1. Introduction
The tremendous increase in the use of heavy metals over the past few decades has eventually resulted in an increased flux of metallic substances in the environment. The heavy metals are of special concern because they are non-degradable and persistent. From the view of environmental protection, heavy metal ions should be removed from the source to avoid pollution of natural waters and subsequent metal accumulation in food chain. Various technological methods, such as precipitation, cementation, sedimentation, filtration, coagulation, flotation, complexing, solvent extraction, membrane separation, electrochemical technique, biological process, reverse osmosis, ion exchange and adsorption have been used for the removal of toxic heavy metals from wastewater. Among these methods, adsorption is a cost-effective technique and simple to operate [1-3]. The adsorption process usually used natural organic or inorganic materials which are particularly abundant and inexpensive. These natural materials include bark/tannin-rich materials, lignin, chitin/chitosan, dead biomass, seaweed/ algae/alginate, xanthate, zeolite, clay, fly ash, peat moss, bone gelatin beads, leaf mould, moss, iron-oxide-coated sand, modified wool and modified cotton [4].
The presence of iron ions as one of the heavy metals in ground and industrial water becomes toxic at high level and then may cause environmental and human health problems [5-8]. Iron ions are attracting wide research attention since they are found in many manufacturing industries such as coatings, car, aeronautic and steel industries [9]. Many absorbents have been reported for removal of iron ions (Fe2+ or Fe3+), such as bacterias [9, 10], chitin [11], palm fruit bunch and maize cob [12], Bengal gram husk [13], tur dal husk [14], eggshells [15], ash [16], sawdust [17], activated carbon [18], natural zeolite [19], bentonite and quartz [20], and apatite [21- 23].
The general formula of apatite is M10(XO4)6Y2 (M = Ca2+, Sr2+, Pb2+, Cd2+, Ba2+, Zn2+, Mg2+, ...; XO4 =
, 
, 
, ...; Y = F−, OH−, Cl−, ...) [23-30]. Hydroxyapatite (Ca10(PO4)6(OH)2, HAP) is a member of apatite mineral family. It is an ideal adsorptive material for long-term containment because of its high sorption capacity for heavy metals, low water solubility, high stability under reducing and oxidizing conditions, availability and low cost [24]. There are many reports on the use of HAP for stabilizing a variety of metals such as Co, Pb, Cu, Zn, Cd, Sb, Cr [23-28]. Different mechanisms for metal cations retention such as ion exchange, adsorption, dissolution/precipitation and formation of surface complexes have been proposed [26,29,30].
Ma et al. [23] have reported the use of hydroxyapatite for removal of heavy metals (Pb2+, Zn2+, Cu2+, Cd2+, Fe2+ and Al3+) in batch experiments. In Fe2+ removal experiments, these authors only observed some retention of Fe2+ on new phases but no Fe2+ solids were detected by XRD. Gschwend and Reynolds [31] have reported the in-situ subsurface formation of colloidal Fe phosphate solids, which were attributed to the interaction of different phosphate species combined with Fe2+ released from the solids of the aquifer. It was suggested that the colloidal solids were vivianite. Oliva et al. [21] have studied the sorption of Fe2+ onto biogenic hydroxyapatite with pH 4.5 and 75 mg/L Fe2+. The results indicated that apatite is an effective absorbent for the removal of Fe2+. The study of iron adsorption by eggshells indicated the reversible and exothermic nature of sorption [15]. However, so far there has been no report on the use of natural apatite for removing Fe3+ from aqueous solution.
In this work, we presented the use of natural apatite for removal of Fe3+ from aqueous solution. The influencing factors such as pH, initial concentration, contact time, dosage of the adsorbent and temperature have been systematically investigated. And the removal mechanism of Fe3+ by natural apatite was discussed.
2. Materials and Methods
2.1. Materials
Natural apatite used here comes from a phosphate rock in Yichang, China. The sample was ground in a ball mill and sized by wet sieve analysis separately for experimental work. In the present studies, rock phosphate of 74 - 150 μm size range was used for the removal of Fe3+ from aqueous solution. Mineralogical analysis of rock phosphate sample reports mainly of apatite and quartz. Dolomite, calcite and iron oxide are the other associated gangue minerals.
2.2. Experimental Procedure
Iron solutions were prepared by dissolving Fe(NO3)3 analytical grade. Adsorption experiments were carried out in the batch reactors (200 mL) containing natural apatite as adsorbent and 50 mL of Fe3+ solutions having different concentrations (10, 50, 100, 150 and 200 mg/L) and pH (1 - 7) was adjusted with 0.1 M HNO3 and 0.1 M NH3·H2O. In order to investigate the effect of the temperature on the adsorption, four adsorption temperatures (273, 293, 303, and 323 K) were studied. The suspensions were stirred with a magnetic stirring bar inside the reactor for different time. Then, the suspensions were filtered through a 0.45 µm membrane filter and the Fe3+ concentration in the filtrate was analyzed using UV/Vis spectrophotometer (Shimadzu, UV-2550) at 510 nm [32].
The morphologies and microstructures of natural apatite and the reaction production were observed by using a JSM-5510 scanning electron microscopy (JEOL, Japan). The crystalline phase of natural apatite and the reaction production were detected by powder X-ray diffraction (Bruker D8 ADVANCE, Germany) using Cu Kα radiation.
3. Results and Discussions
3.1. Effect of Initial Fe3+ Concentration
The effect of initial Fe3+ concentration in the range of 10 to 200 mg/L on adsorption was investigated (Figure 1). It is evident from Figure 1 that the removal efficiency of Fe3+ decreased with the increasing initial Fe3+ concentration. And the equilibrium adsorption capacity increasedwith increasing initial concentration indicating that higher initial concentration of Fe3+ can enhance the adsorption process. The initial Fe3+ concentration provides the necessary driving force to overcome the resistances to the mass transfer of iron between the aqueous phase and the solid phase. The increase in initial Fe3+ concentration also enhances the interaction between iron and apatite powder. Therefore, an increase in initial Fe3+ concentration enhances the adsorptive uptake of Fe3+. This is due to increase in the driving force of concentration gradient, as an increase in the initial Fe3+ concentration. The percentage of Fe3+ removal efficiency was found to be 92.90% for 10 mg/L Fe3+ and 40.03% for 200 mg/L Fe3+. The experimentally derived maximum removal capability of natural apatite was 0.179 mmol/g.
Figure 1. Effect of initial concentration on the removal Fe3+ by natural apatite (the dosage = 8 g/L, initial pH = 2 and contact time = 10 min).
3.2. Effect of Contact Time
The effect of contact time was investigated in the range of 1 min to 17 min. The result (Figure 2) indicated that the removal rate of Fe3+ increased at initial period of contact time and it decreased gradually with time until the removal rate reached an equilibrium point. The equilibrium time was established within 9 min. This is probably due to the number and availability of natural apatite surface active sites, as well as the highest driving force for the mass transfer, caused the rapid Fe3+ uptake at the beginning. Another reason is that the addition of natural apatite increases the pH of solution, which leads to the facile hydrolyzation of Fe3+. As the surface adsorption sites become exhausted, the uptake rate is controlled by the rate at which the adsorbate is transported from the exterior to the interior sites of the natural apatite particles.
3.3. Effect of Natural Apatite Dosage
The effect of apatite dosage was studied in the range of 1 g/L to 28 g/L for the initial Fe3+ concentration of 200 mg/L at pH 2.85. The variation of the removal efficiency of Fe3+ ions and the pH of filter with natural apatite dosage is shown in Figure 3. It can be observed that the removal efficiency increases quickly with the increase in natural apatite dosage initially; but beyond a certain value 20 g/L, the removal efficiency reaches almost a constant value. This trend is expected because the number of adsorbent particles increases with increasing the natural apatite dosage which leads to more Fe3+ attached onto their surfaces. The significant increase in uptake was observed when the dose was increased from 4 to 20 g/L. Any further addition of the adsorbent beyond this did not cause any significant change in the adsorption. A maximum removal efficiency of 98.10% was observed at ad-
Figure 2. Effect of contact time on the removal Fe3+ by natural apatite (initial pH = 2.85, the dosage = 20 g/L, initial Fe3+ concentration = 200 mg/L).
Figure 3. Effect of dosage on the removal efficiency of Fe3+ and filter pH.
sorbent dosage of 28 g/l. The red brown insoluble iron hydroxide increased with increasing the dosage of natural apatite was observed in the filtration residue. That is to say, the hydrolyzation of Fe3+ increased very dramatically with increasing the dosage of natural apatite. The reason is that the addition of natural apatite can increase the pH of solution, as a result of sorption of H+ ions from the acid solution by surface active sites [33]. The result (Figure 3) indicates that the natural apatite dosage has an obvious effect on pH up to 16 g/L. The surface sorption active sites of H+ ions are increased with increasing the dosage of natural apatite. The active sites can adsorb most of H+ ions in the solution at the dosage 16 g/L which leads to the change of filter pH to 7. Any further addition of the natural apatite beyond this did not cause any significant change in the pH. From the discussion above, the adsorption effect and precipitation effect are may be exist together to remove Fe3+ in aqueous solution.
3.4. Effect of Initial pH
One of the most critical parameters in the adsorption process of metal ions from aqueous solutions is the pH of the medium. Hence, the effect of initial pH on removal of Fe3+ ions from aqueous solution on natural apatite was studied. The initial pH values was ranged from 1 to 7 at room temperature (20 ˚C) and the initial concentration of Fe3+ ions was chosen at 200 mg/L. The natural apatite dosage was taken as 20 g/L. The variation of the removal efficiency of Fe3+ ions and removal capability of Fe3+ ions with initial pH is shown in Figure 4. It can be observed that the removal efficiency is increased dramatically with increasing the initial pH, but beyond a certain value 3.0, the removal efficiency reaches almost a constant value. It is known that heavy metal ions convert to insoluble hydroxide and precipitated from the solution at higher value, which lead to the reduction of metal ions. Therefore, both adsorption and precipitation may be re-
Figure 4. Effect of initial pH on the removal Fe3+ by natural apatite (contact time = 10 min).
sponsible to remove Fe3+ in aqueous solution.
To understand the removal mechanism, the effect of pH on the hydrolyzation of Fe3+ was investigated. The initial concentration of Fe3+ ions was 200 mg/L whereas the initial pH values were adjusted to 1.0, 2.0, 3.0, 4.0, 5.0, 6.0, 7.0 by HNO3 and NH3·H2O solutions. The solutions were filtered after stirred 10 min and the Fe3+ concentration in the filtrate was analyzed. The result (Figure 5) indicated that initially the reduction of Fe3+ ions in the solution was not strongly. But beyond pH 3.0, the reduction increased dramatically with increasing pH and almost reached a constant value after the pH beyond 5.0. That is to say, the insoluble iron hydroxide increased with increasing pH. After pH 3.0 the hydrolyzation increased dramatically with increasing pH and almost reached a constant value after the pH beyond 5.0. The SEM images of the natural apatite and the reaction production between natural apatite and aqueous Fe3+ at pH 1 and pH 5 are shown in Figure 6. The XRD pattern of the natural apatite and the reaction production between natural apatite and aqueous Fe3+ at pH 1 and pH 5 are shown in Figure 7. It can be seen that natural apatite dissolved clearly at pH 1 solution (Figure 6(b)), and the XRD pattern of the reaction production (Figure 7) revealed the presence of FePO4.2H2O corresponding to the respective diffraction angles as 17.03˚, 17.91˚, 20.88˚, 22.06˚, 24.27˚, 32.33˚, 63.49˚, 64.07˚ and 65.48˚. The SEM images of the reaction production between natural apatite and aqueous Fe3+ at pH 5 (Figure 6(c)) indicated that there are some gel covered on the natural apatite surface, and the XRD pattern of reaction production (Figure 7) revealed the presence of β-FeOOH corresponding to the respective diffraction angles as 24.13˚, 35.66˚, 56.06˚, 61.97˚ and 64.33˚. Accordingly, we conclude that the pH value clearly affects the removal mechanism of Fe3+ by natural apatite. At low pH value, dissolution/precipitation is the predominant mechanism. The effect of hydrolyza-
Figure 5. Effect of initial pH on the hydrolyzation efficiency and the concentration of Fe3+ in the solution.
Figure 6. SEM images of natural apatite (a) and the reaction production between natural apatite and aqueous Fe3+ at pH 1 (b) and pH 5 (c).
Figure 7. XRD pattern of the natural apatite and the reaction production between natural apatite and aqueous Fe3+ at pH 1 and pH 5.
tion increased with the pH value. After the pH value above 5.0, hydrolyzation is the predominant mechanism.
3.5. Effect of Temperature
The temperature effect on removal of Fe3+ at 273 K, 293 K, 303 K, and 323 K was studied for the initial Fe3+ concentration 200 mg/L at pH 2.85. Figure 8 shows the relationship between the removal efficiency of Fe3+ and temperature. It is found that the adsorption of Fe3+ increases with increasing temperature. The increase in the removal efficiency of Fe3+ at increased temperature indicates that the adsorption of Fe3+ ions onto natural apatite is endothermic in nature.
3.6. Adsorption Isotherms
The capacity of the adsorption isotherm plays an important role in the determination of the maximum capacity of adsorption. It also provides a panorama of course taken by the system under study in a concise form and indicates how efficiently an adsorbent will adsorb and allows an estimate of the economic viability of the adsorbent commercial applications for the specified solute. In order to adapt for the considered system, an adequate model that can reproduce the experimental results obtained. The experimental equilibrium data were fitted using Langmuir, DKR and Freundlich models (Table 1). Sorption parameters are as follows: Ce (mmol/L) is the equilibrium concentration of Fe3+ in the solution, Qe (mmol/g) is the equilibrium concentration of Fe3+ at natural apatite surface, qm (mmol/g) is the maximum sorption capacity, KL (L/mmol) the Langmuir constant related to the energy of adsorption, β (mol2/J2) the DKR constant related to adsorption energy, ε (J/mol) the Polanyi potential, and K (mmol1−n Ln g−1) and n are the Freundlich constants related to the capacity and intensity of the sorption process. The graphical illustration of the linear data fitting is shown in Figure 9, while calculated sorption parameters and the degree of correlation between the sorption data and applied models are listed in Table1
Figure 8. Effect of temperature on the removal efficiency of Fe3+ (the natural apatite dosage = 20 g/L, contact time = 10 min).
(c)
(b)
(c)
Figure 9. Langmuir(a), DKR(b) and Freundlich (c) isotherms for the adsorption of Fe3+ onto natural apatite.
It is obvious from Table 1, that the highest correlation coefficient was obtained using Langmuir model, suggesting that the Fe3+ ions were sorbed by natural apatite
Table 1 . Correlations coefficients and sorption parameters obtained using Langmuir, DKR and Freundlich model.
in a monolayer. Furthermore, the qm value calculated by this equation corresponds well with the experimentally obtained sorption capacity which indicates that the Langmuir equation better fits the experimental data.
Knowing the Langmuir constant, KL, the dimensionless separation factor (RL) can be derived using the expression [34]:
where C0 (mol/dm3) denotes the initial solution concentration. RL values indicate that the sorption process is: unfavorable for RL > 1, linear for RL = 1, favorable for 0 < RL < 1, or irreversible for RL =0.
For initial Fe3+ concentration range from 1.8 × 10−4 to 3.6 × 10−3 mol/dm3, used in this study, calculated RL values were between 0.353 and 0.027, which indicated that the adsorption of Fe3+ by natural apatite was favorable.
4. Conclusion
The present investigation showed that natural apatite was an effective adsorbent for the removal of Fe3+ from aqueous solutions. The removal of Fe3+ by natural apatite was found to be dependent upon pH, contact time, initial Fe3+ concentration, dosage of the adsorbent and temperature. The removal efficiency of Fe3+ increases with the increase of adsorbent dosage and decreases with the increase of initial Fe3+ concentration. The pH value clearly affects the removal mechanism of Fe3+ by natural apatite. At low pH value, dissolution/precipitation is the predominant mechanism. The effect of hydrolyzation increased with the pH value. After the pH value above 5.0, hydrolyzation is the predominant mechanism. The equilibrium data were analyzed using Langmuir, DKR, and Freundlich isotherms. The experimental data yielded excellent fits within the following isotherms order Langmuir > DKR > Freundlich, based on its correlation coefficient values.
[2] M. Uchida, S. Ito, N. Kawasaki, T. Nakamura and S. Tanada, “Competitive Adsorption of Chloroform and Iron Ion onto Activated Carbon Fiber,” Journal of Colloid Interface Science, Vol. 220, No. 2, 1999, pp. 406-409. http://dx.doi.org/10.1006/jcis.1999.6519
[3] C. Huang and W. P. Cheng, “Thermodynamic Parameters of Iron-Cyanide Adsorption onto γ-Al2O3,” Journal of Colloid Interface Science, Vol. 188, No. 2, 1997, pp. 270-274. http://dx.doi.org/10.1006/jcis.1997.4758
[4] M. Pakula, S. Biniak and A. Swiatkowski, “Chemical and Electrochemical Studies of Interactions between Iron(III) Ions and an Activated Carbon Surface,” Langmuir, Vol. 14, No. 5, 1998, pp. 3082-3089. http://dx.doi.org/10.1021/la9705625
[5] S. E. Bailey, T. J. Olin, R, M, Bricka and D. D. Adrian, “A Review of Potentially Low-Cost Sorbents for Heavy Metals,” Water Research, Vol. 33, No. 11, 1999, pp. 2469-2479. http://dx.doi.org/10.1016/S0043-1354(98)00475-8
[6] P. Sarin, V. L. Snoeyink, J. Bebee, K. K. Jim, M. A. Beckett, W. M. Kriven, et al., “Iron Release from Corroded Iron Pipes in Drinking Water Distribution Systems: Effect of Dissolved Oxygen,” Water Research, Vol. 38, No. 5, 2004, pp. 1259-1269. http://dx.doi.org/10.1016/j.watres.2003.11.022
[7] S. Chaturvedi, P. N. Dave, “Removal of Iron for Safe Drinking Water,” Desalination, Vol. 303, No. 1, 2012, pp. 1-11. http://dx.doi.org/10.1016/j.desal.2012.07.003
[8] R. B. Lauffer, “Iron and Human Diseases,” CRC Press, London, 1992.
[9] T. F. Emery, “Iron and Your Health: Facts and Fallacies,” CRC Press, London, 1991.
[10] A. Selatnia, A, Boukazoula, N. Kechid, M. Z. Bakhti and A. Chergui, “Biosorption of Fe3+ from Aqueous Solution by a Bacterial Dead Streptomyces rimosus Biomass,” Process Biochemistry, Vol. 39, No. 11, 2004, pp. 1643- 1651. http://dx.doi.org/10.1016/S0032-9592(03)00305-4
[11] Z. Aksu, A. Calik, A. Y. Dursun and Z. Demircan, “Biosorption of Iron(III)-Cyanide Complex Anions to Rhizopus arrhizus: Application of Adsorption Isotherms,” Process Biochemistry, Vol. 34, No. 5, 1999, pp. 483-491. http://dx.doi.org/10.1016/S0032-9592(98)00115-0
[12] G. Karthikeyan, N. M. Andal and K. Anbalagan, “Adsorption Studies of Iron(III) on Chitin,” Journal of Chemical Sciences, Vol. 117, No. 6, 2005, pp. 663-672. http://dx.doi.org/10.1007/BF02708296
[13] M. M. Nassar, K. T. Ewida, E. E. Ebrahiem, Y. H. Magdy and M. H. Mheaedi, “Adsorption of Iron and Manganese Ions Using Low-Cost Materials as Adsorbents,” Adsorption Science & Technology, Vol. 22, No. 1, 2004, pp. 25- 37. http://dx.doi.org/10.1260/026361704323150971
[14] P. S. Kumar, R. Gayathri and R. P. Arunkumar, “Adsorption of Fe(III) Ions from Aqueous Solution by Bengal gram Husk Powder: Equilibrium Isotherms and Kinetic Approach,” Electronic Journal of Environmental, Agricultural & Food Chemist, Vol. 9, No. 6, 2010, pp. 1047- 1058.
[15] N. Ahalya, R. D. Kanamadi and T. V. Ramachandra, “Cr (VI) and Fe (III) Removal Using Cajanus cajan Husk,” Journal of Environmental Biology, Vol. 28, No. 4, 2007, pp. 765-769.
[16] N. Yeddou and A. Bensmaili, “Equilibrium and Kinetic of Modeling of Iron Adsorption by Eggshells in a Batch System: Effect of Temperature,” Desalination, Vol. 206, No. 1-3, 2006, pp. 127-134. http://dx.doi.org/10.1016/j.desal.2006.04.052
[17] B. Das, P. Hazarika, G. Saikia, H. Kalita, D. C. Goswami, H. B. Das, et al., “Removal of Iron from Groundwater by Ash: A Systematic Study of a Traditional Method,” Journal of Hazardous Materials, Vol. 141, No. 3, 2007, pp. 834-841. http://dx.doi.org/10.1016/j.jhazmat.2006.07.052
[18] S. A. Ahmed, “Batch and Fixed-Bed Column Techniques for Removal of Cu(II) and Fe(III) Using Carbohydrate Natural Polymer Modified Complexing Agents,” Carbohydrate Polymers, Vol. 83, No. 4, 2011, pp. 1470-1478. http://dx.doi.org/10.1016/j.carbpol.2010.09.051
[19] A. A. Hussain, S. R. Mohammed, M. Nallu and S. Arivoli, “Adsorption of Fe(III) from Aqueous Solution by Acanthaceae Activated Carbon,” Journal of Chemical and Pharmaceutical Research, Vol. 4, No. 4, 2012, pp. 2325- 2336.
[20] M. Al-Anber and Z. A. Al-Anber, “Utilization of Natural Zeolite as Ion-Exchange and Sorbent Material in the Removal of Iron,” Desalination, Vol. 225, No. 1-3, 2008, pp. 70-81. http://dx.doi.org/10.1016/j.desal.2007.07.006
[21] M. A. Al-Anber, “Removal of high-level Fe3+ from aqueous solution using natural inorganic materials: Bentonite (NB) and quartz (NQ),” Desalination, Vol. 250, No. 3, 2010, pp. 885-891. http://dx.doi.org/10.1016/j.desal.2009.06.071
[22] J. Oliva, J. D. Pablo, J.-L. Cortina, J. Cama and C. Ayora, “The Use of Apatite IITM to Remove Divalent Metal Ions Zinc(II), Lead(II), Manganese(II) and Iron(II) from Water in Passive Treatment Systems: Column Experiments,” Journal of hazardous materials, Vol. 184, No. 1, 2010, pp. 364-374. http://dx.doi.org/10.1016/j.jhazmat.2010.08.045
[23] G. N. Kousalya, M. R. Gandhi, C. S. Sundaram and S. Meenakshi, “Synthesis of Nano-Hydroxyapatite Chitin/ Chitosan Hybrid Biocomposites for the Removal of Fe(III),” Carbohydrate Polymers, Vol. 82, No. 3, 2010, pp. 594-599. http://dx.doi.org/10.1016/j.carbpol.2010.05.013
[24] Q. Y. Ma, S. J. Tralna, T. J. Logan and J. A. Ryan, “Effects of Aqueous Al, Cd, Cu, Fe(II), Ni, and Zn on Pb Immobilization by Hydroxyapatite,” Environment Science & Technology, Vol. 28, No. 7, 1994, pp. 1219-1228. http://dx.doi.org/10.1021/es00056a007
[25] A. Krestou, A. Xenidis and D. Panias, “Mechanism of Aqueous Uranium(VI) Uptake by Hydroxyapatite,” Minerals Engineering, Vol. 17, No. 3, 2004, pp. 373-381. http://dx.doi.org/10.1016/j.mineng.2003.11.019
[26] A. Corami, S. Mignardi and V. Ferrini, “Copper and Zinc Decontamination from Singleand Binary-Metal Solutions Using Hydroxyapatite,” Journal of Hazardous Materials, Vol. 146, No. 1-2, 2007, pp. 164-170. http://dx.doi.org/10.1016/j.jhazmat.2006.12.003
[27] Y. Xu, F. W. Schwartr and S. J. Traina, “Sorption of Zn2+ and Cd2+ on Hydroxyapatite Surfaces,” Environment Science & Technology, Vol. 28, No. 8, 1994, pp. 1472- 1480. http://dx.doi.org/10.1021/es00057a015
[28] I. Smiciklas, S. Dimovic, I. Plecas and M. Mitric, “Removal of Co2+ from Aqueous Solutions by Hydroxyapatite,” Water Research, Vol. 40, No. 12, 2006, pp. 2267- 2274. http://dx.doi.org/10.1016/j.watres.2006.04.031
[29] B. Sandrine, N. Ange, B.-A. Didierb, C. Eric and S. Patrick, “Removal of Aqueous Lead Ions by Hydroxyapatites: Equilibria and Kinetic Processes,” Journal of Hazardous Materials, Vol. 139. No. 3, 2007, pp. 443-446. http://dx.doi.org/10.1016/j.jhazmat.2006.02.039
[30] M. Srinivasan, C. Ferraris and T. White, “Cadmium and Lead Ion Capture with Three Dimensionally Ordered Macroporous Hydroxyapatite,” Environment Science & Technology, Vol. 40, No. 22, 2006, pp. 7054-7059. http://dx.doi.org/10.1021/es060972s
[31] N. C. da Rocha, R. C. de Campos, A. M. Rossi, E. L. Moreira, A. F. Barbosa and G. T. Moure, “Cadmium Uptake by Hydroxyapatite Synthesized in Different Conditions and Submitted to Thermal Treatment,” Environment Science & Technology, Vol. 36, No. 7, 2002, pp. 1630- 1635. http://dx.doi.org/10.1021/es0155940
[32] P. M. Gschwend and M. D. Reynolds, “Monodisperse Ferrous Phosphate Colloids in an Anoxic Groundwater Plume,” Journal of Contaminant Hydrology, Vol. 1, No. 3, 1987, pp. 309-327. http://dx.doi.org/10.1016/0169-7722(87)90011-8
[33] APHA, “Standard Methods for the Examination of Water and Wastewater,” American Public Health Association, Washington DC, 2005.
[34] K. Skartsila and N. Spanos, “Surface Characterization of Hydroxyapatite: Potentiometric Titrations Coupled with Solubility Measurements,” Journal of Colloid and Interface Science, Vol. 308, No. 2, 2007, pp. 405-412. http://dx.doi.org/10.1016/j.jcis.2006.12.049
[35] K. R. Hall, L. C. Eagleton, A. Acrivos and T. Vermeulen, “Pore and Solid Diffusion Kinetics in Fixed Bed Adsorption under Constant Pattern Conditions,” Industrial & Engineering Chemistry Fundamentals, Vol. 5, No. 2, 1966, pp. 212-223. http://dx.doi.org/10.1021/i160018a011