of these thin nanowires was less than 15 nm. Combined with the XRD results, the as-prepared products were MnCO3 nanowires. Previous reports have demonstrated that polyalcohols are a very important ligand to form coordination complexes with metal cations [33, 34]. Among the polyalcohols, EG is a bidentate ligand, which can react with some cations to from a linear complex [33]. In our previous work, using the similar method, the morphology of CdCO3 7

microcrystals was well controlled by the change in the volume ratio of EG and water [35]. Herein, the influence of the volume ratios of EG and water on the morphology of MnCO3 product was also investigated. The results suggested that the volume ratios of EG and water also have a great influence on the structures of the MnCO3 product. Fig. 3a, b present the SEM images of the products synthesized with a 50 : 50 volume ratios of EG and water. It is cleat that most of the products were nanowires. Some small nanoparticles were also formed, as shown in Fig. 3a. In addition, a few of microtubes were observed, as shown in Fig. 3b. On further increasing the volume of water, more and more microtubes were obtained, as shown in Fig. 3c, d. Furthermore, only microcubes were synthesized with a 10 : 90 volume ratios of EG and water, as shown in Fig. 4a. The products prepared in pure water were similar to the ones synthesized with a 10 : 90 volume ratios of EG and water, as shown in Fig. 4b. The results suggested that the MnCO3 products were transformed from nanowires to microcubes with increasing water volume. By using the ultrasonic irradiation method, Yang et al. reported the synthesis of MnCO3 microtubes in pure water [18]. With the presence of sodium dodecyl sulfate, only ellipsoidal products were obtained. Furthermore, though MnCO3 nanowires were prepared in a complicated detached solvothermal reaction system when using cetyltrimethylammonium bromide as surfactant, but the length of them was only in the range of 0.5 to 1m [14]. In the present work, the structure of the MnCO3 products can be well controlled by the change in the volume ratio of EG and water. The results suggested that the presence of EG was in favor of the growth of the nanowires. However, when further increase the volume of EG, besides of the nanowires, many small nanoparticles were also obtained, as shown in Fig. 5. Thus, the optimum volume ratios of EG and water was 75:25. It is well known that EG can react with 8

metal cations to form linear complexes upon heating [33, 35, 36]. So, herein, EG could serve as a ligand to form long chain-like coordination complexes with Mn2+ ions. At the same time, with the hydrolyzation of urea, the concentration of CO32− was increased. Then, Mn2+ in chain-like complexes reacted with CO32− to form MnCO3 products, which then precipitated out from the reaction medium in the form of nanowires. It is well known that the large surface area of the nanomaterials will in favor of the improvement of their performance. Fig. 6a presents the N2 adsorption– desorption isotherm and BJH adsorption pore size distribution of the MnCO3 nanowires. The isotherm presents a type IV with H1 hysteresis loops characteristic, suggesting the obtained product was mesoporous. The pore size distribution curve shown in Fig. 6b presents a broader pore size distribution around 2-3 nm. These pores may be related to the interspaces between the nanowires. The BET surface area of the obtained MnCO3 nanowires was 57.6 m2 g-1. 3.2. Fluoride removal performance The ion-exchange between the carbonate-like anions of adsorbents and fluoride ions in solution has been demonstrated [26-29]. The surface carbonate-like groups of the obtained MnCO3 products may also take part in the ion-exchange process. Furthermore, the surface structures of MnCO3 were greatly depended on the pH values of the reaction conditions. Under base conditions, MnOOH will form on the surface of MnCO3 owing to the oxidization and hydrolysis of the dissolved Mn2+ ions [37]. The hydroxyl groups on the surface of the generated MnOOH will be in favor of the ion-exchange towards fluoride. Thus, the fluoride removal performance of the as-prepared MnCO3 products was studied. Fig. 7a presents the influence of the synthesis conditions on the fluoride adsorption capacity. It is clear that the obtained 9

MnCO3 nanowires shows the highest adsorption capacity. With the initial concentration of 200 mg L-1, the adsorption capacity was 45.55 mg g-1. With the increasing of water volume, the adsorption capacities were decreased step by step. The adsorption capacity of the MnCO3 microcubes synthesized with a 10 : 90 volume ratios of EG and water was only 12.10 mg g-1, which was equal to the ones prepared in pure water. Thus, the fluoride removal properties of the MnCO3 nanowires were further investigated. To reveal the kinetics of adsorption, the fluoride adsorption onto MnCO3 nanowires was studied as a function of contact time. A plot between time (t) and amount of the adsorbed fluoride with time (qt) is shown in Fig. 7b. The results suggested that the MnCO3 nanowires exhibited an initial rapid uptake of fluoride. With the initial concentration of 5.0 mg L -1, nearly 95% of fluoride anions were removed within 20 min. The time reached to the adsorption equilibrium was about 50 min. The adsorption kinetic experimental data were fitted into a pseudo-first-order [38] and a pseudo-second-order [39] rate kinetic models. The mathematical expressions of the pseudo-first-order and the pseudo-second-order models were showed in the following equations (4) and (5), respectively.

log(qe  qt )=logqe 

k1 t 2.303

(4)

t 1 t =  2 qt k2 qe qe

(5)

where qe and qt (mg/g) are the amount of adsorbed fluoride at equilibrium and at any reaction time t (min); k1 (min-1) and k2 (g mg-1 min-1) are the equilibrium rate

10

constants. During the fitting, normalized standard deviation (Δq%) was used as the objective function to be minimized:

q% = 100

 ((q

exp

 qcal ) / qexp )2 / ( N i  1)

(6)

where qexp (mg/g) is the experimental fluoride uptake, qcal (mg/g) is the amount of the calculated fluoride adsorbed and Ni is the number of data points, respectively [40]. The linear fitting of the two models to the kinetic data are shown in Fig. S1. The corresponding parameters and Δq% are listed in Table 1. The results suggested that the pseudo-second-order model predicted the experimental data satisfactorily well. Thus, the pseudo-second-order model was the suitable ones to describe the adsorption kinetics. Fig. 8 shows the adsorption isotherm, and the adsorption equilibrium data were analyzed using Langmuir, Freundlich, Tempkin and Dubinin–Radushkevich (D-R) isotherm models, respectively. The Langmuir isotherm model [41] is valid for monolayer adsorption, which can be represented by the following equation: Ce 1 Ce  + qe qmb qm

(7)

where qm (mg g-1) is the maximum amount of adsorbate per unit weight of adsorbent at high Ce, and b (L mg-1) is the Langmuir isotherm constant that relates to the adsorption energy. The linear plot of Ce/qe versus Ce suggests that adsorption obeys the Langmuir model, and the constants qm and b can be calculated from slope and intercept, respectively. The obtained linear plot is shown in Fig. S2a, and the calculated values of qm and b are listed in Table 2. The value of R2 was 0.9994. The calculated maximum adsorption capacity was 46.80 mg g-1. Furthermore, according to the Langmuir equation, the adsorption capacity was 11.58 mg g-1 when the 11

equilibrium fluoride concentration was just below the WHO guideline of 1.5 mg L-1. The Freundlich isotherm model [42] can be utilized for simulating the adsorption on the heterogeneous surfaces. The Freundlich model is described as follows: ln qe 

1 ln Ce  ln K F n

(8)

where KF is a constant related to the adsorption capacity, and 1/n is the adsorption intensity. The above constants can be calculated from the intercept and slope of linear plot of lnqe versus lnCe, respectively. The obtained results are shown in Fig. S2b, and the values of KF and 1/n are also listed in Table 2. The value of 1/n lies between 0 and 1 suggests the favorable conditions for fluoride adsorption. However, the value of R2 was 0.9164, which was much lower than the ones of the Langmuir model. In addition, the value of Δq% was about three times higher than that of the Langmuir model. Thus, the fluoride adsorption can be fitted well with the Langmuir model. The Temkin isotherm model considers that the adsorption heat will decrease linearly with coverage owing to the adsorbent/adsorbate interactions, and the mathematical expression of the Temkin isotherm model is given as [43]: qe  RT / b(ln ACe )

(9)

B  RT / b

(10)

where b is the Temkin isotherm constant related to adsorption heat, A (L mg-1) is the equilibrium isotherm binding constant. R is the gas constant (8.314 J/(mol K)) and T is the absolute temperature (K). The Temkin model in linear form is: qe  B(ln A)  B(ln Ce )

(11)

The constants A and B can be determined from the intercept and slope of linear 12

plot of qe versus lnCe, respectively. The obtained results are shown in Fig. S2c and Table 2. The R2 value of the Temkin model was 0.9446, which was lower than the ones of the Langmuir isotherm model. Furthermore, the value of Δq% was also very high. Thus, the Temkin isotherm model was not suitable for representation the experimental data of the complex liquid phase adsorption system, which was coincident with the previous report [44]. Furthermore, it is known that Langmuir and Freundlich isotherms are not sufficient to explain the physical and chemical characteristics of the adsorption process. Thus, the data were also analyzed using the Dubinin–Radushkevich model [45] to determine the type of the adsorption. The linear form of D–R isotherm is expressed as:

ln qe  ln qm   2

(12)

Where qe (mol g-1) has the same meaning as mentioned above, and qm (mol g-1) is the D–R adsorption capacity.  (mol2 KJ-2) is the activity coefficient related to the mean adsorption energy. ε is the Polanyi potential and can be calculated as following equation:

 =RT ln(1 

1 ) Ce

(13)

The values of qm and  can be calculated from the intercept and slope according to the line of the lnqe versus ε2. The mean adsorption energy, E (kJ mol-1), can be obtained using the equation [46]:

E=

1 2

(14)

13

From Fig. S2d and Table 2, The R2 value of the D–R isotherm was the lower than that of Langmuir model, and higher than the ones of Freundlich and Temkin models. It is well known that the adsorption type can be determined by the value of E. If the value is between 1.0 to 8.0 kJ mol-1, the adsorption type can be explained by physical adsorption. If the value of E ranges from 8 and 16 kJ mol-1, the adsorption type corresponds to a chemical ion-exchange process [46]. Herein, the value of E was 11.68 kJ mol-1. The results suggested that the fluoride adsorption on MnCO3 nanowires was an ion-exchange process. Furthermore, a comparison of the fluoride adsorption of various metallic oxides is summarized in Table 3. Accordingly, the results definitely show that the as-prepared MnCO3 nanowires achieved good adsorption performance towards fluoride. 3.3. Fluoride removal mechanism For the chemical adsorption process, FTIR has been proved to be a very useful technique to reveal the structure changes of the adsorbents. From Fig. 9, the peaks centered at 3348 and 1626 cm-1 were related to the stretching and bending vibrations of the adsorbed water, respectively. The peak at 1387 cm-1 was assigned to the stretching bands of carbonates [58]. The peaks at 1456, 1348, and 1064 cm-1 were assigned to the surface monodentate carbonates, and the peaks at 1583 and 864 cm-1 were related to the surface bidentate carbonates[59, 60]. The weak peak at 1156 cm-1 was assigned to the bending vibration of the hydroxyl group of MnOOH [61, 62]. The formation of MnOOH on the surface MnCO3 was consistent with the previous report [37]. The peaks around 640 and 505 cm-1 may be assigned to the Mn-O vibrations of MnOOH [61-64]. The peak at 1036 cm-1 can be assigned to the bending vibration of the surface hydroxyl groups [65-67]. After fluoride removal, the peaks at 1456 and 14

1348 cm-1 turned slightly weak, and the peak at 1064 cm-1 was turned weak obviously. Moreover, the peak at 1583 cm-1 can not be observed. These FTIR peaks were related to the surface carbonate-like groups. Thus, the results suggested that the ion-exchange process were taken place between the surface carbonate-like groups and fluoride. In addition, the peaks at 1156 and 1036 cm-1 were also turned very weak, suggesting that the hydroxyl groups were also participated in the ion-exchange process. To further confirm the proposed fluoride removal mechanism, XPS was utilized to investigate the surface structures of the MnCO3 nanowires, and the results are shown in Fig. 10. From the survey spectra shown in Fig. 10a, the peak at 685 eV was detected after fluoride removal. This peak can be assigned to the F1s spectrum, as shown in Fig. 10b. From the C1s spectra shown in Fig. 10c, before fluoride adsorption, the C1s spectrum suggests four peaks are located at 284.7, 286.3, 287.8, and 289.4 eV, respectively. The peak at 289.4 eV can be assigned to the surface carbonate-like groups of the MnCO3 nanowires [68]. It is clear that the carbon atom content was decreased from 20.11% to 15.77%. Considering the analysis of FTIR, the results further confirmed the presence of the ion-exchange process between the surface carbonate-like groups and fluoride. Furthermore, Fig. 10d presents the O1s spectra before and after fluoride removal. For MnCO3 nanowires before fluoride adsorption, the O1s spectrum suggests four peaks are located around 530.1, 531.4, 532.2, and 533.3 eV, respectively. The peaks at 530.1, 531.4 and 533.3 eV can be assigned to the bulk oxygen (O2-), the surface hydroxyl groups, and the adsorbed water, respectively [69]. The front two peaks should be related to MnOOH. The peak at 532.2 eV can be assigned to the surface carbonate-like groups of MnCO3 [37, 70]. From Fig. 10d, the content of the oxygen atoms in the surface hydroxyl groups was decreased from 26.99% to 24.61%. At the same time, the content of the oxygen atoms in the surface 15

carbonate-like groups was decreased from 20.57% to 18.07%. Thus, the results mean that both of the surface hydroxyl groups and the carbonate-like groups were participated in the ion-exchange process, which was also consistent with the FTIR results. To further confirm the proposed fluoride removal mechanism, in the process of the kinetics experiment, the influence of the contact time on the pH values of the solution were also investigated. The initial concentration of fluoride was chosen as 100 mg L-1, and the pH of the solution was adjusted to 7.0. After the experiment started, the pH value was not adjusted anymore during the adsorption process. From Fig. 11a, the time reached to the adsorption equilibrium was about 3 h. At the same time, the pH values were increased rapidly during the first hour, and then gradually reaches a plateau, as shown in Fig. 11b. The final pH value was 7.45, suggesting the surface hydroxyl groups were released to the solution. Moreover, some coexisting anions such as chloride ions, nitrate, sulfate, bicarbonate, and phosphate may interfere in the fluoride adsorption process. Thus, to evaluate the influence of these coexisting anions on fluoride adsorption, the above anions were added to fluoride solutions at three concentration levels (10, 50 and 100 mg L-1). The fluoride removal rate was studied, and the results are shown in Fig. 11c. It is clear that chloride ions and nitrate have almost no influence on fluoride removal. With the presence of 10 mg L-1 of sulfate, the fluoride removal rate was almost unchanged. However, the removal rate was decreased to 85.3 and 78.6% when the concentration of sulfate increased to 50 and 100 mg L-1, respectively. Furthermore, the presence of bicarbonate and phosphate has a great influence on fluoride removal. When the concentration increased to 100 mg L-1, the removal rates were decreased to 16

62.0% and 45.4%, respectively. The results were similar to the previously reported adsorbents [29, 47, 56]. Based on the above proposed fluoride removal mechanism, the surface hydroxyl groups play the important roles in the fluoride removal process. That is to say the surface hydroxyl groups will be substituted by fluoride. Thus, based on the Le Chatelier's principle, the high pH of solution was unfavorable for the ion-exchange between the surface hydroxyl and fluoride [47, 56]. So, the influence of solution pH on fluoride adsorption was studied. From Fig. 12a, the fluoride adsorption on MnCO3 nanowires was very effective under weak acidic and weak basic conditions. In the pH range of 5.0 to 8.0, the adsorption capacities were larger than 44 mg g-1. At pH 4, the adsorption capacity was decreased to 39.50 mg g-1, which may be related to the low stability of adsorbent in the acid conditions. Furthermore, the adsorption capacities were decreased with the increased pH values. The adsorption capacity was decreased to 23.30 mg g-1 at pH 11. The high concentration of OH- ions will interfere the substitution reaction between the surface hydroxyl groups and fluoride ions, which was coincident with above proposed mechanism. In addition, the strong basic solution will in favor the regeneration of the adsorbent. Thus, the fluoride adsorption performance of the regenerated adsorbent was studied. Five consecutive adsorption– desorption cycles were performed, and the results are shown in Fig. 12b. The adsorption capacities were decreased with the increased adsorption–desorption cycles. After the first cycle, it was decreased to 35.50 mg g-1. Furthermore, the adsorption capacity decreased to 12.10 mg g-1 after five cycles. The results suggested that the asprepared adsorbent could be partly regenerated in strong basic solution, and also confirmed the proposed fluoride mechanism. Thus, based on the above discussion, a reasonable fluoride mechanism is schematically presented in Fig. 13. 17

4. Conclusions In summary, this work demonstrated a simple EG mediated solvothermal method for the controlled synthesis of MnCO3 products. The volume ratios of EG and water have a decisive impact on the morphology of MnCO3. With a decreased water volume, MnCO3 products transformed from microcubes to nanowires. Fluoride removal properties of the obtained MnCO3 nanowires were revealed. Kinetic study results indicated that the adsorption process followed the pseudo-second-order model. The equilibrium data were fitted to the Langmuir model, and the maximum adsorption capacity at neutral pH was 46.80 mg g-1. The coexisting chloride ions and nitrate have no influence on fluoride removal. Sulfate, bicarbonate and phosphate obviously reduces the fluoride removal, especially at high concentrations. The as-prepared MnCO3 adsorbent could be readily regenerated using NaOH solution and be repeatedly used for fluoride removal. More importantly, the fluoride removal mechanism was revealed by FTIR and XPS, and the results suggested that the removal of fluoride was an ion-exchange process based on the surface hydroxyl groups and the surface carbonate-like groups. This work demonstrates the important role of the surface carbonate-like groups of carbonates in fluoride removal process, and also expands the application field of carbonates, which should be of importance for both theoretical investigations and practical water treatment applications.

Acknowledgements This work was supported by the National Natural Science Foundation of China (Grant no.51302102), the Key Natural Science Research Project for Colleges and Universities of Anhui Province (Grant no.KJ2016A638), the Huaibei Scientific Talent Development Scheme(Grant no.20140305), and the 18

exploratory research project of Anhui University of Traditional Chinese Medicine (Grant no.2016ts068).

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28

Figure captions Fig. 1. XRD pattern of the obtained product synthesized with a 75 : 25 volume ratios of EG and water. Fig. 2. SEM (a, b) and TEM (c, d) images of the as-prepared MnCO3 nanowires synthesized with a 75 : 25 volume ratios of EG and water. Fig. 3. SEM images of the MnCO3 products synthesized with a 50 : 50 (a, b) and 25:75 (c, d) volume ratios of EG and water. Fig. 4. SEM images of the MnCO3 products synthesized with a 10 : 90 (a) volume ratios of EG and water, and pure water (b). Fig. 5. Low-magnification (a) and high-magnification (b) SEM images of the MnCO3 products synthesized with a 90 : 10 volume ratios of EG and water. Fig. 6. Nitrogen adsorption–desorption isotherm (a), and the pore-size distribution curve (b) of the obtained MnCO3 nanowires. Fig. 7. (a) Adsorption capacities of fluoride onto MnCO3 products synthesized with different volume ratios of EG and water (C0 = 200 mg L-1, pH = 7.0, temperature = 25 °C, adsorbent dose = 1 g L-1). (b) Adsorption kinetics of fluoride on MnCO3 nanowires (C0 = 5 mg L-1, pH = 7.0, temperature = 25 °C, adsorbent dose = 1 g L-1). Fig. 8.

Adsorption isotherm obtained from fluoride adsorption onto MnCO3

nanowires (pH = 7.0, temperature = 25 °C, adsorbent dose = 1 g L-1). Fig. 9. FTIR spectra of the MnCO3 nanowires before and after fluoride removal. Fig. 10. XPS survey spectra (a) of the MnCO3 nanowires before and after fluoride removal. XPS F1s spectrum (b) of the MnCO3 nanowires after fluoride adsorption. 29

XPS C1s (c) and O1s (d) spectra of the MnCO3 nanowires before and after fluoride removal. Fig. 11. Effects of contact time on adsorption of fluoride (a), and on the pH values (b) of the solution (C0 = 100 mg L-1, temperature = 25 °C, adsorbent dose = 1 g L-1). (c) Effects of competing anions on fluoride adsorption (C0 = 5 mg L-1, temperature = 25 °C, contact time = 24 h, adsorbent dose = 1 g L-1).

Fig. 12. (a) Effects of pH value on fluoride adsorption on MnCO3 nanowires. (b) Adsorption capacities of fluoride onto MnCO3 nanowires in a successive cycle (C0 = 200 mg L-1, temperature = 25 °C, adsorbent dose = 1 g L-1). Fig. 13. Schematic illustration of the fluoride adsorption mechanism on the MnCO3 adsorbent.

Table 1 Comparison of pseudo-first-order and pseudo-second-order models parameters, and calculated qe(cal) and experimental qe(exp) values. Table 2 Comparison of the coefficients isotherm parameters for fluoride adsorption onto MnCO3 nanowires. Table 3 Comparison of the fluoride adsorption properties of various adsorbents.

30

Fig. 1. XRD pattern of the obtained product synthesized with a 75 : 25 volume ratios of EG and water.

31

Fig. 2. SEM (a, b) and TEM (c, d) images of the as-prepared MnCO3 nanowires synthesized with a 75 : 25 volume ratios of EG and water.

32

Fig. 3. SEM images of the MnCO3products synthesized with a 50 : 50 (a, b) and 25:75 (c, d) volume ratios of EG and water.

33

Fig. 4. SEM images of the MnCO3 products synthesized with a 10 : 90 (a) volume ratios of EG and water, and pure water (b).

34

Fig. 5. Low-magnification (a) and high-magnification (b) SEM images of the MnCO3 products synthesized with a 90 : 10 volume ratios of EG and water.

35

Fig. 6. Nitrogen adsorption–desorption isotherm (a) and the pore-size distribution curve (b) of the obtained MnCO3 nanowires.

36

Fig. 7. (a) Adsorption capacities of fluoride onto MnCO3 products synthesized with different volume ratios of EG and water (C0 = 200 mg L-1, pH = 7.0, temperature = 25 °C, adsorbent dose = 1 g L-1). (b) Adsorption kinetics of fluoride on MnCO3 nanowires (C0 = 5 mg L-1, pH = 7.0, temperature = 25 °C, adsorbent dose = 1 g L-1).

37

Fig. 8.

Adsorption isotherm obtained from fluoride adsorption onto MnCO 3

nanowires (pH = 7.0, temperature = 25 °C, adsorbent dose = 1 g L-1).

38

Fig. 9. FTIR spectra of the MnCO3 nanowires before and after fluoride removal.

39

Fig. 10. XPS survey spectra (a) of the MnCO3 nanowires before and after fluoride removal. XPS F1s spectrum (b) of the MnCO3 nanowires after fluoride adsorption. XPS C1s (c) and O1s (d) spectra of the MnCO3 nanowires before and after fluoride removal.

40

Fig. 11. Effects of contact time on adsorption of fluoride (a), and on the pH values (b) of the solution (C0 = 100 mg L-1, temperature = 25 °C, adsorbent dose = 1 g L-1). (c) Effects of competing anions on fluoride adsorption (C0 = 5 mg L-1, temperature = 25 °C, contact time = 24 h, adsorbent dose = 1 g L-1).

41

Fig. 12. (a) Effects of pH value on fluoride adsorption on MnCO3 nanowires. (b) Adsorption capacities of fluoride onto MnCO3 nanowires in a successive cycle (C0 = 200 mg L-1, temperature = 25 °C, adsorbent dose = 1 g L-1).

42

Fig. 13. Schematic illustration of the fluoride adsorption mechanism on the MnCO3 adsorbent.

43

Table 1 Comparison of pseudo-first-order and pseudo-second-order models parameters, and calculated qe(cal) and experimental qe(exp) values. pseudo-first-order model C0 (mg L-1)

qe(exp) (mg g-1)

k1 (min-1)

qe(cal) (mg g-1)

R2

Δq%

5

4.81

0.0168

0.35

0.4631

58.2

Pseudo-second-order model C0 (mg L-1)

qe(exp) (mg g-1)

k2 (g mg-1 min-1)

qe(cal) (mg g-1)

R2

Δq%

5

4.81

0.186

4.83

0.9999

3.58

44

Table 2 Comparison of the coefficients isotherm parameters for fluoride adsorption onto MnCO3 nanowires. Isotherm model

Parameters

Langmuir qm (mg g-1)

46.80

-1

b (L mg ) R

0.2254

2

0.9994

Δq%

8.18

Freundlich 1/n

0.3579

KF(mg g-1) (L mg-1)1/n

9.77

R

2

0.9164

Δq%

22.99

Tempkin A (L mg-1)

6.060

B (mg g-1)

7.075

R2

0.9446

Δq%

40.14

Dubinin–Radushkevich qm (mg g-1)

90.35

 (mol2 KJ-2)

0.00368

E (KJ mol-1)

11.68

R2

0.9508

Δq%

18.67

45

Table 3 Comparison of the fluoride adsorption properties of various adsorbents. Adsorbents

S(m2/g)

pH

Ceq (mg L-1)/qe(mg g-1)

Ceq (mg L-1)/qe(mg g-1)a

Ref.

Ferric hydroxide

250-300

7.0

40/7.0

1/~3.0

47

Hydrous ZrO2

134

7.0

70/68

1.5/8.3

48

Boehmite

/

7.5

50/2.06

5/~0.4

49

Fe-Ti bimetal oxide

/

/

130/47.0

2.5/~7.5

50

Fe-doped TiO2

/

5.0

120/53.22

1.5/9.77

51

CaO/Al2O3

93

6.8

1000/137

1.5/10.2

52

CeO2/Al2O3

/

/

45/37.14

5/~7.5

53

Sulfate-doped Fe 3O4/Al2O3

63.4

7.0

100/70.4

1.5/17.8

54

Alum-impregnated Al2O3

176

6.5

20/40.3

1.5/0.54

55

Natroalunite

206.25

7.0

120/85.84

1.5/19.6

56

Mn3O4

23

6.0

120/2.8

/

57

MnCO3 nanowires

57.6

7.0

150/46.80

1.5/11.58

This work

a

The author reported, and calculated from the adsorption isotherm or Langmuir equation.

46

Graphical abstract

MnCO3 nanowires and microcubes were controlled synthesized. Ionexchange adsorption of fluoride on MnCO3 based on the surface carbonatelike groups and hydroxyl groups were revealed.

47