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Abstract

The silica-based CeO₂ adsorbent (CeO₂/SiO₂) was prepared for removing fluoride from the aqueous solution. The synthesized adsorbent was characterized by scanning electron microscope, energy dispersive spectrum, X-ray diffractometer, Fourier transform infrared spectrometer, and zeta potential measurement analyses. The adsorption batch experiments in the various experimental conditions including solution pH, contact time, initial fluoride concentration, and adsorption temperature were performed and investigated. The maximum adsorption capacity of fluoride into CeO₂/SiO₂ was 2.441 mmol/g at pH 3 and 298 K. The adsorption kinetics and isotherms were well described by the pseudo-second-order model and the Langmuir model, respectively. The fluoride adsorption reached the equilibrium in 15 min from the aqueous solution with the initial fluoride concentration of 400 mg/l at 298 K. In the temperature range of 298–338 K, the maximum adsorption capacity of fluoride decreased from 2.441 mmol/g to 2.109 mmol/l at pH 3. The adsorption thermodynamics study revealed that this process was a spontaneous, exothermic, and entropy-driving adsorption. Furthermore, the mechanism of adsorption was identified as the anion exchange and the electrostatic interaction. The desorption efficiency of fluoride-loaded CeO₂/SiO₂ adsorbent could reach about 95% by 0.1 mol/l NaOH.

Keywords

porous silica fluoride adsorption desorption

Introduction

As is well known, fluorine is one of the essential trace elements in the human body for the formation of dental enamel and normal mineralization of bones (Bell and Ludwig, 1970; Xu et al., 2017). However, the excessive intake of fluoride is harmful to the body that can cause fluorosis (Cai et al., 2015; Ghosh et al., 2013). According to World Health Organization, the maximum acceptable level of fluoride in drinking water is 1.5 mg/l (Lin et al., 2015). In recent years, water pollution has become more and more severe due to the urbanization and industrialization. The wastewater produced from the fertilizer manufacturing plants, aluminum and steel smelteries contained a significant amount of fluoride (Wang et al., 2017a). Moreover, with the development of the nuclear industry, nuclear wastewater containing a significant amount of fluoride from the production of nuclear fuel components also has been increasing year by year (Sawant et al., 2007). The fluoride discharge standard for the industrial wastewater is 4 mg/l according to the United State Environmental Protection Agency regulation. Therefore, it is necessary to reduce the fluorine concentration to the introduced safe limits in drinking water and industrial wastewater discharges.

Some commonly used methods for removing fluoride from wastewater are adsorption, chemical precipitation, membrane separation process, electrolytic defluoridation, and electrodialysis (Cai et al., 2016; Dhillon et al., 2015; Loganathan et al., 2013; Shen et al., 2015). Among them, membrane separation, electrolytic defluoridation, and electrodialysis techniques are relatively expensive. Also, the concentration of fluoride in water cannot be reached to less than 2 mg/l by precipitation of calcium fluoride with lime (Fan et al., 2003; Wang and Reardon, 2001). Adsorption technique, with the advantages such as the fast kinetic, ease of handling, and simple equipment, has been widely used in the research and industrial application. Therefore, various adsorbents for adsorption of fluoride from aqueous solutions, such as activated carbon, bimetal mixed oxide, biomass-based adsorbents, Co–Al hydroxide carbonates, red mud, mesoporous alumina, hydroxyapatite, and so on, have been studied (Biswas et al., 2009; Dayananda et al., 2014; Liang et al., 2014; Ma et al., 2009; Mourabet et al., 2015; Mukherjee and Halder, 2016; Zhao et al., 2015). As known, La(III), Ce(IV), Y(III), and Zr(IV) in the form of oxides showed a good fluoride removal performance (Tokunaga et al., 1995). Among these metal oxides, cerium oxide has exhibited the high selectivity for the fluoride adsorption (Nomura et al., 1990). However, the powder form of cerium oxide could cause the high pressure drop, column blockage, and low flow rate, which is not suitable for the large-scale column operation. To overcome these limitations, the CeO₂–ZrO₂ nanocages adsorbent (Wang et al., 2013) and Ce–La binary hydroxide adsorbent (Zhang et al., 2010) have been synthesized, and both of them showed a good fluoride uptake, but their kinetics study revealed the slow adsorption rate. Therefore, the attempt was made to synthesize a novel adsorbent which has the large specific surface, fast diffusion kinetics, and low pressure drop in a packed adsorption column. A novel porous silica support with the high porosity of 60%, high mechanical strength, and high radiation stability has been developed by Wei et al. (2014) to make silica-based adsorbents suitable for the column operation (Wu et al., 2012; Zhang et al., 2016). Thus, a novel CeO₂/SiO₂ adsorbent was synthesized in this study.

The present study dealt with the preparation and characterization of CeO₂/SiO₂ adsorbent. The adsorption and desorption behavior of fluoride onto CeO₂/SiO₂, such as the kinetics, isotherms, and thermodynamics, have also been investigated. Moreover, the adsorption mechanism was studied by the Fourier transform infrared spectrometer (FT-IR) and zeta potential measurement.

Experimental

Materials

All the reagents are of analytical grade. Cerium nitrate hexahydrate (CeN₃O₉·6H₂O, 99.5%, Aladdin, P.R. China) was used to synthesize the adsorbent. Sodium fluoride (NaF, 98.0%, Sinopharm Chemical Reagent Co., Ltd, P.R. China) was dissolved to prepare the stock solution containing 1000 mg/l of F⁻. Trisodium citrate dehydrate (C₆H₅Na₃O₇·2H₂O, 99.0%, Sinopharm Chemical Reagent Co., Ltd, P.R. China) and sodium nitrate (NaNO₃, 99.0%, Sinopharm Chemical Reagent Co., Ltd, P.R. China) were used to prepared the buffer solution (0.2 mol/l C₆H₅Na₃O₇–1 mol/l NaNO₃, pH = 5∼6). The SiO₂ with the particle size of 60 µm was used as the support material in synthesizing the adsorbent.

Instruments

The surface morphology and energy dispersive spectrum (EDS) analysis of CeO₂/SiO₂ were characterized by the field-emission scanning electron microscopy (LV UHR FE-SEM, NOVA NanoSEM 230, USA). X-ray diffraction pattern was obtained using the polyfunctional X-ray diffractometer (XRD, D8 ADVANCE Da Vinci, Germany). The adsorption mechanism was investigated by means of the FT-IR (Nicolet 6700, USA) and Zeta potential analyzer (ZS90, Malvern Instruments Ltd, UK). The ion meter (PXSJ-216F, Shanghai INESA Scientific Instrument Co., Ltd, P.R. China) was used to determine the concentration of fluoride.

Preparation of silica-based adsorbent

In this study, cerium oxide was immobilized in the porous silica support to prepare the CeO₂/SiO₂ absorbent. The CeO₂/SiO₂ was synthesized by the following steps. At first, cerium nitrate hexahydrate powder was dissolved in methanol to prepare the Ce(NO₃)₃/methanol solution. Then, the SiO₂ particles were added and the mixture was rotated by a rotary evaporator for 1 h at room temperature, so that the Ce(NO₃)₃/methanol solution could be dispersed into the pores of SiO₂. Afterward, the resultant mixture was rotated under reduced pressure at 339 K to remove methanol. Subsequently, the product was dried under vacuum overnight at 343 K to form Ce(NO₃)₃/SiO₂. Finally, Ce(NO₃)₃/SiO₂ was calcinated in the muffle furnace at 523 K for 3 h to decompose Ce(NO₃)₃ into CeO₂, and the CeO₂/SiO₂ absorbent was obtained. The preparation procedure of CeO₂/SiO₂ absorbent is given in Figure 1.

Figure 1.

Preparation procedure of CeO₂/SiO₂ absorbent.

Batch experiments

The batch experiments were used to investigate the effect of pH, contact time, temperature, and the initial fluoride concentration to optimize the adsorption process of fluoride onto the CeO₂/SiO₂ adsorbent. The experiments were performed in a glass vial with 20 ml of the aqueous solution containing fluoride and 0.05 g CeO₂/SiO₂ adsorbent. The glass vial was shaken mechanically at 120 r/min in a thermostatic water bath at the determined adsorption conditions. The adsorbent was then separated from the solution by a filter with the pore size of 0.45 µm. The concentration of fluoride in the aqueous solutions before and after adsorption was determined by an ion meter. Moreover, desorption experiments were accomplished using different concentrations of sodium hydroxide as the desorption agent.

The adsorption capacity (Q, mmol/g), distribution coefficient (K_d, ml/g), and desorption efficiency in a batch experiment were calculated by the following equations

\begin{matrix} K_{d} = \frac{C_{0} - C_{e}}{C_{e}} \times \frac{V}{m} \end{matrix}

(1)

\begin{matrix} Q = \frac{(C_{0} - C_{e}) \times V}{1000 \times m} \end{matrix}

(2)

\begin{matrix} Desorption efficiency = \frac{C_{e} \times V}{1000 \times Q \times m} \times 100 % \end{matrix}

(3)

where C₀ (mmol/l) and C_e (mmol/l) are the initial concentration of fluoride and the concentration of fluoride at equilibrium, respectively. V (ml) and m (g) denote the volume of the aqueous solution and the weight of CeO₂/SiO₂ adsorbent, respectively.

Results and discussion

Characterization of CeO₂/SiO₂

Scanning electron microscope (SEM) images of CeO₂/SiO₂ before and after adsorption are shown in Figure 2. According to Figure 2, the surface morphology of CeO₂/SiO₂ was observed clearly, and its homogeneous surface showed a good porosity. After adsorption, the surface morphology of CeO₂/SiO₂ was almost unchanged and the overall structure of the adsorbent was in good shape, reflecting that CeO₂/SiO₂ had the high mechanical strength. The distribution of elements in CeO₂/SiO₂ after adsorption is shown in Figure 2(d). From Figure 2(d), it was found that the adsorbent contained Ce and F elements, which indicated that cerium oxide was successfully immobilized onto the silica support and fluoride was well adsorbed into CeO₂/SiO₂, respectively. EDS analysis of CeO₂/SiO₂ is displayed in Figure 3. Figure 3(a) showed that there was almost no fluoride peak and the content of fluoride was 0% before adsorption. As shown in Figure 3(b), the apparent fluoride peak appeared in the spectra, and the content of fluoride in the adsorbent was 5.58% after adsorption. Therefore, EDS analysis further confirmed that CeO₂/SiO₂ adsorbent can absorb fluoride from aqueous solution.

Figure 2.

(a) and (b) SEM images of CeO₂/SiO₂ before adsorption, (c) SEM images of CeO₂/SiO₂ after adsorption, and (d) elemental mapping images of CeO₂/SiO₂ after adsorption. SEM: scanning electron microscope.

Figure 3.

EDS of CeO₂/SiO₂ adsorbent (a) before and (b) after adsorption of fluoride. EDS: energy dispersive spectrum.

It can be seen in XRD pattern (Figure 4) that the diffraction peaks appeared at 2θ = 28.44, 32.99, 47.33, and 56.22, which were in agreement with JCPDS Card No. 43-1002. It is noticed that there was no other diffraction peak and the characteristic diffraction peaks of CeO₂ were sharp, revealing that CeO₂ had high purity and was well crystallized. The BET surface area of CeO₂/SiO₂ was measured as 230 m²/g by mercury intrusion method. Above results indicated that CeO₂/SiO₂ adsorbent was successfully synthesized.

Figure 4.

X-ray diffraction pattern of CeO₂/SiO₂.

Effect of pH

The influence of pH on the adsorption capacity was studied under different background electrolytes at 298 K, and the result is presented in Figure 5. It was obvious that the adsorption capacity changed significantly with the variation of pH value, and the adsorption capacity reached the maximum at pH = 2.5. As seen in Figure 5, the adsorption capacity was increased with increasing pH from 1 to 2.5. It might be due to the dominant chemical form of fluorine which is existed as HF at low pH value, and the free fluorine ion (F⁻) increased with an increase in pH (Tang et al., 2009). After reaching the maximum value, the adsorption capacity decreased dramatically. It might be because of the competition of hydroxyl ions in the solution with the fluoride above pH 2.5. In addition, the effect of ionic strength was also studied and the result showed that the ionic strength had no obvious effect on the adsorption.

Figure 5.

Effect of pH on the adsorption of fluoride toward CeO₂/SiO₂ adsorbent (T = 298 K, m/V = 0.05 g/20 ml, t = 5 h, r = 120 r/min, [F⁻]_initial = 100 mg/l).

Adsorption kinetics

The fluoride adsorption kinetics of CeO₂/SiO₂ adsorbent at different initial concentrations of fluoride was studied. As seen in Figure 6(a), the adsorption capacity increased rapidly at the initial stage, which might be caused by a large amount of surface hydroxyl groups on CeO₂/SiO₂ adsorbent at the beginning of the adsorption. The equilibrium was achieved within 15 min with different initial concentrations of fluoride at 298 K. To further analyze the adsorption kinetics, the observed adsorption data were fitted by the pseudo-first-order and pseudo-second-order models.

Figure 6.

(a) Effect of contact time on the adsorption with different initial fluoride concentrations (T = 298 K, m/V = 0.05 g/20 ml, r = 120 r/min, [F⁻]_initial = 100, 400 mg/l, pH = 3); (b) the fitting of adsorption data on the pseudo-first-order kinetic equation; and (c) the fitting of adsorption data on the pseudo-second-order kinetic equation.

The linearized form of the pseudo-first-order model can be written as equation (4) (Javadian et al., 2013)

\begin{matrix} \ln (Q_{e} - Q_{t}) = \ln Q_{e} - k_{1} t \end{matrix}

(4)

where Q_e (mmol/g) and Q_t (mmol/g) are the equilibrium adsorption capacity and the amount of fluoride absorbed at time t (h), respectively, and k₁ (h⁻¹) is the rate constant of adsorption. The plot of ln(Q_e − Q_t) versus t is represented in Figure 6(b), and the relevant parameters are given in Table 1.

Table 1.

The adsorption kinetic models parameters for the adsorption of fluoride toward CeO₂/SiO₂ adsorbent.

C₀ (mg/l)	Pseudo-first-order kinetic model			Pseudo-second-order kinetic model				Q_e,exp (mmol/g)
C₀ (mg/l)	Q_e,cal (mmol /g)	k₁ (1/h)	R ²	Q_e,cal (mmol /g)	k₂ (g/mmol/h)	h (mmol/g/h)	R ²	Q_e,exp (mmol/g)
100	0.134	1.86	0.7531	1.373	71.17	134.23	0.9998	1.366
400	0.0561	0.112	0.1546	2.183	203.73	970.87	0.9991	2.158

The linearized form of the pseudo-second-order model can be given as equation (5) (Mukherjee et al., 2017)

\begin{matrix} \frac{t}{Q_{t}} = \frac{1}{k_{2} {Q_{e}}^{2}} + \frac{t}{Q_{e}} \end{matrix}

(5)

where k₂ (g/mmol/h) is the rate constant of adsorption in the pseudo-second-order kinetic model. The plot of t/Q_t versus t is shown in Figure 6(c), and the pseudo-second-order kinetic model parameters, including Q_{e, cal}, k₂, h (

h = k_{2} Q_{e}^{2}

(Liu et al., 2014)) and the correlation coefficient R², are listed in Table 1.

As shown in Table 1, Q_{e, exp} increased from 1.366 to 2.158 mmol/g with an increase in the initial concentration of fluoride from 100 to 400 mg/l. The correlation coefficient R² for the pseudo-second-order kinetic model was much closer to 1 than that for the pseudo-first-order kinetic model, and Q_{e, cal} for the pseudo-second-order kinetic model was approximately the same as Q_{e, exp} at different initial concentrations. Therefore, the adsorption was fitted better to the pseudo-second-order kinetic model, indicating the fluoride adsorption mechanism of CeO₂/SiO₂ adsorbent was of chemisorption type. The adsorption process included three stages: the external film diffusion, the intraparticle diffusion, and the adsorption of sorbate at the internal active sites. Since CeO₂/SiO₂ adsorbent had the high porosity and small particle size, the intraparticle diffusion had less influence on the kinetics of adsorption. Additionally, the adsorption of sorbate at the internal active sites seemed to be fast considering the achieved equilibrium time. Therefore, the external film diffusion was the rate-controlling step in the fluoride adsorption process. The pseudo-second-order rate constant (k₂) reflecting that the rate of adsorption increased with the increase in the initial concentration of fluoride, which was because of an increase in driving force in the external film diffusion stage (Choy et al., 2004).

Adsorption isotherms

To obtain the fluoride adsorption capacity of CeO₂/SiO₂ adsorbent, the adsorption isotherm study, as given in Figure 7(a), was used to investigate the effect of the initial concentration of fluoride on the adsorption capacity at the various temperatures. The adsorption curve exhibited that the uptake of fluoride increased by increasing the equilibrium concentration of fluoride in the solution and then reached the saturated adsorption capacity. Moreover, the equilibrium adsorption capacity decreased with the increase in temperature, indicating that the adsorption process was favored at lower temperatures. Furthermore, the experimental data were investigated by the Langmuir and Freundlich isotherm models which are the most commonly used in the adsorption isotherm study (Javadian et al., 2013; Özeroğlu and Metin, 2011).

Figure 7.

(a) Adsorption isotherm of fluoride toward CeO₂/SiO₂ adsorbent at different temperatures (m/V = 0.05 g/20 ml, t = 5 h, r = 120 r/min, pH = 3); (b) the plot of the Freundlich isotherm model for the adsorption; and (c) the plot of the Langmuir isotherm model for the adsorption.

The linear expression of the Freundlich isotherm model can be given as equation (6) (Javadian et al., 2013)

\begin{matrix} \log Q_{e} = \log K_{F} + \frac{1}{n_{F}} \log C_{e} \end{matrix}

(6)

where K_F (mmol/g) and n_F are Freundlich constants representing the adsorption capacity and intensity of adsorption, respectively. The Freundlich plot of log Q_e versus log C_e is depicted in Figure 7(b). The obtained parameters including K_F, n_F, and the correlation coefficient R² are listed in Table 2.

Table 2.

The Freundlich and Langmuir adsorption isotherm parameters for the adsorption of fluoride toward CeO₂/SiO₂ adsorbent.

	Freundlich isotherm model			Langmuir isotherm model				Q_m,exp (mmol/g)
T (K)	n _F	K_F (mmol/g)	R ²	Q_m,cal (mmol/g)	K_L (l/mmol)	R _L	R ²	Q_m,exp (mmol/g)
298	3.50	1.17	0.9416	2.489	1.31	0.0280–0.587	0.9937	2.441
318	3.51	1.08	0.9232	2.257	1.36	0.0269–0.586	0.9904	2.259
338	3.55	0.96	0.9382	2.105	1.02	0.0355–0.652	0.9870	2.109

The linear expression of the Langmuir isotherm model can be given as equation (7) (Javadian et al., 2013)

\begin{matrix} \frac{C_{e}}{Q_{e}} = \frac{1}{K_{L} Q_{m}} + \frac{C_{e}}{Q_{m}} \end{matrix}

(7)

where Q_m (mmol/g) and K_L (l/mmol) are the saturated adsorption capacity and Langmuir constant which is related to the binding energy of adsorption of Langmuir (Zhou et al., 2012), respectively. The Langmuir plot of C_e/Q_e versus C_e is displayed in Figure 7(c). Values for Q_m, K_L, the correlation coefficient R², and the dimensionless separation factor or the equilibrium parameter R_L are summarized in Table 2. The parameter R_L is defined by equation (8) (Alkan and Doğan, 2001; Liu et al., 2014)

R_{L} = \frac{1}{1 + K_{L} C_{0}}

(8)

where C₀ (mmol/l) is the initial concentration of fluoride. According to the R_L value, R_L > 1: unfavorable adsorption, R_L = 1: linear adsorption, 0 < R_L < 1: favorable adsorption, R_L = 0: irreversible adsorption (Alkan and Doğan, 2001).

As seen from the Table 2, the values of R² for the Langmuir isotherm model were closer to 1 at different temperatures, and the calculated saturated adsorption capacity Q for the Langmuir model was almost same as the experimental one. For the Freundlich isotherm model, the values of R² were relatively small. Therefore, the Langmuir isotherm model fitted the adsorption isotherms well when compared to the Freundlich model, indicating that the adsorption process was the monolayer adsorption on a homogeneous surface of the CeO₂/SiO₂ adsorbent. When temperature increased from 298 to 338 K, the adsorption capacity of fluoride decreased from 2.441 to 2.109 mmol/g, which meant the adsorption process was affected by the temperature and favored at a lower temperature. K_L represents the adsorption equilibrium constant, and the increase of temperature was unfavorable to forward the adsorption reaction, which led to decrease the K_L at a higher temperature. Moreover, the values of R_L given in Table 2 showed that the adsorption of fluoride toward CeO₂/SiO₂ adsorbent was favorable.

According to Table 3, CeO₂/SiO₂ adsorbent had the higher adsorption capacity than other adsorbents reported previously. Except for Al-modified hydroxyapatite, the BET surface area of CeO₂/SiO₂ was also greater than other adsorbents. Thus, there might be abundant surface hydroxyl groups onto the CeO₂/SiO₂ adsorbent, resulting in the efficient fluoride removal. Furthermore, except for ZrO₂/SiO₂/Fe₃O₄, the adsorption rate of CeO₂/SiO₂ was much faster than that of other adsorbents. Therefore, CeO₂/SiO₂ adsorbent showed the excellent defluorination performance.

Table 3.

Comparison of adsorption capacity of different adsorbents.

Adsorbent	pH	BET (m²/g)	Adsorption equilibrium time	Adsorption capacity (mg/g)	References
CeO₂/SiO₂	2.5	230	15 min	46.36	Present work
ZrO₂/SiO₂/Fe₃O₄	4	4.6	15 min	14.7	Chang et al. (2011)
CTAB-assisted mixed iron oxide	5.75	201.9	7 h	40.5	Mohapatra et al. (2011)
FeOOH/GO	2–10	202	>1 h	17.67	Kuang et al. (2017)
Zirconium-modified-Na-attapulgite	4.13	–	110 min	24.55	Zhang et al. (2012)
Pine wood biochar	2	2.73	2 days	7.66	Mohan et al. (2012)
Al-modified hydroxyapatite	7	258.6	5 h	32.57	Nie et al. (2012)
Polycinnamamide thorium (IV) phosphate	7	101.2	30 min	4.749	Islam et al. (2011)
3D RGO aerogel	6	–	About 200 min	31.3	Wu et al. (2015)
Mg–Al–Zr triple-metal composite	7	200.7	6 h	22.9	Wang et al. (2017b)
Zr-CTS/GO membrane	3–11	–	45 min	29.06	Zhang et al. (2017)
Polypyrrole/TiO₂	7	95.71	30 min	33.18	Chen et al. (2017)
Graphene	7	3.08	20–60 min	35.59	Li et al. (2011)

BET: Brunner?Emmet?Teller surface area; CTAB: Cetyltrimethyl ammonium bromide; CTS: Chitosan; RGO: Reduced-graphene oxide; GO: Graphene oxide.

Adsorption mechanism

The adsorption mechanism of fluoride into the CeO₂/SiO₂ was investigated to understand the interactions between the adsorbent and fluoride. Therefore, IR studies of CeO₂/SiO₂ before and after adsorption and zeta potential measurement were used to explore the fluoride adsorption mechanism of CeO₂/SiO₂. According to the analysis of IR spectra (Figure 8), the broad and strong peak at 3431 cm⁻¹ was due to the stretching vibration of adsorbed water-hydroxyl and surface hydroxyl groups. Also, the peaks at 1631 and 1382 cm⁻¹ were attributed to the bending vibrations of the adsorbed water-hydroxyl and surface hydroxyl groups, respectively. The peaks appeared at 1112, 803, and 474 cm⁻¹ were the characteristic adsorption peaks of SiO₂ (Song et al., 2008).

Figure 8.

IR spectra of SiO₂ and CeO₂/SiO₂ before and after adsorption. IR: infrared.

The main mechanism of the adsorption of anions on the hydrous metal oxides was identified as the anion exchange (Stumm and Werner, 1992). Therefore, the main mechanism of the fluoride adsorption on CeO₂/SiO₂ was inferred as follows

M - OH + F^{-} \Leftrightarrow M - F + O H^{-}

(9)

where M represents CeO₂. From Figure 8, the peak of CeO₂/SiO₂ before adsorption at 1382 cm⁻¹ was obvious and sharp, and it almost disappeared after the adsorption, indicating the fluoride ions in the solution were exchanged with surface hydroxyl groups of CeO₂. As shown in Table 4, the pH value of the solution increased after adsorption, which was because of the replacement of hydroxyl groups with fluoride.

Table 4.

The pH changes in the solution with different background electrolytes during the adsorption process.

[NaCl] (mol/l)	pH
[NaCl] (mol/l)	Before adsorption	After adsorption
0	3.06	3.19
0.001	2.95	2.97
0.01	3.06	3.17
0.1	3	3.15

The result of zeta potential measurement is shown in Figure 9. As can be seen, the zeta potential decreased with increasing pH value, revealing that the electrical charge at the adsorbent surface depended on the pH value of the aqueous solution. Thus, the adsorbent surface had the positive and negative surface charge under acidic and alkaline conditions, respectively. According to the previous report, the protonation and deprotonation of surface hydroxyl groups can be given as follows (Datsko and Zelentsov, 2016)

\begin{matrix} M - O {H_{2}}^{+} \to M - OH \to M - O^{-} \\ Low pH High pH \end{matrix}

(10)

Figure 9.

The variation of the zeta potential with pH of CeO₂/SiO₂.

Therefore, the CeO₂/SiO₂ adsorbent could adsorb the fluoride by the electrostatic interaction, which revealed the mechanism of the fluoride adsorption also included the electrostatic interaction.

By the above inferred adsorption mechanisms, the adsorption capacity in the low pH medium should be larger than that in the high pH medium, and the result of “effect of pH” (Figure 5) was well conformed to the inferred mechanisms.

Adsorption thermodynamics

To obtain adsorption thermodynamic parameters, the batch experiments were performed at different temperatures ranging from 298 to 338 K. The thermodynamic parameters can be determined by the Van't Hoff equation (Palodkar et al., 2017)

\begin{matrix} \ln K_{d} = \frac{Δ S}{R} - \frac{Δ H}{RT} \end{matrix}

(11)

where K_d (ml/g), ΔS (J/mol K), ΔH (J/mol), T (K), and R (8.314 J/mol K) are the distribution coefficient, the entropy change, the enthalpy change, temperature, and the universal gas constant, respectively. The Gibbs free energy change ΔG (kJ/mol) can be determined by the following equation (Palodkar et al., 2017)

Δ G = Δ H - T Δ S

(12)

Figure 10 shows the plot of ln K_d versus 1/T, from which ΔS, ΔH, and ΔG can be obtained, and the thermodynamic parameters are listed in Table 5. From Table 5, the values of ΔH and ΔS were negative and positive, respectively, reflecting that the adsorption process was exothermic and probably irreversible. When the temperature increased, the values of ΔG became more negative, which indicated that the fluoride adsorption of CeO₂/SiO₂ was feasible and spontaneous. Furthermore, the data |ΔH|<|−TΔS| suggested that the adsorption process was dominated by entropy changes rather than enthalpy changes.

Figure 10.

The Van't Hoff plot for the adsorption of fluoride toward CeO₂/SiO₂ adsorbent (m/V = 0.05 g/20 ml, t = 5 h, r = 120 r/min, [F⁻]_initial = 100 mg/l, pH=3).

Table 5.

Adsorption thermodynamic parameters for the adsorption of fluoride toward CeO₂/SiO₂ adsorbent.

ΔH (kJ/mol)	ΔS (J/mol K)	ΔG (kJ/mol)			R ²
ΔH (kJ/mol)	ΔS (J/mol K)	298 K	318 K	338 K	R ²
−7.59	30.28	−16.61	−17.22	−17.82	0.9832

Desorption and reusability study

Studying the regeneration and recycling of the spent adsorbents is an important indicator to evaluate the performance of the synthesized adsorbent. Various concentrations of NaOH were used as the desorption agent. Figure 11 shows that the desorption efficiency increased rapidly in the initial stage and the desorption process reached the equilibrium in 30 min. Furthermore, the higher the concentration of NaOH, the better the desorption efficiency would be. The desorption efficiency can reach about 95% by 0.1 mol/l NaOH solution, indicating CeO₂/SiO₂ adsorbent had an excellent desorption property.

Figure 11.

Desorption of fluoride from loaded CeO₂/SiO₂ with different concentration of NaOH solution (T = 298 K, m/V = 0.05 g/20 ml, r = 120 r/min).

To investigate the reusability of the CeO₂/SiO₂ adsorbent, the adsorption–desorption experiment was repeated four cycles, using 0.1 mol/l NaOH solution as the regeneration agent. As demonstrated in Figure 12, there was a 22% decrease in the adsorption efficiency in the first cycle and then the adsorption efficiency almost kept constant. The reason was that a partial amount of CeO₂ would fall off from the adsorbent during the adsorption process in the first cycle, and the remaining CeO₂ was stably immobilized into the SiO₂. Also, the desorption efficiency showed almost no change with the increase of adsorption–desorption cycles, which further confirmed that the CeO₂/SiO₂ adsorbent had a significant desorption performance with the excellent reusability.

Figure 12.

Effect of cycle number on the adsorption and desorption of CeO₂/SiO₂ (T = 298 K, m/V = 0.05 g/20 ml, r = 120 r/min, t = 5 h).

Conclusion

The silica-based adsorbent (CeO₂/SiO₂) was successfully synthesized to adsorb the fluoride ions from the aqueous solution. The suggested mechanisms of the adsorption included the anion exchange and electrostatic interaction. The optimum pH value of the fluoride adsorption was about 2.5. The fluoride adsorption process was described well by the pseudo-second-order kinetic model and the Langmuir isotherm model, indicating the adsorption process was the chemisorption type and monolayer adsorption on a homogeneous surface, respectively. Also, the obtained thermodynamic parameters showed that the adsorption process was a spontaneous, exothermic, and entropy-driving process. The adsorption process was favored at lower temperature, and the adsorption capacity decreased from 2.441 to 2.109 mmol/g with the increase in temperature from 298 to 338 K at pH 3. Moreover, CeO₂/SiO₂ adsorbent showed the excellent desorption property and reusability. Therefore, the synthesized CeO₂/SiO₂ adsorbent would be useful for the fluoride removal from the aqueous solution in the practical applications.

Footnotes

Declaration of Conflicting Interests

The author(s) declared no potential conflicts of interest with respect to the research, authorship, and/or publication of this article.

Funding

The author(s) disclosed receipt of the following financial support for the research, authorship, and/or publication of this article: This work was supported by the National Natural Science Foundation of China (11405106) and the Scientific Research Foundation for the Returned Overseas Chinese Scholars, State Education Ministry (No. 48 installment).

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Preparation of a novel CeO 2 /SiO 2 adsorbent and its adsorption behavior for fluoride ion

Abstract

Keywords

Introduction

Experimental

Materials

Instruments

Preparation of silica-based adsorbent

Batch experiments

Results and discussion

Characterization of CeO2/SiO2

Effect of pH

Adsorption kinetics

Adsorption isotherms

Adsorption mechanism

Adsorption thermodynamics

Desorption and reusability study

Conclusion

Footnotes

Declaration of Conflicting Interests

Funding

References

Characterization of CeO₂/SiO₂