Synthesis of Polyacrylic Acid Stabilized Amorphous Calcium

Jun 26, 2010 - The exothermic peak at 408 °C indicates the combustion of PAA. Synthesis of Polyacrylic Acid Stabilized ACCs. J. Phys. Chem. C, Vol. 1...
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Synthesis of Polyacrylic Acid Stabilized Amorphous Calcium Carbonate Nanoparticles and Their Application for Removal of Toxic Heavy Metal Ions in Water Guo-Bin Cai,† Gui-Xia Zhao,‡ Xiang-Ke Wang,‡ and Shu-Hong Yu*,† DiVision of Nanomaterials and Chemistry, Hefei National Laboratory for Physical Sciences at Microscale, Department of Chemistry, UniVersity of Science and Technology of China, Hefei, Anhui 230026, People’s Republic of China and Key Laboratory of NoVel Thin Film Solar Cells, Institute of Plasma Physics, Chinese Academy of Sciences, P.O. Box 1126, Hefei, Anhui 230031, People’s Republic of China ReceiVed: April 17, 2010; ReVised Manuscript ReceiVed: June 26, 2010

We report an efficient method for synthesis of poly(acrylic acid) stabilized amorphous calcium carbonate nanoparticles (ACC) and their application for removal of toxic heavy metal ions from aqueous solutions. The maxium removal capacities for Cd2+, Pb2+, Cr3+, Fe3+, and Ni2+ ions were found to be 514.62 mg g-1, 1028.21, 258.85, 320.5, and 537.2 mg g-1, respectively. The distinguishing features of the ACC nanoparticles in water treatment involve not only high removal capacities, but also decontamination of trace ions. Approximately 83.0% on average removal can be achieved in the treatment of polluted water containing trace amount of radioactive Eu3+ ions. A precipitation transformation mechanism is proposed to play a key role in such water treatment. Introduction Environmental problems have recently gained a great deal of attention. Increasing worldwide contamination of freshwater systems has become one of the key environmental problems facing humanity.1 Among all water contaminants, heavy metals are toxic even in relatively low concentrations, as they can be stored, accumulated, and transferred by organisms.2 Strict drinking water standards have been made all over the world. It has been recommended by the World Health Organization (WHO) and the Health Ministry of P. R. China that strict values be followed for drinking water, 0.005 mg Cd/L, 0.01 mg Pb/L, 0.05 mg Cr/L and so on. Various treatment technologies have been used for water contaminated by heavy metals, including chemical precipitation, solvent extraction, reverse osmosis, ionexchange, evaporation, adsorption, filtration, flotation, and so on.3 Economy and efficiency are key factors to be considered in water treatment. Compared with traditional materials, nanomaterials have shown much higher efficiency and faster rates on water treatment. Much research work has been focused on the development of nanomaterials for water decontamination, such as iron-based nanoadsorbents,4 MnO2 hollow structures,5 carboxymethyl cellulose grafted on multiwalled carbon nanotubes,6 magnesium silicate hollow nanostructures,7 and mesoporous F-TiO hollow microspheres.8 All of these materials could adsorb toxic contaminants in a short time span and with a high sorption capacity. However, the cost and inconvenience in product preparation has limited their commercial application. Therefore, a high efficiency and low-cost material is still needed for commercial applications. From an economic perspective, calcium carbonate is one of the cheapest materials in nature. Calcium carbonate has * To whom correspondence should be addressed. Fax: + 86 551 3603040. E-mail: [email protected]. † Division of Nanomaterials and Chemistry, Hefei National Laboratory for Physical Sciences at Microscale, Department of Chemistry, University of Science and Technology of China. ‡ Key Laboratory of Novel Thin Film Solar Cells, Institute of Plasma Physics, Chinese Academy of Sciences.

already been studied for heavy metal ions removal.9 However, natural calcite has only a very low efficiency on heavy metal adsorption. Maximum adsorption capacities are found as 18.52 mg g-1 Cd and 19.92 mg g-1 Pb for natural calcite at 25 °C, respectively.10 The equilibrium adsorption takes a very long time, which makes it inconvenient in water treatment.11 Thus, other forms of calcium carbonate should be considered, and we found that amorphous calcium carbonate can be very efficient in water treatment with heavy metal ions. Recently, amorphous calcium carbonate (ACC) has gained a great deal of attention. ACC is widely used by organisms to serve as temporary storage sites, precursor phases, or in a stabilized form for mechanical purposes.12 It is found in various organisms like ascidian skeletons,13 Mollusk shells,14 American lobsters,15 plants cystoliths,16 and so forth. ACC contains various amounts of water,17 which makes it lower density than other forms of calcium carbonate.18 The solubility product of ACC is reported to be 4.0 × 10-7 at 25 °C,19 which is higher than other forms of calcium carbonate and most heavy metal carbonates like cadmium, lead, zinc, and so on. Therefore, ACC could be a suitable choice for ion exchange with metal ions, especially for precipitation transformation. For years, the preparation and characterization of ACC developed slowly, due to its unstable nature.20 Recently, several additives have been found to stabilize ACC in a shorter time frame, such as magnesium ion,21 poly(acrylic acid) (PAA),22 phytic acid,23 poly(sodium 4-styrene sulfonate),24 and polypeptide.25 Herein, we prepared ACC by using PAA as a stabilizer, and examined its ability to remove heavy metal ions, including Cd2+, Pb2+, Cr3+, Fe3+, Ni2+, and even radioactive Eu3+ ions. Experimental Section Preparation of ACC Bulk Materials. Poly(acrylic acid) (PAA) stabilized ACC was prepared by a modified Wegner’s method.26 Typically, stock aqueous solutions of PAA (0.4 M; PAA, Aldrich, Mw ) 2000), CaCl2 (0.1 M; Shanghai Chemical Reagent Company) and NaOH (0.5 M; Shanghai Chemical Reagent Company) was freshly prepared in deionized water

10.1021/jp103464p  2010 American Chemical Society Published on Web 07/12/2010

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Figure 1. (a) XRD pattern of precipitated ACC sample. (b) FT-IR spectra of ACC sample. The appearance of the broad peaks of ν1 (1072 cm-1) and ν2 (867 cm-1), the split peaks of ν3 (1463 and 1409 cm-1) and the disappearance of the peak ν4 (712 cm-1) are typical FT-IR spectra of ACC. (c) Raman spectra of as-synthesized ACC. The two broad peaks at around 150-300 cm-1 and 1082 cm-1 are characteristic to ACC. (d) Thermogravimetric curve of as-prepared ACC precipitates. (e) Differential scanning calorimetry curve of as-prepared ACC precipitates. The exothermic peak at 408 °C indicates the combustion of PAA.

separately. The PAA concentration was prepared as that of the monomer units. A 20 mL NaOH solution was quickly mixed with a 80 mL aqueous solution, which contained 0.001 mol anhydrous CaCl2 (10 mL 0.1 M solution), 0.005 mol dimethy carbonate (420 µL) and 0.004 mol PAA (625 µL, 0.4 M solution). The reaction solution was kept under stirring for about 2.5 min at room temperature, and the pH value was about 12. The precipitates were separated rapidly by centrifugation and washed using distilled water and anhydrous acetone, respectively. The obtained ACC powder was dried under vacuum at room temperature for 48 h (yield 0.0765 ( 0.0052 mg in one batch).

Experiments for Removal of Heavy Metal Ions. PbCl2, CrCl3 · 6H2O, CdCl2 · 2.5H2O, and FeCl3, NiCl2 · 6H2O (Shanghai Chemical Reagent Company) were used as the sources of Pb2+, Cr3+, Cd2+, Fe3+, and Ni2+ ions, respectively. Solutions containing different concentrations of the above heavy metal ions were prepared. As-prepared ACC powder (0.025 g) was added to 25 mL of the Cd2+ solution under stirring. At different time intervals, the solids and liquids were separated by centrifugation and the concentration of Cd2+ remaining in the solution was measured by inductively coupled plasma atomic emission spectroscopy (ICP-AES) to obtain the adsorption isotherm. For other toxic ion removal processes, the experiments were taken

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under the same conditions, with the exception that the remaining ions were tested after stirring for 2 h. Experiments for Removal of Trace Amount of Metal Ions. The radiotracer 152+154Eu3+ was used in the experiments. Typically, stock aqueous solutions of Eu3+ (2 × 10-5 M) and suspensions of ACC (6 g L-1) were freshly prepared in deionized water separately. Different volumes of Eu3+ solution with 154Eu3+ radiotracer, 1.5 mL ACC suspension, and different amounts of water were added together to make a 6 mL solution. The mixture was shaken for 24 h and then centrifuged for further analysis. The 152+154Eu3+ concentration was measured by liquid scintillation counting (Packard 3100 TR/AB Liquid Scintillation analyzer, PerkinElmer) with an ULTIMA GOLD ABTM (Packard) Scintillation cocktail. Characterization. The collected sample was characterized on an (Philips X’Pert Pro Super) X-ray powder diffractometer with Cu KR radiation (λ ) 1.541874 Å). FTIR spectra were measured on a Bruker Vector-22 FT-IR spectrometer from 4000 to 400 cm-1 at room temperature. A Raman spectrum was conducted on a laser confocal Raman spectrometer (LABRAMHR, JY, France) ranging from 120 to 1800 cm-1 at room temperature. Thermogravimetric analysis (TGA) was carried out on a DTG-60H thermal analyzer (Shimadzu) with a heating rate of 10 K · min-1 from room temperature to 800 °C in an air flow. Differential scanning calorimetry (DSC) was performed on a VP-DSC calorimeter (Microcal) with a heating rate of 10 K min-1 from room temperature to 600 °C in an air flow. SEM images were obtained with a JEOL JSM-6700F scanning electron microscope operating at 10 kV and a Zeis Supra 40 high resolution field emission scanning electron microscope operating at 5 KV. Transmission electron microscope (TEM) was performed on a Hitachi H7650 transmission electron microscope operated at an acceleration voltage of 120 kV equipped with CCD imaging system. N2 adsorption measurements were performed on a ASAP 2020 Accelerated Surface Area and Porosimetry (Micromeritics) at 77 K using BarrettEmmett-Teller (BET) calculations for surface area and BJH calculations for pore size distribution for the adsorption branch of the isotherm. The concentration of metal ions in the aqueous solution was analyzed by the inductively coupled plasma atomicemission spectroscopy (ICP-AES) technique using an Atomscan Advantage (Thermo Jarrell Ash Corporation, U.S.A.) instrument. The concentration of trace Cd2+ ions was examined by inductively coupled plasma mass spectrometer (X Series 2, Thermo Fisher Scientific, U.S.A.). The ζ potential was determined on Delsa Nano C Particle Analyzer (Beckman Coulter). Results and Discussion The resultant products were characterized by powder X-ray diffraction (XRD), Fourier transform infrared spectroscopy (FTIR), Raman spectroscopy, differential scanning calorimetry (DSC), and thermogravimetric analysis (TGA). Two broad peaks on an XRD pattern show no diffraction peak of crystalline CaCO3 (Figure 1a). FT-IR and Raman spectra further confirm the amorphous phase. In the FT-IR spectrum, the in-plane bending of carbonate at 712 cm-1 (ν4) disappears, the out-ofplane bending (ν2) shifts to 867 cm-1, the symmetric stretch at 1072 cm-1 broadens, and the asymmetric stretch (ν3) splits into two parts at around 1409 and 1463 cm-1 (Figure 1b), which is in accordance with the spectra of typical ACC.27 In the Raman spectrum, a broad peak in the region 150-300 cm-1 and the peak at 1082 cm-1correspond to the lattice frequency and the symmetric stretch of carbonate (ν1) of ACC, respectively (Figure 1c). In contrast, other crystalline forms of calcium carbonate all have a series of relatively sharp peaks.27

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Figure 2. SEM and TEM analysis of as-prepared ACC. (a) SEM image of as-obtained ACC sample. Inset shows a photo of as-prepared ACC powders. (b) TEM image of the ACC sample.

TGA is used to analyze the composition of the ACC. The TG curve of the products shows three-step weight loss (Figure 1d). The weight loss before 300 °C is attributed to the water loss.28 PAA decomposed at around 400 °C and CaCO3 decomposed after 500 °C, as indicated by DSC curve (Figure 1e). Then, the composition of CaCO3:H2O:PAA is counted to be 79:16:5 in weight and the mole ratio of CaCO3 to H2O is ∼14: 16. Thus, the possible structure of ACC could be defined as CaCO3 · H2O, which is quite similar to the biogenic ACC investigated by Addadi and co-workers.29 The morphology of as-prepared ACC is characterized by scanning electron microscopy (SEM) and transmission electron microscopy (TEM). A typical SEM image of the obtained ACC is shown in Figure 2a. The sample is composed of aggregated nanoparticles, as shown by the TEM image (Figure 2b). Again, it confirms that these ACC particles are at nanoscale and easy to aggregate and the primary particles have a size of about 40 nm. The Brunauer-Emmett-Teller (BET) surface area and the pore size of the sample were measured. A typical N2 sorption isotherm and the corresponding Barret-Joyner-Halenda (BJH) pore size distribution curve of the as-synthesized ACC are shown in Figure 2, parts a and b, respectively. The sample has a BET surface area of 120 m2 g-1. The N2 isotherm of ACC is a type IV isotherm with a type H2 hysteresis loop (Figure 3a), indicating that a mesoporous structure may exist in such an ACC sample.30 The BJH analyses show that the PAA stabilized ACC exhibit a pore size of ∼8.9 nm (Figure 3b). The high BET surface area, mesoporous structure, and the nanoscale size are beneficial for ion adsorption, exchange, and diffusion. To investigate the removal capacity of the as-obtained ACC samples, toxic metal ions such as Pb2+, Cd2+, Cr3+ ions, and other metal ions (Fe3+ and Ni2+) in water solution were used. A series of systematic experiments have been performed to examine the removal of Cd2+ ions. The curve in Figure 4a shows that the balance will be established in a short time of less than 2 h for removal of Cd2+ ions ([Cd2+] ) 586.3 mg g-1, the ACC dosage was 1 g L-1). Removal of the Cd2+ ions as a function of pH is presented in Figure 4b. The removal percentage of Cd2+ is found to increase as the initial pH increasing from 1.0 to 8.0. The decontaminating effect of Cd2+ ions is quite low at pHin e 1, as 1 g L-1 ACC could not exist at such a low pH value. At pH values from 3.0 to 8.0, the removal percentage is around 85%, and the final pH value after decontamination is found to be around 8.1, which means that ACC could make a basic buffer environment. Thus, there is no need for adjusting initial pH and the applicable pH range is quite extensive. The removal percentage of Cd2+ ions increases with increasing the dosage of ACC sorbent. Under the same experimental conditions applied, the removal percentage of Cd2+ ions from the initial solution increases from ∼51% to ∼99% with increasing the ACC dosage from 0.50 to 2.0 g L-1, as shown in Figure 4c.

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Figure 3. (a) Nitrogen adsorption and desorption isotherms for the PAA stabilized ACC product. (b) The corresponding BJH pore size distribution curve. Prior to measurement, the sample was evacuated at 373 K and 10-6 Torr for 6 h to decrease the influence of water.

While the ACC dosage remains constant (1 g L-1), the removal ability increases with the concentration of Cd2+ ions. Figure 4d shows that the removal capacity of Cd2+ ions increases from 110.7 to 580.0 mg g-1 with the increasing initial concentration of Cd2+ ions. However, the removal percentage of Cd2+ ions decreases almost from 100% to 52%, as seen in Figure 4e. Figure 4e also indicates that Cd2+ ions could nearly completely be removed when the initial concentration of Cd2+ ions is less than 450 mg L-1 by using a dosage of 1 g L-1 ACC. The maximum removal ability for Cd2+ is ∼514.62 mg g-1. Previously, a removal capacity of 18.52 mg g-1 Cd for nature calcite at 25 °C was also reported by Yavuz and co-workers,10 which is relatively quite lower than the as-prepared ACC nanoparticles. Therefore, the quick and efficient decontamination capability can be achieved for the ACC nanoparticles used here. The maximum removal capability for different metal ions was examined. It was measured to be 258.85 mg g-1 for Cr3+, 1028.21 mg g-1 for Pb2+, 320.5 mg g-1 for Ni2+, and 537.2 mg g-1 for Fe3+ ions at room temperature, which is relatively higher compared to the literature.31 To make the decontamination effect of ACC visually, coordination complex was used as indicators. Figure 5 reveals that the concentration of Ni2+ ions in the aqueous solution (115.5 mg L-1) decreased dramatically after absorption with as-obtained ACC samples. The Ni2+ solution after absorption by ACC showed no color change with the addition of dimethylglyoxime, indicating almost complete removal of Ni2+ ions. So do Fe3+ ions (390.9 mg L-1), as indicated by adding SCN- ions. The difference between complete removal and removal capacity here indicates that the trace ions are still difficult to be removed in one time. Trace amount of ions are difficult to remove in solution chemistry. However, the as-prepared ACC still has the capability to remove trace heavy metal ions. A radioactive Eu3+ ion was chosen to check the capacity of ACC on decontamination of trace metal ions. Figure 6 shows that the removal capacity of Eu3+ ions increases from 0.34 to 0.85 mg g-1 with the increasing initial concentration of Eu3+ from 0.61 to 1.52 mg L-1 at an ACC dosage of 1.5 g L-1. However, to check the removal effect, the removal percentage is almost the same, which is ∼83.0% on average. We also examined the effect of ACC on removal of trace amount of Cd2+ ions. The removal percentage decreased from 68.5% to 54.9% as the concentration of Cd2+ ions decreased from 73.29 to 18.44 µg L-1. These results demonstrate that it is not easy to get rid of trace amounts of metal ions in the solution, but ACC has the ability to remove, for the most part, trace ions.

The mechanism of such ACC used for water contaminant scavenging has been investigated. In the present system, the heavy metal ions may be removed via three ways, i.e., adsorption, hydrolyzation, and precipitation transformation. At the solid/liquid interface, the adsorption mechanism is considered to be an ion exchange between metal ions and adsorbents. For PAA stabilized ACC here, the ζ potential was analyzed to be -32.06 mV. The negative charge in surface must promote the adsorption of cations on ACC nanoparticles. For a comparison, the removal of anions for ACC is quite lower than cations. The removal capacity is only 0.262 mg Cr(VI) g-1, which is quite lower than Cr3+ ions. After water treatment with Cd2+ ions, the ζ potential changed to -9.32 mV. The difference here may come from adsorption of cations. However, it cannot exclude the crystal structural change that caused the surface charge ascending. X-ray photoelectron spectroscopy was used to measure the surface atomic composition of precipitations. The result shows that the surface atomic ratio of Ca to Cd is about 3:7. According to ICP-AES results, the molar ratio of Ca to Cd in solids collected after treatment is ∼1:2. The close of surface atomic ratio and total atomic ratio indicates that adsorption may not be the main factor in water treatment with ACC. Furthermore, the removal isotherm (Figure 7) could not match with Langmuir isotherm or Freundlich isotherm, which indicates that the mechanism of ACC on removal of Cd2+ ion is not through a simple adsorption process. Because ACC could not exist in acid condition, we did not make a buffer solution in water treatment. The initial pH of PAA-ACC suspension (1 g L-1) was measured to be around 10.45. After treatment with cadmium ion solution, the pH is around 8.10. In such basic environment, Cd2+ ions could hydrolyze into Cd(OH)2 compound. In control experiments, we adjusted the pH of Cd2+ solution (586.3 mg L-1) between 8.10 and 10.40, then examined the Cd2+ ions left in water after stirring for 2 h. It turned out that the concentration of Cd2+ in water after treatment was from 135.9 mg L-1 (pH 8.10) to 2.82 mg L-1 (pH 10.40). This means that Cd2+ ions could be removed only by adjusting the pH value of solution, and the results were even better than the case using ACC. The basic condition is one of the factors. In the case of ACC, precipitation transformation must be considered, as for the existence of CO32- ions. The solubility of as-synthesized ACC was examined. It is found that in stable solution (1 wt‰ ACC/ water), the concentration of Ca2+ ion is about 92.17 mg L-1 after stirring for two hours, then the solubility product is ∼5.31 × 10-6, which is higher than the ACC prepared without

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Figure 4. (a) Removal rate of Cd2+ by the as-prepared ACC samples. C0 represents the initial concentration of Cd2+ ion; C represents the final concentration after ion exchange. C0 ) 586.3 mg L-1. (b) Effect of pH value on the removal of Cd2+ ions by ACC at 25 °C. The initial concentration of the contaminator is 586.3 mg L-1. (c) Effect of decontaminator concentration on the removal of Cd2+ ions at 25 °C. The initial concentration of the contaminator is 586.3 mg L-1. (d, e) Removal effect on the initial concentration of Cd2+ ions by ACC at concentration of 1 g L-1 at 25 °C.

Figure 5. Color change before and after removal process by indicators. (a) Before and (b) after ion exchange with ACC for Ni2+ indicated by 1 wt % dimethylglyoxime ethanol solutions, [Ni2+]in ) 115.5 mg L-1. (c) Before and (d) after ion exchange with ACC for Fe3+ by adding 5 wt % SCN- ions into Fe3+ aqueous solution, [Fe3+]in ) 390.9 mg L-1.

stabilizers.19 This solubility product is relatively larger than most undissolved compounds, which should definitely promote transformation of precipitation between ACC and heavy metal carbonates. The stability of PAA stabilized ACC here has been examined. The phase of precipitates after treatment was determined by X-ray diffraction. In control experiment, the XRD pattern perfectly agrees with that in literature (JCPDS 05-0586) as shown in Figure 8a, which indicates that the as-synthesized ACC transformed into calcite after being stirred for 2 h in water. To cadmium polluted solution, the XRD pattern of precipitates after treatment just agrees with that in literature (JCPDS 42-1342) as shown in Figure 8b, which proves that a precipitation transformation from ACC to otavite did occur. Moreover, the morphology of precipitates collected after treatment was

Synthesis of Polyacrylic Acid Stabilized ACCs

Figure 6. Removal effect on the initial concentration of Eu3+ ions by use of ACC at concentration of 1.5 g L-1 at 25 °C.

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Figure 9. SEM images of sediments collected after treatment with different precipitators. (a, b) ACC, (c) NaOH, (d) Na2CO3, [Cd2+] ) 586.3 mg L-1.

surface, then inducing precipitation transformation from CaCO3 to CdCO3 and other part of Cd2+ ions directly react with CO32ions. The precipitation transformation equation can be formulated as follows:

CaCO3 + Cd2+ f CdCO3 + Ca2+

Figure 7. Removal isotherm of ACC on the removal of Cd2+ ions at 25 °C.

According to eq 1, the equilibrium concentration of Cd2+ and Ca2+ ions in water should match with following equations theoretically.

Ksp,CaCO3 [Ca2+] 5.31 × 10-6 ) ) ≈ 1.0 × 106 2+ -12 K [Cd ] 5.2 × 10 sp,CdCO3

Figure 8. XRD patterns of precipitates after water treatment. (a) Pure water. (b) Water containing Cd2+ ions.

examined by SEM (Figure 9). The obvious differences here show that a recrystallization process occurred in such treatment process. The remaining sediment (Figure 9a,b) was quite different from Cd(OH)2 (Figure 9c), which proves again the main behavior of precipitation transformation than hydrolization. Also, there is a difference between precipitates here and CdCO3 (Figure 9d) that obtained by direct mixing CdCl2 and Na2CO3 solutions, indicating that the precipitation transformation is not a simple reaction between Cd2+ and CO32- ions that exist in solution. According to the results by XPS and ICP-AES, more cadmium was found than calcium on the solid surface; thus, it can be presumed that some of Cd2+ ions adsorbed on ACC

(1)

(2)

This value again confirms the ability of ACC to remove Cd2+ ions. However, the real molar ratio of Ca2+ and Cd2+ ions left in water is only around 30. After treatment, the concentration of Ca2+ and Cd2+ in solution is 208.1 mg L-1 and 17.48 mg L-1, respectively. The great difference here may come from a lack of time, or from the solubility reduction caused by phase transformation. The diffusion of Cd2+ ions from surface to inside needs more time, and at the same time, the phase transformation from ACC to calcite may slower the diffusion rate and decrease solubility greatly. In comparison, the ACC without using additives was also prepared and used for removal of Cd2+ ions. After treatment, the concentration of Cd2+ ions was 95.74 mg L-1, which is much larger than that treated by PAA-ACC (17.48 mg L-1). As pure ACC could only exist in water for a few minutes, the transformed calcite could delay the diffusion of Cd2+ and reduce the solubility dramatically. Therefore, this emphasizes that PAA-ACC nanoparticles are highly efficient water decontaminators. Conclusions PAA stabilized ACC nanoparticles with negative charge and high surface area synthesized in large quantities can be applied for the removal of a series of toxic heavy metal ions, Pb2+, Cd2+, Cr3+; other metal ions, such as Fe3+ and Ni2+; and radioactive Eu3+ ions from water. It is found that such ACC nanoparticles show powerful capability for fast removal of toxic heavy metal ions and are promising candidates for water

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treatment. However, some disadvantages still exist, such as nonrenewability and further treatment necessity for the sludge after ACC treatment. It is expected that the combination of the as-prepared ACC with another decontaminator could achieve further enhancement and synergistic effects in water treatment, which requires further investigation in the future. Acknowledgment. This work was supported by the National Basic Research Program of China (2010CB934700), the National Science Foundation of China (NSFC) (Grant No. 50732006), and the International Science & Technology Cooperation Program of China (2010DFA41170). References and Notes (1) Schwarzenbach, R. P.; Escher, B. I.; Fenner, K.; Hofstetter, T. B.; Johnson, C. A.; Gunten, U.; Wehrli, B. Science 2006, 313, 1072. (2) Nimick, D. A.; Gammons, C. H.; Cleasby, T. E.; Madison, J. P.; Skaar, D.; Brick, C. M. Water Resour. Res. 2003, 39, 1247. (3) Mohan, D.; Singh, K. P. Water Res. 2002, 36, 2304. (4) (a) Zhong, L. S.; Hu, J. S.; Liang, H. P.; Cao, A. M.; Song, W. G.; Wan, L. J. AdV. Mater. 2006, 18, 2426. (b) Zhong, L. S.; Hu, J. S.; Cao, A. M.; Liu, Q.; Song, W. G.; Wan, L. J. AdV. Mater. 2007, 19, 1648. (5) Fei, J. B.; Cui, Y.; Yan, X. H.; Qi, W.; Yang, Y.; Wang, K. W.; He, Q.; Li, J. B. AdV. Mater. 2008, 20, 452. (6) Shao, D. D.; Jiang, Z. Q.; Wang, X. K.; Li, J. X.; Meng, Y. D. J. Phys. Chem. B 2009, 113, 860. (7) Zhuang, Y.; Yang, Y.; Xiang, G. L.; Wang, X. J. Phys. Chem. C 2009, 113, 10441. (8) Pan, J. H.; Zhang, X.; Du, A. J.; Sun, D. D.; Leckie, J. O. J. Am. Chem. Soc. 2008, 130, 11256. (9) (a) Stipp, S. L.; Hochella, M. F.; Parks, G. A.; Leckie, J. O. Geochim. Cosmochim. Acta 1992, 56, 1941. (b) Al-Degs, Y. S.; ElBarghouthi, M. I.; Issa, A. A.; Khraisheh, M. A.; Walker, G. M. Water Res. 2006, 40, 2645. (c) Shirvani, M.; Kalbasi, M.; Shariatmadari, H.; Nourbakhsh, F.; Najafi, B. Chemosphere 2006, 65, 2178. ¨ .; Guzel, R.; Aydin, F.; Tegin, I.; Ziyadanogullari, R. Pol. (10) Yavuz, O J. EnViron. Stud. 2007, 16, 467. (11) (a) Bailey, E. H.; Mosselmans, J. F. W.; Young, S. D. Mineral. Mag. 2005, 69, 563. (b) Ahmed, I. A. M.; Crout, N. M. J.; Young, S. D. Geochim. Cosmochim. Acta 2009, 72, 1498. (12) (a) Lowenstam, H. A.; Weiner, S. On Biomineralization; Oxford University Press: New York, 1989. (b) Simkiss, K. In Biomineralization; Allemand, D., Ed.; Musee Oceanographique: Monaco, 1993. (c) Meldrum, F. C.; Co¨lfen, H. Chem. ReV. 2008, 108, 4332. (13) Beniash, E.; Aizenberg, J.; Addadi, L.; Weiner, S. Proc. R. Soc. London Ser. B 1997, 264, 461. (14) (a) Weiss, I. M.; Tuross, N.; Addadi, L.; Weiner, S. J. Exp. Zool. 2002, 293, 478. (b) Nassif, N.; Pinna, N.; Gehrke, N.; Antonietti, M.; Ja¨ger,

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