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J. Phys. Chem. C 2008, 112, 17540–17545
Facile Synthesis of Monodisperse Manganese Oxide Nanostructures and Their Application in Water Treatment Hongmin Chen†,‡ and Junhui He*,† Functional Nanomaterials Laboratory, Technical Institute of Physics and Chemistry, Chinese Academy of Sciences (CAS), Beijing 100190, China, and Graduate UniVersity of Chinese Academy of Sciences, Beijing 100049, China ReceiVed: July 12, 2008; ReVised Manuscript ReceiVed: September 17, 2008
Different manganese oxide nanomaterials were prepared by treating their precursor, which had been prepared by mixing KMnO4 solution and oleic acid at room temperature, at low temperatures (e200 °C). While the hierarchical morphology was kept, the phase structure was transformed from layered manganese oxide to tetragonal hausmannite. The manganese oxide nanostructures were characterized by X-ray diffraction (XRD), transmission electron microscopy (TEM), scanning electron microscopy (SEM) and energy dispersive spectroscopy (EDS), Fourier transform infrared spectroscopy (FT-IR), thermogravimetric analysis (TGA), and nitrogen adsorption-desorption measurements. These nanostructures showed better adsorption capacity of organic polluents (methylene blue) than existing MCM-22, Red mud, and other synthesized manganese oxide (including R-, β-, and γ-) materials. The adsorption capacity of the nanomaterials did not largely depend on their surface area. The possible adsorption mechanisms are also discussed. 1. Introduction Health and environmental problems, such as water pollution, represent major challenges facing the global society. Currently, the development of new approaches to control morphologies of nanomaterials has received growing interest in better solving health and environmental problems. A variety of nanomaterials have been adopted for treatment of pollutants in the environment. For example, semoconductor nanomaterials, such as ZnS, Fe2O3, TiO2, and Cu2O, have been used to remove organic molecules in water due to their excellent photoactivity.1-4 Other materials, such as manganese oxide, were also used as effective redox media for removing organic and inorganic pollutants.5 And also, some silica and carbon-based materials with composite structures were used as sorbtants due to the electrostatic attraction between the composite materials and wastes.6 For example, Wang’s group used various silica-based materials (including mesoporous silica, fly ash) as adsorbtants for removing organic dyes with better efficiencies.7 Shi’s and Holmes’s groups used silica-based materials as adsorbtants to remove toxic metal ions with excellent efficiencies.8,9 Although these silica and carbon-based materials had better adsorption abilities of removing organic dyes and toxic metal ions due to their unique large specific surface area, well-defined pore size, and shape, postgrafting with functional groups is usually necessary to modify the pore surface of silica-based mesoporous materials to get better adsorption capacity, and high cost is a major drawback of carbon materials.6a So, it remains a major challenge to develop a facile, economic route to the preparation of alternative sources with high adsorption capacities. As a typical type of transition-metal oxide, manganese oxide has been the subject of intense interest due to its unique properties, including catalytic activity and excellent wastewater * To whom correspondence should be addressed. Tel/fax: +86-10-8254 3535. E-mail:
[email protected]. † Technical Institute of Physics and Chemistry, CAS. ‡ Graduate University of Chinese Academy of Sciences.
treatment ability.5,10 In our previous report, birnessite MnO2 nanostructures, prepared by a soft chemistry route, showed excellent oxidative decomposition of formaldehyde at low temperatures.11 The soft chemistry route, which is based on a solution process, is effective for synthesis of nanostructured materials with well-controlled shape, size, and structure.12 In the current work, we further extended our synthetic method to the preparation of monodisperse manganese oxide nanostructures. First, birnessite MnO2 precursor with 3D nanostructure was synthesized. It was calcined at lower temperatures (e200 °C) into crystalline Mn3O4 nanostructures. The as-synthesized manganese oxide nanostructures kept the morphology of the precursor, and could be used as sorbent for effective removal of organic dyes and heavy metal ions in containinated water. 2. Experimental Section 2.1. Materials. Potassium permanganate (KMnO4), oleic acid, ethanol, and methylene blue (MB) were purchased from Beijing Chemical Reagent Company and used without further purification. Pure water used in all experiments had a resistivity of 18.2 MΩ · cm, and was obtained from a Milli-Q system (Millipore). 2.2. Synthesis of Manganese Oxide Nanospheres. The precursor was prepared according to the procedure described in our previous reports.11,13 In a typical procedure, 1.0 g (6.3 mmol) of KMnO4 was dissolved in 500 mL of water, and the mixture was fleetly stirred for about 30 min. A 10.0 mL sample of oleic acid was added, and a steady emulsion was formed. After the emulsion had been maintained at room temperature for a certain period of time, brown-black products were collected and washed several times with water and alcohol to remove any possible residual reactants. The product was dried in air at 80 °C for 10 h, giving the precursor. Finally, the precursor was calcined in air at 150 and 200 °C for 5 h, respectively. The as-synthesized samples were identified as S0 (dried in air at 80 °C), S1 (calcined in air at 150 °C), and S2 (calcined in air at 200 °C), respectively.
10.1021/jp806160g CCC: $40.75 2008 American Chemical Society Published on Web 10/21/2008
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2.3. Methylene Blue (MB) Adsorption. A 50 mg sample of manganese oxide powder was added under stirring to a flask containing a MB solution of 100 mg L-1. UV-visible absorption spectra were recorded at different intervals to monitor the process at the wavelength of 665 nm. The adsorption capacity was roughly estimated by eq 1 as follows:
q ) (C0 - C)V/W g-1)
(1) L-1)
where q (mg is the adsorption capacity, C0 (mg is the initial concentration of the MB solution, C (mg L-1) is the concentration of the MB solution at different intervals during the adsorption, V (L) is the initial volume of the MB solution, and W (g) is the weight of the adsorbant, i.e., 0.05 g. 2.4. Hg2+ Ion Adsorption. HgCl2 was used for the preparation of Hg2+ ion solution. Hydrochloric acid was used to prevent the precipitation of the metal ions during the adsorption experiment (for 1000 mL of metal ion solution, 5 mL of 2 M HCl was added). A 0.1 g sample of manganese oxide powder was added into 25 mL of the Hg2+ ion solution and the mixture was stirred for 24 h. The suspension was then separated by centrifugation, and the clear solution was measured by ICP. The adsorption capacity was also roughly estimated according to eq 1. 2.5. Characterization. X-ray diffraction (XRD) patterns of the as-prepared products were recorded on a Holand PANalytical X’Pert PRO MPD X-ray diffractometer, using Cu KR radiation (λ ) 0.1542 nm), operated at 40 kV and 40 mA. Scanning electron microscopy (SEM) observations and energy dispersive spectroscopy (EDS) measurements were carried out on a Hitachi S-4300 field emission scanning electron microscope (FESEM). For transmission electron microscopy (TEM) observations, powder samples were added on carbon-coated copper grids and observed on a JEOL JEM-200CX transmission electron microscope at an acceleration voltage of 150 kV. Fourier transform infrared (FT-IR) spectra were recorded on a Varian 3100 (Excalibur series) Fourier transform infrared spectrophotometer, using KBr pellet. The thermal behavior of manganese oxide nanostructures was investigated under nitrogen atmosphere, using a SDT Q600 thermogravimetric analyzer at a heating rate of 10 °C min-1. Nitrogen adsorption-desorption measurements were performed on a Micromeritics ASAP 2010 accelerated surface area analyzer at -196 °C, using the volumetric method. Specific surface areas and pore volumes were calculated by the Brunauer-Emmett-Teller (BET) method. Pore size distributions were estimated from adsorption branches of isotherms by the Barrett, Joyner, and Halenda (BJH) method. UV-vis absorption spectra were recorded on a Shimadzu UV1601 PC spectrophotometer. The Hg2+ concentrations before and after adsorption were estimated by Optima 5300DV inductively coupled plasma atomic emission spectroscopy (ICPAES, Perkin-Elmer). 3. Results and Discussion 3.1. Structure, Morphology, and Composition of the Samples. Figure 1a shows the XRD pattern of S0, indicating its weakly crystallized structure. The structure should be indexed to layered birnessite with a turbostratic structure.11,14 From the (001) reflection, crystallite sizes of the as-synthesized layered MnO2 nanostructures were calculated to be ca. 8.1 nm with use of the Scherrer equation. The increase in the intensity of the (100) peak might be related to the scattering from remaining organics.11 After the precursor was calcined at 150 °C in air for 5 h, the layered structure of birnersite was almost kept (Figure 1b), although the intensity of the (001) differaction
Figure 1. XRD patterns of S0 (a), S1 (b), and S2 (c).
Figure 2. TEM images of S0 (a and b), S1 (c and d), and S2 (e and f). Insets are the SAED patterns of manganese oxide nanostructures.
decreased. By calcining the precursor at 200 °C in air, the crystalline phase changed from the layered MnO2 to hausmannite. On the basis of the XRD pattern in Figure 1c, all diffraction peaks could be indexed to the tetragonal hausmannite structure (space group I41/amd) of Mn3O4 with lattice constants a ) b ) 5.762 Å and c ) 9.470 Å (JCPDS 24-0734).15 From the (211) reflection, crystallite sizes of the as-synthesized hausmannite nanostructures were calculated to be ca. 13.86 nm with use of the Scherrer equation. Clearly, a less crystallized structure (layered MnO2) was formed at low temperatures. It was transformed to a better crystallized phase, and even to a different crystal structure at higher temperatures. These results agree well with those in previous reports.14,16 Figure 2 shows typical TEM images of manganese oxide nanostructures of S0, S1, and S2. The results indicated that S0 (Figure 2a,b), S1 (Figure 2c,d), and S2 (Figure 2e,f) kept the original morphology of the precursor.11 Selected-area electron diffraction (SAED) patterns showed that crystalinity increased from S0 (inset in Figure 2a) to S1 (inset in Figure 2c) to S2 (inset in Figure 2e), and S2 was well-crystallized. Clearly, both XRD patterns and SAED patterns indicated that a phase transformation must have occurred during the drying process.16 This is not surprising since layered MnO2 has been used
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Figure 3. Typical SEM images of S0 (a and c) and S1 (d), and the EDS spectrum of S0 (b). Panel c and the inset in panel d are magnified images of rectangled parts in panels a and d, respectively.
frequently as a precursor for preparation of other manganese oxide nanomaterials by either hydrothermal treatment or thermal decomposition.17 The above results indicate the fact that the crystalline nature of resultant oxides could be controlled by changing the post-treatment temperature of the as-synthesized precursors. In addition, these results demonstrated success in synthesizing hausmannite and layered MnO2 through a simple and low-temperature synthesis route. Figure 3a shows a typical SEM image of the obtained precursor (S0). Clearly, the product consisted of monodisperse nanospheres ca. 97 nm in diameter. Figure 3c is a magnified SEM image of Figure 3a. It shows that the nanosphere in fact had a hierarchical structure that had been formed by the self-assembly of nanoplatelets.18 The calcined manganese oxide nanostrucures (S1 and S2) also kept the hierarchical morphology, as exemplified by Figure 3d for S1. The corresponding EDS results (Figure 3b) confirmed the presence of Mn, O, and K elements. The C element should be from the adsorbed oleic acid. Figure 4 shows FT-IR spectra of manganese oxide nanostructures. The two peaks at about 519 and 613 cm-1 arise from the stretching vibration of the Mn-O and Mn-O-Mn bonds.19 The peaks at 519 cm-1 of S0 and S1 should be attributed to layered MnO2,19a and the two peaks at about 518 and 613 cm-1 of S2 should be attributed to Mn3O4.19b,c The broad peak at 3200-3600 cm-1 could be assigned to the stretching vibration of the water molecule and OH- in the lattice. The peak at 1600 cm-1 is assigned to the bending vibration of H2O and OH-, which implies that hydroxyl groups existed in the as-synthesized nanostructures. The results clearly indicated that the nanostructures were hydroxylated.19 The peaks of C-H (1300-1500 cm-1) and C-C, C-H2 (2800-3000 cm-1) (as guided by dashed lines) are assigned to the characteristic peaks of oleic acid, indicating that oleic acid existed in the structures of S0 and S1, However, most of the absorbed oleic acid were removed by calcination at 200 °C.20 The result is in good agreement with the enlarged (100) peak of the XRD pattern of S0. The absence
Figure 4. FT-IR spectra of as-synthesized manganese oxide nanostructures at room temperature.
of these characteristic peaks of oleic acid in S2 and the presence of the peak at 1600 cm-1 (bending vibration of H2O and OH-) imply that hydroxyl groups existed in the crystal interior of our as-synthesized nanostructures. The results indicated that some of the adsorbed oleic acid should have been removed during the thermal decomposition, and the thermal decomposition at 200 °C had removed nearly all the absorbed oleic acid. The thermogravimetric analysis (TGA) data are consistent with the FT-IR characterization. The measurements were carried out in a nitrogen atmosphere and the sample temperature was raised at a rate of 10 deg min-1 to 600 °C. As shown in Figure 5, the TGA curve of the precursor shows a weight loss of 28.5% in the procedure. The weight loss of 6.5% below 180 °C should be attributed to the removal of surface-adsorbed water and part of the oleic acid.20 Above 180 °C, the adsorbed oleic acid should have been removed completely, which was also suggested by the FT-IR results of the samples. And the weight loss of 22% is believed to correspond to the release of water from the crystallites and the phase transformation.
Manganese Oxide Nanostructures for Water Treatment
Figure 5. TGA curve of the manganese oxide precursor (S0).
Figure 6. N2 adsorption-desorption isotherms (a) and corresponding pore size distributions (b) of S1 and S2, respectively.
The N2 adsorption-desorption isotherms of S1 and S2 shown in Figure 6a could be categorized as type IV,21 with a hysteresis loop in the relative pressure range of 0.5-1.0, indicating the presence of inhomogeneous mesopores. The measurements showed that the BET surface area of S1 was 70.7 m2 g-1 and that the pore size distribution was centered at 3.7, 4.9, 6.5, and 9.0 nm (Figure 5b) according to the BJH method from N2 adsorption isotherms. The single point pore volume of S1 at P/P0 ) 0.98 was 0.213 cm3 g-1. The BET surface area of S2 was estimated to be 53.8 m2 g-1, and the single point pore volume at P/P0 ) 0.994 was 0.171 cm3 g-1. There was only one small pore distribution peak at 2.3 nm (Figure 5b) according to the BJH method from N2 adsorption isotherms. 3.2. Adsorption Measurements of Organic Dyes and Heavy Metal Ions. Fresh water, free of toxic chemicals is vital to the health of human beings. Water treatment has been paid considerable attention in recent years. The development of nanoscience and nanotechnology suggests that many of the current water-treatment problems might be resolved or greatly
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Figure 7. Adsorption rates of MB on 50 mg of (a) S0, (b) S1, (c) S2, and (d) S2. Conditions: (a-c) 40 mL and (d) 20 mL of aqueous MB (100 mg L-1).
ameliorated by using nanosorbents, nanocatalysts, bioactive nanoparticles, catalytic membranes with nanostructures, and so on.1-6,22 Generally, transition metal oxides could be more suitable as adsorbents or catalysts for removal of organic polluents from water by adsorption and subsequent catalytic combustion at relatively low temperatures.23 Many researchers have used iron oxide nanomaterials to remove toxic ions and organic pollutants from water, and these materials showed higher removal capacities than bulk materials.24-26 Despite their excellent removal ability of organic molecules, the pH value of the mixtures had to be low (ca. pH 2-4) to obtain more adsorption sites. As an example of potential applications, the as-obtained manganese oxide nanastructures were used as adsorbent in wastewater treatment in neutral water (pH 7.6). MB, a common cationic dye in textile industry, was chosen as a typical organic polluent. The characteristic absorption of MB at 665 nm was chosen for monitoring the adsorption process with 50 mg of manganese oxide nanostructures. When the initial volume of aqueous MB is 50 mL, the as-obtained manganese oxide of S0 could remove over 44.2% of MB within 120 min without any additives at room temperature, as shown by curve a in Figure 7. We estimated that 1 g of S0 could remove about 35.4 mg of MB. S1 had better adsorption ability than S0, and could remove about 80% of MB within 120 min (Figure 7, curve b) with a removal capacity of 63.36 mg MB (g S1)-1. S2 could remove even more MB, i.e., above 85% of MB within just 30 min without any additives at room temperature, as shown by curve c in Figure 7, with a removal capacity of 68.4 mg MB (g S2)-1. On the basis of the above results, the adsorption rate and ability were on the order of S2 > S1 > S0. When the initial volume of aqueous MB was 20 mL, S2 could remove over 90% of MB within only 5 min without any additives at room temperature (Figure 7, curve d). The details of the removal ability were also shown by UV-vis absorption spectra (Figure 8a,b) and by photographs (Figure 8c) recorded at different times. The adsorption capability and rate were higher than those of MCM22 and Red mud.27 The nanomaterials could be recovered by combustion at 300 °C and reused. The recovered materials (e.g., the second and third regenerations) could still remove similar amounts of MB (ca. 80% within 120 min). Therefore, the removal capability was almost retained after combustion. S2 had much better adsorption ability than S1 and S0, as is further shown in Table 1. Although the specific surface area and pore volume of S1 were larger than those of S2, the adsorption capacity and the adsorption rate of S2 were relatively higher and faster than those of S1. Clearly, the adsorption capacity of these materials did not largely depend on their
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Chen and He control.31 Hg2+ is considered a highly toxic pollutant and its efficient removal is of great importance. In our experiment, we only chose S2 as absorbent because of its excellent removal ability of MB. According to ICP measurement results before and after absorption, the adsorption capacity was measured as 0.84 mg g-1. The capacity is smaller than that for MB. The manganese oxide nanomaterials prepared in the current work showed better adsorption ability for MB than other synthesized manganese oxide (including R-, β-, and γ-) materials.29 As revealed by the FT-IR results, all the manganese oxide nanomaterials (in nearly neutral (pH ∼7.0) solutions) have abundant hydroxyl groups on their surfaces. The better adsorption ability should be attributed to the strong electrostatic attraction between the surface hydroxyl groups and the cationic groups (R-S+) of MB as well as to hydrogen bonding between the hydroxyl groups and the nitrogen atoms of MB. The combination of electrostatic attraction and hydrogen bonding leads to strong binding of MB by the manganese oxide nanostructures. Although S2 has a smaller specific surface area than S1, it has higher adsorptivity. Thus, the crystal structure might also play an important role in the adsorption of MB. In the case of Hg2+ ions, however, only electrostatic attraction occurs between Hg2+ ions and the manganese oxide nanomaterials, and thus the adsorption ability for Hg2+ is lower than that for MB. 4. Conclusions
Figure 8. Absorption spectra of an aqueous solution of MB (100 mg L-1, 20 mL) in the presence of S2 (50 mg) (a), enlarged spectra of the rectangled area in panel a (b), and photographs (c) of aqueous MB with S2 (50 mg) at different times.
TABLE 1: Comparison of the Specific Surface Area and Single Point Pore Volume with the Adsorption Capacity of the Samples sample
SBET (cm3 g-1 STP)a
pore vol (cm3 g-1)b
adsorption capacity of MB (mg g-1)
S0 S1 S2
70.7 53.8
0.213 0.171
35.4 63.36 68.4
a
Calculated from the N2 adsorption branch using the BET method. b Estimated from the single point amount adsorbed at P/P0 ) 0.98 (S1) and 0.994 (S2), respectively.
specific surface area. On the basis of the FT-IR analysis, calcination of the as-synthesized sample at 150 and 200 °C led to remove of oleic acid to different extents, i.e., higher temperature led to removal of more oleic acid. So, there were more electrostatic attraction sites on the surface of S1 and S2 than S0. Due to the hydrophobic characteristic of adsorbed oleic acid, S0 and S1 had slower adsorption rates and lower adsorption capacity.28 To our best knowledge, for organic wastes manganese oxide usually has its lowest adsorption capacity in nearly neutral solutions (pH ∼7.0) but its highest adsorption ability in the pH value range of 2.6-4.5.29,30 The electrostatic attraction between the manganese oxide surface and MB species in solution at pH 7.6 was responsible for the dye removal. Removal of toxic heavy metal ions from wastewater is also a continuing research objective of environmental pollution
We have successfully fabricated a series of porous manganese oxide nanostructures with the same morphology. The phase structure of manganese oxide nanostructure was controlled by treating the precursor at different temperatures, transforming the precursor from layered manganese oxide to tetragonal hausmannite structure. Water treatment experiments indicated that the prepared manganese oxide nanomaterials exhibited higher removal capacity and rate of organic polluents in neutral solutions compared with the present MCM-22, Red mud, and manganese oxide materials (R-, β-, and γ-). The order of adsorption capacity and rate of organic polluents (methylene blue) was manganese oxide (200 °C) > manganese oxide (150 °C) > manganese oxide (80 °C), which was not largely depent on their specific surface areas. Acknowledgment. This work was supported by the National Natural Science Foundation of China (Grant Nos. 10776034 and 20871118), the Knowledge Innovation Program of the Chinese Academy of Sciences (Grant No. KGCX2-YW-111-5), and the “Hundred Talents Program” of CAS. References and Notes (1) Hu, J. S.; Ren, L. L.; Guo, Y. G.; Liang, H. P.; Cao, A. M.; Wan, L. J.; Bai, C. L. Angew. Chem., Int. Ed. 2005, 44, 1269. (2) Cao, S.; Zhu, Y. J. Phys. Chem. C 2008, 112, 6253. (3) Li, Y. Z.; Kunitake, T.; Fujikawa, S. J. Phys. Chem. B 2006, 110, 13000. (4) Xu, Y.; Chen, D.; Jiao, X.; Xue, K. J. Phys. Chem. C 2007, 111, 16284. (5) (a) Ahn, M.; Fillry, T. R.; Jafvert, C. T.; Nies, L.; Hua, I.; Cruz, J. EnViron. Sci. Technol. 2006, 40, 215. (b) Barrett, K. A.; Mcbride, M. B. EnViron. Sci. Technol. 2005, 39, 9223. (c) Sekine, Y. Atmos. EnViron. 2002, 36, 5543. (6) (a) Hernandez-Ramirez, O.; Holmes, S. M. J. Chem. Mater. 2008, 18, 2751. (b) Liu, A.; Hidajat, K.; Kawi, S.; Zhao, D. Chem. Commun. 2000, 114. (c) Lam, K.; Yeung, K.; Nckay, G. EnViron. Sci. Technol. 2007, 41, 3329.
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