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Enhanced adsorptive removal of fluoride using mesoporous alumina
Microporous and Mesoporous Materials 127 (2010) 152–156
Contents lists available at ScienceDirect
Microporous and Mesoporous Materials journal homepage: www.elsevier.com/locate/micromeso
Enhanced adsorptive removal of fluoride using mesoporous alumina Grace Lee, Chao Chen, Seung-Tae Yang, Wha-Seung Ahn * Department of Chemical Engineering, Inha University, Incheon 402-751, Republic of Korea
a r t i c l e
i n f o
Article history: Received 29 October 2008 Received in revised form 1 July 2009 Accepted 4 July 2009 Available online 10 July 2009 Keywords: Mesoporous alumina Aluminum tri-sec-butoxide Cetyltrimethylammonium bromide Stearic acid Fluoride adsorption
a b s t r a c t Two different kinds of mesoporous alumina samples were prepared using aluminum tri-sec-butoxide in the presence of either cetyltrimethylammonium bromide (MA-1) or stearic acid (MA-2) as a structuredirecting agent, and tested for adsorptive removal of fluoride in water. Both materials contain a wormhole-like mesopore structure, but exhibited different textural properties: surface area (421 or 650 m2/g) and pore volume (0.96 or 0.59 cm3/g). These mesoporous aluminas demonstrated significantly improved adsorption capacity and faster kinetics to those of the commercial activated aluminas in fluoride removal by adsorption process. MA-2 prepared using stearic acid, in particular, demonstrated an adsorption capacity (14.26 mg/g) and initial adsorption rate (14.6 mg/g min) that were respectively 2.2 and 45 times higher than those of a commercial gamma alumina. The textural features of larger surface area and relatively smaller pore size in MA-2 compared to the activated aluminas are believed to be responsible for this enhancement in adsorption process. Ó 2009 Elsevier Inc. All rights reserved.
1. Introduction The presence of excess fluoride ions in water causes serious problems such as dental and skeletal fluorosis, and World Health Organization (WHO) had set a limiting value of fluoride ions in drinking water to be 1.5 mg l1 [1,2]. The concentration of fluoride ions in groundwater of many places in the world, however, exceeds the permitted level; increasing amount of wastewater containing fluoride is being released from various engineering processes, such as semiconductor manufacturing, coal power plants, electroplating, and rubber and fertilizer production to name a few. Therefore, the necessity to remove the excess fluoride from aquatic environment is high. Many methods, i.e. adsorption [3,4], ion exchange [5,6], precipitation with calcium and aluminum salts [4,7], dialysis [2,8,9], and membrane filtration [10] have been adopted to remove excessive fluoride ion from water. Among these, adsorption is less expensive than membrane filtration, easier and safer to handle compared to the contaminated sludge produced by precipitation, and more versatile than ion exchange [11–13]. In particular, adsorption process is generally known to be more effective and economical at low pollutant concentration levels [14]. Diverse kinds of adsorbents such as montmorillonite [15], Al2O3/CNT [16], activated alumina [17], fly ash [18], silica gel [19], bone char [20], and zeolite [21] have been reported so far for fluoride removal. Their low adsorption capacities and poor adsorption kinetics, however, still need improvement. An ideal * Corresponding author. Tel.: +82 32 866 0143; fax: +82 32 872 0959. E-mail address: [email protected] (W.-S. Ahn). 1387-1811/$ - see front matter Ó 2009 Elsevier Inc. All rights reserved. doi:10.1016/j.micromeso.2009.07.007
adsorbent should have uniformly accessible pores, high surface area, and physical and chemical stability [22]. Recently, aluminaimpregnated mesoporous silica SBA-15 [23,24] and mesoporous alumina prepared using an anionic surfactant [25] were reported to exhibit improved adsorption for several anionic species over commercial activated alumina. We also confirmed the effectiveness of mesoporous alumina in adsorptive removal of arsenate and orthophosphate anions [26]. In this work, two representative samples of mesoporous alumina materials, one previously reported by Ray et al. [27] and ˇ ejka et al. [28] were synthesized and tested for rethe other by C moval of fluoride in water. It is meaningful to measure the adsorption performances over different mesoporous alumina samples to draw a correct relationship between adsorption capacity and textural parameters of an alumina, since a range of mesoporous alumina materials prepared using different kinds of surfactant species with different textural properties have been claimed [29– 32]. An extensive review on mesoporous alumina synthesis and characterization was published recently [33]. Fluoride ion adsorption capacities and adsorption kinetics were measured and their effectiveness was critically evaluated against commercial activated alumina samples – Catapal B (Condea Vista) in pseudo-boehmite form and gamma alumina (Condea Chemie). 2. Experimental 2.1. Synthesis and characterization of mesoporous alumina samples MA-1 [27]: Aluminum tri-sec-butoxide (ASB, Aldrich) and cetyltrimethylammonium bromide (CTAB, Aldrich) were used as
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G. Lee et al. / Microporous and Mesoporous Materials 127 (2010) 152–156
2.2. Adsorption experiments 2.2.1. Equilibrium adsorption isotherms A series of flasks with 100 ml of solution having 20–250 mg/L fluoride ion concentration was prepared by dissolving NaF in RO (reverse osmosis)-deionized water. The initial pH of the solutions was adjusted to 6 with 0.01 N NaOH and HCl solutions. Mesoporous alumina (1.2 g) dried at 200 °C for 4 h prior to the adsorption experiment was then put into the flasks. The flasks were magnetically stirred at a constant speed of 200 rpm for 24 h immersed in a water bath maintained at 30 °C. Upon completion, the solutions were filtered immediately using a 0.45 lm microfilter and the liquid was separated from the residual powder by centrifugation. The concentration of fluoride in the filtrate was determined by ion chromatograph (IC, Dionex, ICS-3000). 2.2.2. Adsorption kinetics The fluoride ion adsorption kinetics was monitored by adding 12 g of alumina sample into a 1 L solution. The initial fluoride ion concentration in the solution was 150 mg/L, and the solution pH was adjusted to 6.0 using acidic and alkaline solutions. The sample was agitated at 200 rpm with a magnetic bar for 8 h at 30 °C. A portion of solution was taken at predetermined time intervals and centrifuged followed by filtration. Total sampling volumes were kept below 5% of the total solution volume [17]. The fluoride concentration of the filtered solutions was again analyzed by ion chromatography. 3. Results and discussion 3.1. Characterization of mesoporous alumina and activated alumina Powder X-ray diffraction (XRD) patterns of the calcined mesoporous alumina samples designated as MA-1 and MA-2 are shown in Fig. 1a. Both MA-1 and MA-2 exhibited only single XRD peak at 2h values close to 0.9 and 1.5, respectively, which indicates a disordered mesoporous structure without a long range order, but with a fairly uniform pore size; it is worth mentioning that the single low angle XRD peak in mesoporous alumina does not necessarily
(a)
MA-2
Intensity (a.u.)
as- syn
calcined
MA-1 20
MA-2
2
40
4
60
6
80
8
Two theta (degree)
(b) Boehmite
Intensity (a.u.)
an aluminum source and a cationic structure-directing agent, respectively. 2-Butanol (Aldrich) was employed as an organic solvent. CTAB (4.38 g) was dissolved first in a solution of 10.18 g ASB and 40.0 g of 2-butanol with stirring. To this mixture, a solution of 1.44 g of distilled water in 5.0 g of 2-butanol was then added drop wise with stirring to make a molar composition of the solution mixture in Al:surfactant:water:2-butanol = 1:0.3:2: 15. This solution gradually turned into a gel, which was aged at 100 °C for 24 h in a teflon-lined autoclave at static conditions under autogenous pressure. The resulting product was filtered, extensively washed with 99% absolute ethanol, and dried at 100 °C for 5 h. Finally, the dried sample was calcined in air at 500 °C for 4 h. MA-2 [28]: 5.1 g of stearic acid (Aldrich) was dissolved in 100 ml of 1-propanol (Fluka). Water (3.1 ml) was then added to this solution and stirred for about 30 min. After addition of 13.7 g of ASB, the mixture was stirred for additional 20 min. The gel prepared was heated at 100 °C for 50 h under static conditions. The resulting product was filtered, washed with 99% absolute ethanol, and dried at 50 °C overnight. The powder was subsequently calcined at 410 °C in nitrogen and then at 420 °C in air. X-ray diffraction patterns of the products were obtained with a Rigaku DMAX 2500 using Ni-filtered CuKa radiation. The BET surface area, pore volume, and mean pore diameters were measured by N2 physisorption at 196 °C using a Micromeritics ASAP 2000 automatic analyzer. The morphology of the samples was examined by TEM (Philips, CM 200).
γ -Al2O3 20
40
60
80
γ -Al2O3 Boehmite 2
4
6
8
Two theta (degree) Fig. 1. XRD patterns of alumina samples: (a) mesoporous alumina and (b) commercial activated alumina (inset: corresponding wide angle reflections).
prove well-ordered mesoporous material with uniform pores [34]. Low angle peak remained intact after calcinations at 500 °C, which implies the structural stability of the materials during and after the removal of surfactant. As shown by the wide angle XRD patterns of MA-2 shown in the inset of the Fig. 1a, pseudo-boehmite phase of the as-synthesized MA-2 transformed into gamma phase after calcination as reported by Pinnavaia et al. [32]. XRD patterns of MA-1 in the wide angle region also exhibited the same structural transformation upon calcination (not shown). Activated alumina, on the other hand, did not show any XRD peaks in the small-angle region as shown in Fig. 1b, but exhibited diffraction lines in the wide angle region corresponding to either boehmite or gamma alumina phase depending on the sample. Fig. 2 shows TEM images of mesoporous alumina samples used in this work. Both MA-1 and MA-2 show a typical wormhole-like morphology known for mesoporous aluminas without any distinguishable features. Fig. 3 presents nitrogen adsorption–desorption isotherms and the corresponding pore size distributions of alumina samples. All the isotherms were of the type IV with hysteresis loops in different sizes. The isotherms on mesoporous alumina samples MA-1 and MA-2 are characterized by a relatively steep region in the desorption branch of the isotherms at P/P0 = 0.6–0.8 and 0.3–0.5, respectively, reflecting the capillary condensation in uniform mesopores. As shown by the pore size distribution in the inset of Fig. 3, MA-1 has significantly larger pore diameter and pore volume than MA-2. Textural properties of the corresponding mesoporous alumina and the commercial alumina are presented in Table 1. While the
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Fig. 2. TEM images of mesoporous alumina samples: (a) MA-1 and (b) MA-2.
600
dV/dlogD
Vol. adsorbed (cc/g, STP)
800
400
10
100
Pore diameter (A)
200
0 0.0
0.2
0.4
0.6
0.8
1.0
Relative pressure (P/P0) Fig. 3. N2 adsorption–desorption isotherms of alumina samples (inset: pore size distributions of the corresponding samples from adsorption branch).
Table 1 Textural properties of mesoporous and activated alumina samples. Adsorbent
BET surface area (m2/g)
Pore volume (cc/g)
Average BJH pore diameter (nm)
MA-1 MA-2 Boehmite Gamma alumina
421 650 253 220
0.96 0.59 0.38 0.50
6.3 2.9 5.1 6.6
which a solid submerged in an electrolyte exhibits zero net electrical charge on the surface. Typically ZPC values for commercial aluminas fall in a range of pH 8–10, depending on the grades [35]. The alumina surface is charged positively at pH below the ZPC whereas the surface is negatively charged at pH values above the ZPC. Accordingly, at pH below the ZPC, the adsorption sites on the surface of an activated alumina will adsorb anionic species. In accordance with this, the optimum pH in removal of fluoride by activated alumina was reported to be in the range of 5–7 [17]. In this work, pH of the fluoride solution was fixed to 6 for all adsorption experiments to make a condition close to that of drinking water. Fig. 4 shows the equilibrium fluoride ion adsorption data over alumina samples obtained at 30 °C at initial fluoride concentrations ranging from 20 to 250 mg/L. All the fluoride adsorption data were found well-fitted with the Langmuir isotherm model, which has been successfully applied to many other real sorption processes. Langmuir adsorption isotherm can be given as follows:
qe ¼
qm bC e 1 þ bC e
where qe is the amount of fluoride adsorbed per unit weight of adsorbents at equilibrium (mg/g), qm is the maximum capacity of the adsorbent (mg/g), C e is the equilibrium concentration of fluoride (mg/L), and b is the Langmuir constant related to the energy of adsorption. The maximum adsorption capacities for the MA-1 and MA-2 adsorbents were estimated to be 7.51 and 14.26 mg/g, respectively, while the corresponding values for boehmite and gamma alumina were 6.13 mg/g and 6.36 mg/g, respectively. The surface
16 14
Fluoride uptake (mg/g)
average pore diameter of mesoporous alumina MA-1 is close to those by activated aluminas, it has a surface area and pore volume of 421 m2/g and 0.96 cm3/g, respectively, which are larger than those of commercial gamma alumina by a factor close to 2.0. MA-2, on the other hand, has a significantly smaller pore diameter (2.9 nm) than others (5.1–6.6 nm), but its surface area is the biggest among them (650 m2/g), which is almost 3 times larger than that of commercial gamma alumina.
12 10 8 6 4
3.2. Batch adsorption studies 2
The fluoride adsorption by adsorbents shows distinctive behaviors depending on the bonding between fluoride species and adsorption sites on the surface of an adsorbent [17]. In aqueous solution, the adsorption process is primarily governed by the zero point charge (ZPC) of an adsorbent. The ZPC is the pH value at
0 0
50
100
150
Equilibrium concentration (mg/L) Fig. 4. Adsorption isotherms of fluoride over various alumina samples tested.
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G. Lee et al. / Microporous and Mesoporous Materials 127 (2010) 152–156 Table 2 Kinetic parameters for fluoride ion adsorption.
10
(a)
Adsorbent
Fluoride uptake (mg/g)
8
MA-1 MA-2 Boehmite Gamma alumina
6
2
4.2 10 25.0 102 3.9 102 1.7 102
qe (mg/g)
v0 (mg/g min)
R2
6.23 7.58 5.23 4.42
1.61 14.6 1.06 0.324
0.9970 1.0000 0.9996 0.9989
4
Eqs. (2) and (3) become: 2
t 1 1 ¼ þ t qt v0 qe
0 0
50
100
150
200
250
300
200
250
300
Time (min) 100
(b) 80 -1
t/qt (mingmg )
kad (g/mg min)
60
40
20
0 0
50
100
150
Time (min) Fig. 5. (a) Adsorption kinetics of fluoride ion removal and (b) pseudo-second-order adsorption rates of fluoride ion over various alumina samples tested.
area of MA-2 is about 2.6 times of that of boehmite (650/253 m2/g), while the maximum sorption capacity is about 2.3 times of that of boehmite (14.26/6.13 mg/g). It clearly shows that surface area of the adsorbent is the major factor governing fluoride adsorption capacity in aluminas. Pore size seems to be not critical as long as the pores can accommodate fluoride ions inside freely. In fact, it can be envisaged that large surface area, relatively small mesopores, and small pore volume in an alumina will lead to more close contact between fluoride ions and adsorption sites on the surface. Fig. 5 shows the adsorption kinetics data of fluoride on different alumina samples. The initial fluoride ion concentrations were kept at 150 mg/L. The data were also fitted to a pseudo-second-order kinetic model based on surface reaction [7], which has been suggested recently as more appropriate for describing ionic-type adsorption [36,37]
dqt ¼ kad ðqe qt Þ2 dt
ð1Þ
In Eq. (1), qe is the sorption capacity at equilibrium, qt is the uptake at time t, and kad is the rate constant of pseudo-second-order adsorption (g/mgmin). By integrating Eq. (1) with the boundary conditions of qt = 0 (at t = 0) and qt = qt (at t = t), the following equation can be obtained:
t 1 1 ¼ þ t qt kad q2e qe
ð2Þ
Taking into account the initial sorption rate v0 (mg/gmin)
v0 ¼ kad q2e
ð3Þ
ð4Þ
The constants can be determined experimentally from the slope and intercept of a plot of t/qt versus t shown in Fig. 5b. As shown in Fig. 5a, MA-1 and MA-2 reached their respective adsorption equilibrium after 45 or 15 min, respectively. On the other hand, the adsorption for boehmite and gamma alumina took much longer time to reach an equilibrium; adsorption progressed smoothly for 60 (boehmite) to 100 (gamma alumina) min, respectively, and then leveled off more slowly as the adsorption approaches respective equilibrium. The kinetic parameters – the pseudo-second-order rate constants (kad) and the initial sorption rates (v0) of the adsorbent materials acquired from fitting results are summarized in Table 2. The kinetics data of fluoride could be successfully simulated by the pseudo-second-order rate equation with a correlation coefficient higher than 0.99. These high correlation coefficient (R2) values for a pseudo-second-order equation for all the samples implied that the adsorption of fluoride on those samples was chemisorption in nature [36]. While the rate constant (kad) of MA-2 was 25 102 g/mg min, that of gamma alumina was the smallest as 1.7 102 g/mg min. Also, the sorption capacity at equilibrium (qe) of MA-2 was ca. 1.7 times higher than that of gamma alumina. As mentioned earlier, this difference in sorption capacity is due to the differences in surface areas among the adsorbents, but other factors such as differences in ZPC values among the samples could also influence the sorption behavior as well. The initial sorption rate (v0) is represented by a multiple of kad and q2e , and MA-2 had a v0 of 14.6 mg/g min, which is significantly faster than that of gamma alumina. It is noteworthy in that the sorption capacity values at equilibrium (qe) of mesoporous aluminas and activated aluminas estimated by the model were highly consistent with the corresponding fluoride uptake values shown in the adsorption isotherm plots shown in Fig. 4; equilibrium concentrations in solution measured were 51 (MA-2), 72 (MA-1), 95 (boehmite), and 101 mg/L (gamma alumina), respectively. These adsorption kinetic experimental results indicate that the mesoporous alumina MA-2 with the highest surface area and narrow pore size has the fastest response time for fluoride adsorption. 4. Conclusions Two kinds of mesoporous alumina were prepared using an aluminum alkoxide precursor via either cationic or anionic surfactantmediated synthesis. The effectiveness of these materials for removal of fluoride ions in aqueous solution was evaluated in batch adsorption experiments by measuring adsorption capacities and kinetic parameters. Measured equilibrium adsorption data were fitted to the Langmuir model, and kinetics data were fitted to a pseudo-second-order model with excellent agreement. Mesoporous aluminas demonstrated superior adsorption performances to commercial boehmite or gamma alumina in both sorption capacity and initial sorption rates. The mesoporous alumina prepared using anionic surfactant (MA-2), in particular, demonstrated an adsorption capacity and initial adsorption rate that were significantly
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higher than those of a commercial gamma alumina, mainly attributed to its large surface area. Acknowledgment This study was sponsored by the Korea Energy Management Corporation (KEMCO) through the Energy & Resources Development Program (2007–2008). References [1] K.M. Popat, P.S. Anand, B.D. Dasare, React. Polym. 23 (1994) 23. [2] M. Hichour, F. Persin, J. Sandeaux, C. Gavach, Sep. Purif. Technol. 18 (2000) 1. [3] R. Weerasooriya, H.U.S. Wickramarathne, H.A. Dharmagunawardhane, Coll. Surf. A: Physicochem. Eng. Aspects 144 (1998) 267. [4] H. Lounici, L. Addour, D. Belhocine, H. Grib, S. Nicolas, B. Bariou, N. Mameri, Desalination 114 (1997) 241. [5] C. Castel, M. Schweizer, M.O. Simonnot, M. Sardin, Chem. Eng. Sci. 55 (2000) 3341. [6] E.J. Reardon, Y. Wong, Environ. Sci. Technol. 15 (2000) 3247. [7] P.L. Bishop, G. Sansoucy, J. AWWA 70 (1978) 554. [8] M. Hichour, F. Persin, J. Molenat, J. Sandeaux, C. Gavach, Desalination 122 (1999) 53. [9] Z. Amor, B. Bariou, N. Mameri, M. Taky, S. Nicolas, A. Elmidaoui, Desalination 133 (2001) 215. [10] K. Hu, J.M. Dickson, J. Membr. Sci. 279 (2006) 529. [11] T.F. Lin, J.K. Wu, Water Res. 35 (8) (2001) 2049. [12] Y. Xu, T. Nakajima, A. Ohki, J. Hazard. Mater. 92 (2002) 275. [13] J.L. Schnoor (Ed.), Environmental Modeling: Fate and Transport of Pollutants in Water, Air, and Soil, John Wiley & Sons, New York, 1996 (Chapters 8 and 9).
[14] N. Öztürk, D. Kavak, J. Hazard. Mater. 127 (2005) 81. [15] A. Tor, Desalination 201 (2006) 267. [16] Y.H. Li, S. Wang, A. Cao, D. Zhao, X. Zhang, C. Xu, Z. Luan, D. Ruan, J. Liang, D. Wu, B. Wei, Chem. Phys. Lett. 350 (2001) 412. [17] Y. Ku, H.M. Chiou, Water Air Soil Pollution 133 (2002) 349. [18] A.K. Chaturvedi, K.P. Yadava, K.C. Pathak, V.N. Singh, Water Air Soil Pollution 49 (1990) 51. [19] R. Wang, H. Li, Y. Wang, Water Qual. Res. J. Can. 30 (1995) 81. [20] D.S. Bhargava, D.J. Killedar, Water Res. 26 (1992) 781. [21] S. Mayadevi, Ind. Chem. Eng. A 38 (1996) 155. [22] H. Yoshitake, T. Yokoi, T. Tatsumi, Chem. Mater. 14 (2002) 4603. [23] M. Jang, E.W. Shin, J.K. Park, S.I. Choi, Environ. Sci. Technol. 37 (2003) 5062. [24] E.W. Shin, J.S. Han, M. Jang, S.H. Min, J.K. Park, R.M. Rowell, Environ. Sci. Technol. 38 (2004) 912. [25] Y.H. Kim, C.M. Kim, I.H. Choi, S. Rengaraj, J.H. Yi, Environ. Sci. Technol. 38 (2004) 924. [26] M.J. Yu, W. Li, W.S. Ahn, Micropor. Mesopor. Mater. 113 (2008) 197. [27] J.C. Ray, K.S. You, J.W. Ahn, W.S. Ahn, Micropor. Mesopor. Mater. 100 (2007) 183. [28] J. Cˇejka, N. Zˇilková, J. Rathousky´, A. Zukal, Phys. Chem. Chem. Phys. 3 (2001) 5076. ˇ ejka, Appl. Catal. A: Gen. 254 (2003) 327. [29] J. C [30] A. Corma, Chem. Rev. 97 (1997) 2373. ˇ ejka, N. Zˇilková, J. Rathousky´, A. Zukal, Langmuir 20 (2004) 7532. [31] J. C [32] T.J. Pinnavaia, Z. Zhang, R.W. Hicks, Stud. Surf. Sci. Catal. 156 (2005) 1. ˇ ejka, Catal. Rev. 50 (2008) [33] C. Márquez-Alvarez, N. Zˇilková, J. Pérez-Pariente, J. C 222. [34] N. Zˇilková, A. Zukal, J. Cˇejka, Micropor. Mesopor. Mater. 95 (2006) 176. [35] R.T. Yang, Adsorbents: Fundamentals and Applications, Wiley-Interscience, Hoboken, New Jersey, 2003, p. 168. [36] Y.S. Ho, G. McKay, Water Res. 34 (2000) 735. [37] Z. Reddad, C. Gerente, Y. Andres, P.L. Cloirec, Environ. Sci. Technol. 36 (2002) 2067.
Contents lists available at ScienceDirect
Microporous and Mesoporous Materials journal homepage: www.elsevier.com/locate/micromeso
Enhanced adsorptive removal of fluoride using mesoporous alumina Grace Lee, Chao Chen, Seung-Tae Yang, Wha-Seung Ahn * Department of Chemical Engineering, Inha University, Incheon 402-751, Republic of Korea
a r t i c l e
i n f o
Article history: Received 29 October 2008 Received in revised form 1 July 2009 Accepted 4 July 2009 Available online 10 July 2009 Keywords: Mesoporous alumina Aluminum tri-sec-butoxide Cetyltrimethylammonium bromide Stearic acid Fluoride adsorption
a b s t r a c t Two different kinds of mesoporous alumina samples were prepared using aluminum tri-sec-butoxide in the presence of either cetyltrimethylammonium bromide (MA-1) or stearic acid (MA-2) as a structuredirecting agent, and tested for adsorptive removal of fluoride in water. Both materials contain a wormhole-like mesopore structure, but exhibited different textural properties: surface area (421 or 650 m2/g) and pore volume (0.96 or 0.59 cm3/g). These mesoporous aluminas demonstrated significantly improved adsorption capacity and faster kinetics to those of the commercial activated aluminas in fluoride removal by adsorption process. MA-2 prepared using stearic acid, in particular, demonstrated an adsorption capacity (14.26 mg/g) and initial adsorption rate (14.6 mg/g min) that were respectively 2.2 and 45 times higher than those of a commercial gamma alumina. The textural features of larger surface area and relatively smaller pore size in MA-2 compared to the activated aluminas are believed to be responsible for this enhancement in adsorption process. Ó 2009 Elsevier Inc. All rights reserved.
1. Introduction The presence of excess fluoride ions in water causes serious problems such as dental and skeletal fluorosis, and World Health Organization (WHO) had set a limiting value of fluoride ions in drinking water to be 1.5 mg l1 [1,2]. The concentration of fluoride ions in groundwater of many places in the world, however, exceeds the permitted level; increasing amount of wastewater containing fluoride is being released from various engineering processes, such as semiconductor manufacturing, coal power plants, electroplating, and rubber and fertilizer production to name a few. Therefore, the necessity to remove the excess fluoride from aquatic environment is high. Many methods, i.e. adsorption [3,4], ion exchange [5,6], precipitation with calcium and aluminum salts [4,7], dialysis [2,8,9], and membrane filtration [10] have been adopted to remove excessive fluoride ion from water. Among these, adsorption is less expensive than membrane filtration, easier and safer to handle compared to the contaminated sludge produced by precipitation, and more versatile than ion exchange [11–13]. In particular, adsorption process is generally known to be more effective and economical at low pollutant concentration levels [14]. Diverse kinds of adsorbents such as montmorillonite [15], Al2O3/CNT [16], activated alumina [17], fly ash [18], silica gel [19], bone char [20], and zeolite [21] have been reported so far for fluoride removal. Their low adsorption capacities and poor adsorption kinetics, however, still need improvement. An ideal * Corresponding author. Tel.: +82 32 866 0143; fax: +82 32 872 0959. E-mail address: [email protected] (W.-S. Ahn). 1387-1811/$ - see front matter Ó 2009 Elsevier Inc. All rights reserved. doi:10.1016/j.micromeso.2009.07.007
adsorbent should have uniformly accessible pores, high surface area, and physical and chemical stability [22]. Recently, aluminaimpregnated mesoporous silica SBA-15 [23,24] and mesoporous alumina prepared using an anionic surfactant [25] were reported to exhibit improved adsorption for several anionic species over commercial activated alumina. We also confirmed the effectiveness of mesoporous alumina in adsorptive removal of arsenate and orthophosphate anions [26]. In this work, two representative samples of mesoporous alumina materials, one previously reported by Ray et al. [27] and ˇ ejka et al. [28] were synthesized and tested for rethe other by C moval of fluoride in water. It is meaningful to measure the adsorption performances over different mesoporous alumina samples to draw a correct relationship between adsorption capacity and textural parameters of an alumina, since a range of mesoporous alumina materials prepared using different kinds of surfactant species with different textural properties have been claimed [29– 32]. An extensive review on mesoporous alumina synthesis and characterization was published recently [33]. Fluoride ion adsorption capacities and adsorption kinetics were measured and their effectiveness was critically evaluated against commercial activated alumina samples – Catapal B (Condea Vista) in pseudo-boehmite form and gamma alumina (Condea Chemie). 2. Experimental 2.1. Synthesis and characterization of mesoporous alumina samples MA-1 [27]: Aluminum tri-sec-butoxide (ASB, Aldrich) and cetyltrimethylammonium bromide (CTAB, Aldrich) were used as
153
G. Lee et al. / Microporous and Mesoporous Materials 127 (2010) 152–156
2.2. Adsorption experiments 2.2.1. Equilibrium adsorption isotherms A series of flasks with 100 ml of solution having 20–250 mg/L fluoride ion concentration was prepared by dissolving NaF in RO (reverse osmosis)-deionized water. The initial pH of the solutions was adjusted to 6 with 0.01 N NaOH and HCl solutions. Mesoporous alumina (1.2 g) dried at 200 °C for 4 h prior to the adsorption experiment was then put into the flasks. The flasks were magnetically stirred at a constant speed of 200 rpm for 24 h immersed in a water bath maintained at 30 °C. Upon completion, the solutions were filtered immediately using a 0.45 lm microfilter and the liquid was separated from the residual powder by centrifugation. The concentration of fluoride in the filtrate was determined by ion chromatograph (IC, Dionex, ICS-3000). 2.2.2. Adsorption kinetics The fluoride ion adsorption kinetics was monitored by adding 12 g of alumina sample into a 1 L solution. The initial fluoride ion concentration in the solution was 150 mg/L, and the solution pH was adjusted to 6.0 using acidic and alkaline solutions. The sample was agitated at 200 rpm with a magnetic bar for 8 h at 30 °C. A portion of solution was taken at predetermined time intervals and centrifuged followed by filtration. Total sampling volumes were kept below 5% of the total solution volume [17]. The fluoride concentration of the filtered solutions was again analyzed by ion chromatography. 3. Results and discussion 3.1. Characterization of mesoporous alumina and activated alumina Powder X-ray diffraction (XRD) patterns of the calcined mesoporous alumina samples designated as MA-1 and MA-2 are shown in Fig. 1a. Both MA-1 and MA-2 exhibited only single XRD peak at 2h values close to 0.9 and 1.5, respectively, which indicates a disordered mesoporous structure without a long range order, but with a fairly uniform pore size; it is worth mentioning that the single low angle XRD peak in mesoporous alumina does not necessarily
(a)
MA-2
Intensity (a.u.)
as- syn
calcined
MA-1 20
MA-2
2
40
4
60
6
80
8
Two theta (degree)
(b) Boehmite
Intensity (a.u.)
an aluminum source and a cationic structure-directing agent, respectively. 2-Butanol (Aldrich) was employed as an organic solvent. CTAB (4.38 g) was dissolved first in a solution of 10.18 g ASB and 40.0 g of 2-butanol with stirring. To this mixture, a solution of 1.44 g of distilled water in 5.0 g of 2-butanol was then added drop wise with stirring to make a molar composition of the solution mixture in Al:surfactant:water:2-butanol = 1:0.3:2: 15. This solution gradually turned into a gel, which was aged at 100 °C for 24 h in a teflon-lined autoclave at static conditions under autogenous pressure. The resulting product was filtered, extensively washed with 99% absolute ethanol, and dried at 100 °C for 5 h. Finally, the dried sample was calcined in air at 500 °C for 4 h. MA-2 [28]: 5.1 g of stearic acid (Aldrich) was dissolved in 100 ml of 1-propanol (Fluka). Water (3.1 ml) was then added to this solution and stirred for about 30 min. After addition of 13.7 g of ASB, the mixture was stirred for additional 20 min. The gel prepared was heated at 100 °C for 50 h under static conditions. The resulting product was filtered, washed with 99% absolute ethanol, and dried at 50 °C overnight. The powder was subsequently calcined at 410 °C in nitrogen and then at 420 °C in air. X-ray diffraction patterns of the products were obtained with a Rigaku DMAX 2500 using Ni-filtered CuKa radiation. The BET surface area, pore volume, and mean pore diameters were measured by N2 physisorption at 196 °C using a Micromeritics ASAP 2000 automatic analyzer. The morphology of the samples was examined by TEM (Philips, CM 200).
γ -Al2O3 20
40
60
80
γ -Al2O3 Boehmite 2
4
6
8
Two theta (degree) Fig. 1. XRD patterns of alumina samples: (a) mesoporous alumina and (b) commercial activated alumina (inset: corresponding wide angle reflections).
prove well-ordered mesoporous material with uniform pores [34]. Low angle peak remained intact after calcinations at 500 °C, which implies the structural stability of the materials during and after the removal of surfactant. As shown by the wide angle XRD patterns of MA-2 shown in the inset of the Fig. 1a, pseudo-boehmite phase of the as-synthesized MA-2 transformed into gamma phase after calcination as reported by Pinnavaia et al. [32]. XRD patterns of MA-1 in the wide angle region also exhibited the same structural transformation upon calcination (not shown). Activated alumina, on the other hand, did not show any XRD peaks in the small-angle region as shown in Fig. 1b, but exhibited diffraction lines in the wide angle region corresponding to either boehmite or gamma alumina phase depending on the sample. Fig. 2 shows TEM images of mesoporous alumina samples used in this work. Both MA-1 and MA-2 show a typical wormhole-like morphology known for mesoporous aluminas without any distinguishable features. Fig. 3 presents nitrogen adsorption–desorption isotherms and the corresponding pore size distributions of alumina samples. All the isotherms were of the type IV with hysteresis loops in different sizes. The isotherms on mesoporous alumina samples MA-1 and MA-2 are characterized by a relatively steep region in the desorption branch of the isotherms at P/P0 = 0.6–0.8 and 0.3–0.5, respectively, reflecting the capillary condensation in uniform mesopores. As shown by the pore size distribution in the inset of Fig. 3, MA-1 has significantly larger pore diameter and pore volume than MA-2. Textural properties of the corresponding mesoporous alumina and the commercial alumina are presented in Table 1. While the
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Fig. 2. TEM images of mesoporous alumina samples: (a) MA-1 and (b) MA-2.
600
dV/dlogD
Vol. adsorbed (cc/g, STP)
800
400
10
100
Pore diameter (A)
200
0 0.0
0.2
0.4
0.6
0.8
1.0
Relative pressure (P/P0) Fig. 3. N2 adsorption–desorption isotherms of alumina samples (inset: pore size distributions of the corresponding samples from adsorption branch).
Table 1 Textural properties of mesoporous and activated alumina samples. Adsorbent
BET surface area (m2/g)
Pore volume (cc/g)
Average BJH pore diameter (nm)
MA-1 MA-2 Boehmite Gamma alumina
421 650 253 220
0.96 0.59 0.38 0.50
6.3 2.9 5.1 6.6
which a solid submerged in an electrolyte exhibits zero net electrical charge on the surface. Typically ZPC values for commercial aluminas fall in a range of pH 8–10, depending on the grades [35]. The alumina surface is charged positively at pH below the ZPC whereas the surface is negatively charged at pH values above the ZPC. Accordingly, at pH below the ZPC, the adsorption sites on the surface of an activated alumina will adsorb anionic species. In accordance with this, the optimum pH in removal of fluoride by activated alumina was reported to be in the range of 5–7 [17]. In this work, pH of the fluoride solution was fixed to 6 for all adsorption experiments to make a condition close to that of drinking water. Fig. 4 shows the equilibrium fluoride ion adsorption data over alumina samples obtained at 30 °C at initial fluoride concentrations ranging from 20 to 250 mg/L. All the fluoride adsorption data were found well-fitted with the Langmuir isotherm model, which has been successfully applied to many other real sorption processes. Langmuir adsorption isotherm can be given as follows:
qe ¼
qm bC e 1 þ bC e
where qe is the amount of fluoride adsorbed per unit weight of adsorbents at equilibrium (mg/g), qm is the maximum capacity of the adsorbent (mg/g), C e is the equilibrium concentration of fluoride (mg/L), and b is the Langmuir constant related to the energy of adsorption. The maximum adsorption capacities for the MA-1 and MA-2 adsorbents were estimated to be 7.51 and 14.26 mg/g, respectively, while the corresponding values for boehmite and gamma alumina were 6.13 mg/g and 6.36 mg/g, respectively. The surface
16 14
Fluoride uptake (mg/g)
average pore diameter of mesoporous alumina MA-1 is close to those by activated aluminas, it has a surface area and pore volume of 421 m2/g and 0.96 cm3/g, respectively, which are larger than those of commercial gamma alumina by a factor close to 2.0. MA-2, on the other hand, has a significantly smaller pore diameter (2.9 nm) than others (5.1–6.6 nm), but its surface area is the biggest among them (650 m2/g), which is almost 3 times larger than that of commercial gamma alumina.
12 10 8 6 4
3.2. Batch adsorption studies 2
The fluoride adsorption by adsorbents shows distinctive behaviors depending on the bonding between fluoride species and adsorption sites on the surface of an adsorbent [17]. In aqueous solution, the adsorption process is primarily governed by the zero point charge (ZPC) of an adsorbent. The ZPC is the pH value at
0 0
50
100
150
Equilibrium concentration (mg/L) Fig. 4. Adsorption isotherms of fluoride over various alumina samples tested.
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G. Lee et al. / Microporous and Mesoporous Materials 127 (2010) 152–156 Table 2 Kinetic parameters for fluoride ion adsorption.
10
(a)
Adsorbent
Fluoride uptake (mg/g)
8
MA-1 MA-2 Boehmite Gamma alumina
6
2
4.2 10 25.0 102 3.9 102 1.7 102
qe (mg/g)
v0 (mg/g min)
R2
6.23 7.58 5.23 4.42
1.61 14.6 1.06 0.324
0.9970 1.0000 0.9996 0.9989
4
Eqs. (2) and (3) become: 2
t 1 1 ¼ þ t qt v0 qe
0 0
50
100
150
200
250
300
200
250
300
Time (min) 100
(b) 80 -1
t/qt (mingmg )
kad (g/mg min)
60
40
20
0 0
50
100
150
Time (min) Fig. 5. (a) Adsorption kinetics of fluoride ion removal and (b) pseudo-second-order adsorption rates of fluoride ion over various alumina samples tested.
area of MA-2 is about 2.6 times of that of boehmite (650/253 m2/g), while the maximum sorption capacity is about 2.3 times of that of boehmite (14.26/6.13 mg/g). It clearly shows that surface area of the adsorbent is the major factor governing fluoride adsorption capacity in aluminas. Pore size seems to be not critical as long as the pores can accommodate fluoride ions inside freely. In fact, it can be envisaged that large surface area, relatively small mesopores, and small pore volume in an alumina will lead to more close contact between fluoride ions and adsorption sites on the surface. Fig. 5 shows the adsorption kinetics data of fluoride on different alumina samples. The initial fluoride ion concentrations were kept at 150 mg/L. The data were also fitted to a pseudo-second-order kinetic model based on surface reaction [7], which has been suggested recently as more appropriate for describing ionic-type adsorption [36,37]
dqt ¼ kad ðqe qt Þ2 dt
ð1Þ
In Eq. (1), qe is the sorption capacity at equilibrium, qt is the uptake at time t, and kad is the rate constant of pseudo-second-order adsorption (g/mgmin). By integrating Eq. (1) with the boundary conditions of qt = 0 (at t = 0) and qt = qt (at t = t), the following equation can be obtained:
t 1 1 ¼ þ t qt kad q2e qe
ð2Þ
Taking into account the initial sorption rate v0 (mg/gmin)
v0 ¼ kad q2e
ð3Þ
ð4Þ
The constants can be determined experimentally from the slope and intercept of a plot of t/qt versus t shown in Fig. 5b. As shown in Fig. 5a, MA-1 and MA-2 reached their respective adsorption equilibrium after 45 or 15 min, respectively. On the other hand, the adsorption for boehmite and gamma alumina took much longer time to reach an equilibrium; adsorption progressed smoothly for 60 (boehmite) to 100 (gamma alumina) min, respectively, and then leveled off more slowly as the adsorption approaches respective equilibrium. The kinetic parameters – the pseudo-second-order rate constants (kad) and the initial sorption rates (v0) of the adsorbent materials acquired from fitting results are summarized in Table 2. The kinetics data of fluoride could be successfully simulated by the pseudo-second-order rate equation with a correlation coefficient higher than 0.99. These high correlation coefficient (R2) values for a pseudo-second-order equation for all the samples implied that the adsorption of fluoride on those samples was chemisorption in nature [36]. While the rate constant (kad) of MA-2 was 25 102 g/mg min, that of gamma alumina was the smallest as 1.7 102 g/mg min. Also, the sorption capacity at equilibrium (qe) of MA-2 was ca. 1.7 times higher than that of gamma alumina. As mentioned earlier, this difference in sorption capacity is due to the differences in surface areas among the adsorbents, but other factors such as differences in ZPC values among the samples could also influence the sorption behavior as well. The initial sorption rate (v0) is represented by a multiple of kad and q2e , and MA-2 had a v0 of 14.6 mg/g min, which is significantly faster than that of gamma alumina. It is noteworthy in that the sorption capacity values at equilibrium (qe) of mesoporous aluminas and activated aluminas estimated by the model were highly consistent with the corresponding fluoride uptake values shown in the adsorption isotherm plots shown in Fig. 4; equilibrium concentrations in solution measured were 51 (MA-2), 72 (MA-1), 95 (boehmite), and 101 mg/L (gamma alumina), respectively. These adsorption kinetic experimental results indicate that the mesoporous alumina MA-2 with the highest surface area and narrow pore size has the fastest response time for fluoride adsorption. 4. Conclusions Two kinds of mesoporous alumina were prepared using an aluminum alkoxide precursor via either cationic or anionic surfactantmediated synthesis. The effectiveness of these materials for removal of fluoride ions in aqueous solution was evaluated in batch adsorption experiments by measuring adsorption capacities and kinetic parameters. Measured equilibrium adsorption data were fitted to the Langmuir model, and kinetics data were fitted to a pseudo-second-order model with excellent agreement. Mesoporous aluminas demonstrated superior adsorption performances to commercial boehmite or gamma alumina in both sorption capacity and initial sorption rates. The mesoporous alumina prepared using anionic surfactant (MA-2), in particular, demonstrated an adsorption capacity and initial adsorption rate that were significantly
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higher than those of a commercial gamma alumina, mainly attributed to its large surface area. Acknowledgment This study was sponsored by the Korea Energy Management Corporation (KEMCO) through the Energy & Resources Development Program (2007–2008). References [1] K.M. Popat, P.S. Anand, B.D. Dasare, React. Polym. 23 (1994) 23. [2] M. Hichour, F. Persin, J. Sandeaux, C. Gavach, Sep. Purif. Technol. 18 (2000) 1. [3] R. Weerasooriya, H.U.S. Wickramarathne, H.A. Dharmagunawardhane, Coll. Surf. A: Physicochem. Eng. Aspects 144 (1998) 267. [4] H. Lounici, L. Addour, D. Belhocine, H. Grib, S. Nicolas, B. Bariou, N. Mameri, Desalination 114 (1997) 241. [5] C. Castel, M. Schweizer, M.O. Simonnot, M. Sardin, Chem. Eng. Sci. 55 (2000) 3341. [6] E.J. Reardon, Y. Wong, Environ. Sci. Technol. 15 (2000) 3247. [7] P.L. Bishop, G. Sansoucy, J. AWWA 70 (1978) 554. [8] M. Hichour, F. Persin, J. Molenat, J. Sandeaux, C. Gavach, Desalination 122 (1999) 53. [9] Z. Amor, B. Bariou, N. Mameri, M. Taky, S. Nicolas, A. Elmidaoui, Desalination 133 (2001) 215. [10] K. Hu, J.M. Dickson, J. Membr. Sci. 279 (2006) 529. [11] T.F. Lin, J.K. Wu, Water Res. 35 (8) (2001) 2049. [12] Y. Xu, T. Nakajima, A. Ohki, J. Hazard. Mater. 92 (2002) 275. [13] J.L. Schnoor (Ed.), Environmental Modeling: Fate and Transport of Pollutants in Water, Air, and Soil, John Wiley & Sons, New York, 1996 (Chapters 8 and 9).
[14] N. Öztürk, D. Kavak, J. Hazard. Mater. 127 (2005) 81. [15] A. Tor, Desalination 201 (2006) 267. [16] Y.H. Li, S. Wang, A. Cao, D. Zhao, X. Zhang, C. Xu, Z. Luan, D. Ruan, J. Liang, D. Wu, B. Wei, Chem. Phys. Lett. 350 (2001) 412. [17] Y. Ku, H.M. Chiou, Water Air Soil Pollution 133 (2002) 349. [18] A.K. Chaturvedi, K.P. Yadava, K.C. Pathak, V.N. Singh, Water Air Soil Pollution 49 (1990) 51. [19] R. Wang, H. Li, Y. Wang, Water Qual. Res. J. Can. 30 (1995) 81. [20] D.S. Bhargava, D.J. Killedar, Water Res. 26 (1992) 781. [21] S. Mayadevi, Ind. Chem. Eng. A 38 (1996) 155. [22] H. Yoshitake, T. Yokoi, T. Tatsumi, Chem. Mater. 14 (2002) 4603. [23] M. Jang, E.W. Shin, J.K. Park, S.I. Choi, Environ. Sci. Technol. 37 (2003) 5062. [24] E.W. Shin, J.S. Han, M. Jang, S.H. Min, J.K. Park, R.M. Rowell, Environ. Sci. Technol. 38 (2004) 912. [25] Y.H. Kim, C.M. Kim, I.H. Choi, S. Rengaraj, J.H. Yi, Environ. Sci. Technol. 38 (2004) 924. [26] M.J. Yu, W. Li, W.S. Ahn, Micropor. Mesopor. Mater. 113 (2008) 197. [27] J.C. Ray, K.S. You, J.W. Ahn, W.S. Ahn, Micropor. Mesopor. Mater. 100 (2007) 183. [28] J. Cˇejka, N. Zˇilková, J. Rathousky´, A. Zukal, Phys. Chem. Chem. Phys. 3 (2001) 5076. ˇ ejka, Appl. Catal. A: Gen. 254 (2003) 327. [29] J. C [30] A. Corma, Chem. Rev. 97 (1997) 2373. ˇ ejka, N. Zˇilková, J. Rathousky´, A. Zukal, Langmuir 20 (2004) 7532. [31] J. C [32] T.J. Pinnavaia, Z. Zhang, R.W. Hicks, Stud. Surf. Sci. Catal. 156 (2005) 1. ˇ ejka, Catal. Rev. 50 (2008) [33] C. Márquez-Alvarez, N. Zˇilková, J. Pérez-Pariente, J. C 222. [34] N. Zˇilková, A. Zukal, J. Cˇejka, Micropor. Mesopor. Mater. 95 (2006) 176. [35] R.T. Yang, Adsorbents: Fundamentals and Applications, Wiley-Interscience, Hoboken, New Jersey, 2003, p. 168. [36] Y.S. Ho, G. McKay, Water Res. 34 (2000) 735. [37] Z. Reddad, C. Gerente, Y. Andres, P.L. Cloirec, Environ. Sci. Technol. 36 (2002) 2067.