Priority pollutants (Hg2+ and Cd2+) removal from water by ETS-4 titanosilicate

Desalination 249 (2009) 742–747

Contents lists available at ScienceDirect

Desalination j o u r n a l h o m e p a g e : w w w. e l s ev i e r. c o m / l o c a t e / d e s a l

Priority pollutants (Hg2+ and Cd2+) removal from water by ETS-4 titanosilicate M. Otero a, C.B. Lopes a, J. Coimbra a, T.R. Ferreira b, C.M. Silva b, Z. Lin b, J. Rocha b, E. Pereira a,⁎, A.C. Duarte a a b

CESAM, Department of Chemistry, University of Aveiro, 3810-193, Aveiro, Portugal CICECO, Department of Chemistry, University of Aveiro, 3810-193, Aveiro, Portugal

a r t i c l e

i n f o

Article history: Accepted 22 April 2009 Available online 4 October 2009 Keywords: Mercury Cadmium Titanosilicate Sorption Purification

a b s t r a c t The synthetic microporous titanosilicate material ETS-4 has been used for the removal of Hg2+ and Cd2+ from water. Batch stirred experiments were carried out to study the equilibrium and the kinetics of the removal of Hg2+ and Cd2+ from water. It has been demonstrated that ETS-4 has a great affinity for both these metal cations even when their initial concentrations are low. The uptake rates for both Hg2+ and Cd2+ were well described by the pseudo-second order model which constants confirmed that the kinetics of the removal of Cd2+ is faster than that of Hg2+. However, at the equilibrium, ETS-4 has a higher capacity to remove Hg2+ than Cd2+. Adsorption isotherms for both Hg2+ and Cd2+ were well fitted to Langmuir isotherm and the corresponding monolayer capacities of ETS-4 are 0.43 and 0.24 µmol mg- 1, respectively, which are quite consistent with those predicted by the pseudo-second order kinetic equation. Hence, the contribution of this work is to support the use of this material for the removal of Hg2+ and Cd2+ from water. © 2009 Elsevier B.V. All rights reserved.

1. Introduction The Water Framework Directive, adopted in 2000, is designed to protect European rivers and water basins. One part of the Directive concerns emissions of chemicals that can be harmful for Humans and environment; this is regulated in article 16. The final decision on priority hazardous substances (2455/2001/EC) was adopted on 20 November 2001. The priority substances are given different status regarding reduction or phase out. Both mercury and cadmium have been classified as “priority hazardous substances” and thus, they will be subject to cessation or phasing out of discharges, emissions and losses within an appropriate timetable that shall not exceed 20 years (http://ec.europa. eu/environment/water/water-framework/index_en.html). However both these metals are priority substances, each of them has its own characteristics. Mercury (Hg) is a widely used toxic heavy metal. Globally, mercury has been used in dentistry, measuring and control equipment, batteries and lamps [1]. The chloro-alkaline industry is also known to use large amounts of this metal. Mercury has also been used as a pesticide and biocide on grain and in paper industry. Currently, mercury is found in old electrical appliances, amalgam in teeth and in lights (tubes, energy saving light bulbs and headlamps on cars). Mercury causes damage on inner organs and is very toxic to aquatic organisms, which, together with mammals, accumulate the metal. Both short-term and long-term exposure to mercury in Humans may result in central nervous system (CNS) effects. Several important

⁎ Corresponding author. Tel.: + 351 234 370721; fax: +351 234 370 084. E-mail address: [email protected] (E. Pereira). 0011-9164/$ – see front matter © 2009 Elsevier B.V. All rights reserved. doi:10.1016/j.desal.2009.04.008

ecological accidents caused by mercury, notably between 1953 and 1956 in Minamata Bay in Japan, are well known [2].Thus, mercury pollution risks on a global scale can only be crossed over through strong political will and public involvement in addition to International cooperation [3]. Cadmium (Cd) is a toxic heavy metal which is not found in pure state in nature. It is associated with zinc and non ferrous ores, and usually present in cement and phosphate fertilizer, as well as in fossil fuels [1]. Cadmium has been widely used in pigments and is still in widespread use in electronics metallurgy and for corrosive protection. Its main use has been in batteries. Cadmium and its compounds are toxic (acute and chronic) for mammals. They are listed as probable Human carcinogens. Cadmium compounds are also very toxic to aquatic life forms, especially in fresh water. Cadmium accumulates in soil, and plant uptake is a problem. In water, cadmium is adsorbed to particulate matter. Cadmium accumulates in liver and kidneys in Humans, while uptake of cadmium by daphnia, aquatic insects, molluscs, and crayfish is appreciable. Due to their toxicology and tendency to accumulate in the living bodies, relatively low concentrated discharges in the environment may have disastrous effects. It is known that heavy metal concentrations in the range of mg L- 1 can be reduced by electrochemical treatment [4,5], reverse osmosis [6], or chemical treatment [7] of the wastewater. However, these conventional technologies may be inadequate, expensive or originate secondary problems [8]. Moreover, they are usually not able to reduce economically heavy metal concentration to extremely low levels as required by environmental regulations, which limits its application in purification of drinking waters. For instance, chemical precipitation, mostly used to separate metal ions from solution, has played a major role in water treatment for a

M. Otero et al. / Desalination 249 (2009) 742–747

long time, being widely used for the treatment of many wastewaters before discharge [9]. Nonetheless, the removal of small trace compounds and high removal efficiencies usually imply the formation of a carrier precipitate to incorporate or adsorb the contaminant. Unfortunately, the waste volume from carrier precipitation may be considerable even for trace concentrations, which has encouraged the selection of alternative processes such as adsorption. Adsorption processes are probably among the most attractive and efficient methods for metal decontamination of waters [9]. Yet, the cost of the adsorbent may represent an important limiting factor [10], especially when it does not have a high and specific affinity for the target pollutant and a great amount of adsorbent is necessary to get the required level of purification. Attention has been focused on biopolymers, natural zeolites, clays, natural oxides, carbonaceous wastes, which are cheap and able to remove heavy metals from contaminated waters [11–14]. However making use of this kind of materials, their disordered and inconstant structure, the lack of model predictability and their inability to remove the pollutants down to very low levels may be, for very exigent applications, a drawback of these materials. Thus, research on new technical and highly specific materials is of great concern. Some synthetic zeolites have well-defined structures and high ion-exchange capacity, elevated selectivity and environmental compatibility [8,10]. Therefore, they present great potential for the purification of metal-polluted waters. Environmental application is not usually studied in parallel with materials research and after their production and characterization only a few microporous zeolite- type materials have been used for the removal of heavy metals such as copper, lead and cadmium from water [15,16]. Still, in the case of mercury, there are very few works and most of them deal with relatively high and environmentally unrealistic mercury concentrations (≥ 20 mg L- 1) [17,18] which contrasts with mercury high toxicity to aquatic organisms even in trace levels [19]. Only very recently, it has been proved that synthetic titanosilicates have a great potential for mercury removal from water even under competition by the major fresh-water cations Ca2+, Na+ and Mg2+ [20]. The objective of this work was to produce synthetic microporous ETS-4 material and to evaluate and compare its ability for the removal from waters of two priority substances, Hg2+ and Cd2+. Through reporting on the kinetics and equilibrium uptake of Hg2+ and Cd2+ even at trace levels when using a small amount of ETS-4, it is aimed to show the applicability and high potential of this sort of material at the tertiary treatment of industrial effluents. 2. Experimental

743

room temperature with distilled water, and dried at 70 °C overnight, the final product being an off-white microcrystalline powder. Although the production costs of ETS-4 will obviously depend on the scale and amount produced, at an industrial scale they can be estimated to be 400-600 $ per ton. Table 1 depicts relevant information on ETS-4 titanosilicate. 2.2. Analytic procedures Mercury concentration in solution was measured by atomic fluorescence spectroscopy (AFS), which allows its detection at very low levels (1 ng L- 1), ca. 100 times lower than atomic absorption spectroscopy [21,22]. Analyses were carried out on a flow-injectioncold vapour-atomic fluorescence spectrometer (Hydride/vapour generator PS Analytical Model 10.003, coupled to a PS Analytical Model 10.023 Merlin atomic fluorescence spectrometer) (PS Analytical, Orpington, Kent, England). Cadmium analysis was performed by Inductively Coupled Plasma Mass Spectrometry (ICP-MS), on a Thermo ICP-MS X Series equipped with a Burgener nebuliser. All glassware and plastic containers used in this work were acidwashed prior to use (nitric acid 25%, 12 hours for mercury followed by hydrochloric acid 25% another 12 hours for cadmium). 2.3. Sorption studies All experiments were isothermally (298 K ± 1) carried out in batch conditions in 2 L volumetric flasks. Closed volumetric flasks were used to avoid losses of metals. Mercury and cadmium solutions were prepared daily by diluting the corresponding stock solution to the desired initial concentrations ([Hg2+] ~ 0.25 µmol L- 1 and [Cd2+] ~ 5.75 µmol L- 1) in high purity water (18 MΩcm). Known masses of ETS-4 (between 1.5 and 100 mg) were added to the corresponding metallic solution and this time was considered the starting point of the experiment. The ETS-4 suspension was magnetically stirred and samples were withdrawn at increasing times, filtered through an acidwashed Millipore membrane (0.45 µm) and immediately analyzed. A control Hg2+ or Cd2+ solution was always run in parallel under the same experimental conditions. The pH of the solution was measured before and after stirring and it kept between 4 and 5 along all the experiments carried out. Hg2+ and Cd2+ sorption from solution were compared both from the kinetic and the equilibrium points of view. For the kinetic comparison, experiments were carried out for three different ratios of 1 . Equilibrium points metal/ETS-4: 0.1, 0.2 and 0.3 µmolmetal mg-ETS-4 were found at different ratios for each of these two metals.

2.1. Chemicals and materials 3. Calculation All reagents used in this work were of analytical grade. They were obtained from chemical commercial suppliers and were used without further purification. Mercury (II) nitrate standard solution (1000 ± 2 mg L- 1) was purchased from BDH Chemicals Ltd, while the certified standard solution of cadmium (II) (1001 ± 2 mg L- 1) was purchased from Merck. All working solutions were obtained by diluting the corresponding stock solutions. ETS-4 (Engelhard titanosilicates number 4) is the synthetic analogue of the mineral umbite and it was produced and used as cation exchanger to remove Hg2+ and Cd2+ ions from water. The synthesis of ETS-4 was performed as follows: an alkaline solution was made by dissolving 33.16 g of metasilicate (BDH), 2.00 g NaOH (Merck), and 3.00 g KCl (Merck) into 25.40 g H2O. 31.88 g of TiCl3 (15 % m/m TiCl3 and 10 % m/m HCl, Merck) were added to this solution and stirred thoroughly. This gel, with a molar composition 5.9 Na2O : 0.7 K2O : 5.0 SiO2 : 1.0 TiO2 : 114 H2O, was transferred to a Teflon-lined autoclave and treated at 230 °C for 17 hours under autogenous pressure without agitation. The product was filtered off, washed at

For each metal, the corresponding concentration C = C(t) was determined along with time for the residual ion in solution. The amount of metal sorbed by the titanosilicate at a given time, qt (µmol mg- 1), was calculated by the mass balance: qt = ðC0 −Ct Þ

V W

ð1Þ

Table 1 Characteristics of ETS-4 titanosilicate synthesized and used for Hg2+ and Cd2+ removal from water. Zeolitic material

ETS-4

Formula Physical form Density (kg m- 3) Particle diameter (10- 6 m) Pore diameter (10- 10 m)

(Na9Ti5Si12O38(OH)·12H2O white powder 2200 0.5-0.9 3-4

744

M. Otero et al. / Desalination 249 (2009) 742–747

where C0 (µmol L- 1) is the initial concentration of metal in the liquidphase and confirmed by the control solution, Ct (µmol L- 1) is the instantaneous liquid-phase concentration, V the volume of the solution (L) and W the titanosilicate dry weight (mg). Metal sorbed at the equilibrium, qe (µmol mg- 1), was calculated as [23]: qe = ðC0 −Ce Þ

V W

ð2Þ

where Ce (µmol L- 1) is the metal concentration in the liquid-phase at the equilibrium. In order to compare the kinetic performance of ETS-4 to remove mercury and cadmium from solution, the experimental results were fitted to the pseudo first-order Lagergren and the pseudo secondorder kinetic equations, which have already been used to describe the sorption of other metallic cations onto titanosilicates [15]. These are empirical rate equations commonly used to describe the kinetics of sorption, based on the overall sorption rate. The Lagergren rate expression is generally expressed as follows [24]: dqt = k1 ðqe −qt Þ dt

where KF ((μmol mg- 1)(L μmol- 1)1/n) and n are the Freundlich parameters. 4. Discussion 4.1. Sorption kinetics The experimental concentrations of each Hg2+ and Cd2+ as a function of time are shown in Fig. 1. For both metals concentration decreased along time until the equilibrium. The first-order rate constant (k1) was determined from the slopes and intercepts of plots of log(qe – qt) versus t. On the other hand, using Eq. (6) and plotting t/qt versus t, the second-order rate constant (k2) and qe values were determined from the slope and intercepts of the plots. These plots and the corresponding linear fittings are shown in Figs. 2 and 3 for the first and second order models, respectively. The kinetic sorption rate constants k1 and k2 are shown in Table 2, together with the experimental and calculated qe. As it may be seen, the removal of mercury, under the experimental conditions used, may be described both by the first and the second-order kinetic models. However, the qe obtained from the fitting of the first order equation

ð3Þ

where qt (µmol mg- 1) is the amount of solute sorbed onto the material at time t (h), and kl (h- 1) is the rate constant of first-order sorption. After integrating and applying boundary conditions, t = 0 to t = t and qt = 0 to qt = qe, one obtains: logðqe −qt Þ = log qe −

k1 t 2:303

ð4Þ

A straight line log (qe – qt) versus t indicates the applicability of this pseudo-first order kinetic model. In order to fit Eq. (4) to experimental data, the sorption capacity at the equilibrium, qe, must be known. Often, the pseudo first-order Lagergren equation does not fit well the full contact time range [24–26]. The pseudo second-order equation is also based on the sorption capacity of the solid phase and, in contrast with the first model, it usually predicts the behaviour over the whole range of adsorption [25,27]: dqt 2 = k2 ðqe −qt Þ dt

ð5Þ

where k2 (mg µmol - 1 h - 1) is the second-order sorption rate constant. For the boundary conditions, t = 0 to t = t and qt = 0 to qt = qe, the integrated form of Eq. (5) is: t 1 1 = + t qt qe k2 q2e

ð6Þ

If the second-order kinetics model is applicable, the plot of t/qt versus t should be linear and no parameter has to be known beforehand. Experimental results on the sorption equilibrium of each Hg2+ and Cd2+ onto ETS-4 were fitted to Langmuir (Eq. (7)) and Freundlich (Eq. (8)) isotherm equations: qe =

QKL Ce 1 + KL Ce

ð7Þ

KL (L μmol- 1) and Q (μmol mg- 1) are the Langmuir sorption equilibrium constant and maximum loading of ETS-4, respectively; 1= n

qe = KF Ce

ð8Þ

Fig. 1. Metal sorbed onto ETS-4 (qt (μmol mg- 1)) as a function of time (295 K ± 1). Three initial ratios of metal to zeolite (0.1, 0.2 and 0.3 μmol mg- 1) were used both for mercury (a) and for cadmium (b).

M. Otero et al. / Desalination 249 (2009) 742–747

Fig. 2. Pseudo first-order Lagergren plots for the removal of mercury (a) and cadmium (b) from solution by ETS-4.

better approached the experimentally obtained qe. In the case of cadmium, the qe values assuming first-order kinetics underestimated the experimentally obtained values while those derived from the second-order kinetics well agree with the experimental values. Other authors have also found that cadmium removal from water by using different zeolitic materials followed second-order kinetics [28,29]. Plots shown in Fig. 1 evidenced that the removal of Hg2+ by ETS-4 was slower than that of Cd2+. Anyway, this is confirmed by the kinetic constants shown in Table 2.

4.2. Sorption equilibrium The experimental equilibrium data are shown in Fig. 4 along with fittings to Langmuir and Freundlich isotherms. As it may be seen, at the equilibrium, the residual concentration of Hg2+ and Cd2+ in solution (Ce) was remarkably low. It must be highlighted that residual concentrations reached for Hg2+ and Cd2+ by the use of ETS-4 are below the limits established for water of drinking quality ([Hg2+] ≤ 2 µg L- 1 = 0.010 µmol L- 1; [Cd2+] ≤ 10 µg L- 1 = 0.089 µmol L- 1) by EPA [http://www.epa.gov/safewater/

745

Fig. 3. Pseudo second-order plots for the removal of mercury (a) and cadmium (b) from solution by ETS-4.

contaminants/index.html]. Differently from other materials used for the removal of Hg2+ and Cd2+ from water, ETS-4 has so a high affinity for both Hg2+ and Cd2+ that even when the initial concentration is relatively low, it is able to uptake these cations attaining so low residual concentrations in solution that water is suitable for drinking. This evidences that ETS-4 is a very appropriate material for the treatment of industrial effluents attaining the requirements by the Water Framework Directive for the next future. Furthermore, due to its capacity, a very small amount of ETS-4 is required for the purification of Hg2+ or Cd2+ polluted waters, which significantly reduces the application costs. Fig. 4 shows that both the Freundlich and the Langmuir isotherms suitably describe Hg2+ sorption equilibrium onto ETS-4. In the case of Cd2+ and under the conditions here studied, Langmuir isotherm equation seems to describe equilibrium results better than the Freundlich one. The corresponding model parameters obtained by linear fitting may be seen in Table 3. As it may be seen, the QL corresponding to Hg2+ (0.43 µmol mg- 1) is higher than the one corresponding to Cd2+ (0.24 µmol mg- 1), which indicates that ETS-4 has a higher capacity for Hg2+ removal. This is in agreement with the results obtained by Popa et al. [30] for the sorption of the radioactive

746

M. Otero et al. / Desalination 249 (2009) 742–747

Table 2 Kinetic sorption rate constants, k1 (pseudo first-order Lagergren) and k2 (pseudo second-order), together with experimental and calculated qe corresponding to the removal of Hg2+ and Cd2+ from water by ETS-4. Metal cation

Hg2+

Cd2+

Metal/ETS-4 ratio (µmol mg- 1)

First-order kinetics k1 (h- 1)

qe (µmol mg- 1) experimental

qe (µmol mg- 1) fitted

k2 (mg μmol- 1 h- 1)

qe (µmol mg- 1) experimental

qe (µmol mg- 1) fitted

0.1 0.2 0.3 0.1 0.2 0.3

0.15 0.08 0.06 0.95 0.18 0.09

0.10 0.20 0.26 0.11 0.19 0.22

0.09 0.19 0.25 0.10 0.16 0.19

1.60 0.31 0.18 9.01 2.88 2.59

0.10 0.20 0.26 0.11 0.19 0.22

0.11 0.26 0.34 0.12 0.18 0.22

forms of Cd and Hg onto ETS-4. Anyway, as it is shown in Table 4, the monolayer capacities here displayed by ETS-4 are similar or higher than most of those found in the literature for Hg2+ and Cd2+ adsorption by diverse materials. This is also true even when comparing ETS-4 capacity with that of different zeolitic materials [15,17].

Second-order kinetics

5. Conclusions ETS-4 is a low-cost zeolite-type material that has a great potential for being used in the removal of high priority pollutants such as Hg2+ and Cd2+ from contaminated water. It has been proved that this titanosilicate is able to remove both pollutants from water to a very high extent so to give water suitable for drinking, even when using a very small amount of ETS-4. It must be highlighted the possibility of achieving the desired "Zero Discharge or Totally Effluent Free (TEF)” for Hg2+ and Cd2+ by the use of ETS-4, which will be required for both pollutants by the Water Framework Directive in the next future. ETS-4 has displayed different uptake capacity for each of Hg2+ and Cd2+. Hg2+ was taken from solution more slowly than Cd2+, but the capacity of Hg2+ removal was higher than that of Cd2+. ETS-4 was able to uptake more than 98%, when starting from mercury concentrations below 0.25 μmol L- 1, attaining equilibrium concentrations lower than 9.5 10- 3 μmol L- 1. The sorption of Hg2+ and Cd2+ along time is well described by the pseudo second-order kinetic model adopted. Both Langmuir and Freundlich equations give reliable equilibrium representations for Hg2+ sorption onto ETS-4 while, under the studied conditions, the Langmuir model fits Cd2+ equilibrium results better than the

Table 3 Langmuir and Freundlich parameters corresponding to the fittings of the equilibrium data on the removal of Hg2+ and Cd2+ from water by ETS-4. Metal cation KL (L µmol- 1) QL (µmol mg- 1) KF (µmol mg- 1) (L μmol- 1)1/n n Hg2+ Cd2+

34.14 7.46

0.43 0.24

2.54 0.16

1.43 5.35

Table 4 Maximum Langmuir loading (QL) available in the literature for the sorption of Hg2+ and Cd2+ onto different materials. Adsorbent

Fig. 4. Experimental equilibrium data corresponding to the removal of mercury (a) and cadmium (b) from solution by ETS-4 together with the corresponding fittings to Langmuir and Freundlich models.

Rice husk ash Bacillus sp. Zeolitic mineral Fly ash Activated carbon Eucalyptus bark Carbon aerogel Activated carbon Areca Corn cob Ground wheat stems Activated carbon Na2CO3 treated rice husk Bentonite Zeolite 13X Zeolite 4A

Cation Hg

2+

Cd2+

QL (µmol mg- 1)

Reference

0.03 0.04 0.05 0.07 0.13 0.17 0.17 0.22 0.01 0.10 0.10 0.14 0.18 0.25 0.26 0.27

[31] [32] [33] [34] [35] [36] [37] [38] [39] [40] [41] [42] [43] [44] [44] [44]

M. Otero et al. / Desalination 249 (2009) 742–747

Freundlich one. The Langmuir maximum capacities of ETS-4 are 0.43 µmol mg- 1 Hg2+ and 0.24 µmol mg- 1 Cd2+. ETS-4 titanosilicate may be a very promising material for environmental remediation. Furthermore, very recent work has shown that complete elution of Hg2+ and Cd2+ from ETS-4 may be obtained by NaNO3 (10- 3 M). Fixed-bed operation and sorption-desorption-reuse cycles are being carried out in order to prove the practical application of ETS-4 on real wastewater purification systems. Acknowledgements We thank the Fundação para a Ciência e a Tecnologia (FCT) and FEDER for financial support. This research was supported by a FCT PhD grant (SFRH/BD/19098/2004) funding C.B. Lopes. References [1] T.E. Økland, E. Wilhelmsen, Ø. Solevåg (Eds.), A study of the priority substances of the Water Framework Directive, Monitoring and need for screening, Bergfald & Co, Norway, 2005. [2] F.S. Zhang, J.O. Nriagu, H. Itoh, Water Res. 39 (2005) 389–395. [3] S.P. Mohapatra, I. Nikolova, A. Mitchell, J. Environ. Manag. 83 (2007) 80–92. [4] R. Fischer, H. Seidel, P. Morgenstern, H.-J. Förster, W. Thiele, P. Krebs, Eng. Life Sci. 5 (2005) 163–168. [5] M. Hunsom, K. Pruksathorn, S. Damronglerd, H. Vergnes, P. Duverneuil, Electrochemical treatment of heavy metals (Cu2+, Cr6+, Ni2+) from industrial effluent and modeling of copper reduction, Water Res. 39 (2005) 610–616. [6] T.A. Kurniawan, G.Y.S. Chan, W.-H. Lo, S. Babel, Chem. Eng. J. 118 (2006) 83–98. [7] N. Meunier, P. Drogui, C. Montane, R. Hausler, G. Mercier, J.F. Blais, J. Hazard. Mater. 137 (2006) 581–590. [8] M.I. Panayotova, Waste Manage. 21 (2001) 671–676. [9] Jack S. Watson, Separation methods for waste and environmental applications, Marcel Dekker Ink, New York, 1999. [10] R. Petrus, J. Warchol, Microporous Mesoporous Mater. 61 (2003) 137–146. [11] S. Babel, T. Kurniawan, J. Hazard. Mater. 97 (2003) 219–243. [12] D. Touaibia, B. Benayada, Desalination 186 (2005) 75–80. [13] S.M. Hasany, R. Ahmad, J. Environ. Manag. 81 (2006) 286–295. [14] K. Bedoui, I. Bekri-Abbes, E. Srasra, Desalination 223 (2008) 269–273. [15] J.H. Choi, S.D. Kim, S.H. Noh, S.J. Oh, W.J. Kim, Microporous Mesoporous Mater. 87 (2006) 163–169.

747

[16] A. Cincotti, A. Mameli, A.M. Locci, R. Orru, G. Cão, Ind. Eng. Chem. Res. 45 (2006) 1074–1084. [17] D. Pérez-Quintanilla, I. Del Hierro, M. Fajardo, I. Sierra, Microporous Mesoporous Mater. 89 (2006) 58–68. [18] X. Zhu, S.D. Alexandratos, Ind. Eng. Chem. Res. 44 (2005) 8605–8610. [19] T.A. Davis, B. Volesky, H.S.F. Vieira, Water Res. 34 (2000) 4270–4278. [20] C.B. Lopes, M. Otero, J. Coimbra, E. Pereira, J. Rocha, Z. Lin, A. Duarte, Microporous Mesoporous Mater. 103 (2007) 325–332. [21] P. Pichet, K. Morrison, I. Rheault, A. Tremblay, Analysis of total mercury and methyl-mercury in environmental samples, Mercury in the biogeochemical cycle: natural environments and hydroelectric reservoirs of northern Québec (Canada), Springer, Berlin, Germany, 1999, pp. 41–52. [22] B. Rood, C. Sanford, Mercury determination in the environment, Encyclopedia of environmental analysis and remediation, John Wiley & Sons, New York, 1998, pp. 2737–2744. [23] C.C. Pavel, K. Popa, N. Bilba, A. Cecal, D. Cozma, A. Pui, J. Radioanal. Nucl. Ch. 258 (2003) 243–248. [24] S. Lagergren, Handlingar 24 (1898) 1–39. [25] Y.S. Ho, G. McKay, Process Biochem. 34 (1999) 451–465. [26] Z. Aksu, Process Biochem. 40 (2005) 997–1026. [27] C. Namasivayam, S. Senthilkumar, Ind. Eng. Chem. Res. 37 (1998) 4816–4822. [28] G.P.C. Rao, S. Satyaveni, A. Ramesh, K. Seshaiah, K.S.N. Murthy, N.V. Choudary, J. Environ. Manag. 81 (2006) 265–272. [29] S. Kocaoba, Y. Orhan, T. Akyuz, Desalination 214 (2007) 1–10. [30] K. Popa, C.C. Pavel, N. Bilba, A. Cecal, J. Radioanal. Nucl. Chem. 269 (2006) 155–160. [31] Q. Feng, Q. Lin, F. Gong, S. Sugita, M. Shoya, J. Colloid Interface Sci. 278 (2004) 1–8. [32] C. Green-Ruiz, Bioresour. Technol. 97 (2006) 1907–1911. [33] T. Gebremedhin-Haile, M.T. Olguín, M. Solache-Ríos, Water Air Soil Pollut. 148 (2003) 179–200. [34] S.S. Banerjee, M.V. Joshi, R.V. Jayaram, Sep. Sci. Technol. 39 (2005) 1611–1629. [35] M.M. Rao, D.H.K.K. Reddy, P. Venkateswarlu, K. Seshaiah, J. Environ. Manag. 90 (2009) 634–643. [36] I. Ghodbane, O. Hamdaoui, J. Hazard. Mater. 160 (2008) 301–309. [37] J. Goel, K. Kadirvelu, C. Rajagopal, V.K. Garg, Carbon 43 (2005) 197–200. [38] K. Ranganathan, Carbon 41 (2003) 1087–1092. [39] W. Zheng, X. Li, F. Wang, Q. Yang, P. Deng, G. Zeng, J. Hazard. Mater. 157 (2008) 490–495. [40] M. Šćiban, M. Klašnja, B. Škrbić, Desalination 229 (2008) 170–180. [41] G. Tan, D. Xiao. J. Haz. Mat. 164 (2009) 1359-1363. [42] M.M. Rao, D.K. Ramana, K. Seshaiah, M.C. Wang, S.W.C. Chien. J. Haz. Mat. 166 (2009) 1006-1013. [43] U. Kumar, M. Bandyopadhyay, Bioresour. Technol. 97 (2006) 104–109. [44] G.P.C. Rao, S. Satyaveni, A. Ramesh, K. Seshaiah, K.S.N. Murthy, N.V. Choudary, J. Environ. Manag. 81 (2006) 265–272.