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Analysis of F− removal from aqueous solutions using MgO
Journal of Water Process Engineering 25 (2018) 54–57
Contents lists available at ScienceDirect
Journal of Water Process Engineering journal homepage: www.elsevier.com/locate/jwpe
Analysis of F− removal from aqueous solutions using MgO ⁎
T
Tomohito Kameda , Yusuke Yamamoto, Shogo Kumagai, Toshiaki Yoshioka Graduate School of Environmental Studies, Tohoku University, 6-6-07 Aoba, Aramaki, Aoba-ku, Sendai, Japan
A R T I C LE I N FO
A B S T R A C T
Keywords: MgO F−Water treatment Kinetics Equilibrium
MgO was found to take up F− from NaF aqueous solutions due to the electrostatic attractive force between the positively charged MgO and F−. The F− adsorption by MgO can be represented by pseudo first-order reaction kinetics, and the magnitude of the apparent activation energy (71.6 kJ mol−1) confirms that this is a chemisorption process. The thermodynamic behavior of this process follows Langmuir-type adsorption, with the maximum adsorption amount of 5.6 mmol g–1. The F− can be desorbed from MgO using NaOH solution. The regenerated MgO can still take up F− from the solution despite the lowered adsorption capacity. Therefore, it is possible to recycle the MgO for F− adsorption.
1. Introduction F−-containing wastewater originates from the electronics industry, the glass industry, and from etching processes in general. The effluent standard in Japan for F− is 8 mg L–1. The primary treatment involves capturing the F− ions in the form of slightly soluble CaF2, by adding calcium salts such as Ca(OH)2 and CaCl2 to the wastewater. In a second step, aluminum salts, such as polyaluminum chloride, are added to the wastewater, producing gelatinous aluminum hydroxide to capture the remaining F− ions. However, this two-step process is very cumbersome, and a single-step treatment for F−-containing wastewater is therefore in keen demand. We have examined several new treatment methods for F−-containing wastewater using adsorbents. For example, we studied the possibility of recycling Mg − Al layered double hydroxide (Mg − Al LDH) and its calcination product (Mg − Al oxide) for F− treatment [1,2], and tested them in removing F− in real wastewater [3]. Furthermore, we found that MgO could remove F− in aqueous solutions [4]. In several subsequent studies, Li et al. examined F− removal by porous hollow MgO microspheres [5], and Jin et al. investigated the effective removal of F− by porous MgO nanoplates and the adsorption mechanism [6]. Recently, Lee et al. examined the synthesis of pillarand microsphere-like MgO particles, and their F− adsorption performance in aqueous solutions [7]. However, the desorption and recycling of these materials after they are used for F− removal have not been considered. In this study, we tested the reusability of MgO for removing aqueous F−, as shown in Fig. 1. The MgO adsorbed with F− is then treated with an aqueous solution of NaOH. After the desorption of F−, the MgO is
⁎
Corresponding author. E-mail address: [email protected] (T. Kameda).
https://doi.org/10.1016/j.jwpe.2018.06.009 Received 26 March 2018; Received in revised form 15 June 2018; Accepted 26 June 2018 2214-7144/ © 2018 Elsevier Ltd. All rights reserved.
regenerated and reused. We systematically examined the kinetic and thermodynamic aspects of F− removal by MgO. The effects of the amount of MgO and the temperature were investigated. Finally, the effect of ionic strength on the adsorption of F− by MgO was examined, which helps to reveal the adsorption mechanism. 2. Experimental All reagents were of chemical reagent grade and used without further purification. MgO is purchased from KANTO CHEMICAL CO., INC.. The average particle size is 28.5 μm, and the specific surface area is 4.8 m2/g. 2.1. Removal of F− from aqueous solution NaF solutions were prepared by dissolving NaF in deionized water. For the kinetic measurements, MgO was added to 100 mg L–1 NaF solution (500 mL) without initial pH control, and the resultant suspension was stirred at 10, 30, and 60 °C for 100 h. Samples were withdrawn from the suspension at different time intervals and immediately filtered, and then the filtrates were analyzed for residual F−. To study the thermodynamics of the adsorption process, MgO was added to NaF solutions with the molar ratio of Mg/F = 1, 5, 10, and 20. To determine the adsorption isotherm, 20 mL of a NaF solution (0.005–0.1 mol L–1) and 0.2 g of MgO were placed in a 50-mL screw-top tube and shaken at 30 °C for 1 week.
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Fig. 1. Scheme of the proposed method to remove aqueous F− by MgO. . Fig. 2. Change in the F− concentration over time in the MgO suspension in NaF solution at various Mg/F molar ratios at 30 °C.
2.2. Desorption of F− from MgO The desorption of F− from MgO was carried out using NaOH solution. After a MgO suspension in NaF (Mg/F = 10) was kept at 30 °C for 48 h, the MgO was found to contain 0.8 wt% of F−. This MgO adsorbed with F− (0.1 g) and 20 mL of NaOH solution (1.0 M) were placed in a 50-mL screw-top tube and shaken at 30 °C for 24 h.
XRD patterns for the solids from MgO suspension in NaF at Mg/F = 10 at 30 at 30 °C for 0–48 h. After 12 h, the XRD peak ascribed to Mg(OH)2 was hardly observed; but its intensity increased with time, suggesting hydration on the surface of MgO particles, which caused the dissolution of Mg2+. The XRD patterns (Fig. S3) did not show the peak ascribed to the products composed of Mg and F−. The zero point of charge (ZPC) of MgO is known to be 12.4 [8]. Since the pH values in Fig. S1 are all lower than this, the surface of MgO particles should be positively charged. Therefore, the uptake of F− is attributed to the electrostatic attractive force between the positively charged MgO and F−, implying chemical adsorption. Fig. 3 shows the change in F− removal at Mg/ F = 10 at various temperatures. For all temperatures, the F− adsorption increased with time, particularly at 60 °C. At any time, the F− adsorption also increased with temperature. The kinetics of F− removal by MgO was examined by first-order kinetics, which depend on the concentration of F− as
2.3. Removal of F− from aqueous solution by regenerated MgO MgO regenerated from the MgO adsorbed with F− in 1.0 M NaOH solution at 30 °C for 12 h was suspended in 100 mg L–1 NaF solution at Mg/F = 10 and 30 °C. For comparison, Mg(OH)2 was suspended in 100 mg L–1 NaF solution at the same Mg/F molar ratio and temperature. 2.4. Effects of ionic strength on adsorption The effect of ionic strength on the adsorption of F− by MgO was examined in 100 mg L–1 NaF solutions prepared using 0, 0.001, or 0.01 M NaCl solution with initial pH of 2 − 12. MgO was suspended in this solution at Mg/F = 5 and 30 °C.
–ln (1–x ) = kt
(1)
where x is the degree of F− adsorption, t (h) is the reaction time, and k (h−1) is the rate constant. In the plots in Fig. 4 (Mg/F = 10), good linearity was obtained at each temperature, indicating that F− adsorption can be represented by pseudo first-order reaction kinetics. The respective apparent rate constants at 10, 30, and 60 °C were 8.0 × 10−3, 7.1 × 10−2, and 7.8 × 10−1 h−1, clearly increasing with increasing temperature. Fig. S4 shows the corresponding Arrhenius plot of the apparent rate constant, yielding an apparent activation energy of 71.6 kJ mol−1. This value confirms that the uptake of F− by MgO proceeded under chemical reaction control. Fig. 5 shows the adsorption isotherm, in which the equilibrium adsorption amount increased with increasing equilibrium concentration. These isotherms showed Langmuir-type behavior, which was confirmed by fitting to the Langmuir equation expressed as
2.5. Analytical methods MgO before and after F− adsorption was analyzed using X-ray diffraction (XRD) measurements with Cu Kα radiation. For the adsorption experiments, the residual F− and Mg2+ dissolved from MgO in the filtrates were separated using a Dionex DX-120 ion chromatograph and a Dionex model AS-12 A column (eluent: 2.7 mM Na2CO3 and 0.3 mM NaHCO3; flow rate: 1.3 mL min−1). Their concentrations were measured by inductively coupled plasma-atomic emission spectrometry (ICP-AES). The pH after the adsorption was also measured. 3. Results and discussion 3.1. Removal of F− from aqueous solution Fig. 2 shows change in the F− concentration over time in various MgO suspensions in NaF at 30 °C. For all Mg/F molar ratios, the concentration of F− decreased with time, and the concentration after a given time decreased with increasing Mg/F ratio. When Mg/F = 20, the concentration of F− after 24-h treatment was less than the effluent standard in Japan (8 mg L–1). Therefore, MgO was confirmed to effectively take up F− from NaF solutions. Fig. S1 shows the change in the pH over time in the various suspension at 30 °C. For all Mg/F molar ratios, the pH value first increased rapidly and then remained constant at around 11. This is attributed to the buffer action of Mg2+. Fig. S2 shows that at all Mg/F ratios and 30 °C, the Mg2+ concentration in the solution increased rapidly at first and then decreased with time. The maximum amount of Mg2+ dissolved was around 3% for Mg/F = 10, and this low value indicates that the dissolution of Mg2+ from MgO had little effect on the uptake of F− from the solution. Fig. S3 shows the
Fig. 3. Change in the F− removal over time in the MgO suspension in NaF solution at Mg/F = 10 and various temperatures. 55
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Fig. 6. Change in the F− desorption over time from MgO adsorbed with F− by using a NaOH solution. MgO quantity: 0.1 g; temperature: 30 °C; NaOH solution: 1.0 M.
Fig. 4. Pseudo first-order plot for F− removal by MgO in NaF solution at Mg/ F = 10 and various temperatures.
Fig. 7. Change in the concentration of F− in NaF solution over time for MgO before and after regeneration. Mg/F molar ratio: 10; temperature: 30 °C. Fig. 5. Adsorption isotherm for the adsorption of F− by MgO. MgO quantity: 0.2 g; temperature: 30 °C; time: 1 week.
qe = Ce qm KL/(1 + Ce KL)
3.3. Removal of F− from aqueous solution by the MgO after regeneration Fig. 7 shows F− removal over time using MgO before or after regeneration, at Mg/F = 10 and 30 °C. For the regenerated MgO, the concentration of F− in the solution decreased gradually with time. After 48 h, the concentrations of F− were 5.8 and 30.9 mg L–1 for MgO before and after regeneration, respectively. Hence, the regenerated MgO had lower ability for F− adsorption than the original MgO, due to the progressive hydration of the surface of MgO particles into Mg(OH)2. For reference, Fig. S6 compares the F− concentration over time in NaF suspensions with MgO or Mg(OH)2 at Mg/F = 10 and 30 °C. The results indicate that Mg(OH)2 could only take up very little F− from the solution. In other words, non-hydrolyzed MgO surface is required for F− adsorption. Nevertheless, there was non-negligible uptake of F− by the regenerated MgO, indicating that it is possible to recycle the MgO for F− adsorption, as shown in Fig. 1. By the way, the adsorption rate for regenerated MgO was faster than the original MgO while in the initial adsorption (0–10 h), as shown in Fig. 7. This behavior is similar to the adsorption behavior for Mg(OH)2, as shown in Fig.S6. Fig.S6 presents that the adsorption rate for Mg (OH)2 was faster than MgO while in the initial adsorption (0–4 h). Therefore, the above phenomenon is attributed to the hydration of the surface of MgO into Mg(OH)2 for regenerated MgO.
(2)
where qe (mmol g–1) is the equilibrium adsorption, Ce (mmol L–1) is the equilibrium concentration, qm (mmol g–1) is the maximum adsorption, and KL is the equilibrium adsorption constant. This equation can also be expressed as
Ce / qe = 1/ qm KL + Ce / qm
(3)
Fig. S5 plots Ce/qe versus Ce for the adsorption isotherms. Good linearity was obtained, indicating that this process followed Langmuirtype adsorption. This also suggests that the adsorption of F− by MgO is chemical adsorption. The value of qm, determined from the slope of the straight line in Fig. S5, was 5.6 mmol g–1. This value (i.e. capacity for F− removal) is much higher than our previously reported maximum adsorption for related adsorbents: 3.3 mmol g–1 for NO3•Mg − Al LDH and 3.2 mmol g–1 for Cl•Mg − Al LDH [1], and 3.0 mmol g−1 for Mg–Al oxide (Mg/Al = 2) [2].
3.2. Desorption of F− from MgO Fig. 6 shows the F− desorption from MgO in NaOH solution over time. The amount of desorbed F− first increased rapidly with time and then remained constant. After 24 h, 11.8% of the adsorbed F− was desorbed. The pH value during the reaction (13.0–14.5) is higher than the ZPC of MgO, suggesting that the surface of MgO particles is now negatively charged, and the resulting electrostatic repulsion caused the F− to desorb from MgO.
3.4. Adsorption mechanism Figs. 8 and 9 show the change in the F− adsorption and the final pH as a function of initial pH in MgO suspension in NaF solution (100 mg/ L) with various NaCl concentrations at Mg/F = 5 and 30 °C. Upon decreasing the initial pH, the F− adsorption increased for all NaCl concentrations, and the final pH also decreased. Since all the final pH 56
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pseudo first-order reaction kinetics. The apparent rate constants at 10, 30, and 60 °C were 8.0 × 10−3, 7.1 × 10−2, and 7.8 × 10−1 h−1, respectively, and the apparent activation energy was 71.6 kJ mol−1. This value confirms that the uptake of F− by MgO is chemical adsorption. This process was found to follow a Langmuir-type adsorption, with the maximum adsorption being 5.6 mmol g–1. The adsorbed F− could be desorbed from the MgO using NaOH solution, by creating electrostatic repulsion between the F− and negatively charged MgO. While the regenerated MgO showed a lower ability for F− adsorption (possibly due to the progressive hydration of the MgO surface to form Mg(OH)2), it could still take up an appreciable amount of F-, indicating that MgO can be reused as an adsorbent to remove F-. Finally, the observed dependence of adsorption on the ionic strength implies that the F− adsorption by MgO is driven by the electrostatic attractive force between the positively charged MgO and F−, instead of the direct combination of F− onto the surface of MgO.
Fig. 8. Change in F− adsorption as a function of initial pH in MgO suspension in 100 mg/L NaF solution with various NaCl concentrations. Mg/F molar ratio: 5; temperature: 30 °C.
Acknowledgments This research was supported by the Environment Research and Technology Development Fund (5RFb-1201) of the Ministry of Environment, Japan. Appendix A. Supplementary data Supplementary material related to this article can be found, in the online version, at doi:https://doi.org/10.1016/j.jwpe.2018.06.009. References [1] T. Kameda, J. Oba, T. Yoshioka, Recyclable Mg–Al layered double hydroxides for fluoride removal: kinetic and equilibrium studies, J. Hazard. Mater. 300 (2015) 475–482. [2] T. Kameda, J. Oba, T. Yoshioka, Kinetics and equilibrium studies on Mg–Al oxide for removal of fluoride in aqueous solution and its use in recycling, J. Environ. Manage. 156 (2015) 252–256. [3] T. Kameda, J. Oba, T. Yoshioka, Removal of boron and fluoride in wastewater using Mg−Al layered double hydroxide and Mg-Al oxide, J. Environ. Manag. 188 (2017) 58–63. [4] J. Ishikawa, T. Kameda, Y. Umetsu, Proc. MMIJ Annual Meeting Vol.(II (2005), pp. 121–122. [5] L.-X. Li, D. Xu, X.-Q. Li, W.-C. Liu, Y. Jia, Excellent fluoride removal properties of porous hollow MgO microspheres, New J. Chem. 38 (2014) 5445–5452. [6] Z. Jin, Y. Jia, K.-S. Zhang, L.-T. Kong, B. Sun, W. Shen, F.-L. Meng, J.-H. Liu, Effective removal of fluoride by porous MgO nanoplates and its adsorption mechanism, J. Alloy Compd. 675 (2016) 292–300. [7] S.G. Lee, J.-W. Ha, E.-H. Sohn, I.J. Park, S.-B. Lee, Synthesis of pillar and microsphere-like magnesium oxide particles and their fluoride adsorption performance in aqueous solutions, Korean J. Chem. Eng. 34 (2017) 2738–2747. [8] G.A. Parks, The isoelectric points of solid oxides, solid hydroxides, and aqueous hydroxo complex systems, Chem. Rev. 65 (1965) 177–198. [9] K.F. Hayes, G. Redden, W. Ela, J.O. Leckie, Surface complexation models: an evaluation of model parameter estimation using FITEQL and oxide mineral titration data, J. Colloid Interface Sci. 142 (1991) 448–469. [10] S. Goldberg, H.S. Forster, E.L. Heick, Boron adsorption mechanisms on oxides, clay minerals, and soils inferred from ionic strength effects, Soil Sci. Soc. Am. J. 57 (1993) 704–708. [11] S. Goldberg, L.J. Criscenti, D.R. Turner, J.A. Davis, K.J. Cantrell, Adsorption-desorption processes in subsurface reactive transport modeling, Vadose Zone J. 6 (2007) 407–425. [12] S. Goldberg, Application of surface complexation models to anion adsorption by natural materials, Environ. Toxicol. Chem. 33 (2014) 2172–2180.
Fig. 9. Change in the final pH as a function of initial pH in MgO suspension in 100 mg/L NaF.
values were lower than the ZPC of MgO, the MgO surface should carry increasing positive charge as the initial pH is lowered, resulting in more efficient F− adsorption. In this case, the NaCl was used to adjust the ionic strength. When the NaCl concentration changed from 0.001 to 0.01 M, the F− adsorption decreased. Here, a higher ionic strength means more abundant Cl−, which could compete with F− for adsorption on the positively charged MgO surface. Generally, ionic strengthdependent adsorption suggests the formation of outer-sphere complexes [9–12]. Therefore, the results in Figs. 8 and 9 confirm that the F− adsorption by MgO is not due to the direct combination of F− onto the surface of MgO, but to the electrostatic attractive force between the positively charged MgO and the F-. 4. Conclusions MgO was confirmed to take up F− from NaF solutions. The pH during the reaction remains lower than the ZPC of MgO, suggesting that the surface of MgO particle is positively charged, and therefore the uptake of F− is attributed to the electrostatic attractive force between MgO and F−. When the Mg/F molar ratio was 20 and the reaction time exceeded 24 h, the final concentration of F− was less than the effluent standards in Japan (8 mg L–1). The F− adsorption can be represented by
57
Contents lists available at ScienceDirect
Journal of Water Process Engineering journal homepage: www.elsevier.com/locate/jwpe
Analysis of F− removal from aqueous solutions using MgO ⁎
T
Tomohito Kameda , Yusuke Yamamoto, Shogo Kumagai, Toshiaki Yoshioka Graduate School of Environmental Studies, Tohoku University, 6-6-07 Aoba, Aramaki, Aoba-ku, Sendai, Japan
A R T I C LE I N FO
A B S T R A C T
Keywords: MgO F−Water treatment Kinetics Equilibrium
MgO was found to take up F− from NaF aqueous solutions due to the electrostatic attractive force between the positively charged MgO and F−. The F− adsorption by MgO can be represented by pseudo first-order reaction kinetics, and the magnitude of the apparent activation energy (71.6 kJ mol−1) confirms that this is a chemisorption process. The thermodynamic behavior of this process follows Langmuir-type adsorption, with the maximum adsorption amount of 5.6 mmol g–1. The F− can be desorbed from MgO using NaOH solution. The regenerated MgO can still take up F− from the solution despite the lowered adsorption capacity. Therefore, it is possible to recycle the MgO for F− adsorption.
1. Introduction F−-containing wastewater originates from the electronics industry, the glass industry, and from etching processes in general. The effluent standard in Japan for F− is 8 mg L–1. The primary treatment involves capturing the F− ions in the form of slightly soluble CaF2, by adding calcium salts such as Ca(OH)2 and CaCl2 to the wastewater. In a second step, aluminum salts, such as polyaluminum chloride, are added to the wastewater, producing gelatinous aluminum hydroxide to capture the remaining F− ions. However, this two-step process is very cumbersome, and a single-step treatment for F−-containing wastewater is therefore in keen demand. We have examined several new treatment methods for F−-containing wastewater using adsorbents. For example, we studied the possibility of recycling Mg − Al layered double hydroxide (Mg − Al LDH) and its calcination product (Mg − Al oxide) for F− treatment [1,2], and tested them in removing F− in real wastewater [3]. Furthermore, we found that MgO could remove F− in aqueous solutions [4]. In several subsequent studies, Li et al. examined F− removal by porous hollow MgO microspheres [5], and Jin et al. investigated the effective removal of F− by porous MgO nanoplates and the adsorption mechanism [6]. Recently, Lee et al. examined the synthesis of pillarand microsphere-like MgO particles, and their F− adsorption performance in aqueous solutions [7]. However, the desorption and recycling of these materials after they are used for F− removal have not been considered. In this study, we tested the reusability of MgO for removing aqueous F−, as shown in Fig. 1. The MgO adsorbed with F− is then treated with an aqueous solution of NaOH. After the desorption of F−, the MgO is
⁎
Corresponding author. E-mail address: [email protected] (T. Kameda).
https://doi.org/10.1016/j.jwpe.2018.06.009 Received 26 March 2018; Received in revised form 15 June 2018; Accepted 26 June 2018 2214-7144/ © 2018 Elsevier Ltd. All rights reserved.
regenerated and reused. We systematically examined the kinetic and thermodynamic aspects of F− removal by MgO. The effects of the amount of MgO and the temperature were investigated. Finally, the effect of ionic strength on the adsorption of F− by MgO was examined, which helps to reveal the adsorption mechanism. 2. Experimental All reagents were of chemical reagent grade and used without further purification. MgO is purchased from KANTO CHEMICAL CO., INC.. The average particle size is 28.5 μm, and the specific surface area is 4.8 m2/g. 2.1. Removal of F− from aqueous solution NaF solutions were prepared by dissolving NaF in deionized water. For the kinetic measurements, MgO was added to 100 mg L–1 NaF solution (500 mL) without initial pH control, and the resultant suspension was stirred at 10, 30, and 60 °C for 100 h. Samples were withdrawn from the suspension at different time intervals and immediately filtered, and then the filtrates were analyzed for residual F−. To study the thermodynamics of the adsorption process, MgO was added to NaF solutions with the molar ratio of Mg/F = 1, 5, 10, and 20. To determine the adsorption isotherm, 20 mL of a NaF solution (0.005–0.1 mol L–1) and 0.2 g of MgO were placed in a 50-mL screw-top tube and shaken at 30 °C for 1 week.
Journal of Water Process Engineering 25 (2018) 54–57
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Fig. 1. Scheme of the proposed method to remove aqueous F− by MgO. . Fig. 2. Change in the F− concentration over time in the MgO suspension in NaF solution at various Mg/F molar ratios at 30 °C.
2.2. Desorption of F− from MgO The desorption of F− from MgO was carried out using NaOH solution. After a MgO suspension in NaF (Mg/F = 10) was kept at 30 °C for 48 h, the MgO was found to contain 0.8 wt% of F−. This MgO adsorbed with F− (0.1 g) and 20 mL of NaOH solution (1.0 M) were placed in a 50-mL screw-top tube and shaken at 30 °C for 24 h.
XRD patterns for the solids from MgO suspension in NaF at Mg/F = 10 at 30 at 30 °C for 0–48 h. After 12 h, the XRD peak ascribed to Mg(OH)2 was hardly observed; but its intensity increased with time, suggesting hydration on the surface of MgO particles, which caused the dissolution of Mg2+. The XRD patterns (Fig. S3) did not show the peak ascribed to the products composed of Mg and F−. The zero point of charge (ZPC) of MgO is known to be 12.4 [8]. Since the pH values in Fig. S1 are all lower than this, the surface of MgO particles should be positively charged. Therefore, the uptake of F− is attributed to the electrostatic attractive force between the positively charged MgO and F−, implying chemical adsorption. Fig. 3 shows the change in F− removal at Mg/ F = 10 at various temperatures. For all temperatures, the F− adsorption increased with time, particularly at 60 °C. At any time, the F− adsorption also increased with temperature. The kinetics of F− removal by MgO was examined by first-order kinetics, which depend on the concentration of F− as
2.3. Removal of F− from aqueous solution by regenerated MgO MgO regenerated from the MgO adsorbed with F− in 1.0 M NaOH solution at 30 °C for 12 h was suspended in 100 mg L–1 NaF solution at Mg/F = 10 and 30 °C. For comparison, Mg(OH)2 was suspended in 100 mg L–1 NaF solution at the same Mg/F molar ratio and temperature. 2.4. Effects of ionic strength on adsorption The effect of ionic strength on the adsorption of F− by MgO was examined in 100 mg L–1 NaF solutions prepared using 0, 0.001, or 0.01 M NaCl solution with initial pH of 2 − 12. MgO was suspended in this solution at Mg/F = 5 and 30 °C.
–ln (1–x ) = kt
(1)
where x is the degree of F− adsorption, t (h) is the reaction time, and k (h−1) is the rate constant. In the plots in Fig. 4 (Mg/F = 10), good linearity was obtained at each temperature, indicating that F− adsorption can be represented by pseudo first-order reaction kinetics. The respective apparent rate constants at 10, 30, and 60 °C were 8.0 × 10−3, 7.1 × 10−2, and 7.8 × 10−1 h−1, clearly increasing with increasing temperature. Fig. S4 shows the corresponding Arrhenius plot of the apparent rate constant, yielding an apparent activation energy of 71.6 kJ mol−1. This value confirms that the uptake of F− by MgO proceeded under chemical reaction control. Fig. 5 shows the adsorption isotherm, in which the equilibrium adsorption amount increased with increasing equilibrium concentration. These isotherms showed Langmuir-type behavior, which was confirmed by fitting to the Langmuir equation expressed as
2.5. Analytical methods MgO before and after F− adsorption was analyzed using X-ray diffraction (XRD) measurements with Cu Kα radiation. For the adsorption experiments, the residual F− and Mg2+ dissolved from MgO in the filtrates were separated using a Dionex DX-120 ion chromatograph and a Dionex model AS-12 A column (eluent: 2.7 mM Na2CO3 and 0.3 mM NaHCO3; flow rate: 1.3 mL min−1). Their concentrations were measured by inductively coupled plasma-atomic emission spectrometry (ICP-AES). The pH after the adsorption was also measured. 3. Results and discussion 3.1. Removal of F− from aqueous solution Fig. 2 shows change in the F− concentration over time in various MgO suspensions in NaF at 30 °C. For all Mg/F molar ratios, the concentration of F− decreased with time, and the concentration after a given time decreased with increasing Mg/F ratio. When Mg/F = 20, the concentration of F− after 24-h treatment was less than the effluent standard in Japan (8 mg L–1). Therefore, MgO was confirmed to effectively take up F− from NaF solutions. Fig. S1 shows the change in the pH over time in the various suspension at 30 °C. For all Mg/F molar ratios, the pH value first increased rapidly and then remained constant at around 11. This is attributed to the buffer action of Mg2+. Fig. S2 shows that at all Mg/F ratios and 30 °C, the Mg2+ concentration in the solution increased rapidly at first and then decreased with time. The maximum amount of Mg2+ dissolved was around 3% for Mg/F = 10, and this low value indicates that the dissolution of Mg2+ from MgO had little effect on the uptake of F− from the solution. Fig. S3 shows the
Fig. 3. Change in the F− removal over time in the MgO suspension in NaF solution at Mg/F = 10 and various temperatures. 55
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Fig. 6. Change in the F− desorption over time from MgO adsorbed with F− by using a NaOH solution. MgO quantity: 0.1 g; temperature: 30 °C; NaOH solution: 1.0 M.
Fig. 4. Pseudo first-order plot for F− removal by MgO in NaF solution at Mg/ F = 10 and various temperatures.
Fig. 7. Change in the concentration of F− in NaF solution over time for MgO before and after regeneration. Mg/F molar ratio: 10; temperature: 30 °C. Fig. 5. Adsorption isotherm for the adsorption of F− by MgO. MgO quantity: 0.2 g; temperature: 30 °C; time: 1 week.
qe = Ce qm KL/(1 + Ce KL)
3.3. Removal of F− from aqueous solution by the MgO after regeneration Fig. 7 shows F− removal over time using MgO before or after regeneration, at Mg/F = 10 and 30 °C. For the regenerated MgO, the concentration of F− in the solution decreased gradually with time. After 48 h, the concentrations of F− were 5.8 and 30.9 mg L–1 for MgO before and after regeneration, respectively. Hence, the regenerated MgO had lower ability for F− adsorption than the original MgO, due to the progressive hydration of the surface of MgO particles into Mg(OH)2. For reference, Fig. S6 compares the F− concentration over time in NaF suspensions with MgO or Mg(OH)2 at Mg/F = 10 and 30 °C. The results indicate that Mg(OH)2 could only take up very little F− from the solution. In other words, non-hydrolyzed MgO surface is required for F− adsorption. Nevertheless, there was non-negligible uptake of F− by the regenerated MgO, indicating that it is possible to recycle the MgO for F− adsorption, as shown in Fig. 1. By the way, the adsorption rate for regenerated MgO was faster than the original MgO while in the initial adsorption (0–10 h), as shown in Fig. 7. This behavior is similar to the adsorption behavior for Mg(OH)2, as shown in Fig.S6. Fig.S6 presents that the adsorption rate for Mg (OH)2 was faster than MgO while in the initial adsorption (0–4 h). Therefore, the above phenomenon is attributed to the hydration of the surface of MgO into Mg(OH)2 for regenerated MgO.
(2)
where qe (mmol g–1) is the equilibrium adsorption, Ce (mmol L–1) is the equilibrium concentration, qm (mmol g–1) is the maximum adsorption, and KL is the equilibrium adsorption constant. This equation can also be expressed as
Ce / qe = 1/ qm KL + Ce / qm
(3)
Fig. S5 plots Ce/qe versus Ce for the adsorption isotherms. Good linearity was obtained, indicating that this process followed Langmuirtype adsorption. This also suggests that the adsorption of F− by MgO is chemical adsorption. The value of qm, determined from the slope of the straight line in Fig. S5, was 5.6 mmol g–1. This value (i.e. capacity for F− removal) is much higher than our previously reported maximum adsorption for related adsorbents: 3.3 mmol g–1 for NO3•Mg − Al LDH and 3.2 mmol g–1 for Cl•Mg − Al LDH [1], and 3.0 mmol g−1 for Mg–Al oxide (Mg/Al = 2) [2].
3.2. Desorption of F− from MgO Fig. 6 shows the F− desorption from MgO in NaOH solution over time. The amount of desorbed F− first increased rapidly with time and then remained constant. After 24 h, 11.8% of the adsorbed F− was desorbed. The pH value during the reaction (13.0–14.5) is higher than the ZPC of MgO, suggesting that the surface of MgO particles is now negatively charged, and the resulting electrostatic repulsion caused the F− to desorb from MgO.
3.4. Adsorption mechanism Figs. 8 and 9 show the change in the F− adsorption and the final pH as a function of initial pH in MgO suspension in NaF solution (100 mg/ L) with various NaCl concentrations at Mg/F = 5 and 30 °C. Upon decreasing the initial pH, the F− adsorption increased for all NaCl concentrations, and the final pH also decreased. Since all the final pH 56
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pseudo first-order reaction kinetics. The apparent rate constants at 10, 30, and 60 °C were 8.0 × 10−3, 7.1 × 10−2, and 7.8 × 10−1 h−1, respectively, and the apparent activation energy was 71.6 kJ mol−1. This value confirms that the uptake of F− by MgO is chemical adsorption. This process was found to follow a Langmuir-type adsorption, with the maximum adsorption being 5.6 mmol g–1. The adsorbed F− could be desorbed from the MgO using NaOH solution, by creating electrostatic repulsion between the F− and negatively charged MgO. While the regenerated MgO showed a lower ability for F− adsorption (possibly due to the progressive hydration of the MgO surface to form Mg(OH)2), it could still take up an appreciable amount of F-, indicating that MgO can be reused as an adsorbent to remove F-. Finally, the observed dependence of adsorption on the ionic strength implies that the F− adsorption by MgO is driven by the electrostatic attractive force between the positively charged MgO and F−, instead of the direct combination of F− onto the surface of MgO.
Fig. 8. Change in F− adsorption as a function of initial pH in MgO suspension in 100 mg/L NaF solution with various NaCl concentrations. Mg/F molar ratio: 5; temperature: 30 °C.
Acknowledgments This research was supported by the Environment Research and Technology Development Fund (5RFb-1201) of the Ministry of Environment, Japan. Appendix A. Supplementary data Supplementary material related to this article can be found, in the online version, at doi:https://doi.org/10.1016/j.jwpe.2018.06.009. References [1] T. Kameda, J. Oba, T. Yoshioka, Recyclable Mg–Al layered double hydroxides for fluoride removal: kinetic and equilibrium studies, J. Hazard. Mater. 300 (2015) 475–482. [2] T. Kameda, J. Oba, T. Yoshioka, Kinetics and equilibrium studies on Mg–Al oxide for removal of fluoride in aqueous solution and its use in recycling, J. Environ. Manage. 156 (2015) 252–256. [3] T. Kameda, J. Oba, T. Yoshioka, Removal of boron and fluoride in wastewater using Mg−Al layered double hydroxide and Mg-Al oxide, J. Environ. Manag. 188 (2017) 58–63. [4] J. Ishikawa, T. Kameda, Y. Umetsu, Proc. MMIJ Annual Meeting Vol.(II (2005), pp. 121–122. [5] L.-X. Li, D. Xu, X.-Q. Li, W.-C. Liu, Y. Jia, Excellent fluoride removal properties of porous hollow MgO microspheres, New J. Chem. 38 (2014) 5445–5452. [6] Z. Jin, Y. Jia, K.-S. Zhang, L.-T. Kong, B. Sun, W. Shen, F.-L. Meng, J.-H. Liu, Effective removal of fluoride by porous MgO nanoplates and its adsorption mechanism, J. Alloy Compd. 675 (2016) 292–300. [7] S.G. Lee, J.-W. Ha, E.-H. Sohn, I.J. Park, S.-B. Lee, Synthesis of pillar and microsphere-like magnesium oxide particles and their fluoride adsorption performance in aqueous solutions, Korean J. Chem. Eng. 34 (2017) 2738–2747. [8] G.A. Parks, The isoelectric points of solid oxides, solid hydroxides, and aqueous hydroxo complex systems, Chem. Rev. 65 (1965) 177–198. [9] K.F. Hayes, G. Redden, W. Ela, J.O. Leckie, Surface complexation models: an evaluation of model parameter estimation using FITEQL and oxide mineral titration data, J. Colloid Interface Sci. 142 (1991) 448–469. [10] S. Goldberg, H.S. Forster, E.L. Heick, Boron adsorption mechanisms on oxides, clay minerals, and soils inferred from ionic strength effects, Soil Sci. Soc. Am. J. 57 (1993) 704–708. [11] S. Goldberg, L.J. Criscenti, D.R. Turner, J.A. Davis, K.J. Cantrell, Adsorption-desorption processes in subsurface reactive transport modeling, Vadose Zone J. 6 (2007) 407–425. [12] S. Goldberg, Application of surface complexation models to anion adsorption by natural materials, Environ. Toxicol. Chem. 33 (2014) 2172–2180.
Fig. 9. Change in the final pH as a function of initial pH in MgO suspension in 100 mg/L NaF.
values were lower than the ZPC of MgO, the MgO surface should carry increasing positive charge as the initial pH is lowered, resulting in more efficient F− adsorption. In this case, the NaCl was used to adjust the ionic strength. When the NaCl concentration changed from 0.001 to 0.01 M, the F− adsorption decreased. Here, a higher ionic strength means more abundant Cl−, which could compete with F− for adsorption on the positively charged MgO surface. Generally, ionic strengthdependent adsorption suggests the formation of outer-sphere complexes [9–12]. Therefore, the results in Figs. 8 and 9 confirm that the F− adsorption by MgO is not due to the direct combination of F− onto the surface of MgO, but to the electrostatic attractive force between the positively charged MgO and the F-. 4. Conclusions MgO was confirmed to take up F− from NaF solutions. The pH during the reaction remains lower than the ZPC of MgO, suggesting that the surface of MgO particle is positively charged, and therefore the uptake of F− is attributed to the electrostatic attractive force between MgO and F−. When the Mg/F molar ratio was 20 and the reaction time exceeded 24 h, the final concentration of F− was less than the effluent standards in Japan (8 mg L–1). The F− adsorption can be represented by
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