Adsorption of Hg(II) from aqueous solutions using TiO2 and titanate nanotube adsorbents

Accepted Manuscript Title: Adsorption of Hg(II) from aqueous solutions using TiO2 and titanate nanotube adsorbents Author: Mar´ıa-Jos´e L´opez-Mu˜noz Amaya Arencibia Luis ´ Cerro Raquel Pascual Alvaro Melgar PII: DOI: Reference:

S0169-4332(16)00146-X http://dx.doi.org/doi:10.1016/j.apsusc.2016.01.109 APSUSC 32341

To appear in:

APSUSC

Received date: Revised date: Accepted date:

11-11-2015 11-1-2016 12-1-2016

Please cite this article as: M.-J. L´opez-Mu˜noz, A. Arencibia, L. Cerro, R. ´ Melgar, Adsorption of Hg(II) from aqueous solutions using Pascual, A. TiO2 and titanate nanotube adsorbents, Applied Surface Science (2016), http://dx.doi.org/10.1016/j.apsusc.2016.01.109 This is a PDF file of an unedited manuscript that has been accepted for publication. As a service to our customers we are providing this early version of the manuscript. The manuscript will undergo copyediting, typesetting, and review of the resulting proof before it is published in its final form. Please note that during the production process errors may be discovered which could affect the content, and all legal disclaimers that apply to the journal pertain.

Graphical Abstract Aqueous Hg(II) adsorption

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TNT TiO2-SG TiO2-P25

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TiO2-P25, TiO2-SG and titanate nanotubes were compared as Hg adsorbents Mercury uptake was significantly enhanced as pH increased up to pH 10 Porosity is not a crucial variable in the global Hg(II) adsorption rate Titanate nanotubes exhibited net adsorption capacities higher than titania samples TNT materials showed the smallest capacities normalized to surface area

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Adsorption of Hg(II) from aqueous solutions using TiO2 and titanate nanotube adsorbents

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María-José López-Muñoz, Amaya Arencibia*[email protected], Luis Cerro, Raquel Pascual, Álvaro Melgar

Department of Chemical and Energy Technology, Chemical and Environmental Technology, Mechanical Technology and Analytical Chemistry, ESCET, Universidad Rey Juan Carlos, C/ Tulipán s/n, 28933 Móstoles, Madrid, Spain Escuela Superior de Ciencias Experimentales y Tecnología, Universidad Rey Juan Carlos., Campus de Móstoles. C/ Tulipán s/n 28933 Móstoles (Madrid) SPAIN. Phone: +34 91 488 70 85, fax: +34 91 488 7068

Abstract Titania and titanate nanotubes were evaluated as adsorbents for the removal of Hg(II) from aqueous solution. Commercial titanium dioxide (TiO2-P25, Evonik), a synthesized anatase sample obtained by the sol-gel method (TiO2-SG) and titanate nanotubes prepared via hydrothermal treatment (TNT) were compared. Mercury adsorption was analysed by kinetic and equilibrium experiments, studying the influence of pH and the type of adsorbents. The kinetics of Hg(II) adsorption on titania and titanate nanotubes could be well described by the pseudosecond order model. It was found that the process is generally fast with small differences between adsorbents, which cannot be explained by their dissimilarities in textural properties.

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Keywords Hg(II) removal; Adsorption; TiO2; Titanate nanotubes; water treatment

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1. INTRODUCTION

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Equilibrium isotherm data were best fitted with the Sips isotherm model. The maximum adsorption capacities of Hg(II) were achieved with titanate nanotubes sample, whereas between both titania samples, TiO2-SG exhibited the highest mercury uptake. For all adsorbents, adsorption capacities were enhanced as pH was increased, achieving at pH 10 Hg(II) adsorption capacities of 100, 121, and 140 mg g−1 for TiO2-P25, TiO2-SG, and TNT, respectively. Differences between samples were discussed in terms of their crystalline phase composition and chemical nature of both, mercury species and surface active sites.

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Pollution of water resources with heavy metals is a worldwide issue of major concern due to their great potential threat to the environment and human health. Among heavy metals, mercury is highly toxic to living organisms. It is well documented that exposures to mercury can affect the human nervous system and harm the brain, heart, kidneys, lungs, and immune system [1]. Mercury and its compounds may enter aquatic resources through weathering of soils and rocks, atmospheric deposition or as a result of a variety of human activities mainly related to coal use, mining of metals and subsequent production (e.g. cement, chlorine, caustic soda). Once in surface water, mercury can undergo a number of chemical transformations, among them conversion by naturally present bacteria to methylmercury, which is considered the more toxic mercury species [2]. Moreover, contaminated sediments at the bottom of surface waters serve as mercury reservoirs, recycling mercury back into the aquatic ecosystem [3].

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It is therefore important to develop effective technologies for the treatment of effluents containing mercury before its disposal to the environment. So far, different physicochemical and biological approaches have been investigated [4] among which, adsorption has been widely used. A variety of materials have been proposed as mercury adsorbents, such as activated carbons [5,6, ], silica [7], alumina and metal hydroxides [8,9, ] clay minerals [10,11], or biosorbent materials [12,13], being the adsorbents prepared from organic sulphur functionalization of inorganic matrix underlined for their very high adsorption capacities [14,15,16].

Taking into account the current interest of finding suitable and low-cost adsorbents for the removal of inorganic mercury in aqueous systems, the present study aimed at investigating the feasibility of titanium dioxide and titanate nanotubes for such purpose. Titanium dioxide appears as an interesting material to be explored as mercury adsorbent. On one side, it has the advantage of being very stable in aqueous solution and its point of zero charge (pHpzc) near neutral pH makes it possible to study Hg(II) sorption on positively or negatively charged surfaces of TiO2 over a broad range of pH [17]. On the other side, it has been previously shown that the use of TiO2 as a photocatalyst provides an effective procedure to achieve the uptake of Hg(II) from aqueous solutions [18,19,20,21,22,23] and the extent of pollutant sorption on the semiconductor can be a determining factor in the overall photocatalytic process.

With regard to titanate nanotubes synthesized by hydrothermal treatment of crystalline TiO2 in highly concentrated alkaline solutions, they have recently received a great deal of attention since the first reports by Kasuga et al. [24,25]. Their high specific surface area, pore volume, and ion-exchange capacity make them materials of great potential as effective adsorbents to metals.

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The removal of a variety of species such as Cu(II), Cd(II), Pb(II), Cd(II), Cr(III) or Pd(II) by adsorption on titanate nanotubes has been previously investigated [26,27,28,29]. However, to the best of our knowledge, the adsorption capacity of titanate nanotubes for aqueous Hg(II) has not been yet explored in detail.

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In the present paper, titania and titanate nanotubes have been evaluated as adsorbents for the uptake of Hg(II) from aqueous solutions. The influence of some operating conditions such as initial Hg(II) concentration and pH on the adsorption process was investigated by analyzing the equilibrium and kinetics in batch adsorption experiments. The Langmuir, Freundlich, and Sips isotherm models were examined for their ability to model the equilibrium sorption experimental data whereas pseudo first order and pseudo second order equations were tested for reproducing the kinetic results.

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2. EXPERIMENTAL 2.1. Materials Two different titania samples were evaluated as Hg(II) sorbents: i) a commercial titanium dioxide, (TiO2-P25, Evonik), which was used as provided; and ii) an anatase sample, which was synthesized by a sol-gel method based in a procedure reported by Zhu et al [30], in which the hydrolysis of titanium butoxide is controlled through an esterification reaction between acetic acid and ethanol. Briefly, 37 mL of titanium n-butoxide (TNB, ≥ 99%, Acros Organic) were dissolved in 75 mL of absolute ethanol (≥ 99.9%, Scharlau), and then 60 mL of acetic acid (CH3COOH, ≥ 99.8%, Scharlau) and 3 mL of sulphuric acid (H2SO4, ≥ 95%, Scharlau) were added. The reaction mixture was then stirred for 2 h at 328 K and thereafter, aged for 2 h. The resultant precipitate was separated by filtration, rinsed with ethanol and deionized water and dried overnight at 363 K. The dried gel was then grounded to fine powders in agate mortar, and calcined in air at 1023 K to yield a TiO2 sample which hereinafter will be named TiO2-SG.

Titanate nanotubes were prepared via a hydrothermal treatment of TiO2 following a procedure based on that reported by Kasuga et al. [Error! Bookmark not defined., 25]. In a typical synthesis, 10 g of TiO2-P25 were added to 100 mL of 10 M NaOH solution. After stirring, the mixture was placed in a Teflon-lined autoclave and heated at 403 K for 20 hours. The white precipitate obtained was filtered and repeatedly washed with deionized water and a 0.01 M HCl solution until the pH was around 7, in order to achieve the ion exchange of surface Na+ by H+. Finally the solid was dried overnight at 373 K to yield the as-prepared titanate nanotubes, hereinafter named as TNT.

2.2 Adsorption experiments The adsorption study was carried out at 293 K in the batch mode. Mercury stock solutions were prepared by dissolving mercuric chloride (HgCl2) in Milli-Q deionized water (resistivity 18.2 MΩ·cm).

In each individual run a proper amount of TiO2 or titanate nanotubes samples was dispersed in HgCl2 aqueous solution to get a 2 g L−1 suspension. A small amount of HCl or NaOH 0.1 M solutions was used to adjust the pH at different values: 2, 4.5 (natural pH), 7, and 10. After

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removing the adsorbent with a syringe filter of 0.22 µm the filtered solutions were collected for analysis. The amount of mercury adsorbed was determined by difference between initial and final metal concentrations in the solution.

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For kinetic adsorption experiments 1 L of the suspension was stirred at 200 r.p.m. during 3 hours, taking the appropriate aliquots at established times. The initial concentration of Hg(II) was fixed at 100 mg L−1.

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2.3 Characterization and analytical techniques

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Equilibrium adsorption experiments to obtain the aqueous mercury isotherms were carried out by modifying the solution concentration of Hg(II) from 5 to 200 mg L−1. Individual runs were performed by maintaining 45 mL suspension of TiO2 or titanate nanotubes samples under stirring until equilibrium was certainly achieved. Thus, experiments lasted 1 hour at pH 7 and 10, and 2.5 hours for pH 2 and pH 4.5.

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X-ray diffraction (XRD) patterns of the commercial and synthesized materials were acquired on a Philips X’PERTMPD equipment, using CuKα radiation (λ = 1.54059 Å). Scans were made in the 2θ range 5–70° with a step size of 0.04 and a step time of 1 s, enough to obtain a good signal-to-noise ratio in all the studied reflections.

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Nitrogen adsorption-desorption isotherms at 77 K were measured on a Micrometrics Tristar 3000 sorptometer (Micrometrics Instruments Corp. USA). The surface areas were determined by using the Brunauer-Emmet-Teller (BET) equation in the pressure range 0.05-0.2 and the pore volumes were taken at P/P0 = 0.97. Prior to the analysis, samples were degassed in two isothermic steps: at 90°C for 1 h and then at 130°C for 12 h.

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Concentrations of Hg(II) in the aqueous solutions were determined by Cold Vapour Atomic Fluorescence Spectroscopy (CV-AFS) using a PSA Analytical Millennium Merlin. The analyses were performed following the EN 13506 standard procedure [31].

Scanning electron microscopy (SEM) images were acquired with two equipments, an XL30 ESEM microscope and an UHR Nova NanoSEM 230 device.

Transmission electron micrographs (TEM) were taken on a JEOL JEM-2000 FX instrument, working at 200 kV. Prior to the analysis the samples were dispersed in acetone, stirred in an ultrasonic bath and placed on a carbon-coated copper grid.

2.4 Mercury speciation modelling The distribution of mercury species in aqueous solution was theoretically obtained as a function of pH by using the chemical equilibrium modelling software MINEQL+[32]. Stability constants of mercury complexes used in these calculations are shown in Table 1.

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1. RESULTS AND DISCUSSION

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3.1 Characterization of adsorbent materials The XRD patterns in Fig. 1 illustrate the crystalline phases of the three evaluated samples. TiO2P25 consists of anatase and rutile in a ratio of ca. 4:1, according to the relative intensity of the most intense diffraction peaks for each pure phase, e.g. (101) anatase (2θ~25.3°) and (110) rutile (2θ~27.4°). The diffraction peaks for TiO2-SG sample demonstrate it is pure anatase. The XRD pattern of TNT sample shows diffractions at scattering angles (2θ) 9.6°, 24.1°, 28.5°, 48.3° and 62.4°, similar to those previously ascribed to trititanate nanotubes which are consistent with a well-defined, layered structure [33,34]. In particular, the reflection at 9.6° is usually attributed to the interlayer distance in NaxH2-xTi3O7·nH2O where the values of x and n depend on the degree of sodium proton exchange during the washing procedure [35].

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Fig. 2 displays SEM and TEM micrographs acquired for titania and TNT samples respectively. As it is well-known, TiO2-P25 material (Fig 2a) consists of particles with size around 20-30 nm that form agglomerates in solution. TiO2-SG (Fig 2b) is also made of nanometric particles that seem slightly rounded, with sizes ranging from 70 to 90 nm. The nanotubular structure of TNT sample was confirmed by transmission electron microscopy (Fig 2c). As can be seen, the material consists of multi-walled tubes with inner and external diameters of around 4-5 nm and 8-10 nm, respectively. Most tubes were observed to have open ends with lengths reaching more than 300 nm as previously described [33].

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Nitrogen adsorption-desorption isotherms at 77 K obtained for the three materials are displayed in Fig. 3. The commercial TiO2-P25 and the sol-gel synthesized TiO2–SG material exhibit similar adsorption properties as shown by their overlapping isotherms, both of type III according to the IUPAC classification, thus corresponding to slightly porous structures with low values for specific surface area (47 and 38 m2 g-1 respectively) and pore volume (0.1 cm3 g-1) (Table 2). Nitrogen adsorbed volume is just important at high relative pressure revealing interparticular porosity. Titanate nanotubes displayed a type IV isotherm, showing a sharp increment in the adsorption branch and a clear hysteresis cycle, indicative of the presence of mesopores that are formed between nanotubes [36]. The BET surface area (154 m2 g-1) is significantly higher than that obtained for TiO2-based materials.

3.2. Aqueous Hg(II) adsorption 3.2.1. Mercury adsorption kinetics A brief kinetic study of the mercury adsorption from water was performed in order to obtain the equilibrium time and get information about the overall kinetic behavior of the TiO2 and TNT samples. Fig. 4 shows the experimental data as Hg(II) concentration against contact time for TiO2-P25 at different pH values (Fig 4a), and for every adsorbent at pH 7 (Fig 4b). For TiO2-P25 and 100 mg L−1 of Hg(II), the mercury sorption sharply increased in the early stages of the process, to slowly approach the equilibrium thereafter. The adsorption kinetics clearly depends on the pH since the amount of mercury adsorbed after 20 min was 68%, 88%, and 96% of the final value obtained at pH 2 (3.8 mg g−1), 4.5 (12.5 mg g−1) and 7 (32.0 mg g−1) respectively, therefore indicating that the adsorption process speeds up as the pH increases. At neutral pH, the amount of mercury adsorbed is higher and a significant change in the kinetics

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profile is observed in comparison to acidic conditions, with the increase of the Hg(II) uptake being much more pronounced during the first 10 min. Nevertheless, it is worthy to notice that despite the differences found in the overall behavior, a constant concentration was achieved at 1 h in all cases, matching the values determined after 24 h of stirring.

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The comparison of the results obtained for the different titania and TNT samples at the same pH (Fig 4b), also reveals the dependence of the adsorption kinetics on the nature of the surface adsorbent. The kinetic profiles show that mercury adsorption rate at pH 7 increased in the order TiO2-SG < TiO2-P25 < TNT since the time needed to get the 95% of the final concentration value was 25, 15 and 10 min, respectively. Thus, while the synthesized TiO2-SG yielded a slower process, less significant differences are revealed between TNT and the commercial titania sample. As it can be observed, the Hg(II) concentration was almost constant at around 30-40 minutes, being 1 hour time enough to completely achieve the equilibrium state for every adsorbent.

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To the best of our knowledge no studies on aqueous Hg(II) adsorption onto titanates nanotubes have been reported up-to-date. However, the equilibrium time found for the adsorption of mercury on TNT sample is analogous or even somewhat smaller than found for the adsorption of some other metals onto titanate nanotubes synthesized by different methods. As an example, Chen et al. [37] described that the adsorption of Pb(II) onto nanotubes prepared by microwave hydrothermal method reached the equilibrium within 30 minutes. Also, with titanate nanotubes synthesized by an alkaline hydrothermal method analogous to that used in the present work, Wang et al. [Error! Bookmark not defined.] reported that the adsorption of Pb(II), Cd(II) and Cr(III), was arrived at equilibrium within 60 minutes, whereas Xiong et al. [Error! Bookmark not defined.] found that Pb(II) and Cd(II) concentration was constant after 80 minutes although they stated 180 min for equilibrium time at any condition.

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In contrast, both TiO2 materials yielded equilibrium times significantly shorter than other titania samples reported in the literature. TiO2 containing 87% rutile phase and 13% anatase phase at similar Hg(II) level and high adsorbent concentration (5 g L-1) needed at least 6 hours for getting the adsorption equilibrium of mercury [38], copper and lead [39]. Hg(II) removal by adsorption with anatase TiO2 nanoparticles was also reported to be a slower process [40] and even higher equilibrium times, longer than 10 hours, were determined for the adsorption of Cu(II) [41] or Se (IV) [42] over some commercial anatase TiO2. Aqueous arsenate adsorption on TiO2 was also reported with shorter equilibrium times [43].

Two semi-empirical kinetic models based on adsorption equilibrium capacity were used to analyze the overall kinetics: a pseudo-first order proposed by Lagergren [44], and a pseudosecond order reported by Ho and McKay [45], given respectively by equations 1 and 2. The adsorption rate is related to the amount of Hg(II) adsorbed at equilibrium (qe) and that at any time t (qt), both expressed as mg g−1 according to:

dqt = k1 ( qe − qt ) dt dqt 2 = k2 ( qe − qt ) dt

(1) (2)

where k1 is the pseudo-first order kinetic constant , in min−1 units, and k2 is the pseudo-second order kinetic constant, expressed as g mg−1 min−1. Equation integration for the boundary conditions t = 0 and qt = 0 to t =t and qt=qt and rearrangement yield the corresponding lineal

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forms (eq. 3 and 4) that are commonly used to check the validity of these models and to obtain the kinetic parameters: (3) ln ( qe − qt ) = ln qe − k1 t

t 1 t = + qt k2 qe 2 qe

(4)

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Fig. 5 shows the corresponding linear plots for the kinetic data. While pseudo-first order model is not suitable to describe the kinetic profiles (Fig. 5a), rates of aqueous Hg(II) adsorption over both TiO2 and TNT samples are accurately described by the pseudo-second-order equation (Fig. 5b and 5c). Table 3 summarizes the calculated parameters qe, k2 and v0 (estimated as v0 = k2·qe2) for every adsorbent at each initial conditions. Regression coefficients R2 close to 1 and SD < 1.5 obtained from linear fits indicate a regular good correlation. Experimental results of adsorbed Hg(II) amount at equilibrium obtained after stirring 24 hours, qeexp, are also included in Table 3. As seen, these values are in very good agreement with those calculated from the kinetic model.

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For TiO2-P25, both the pseudo-second order constant and the amount of mercury adsorbed significantly increase as the solution pH rises from 2 to 7 at 100 mg Hg(II) L-1. As a result, the initial adsorption rate is strongly enhanced from 0.27 to 83.3 mg g−1min−1 with the increase of pH. The results evidence that mercury is more readily adsorbed on the titania surface at the higher pH value, what can be attributed to the dependence with pH of HgCl2 speciation in aqueous solution, as it will be discussed in the equilibrium section.

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The comparison between adsorbents shows that k2 increases in the order TiO2-SG
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Langmuir

qe =

Q0 b Ce 1 + b Ce

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Freundlich

q e = K F ⋅ Ce

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Q 0 ⋅ b ⋅ Ce n 1 + b ⋅ Ce

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reduces to the Freundlich isotherm at low sorbate concentrations whereas at high sorbate concentrations, it predicts a monolayer sorption capacity characteristic of the Langmuir isotherm. The three isotherms are expressed by the following equations for non-linear fit:

n

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where Q0 denotes the maximum Langmuir adsorption capacity, that is, the amount of metal to form a complete monolayer and b is the Langmuir constant related to the intensity of adsorption interaction. KF, and n, are the empirical constants of Freundlich equation, which are associated respectively to the adsorption capacity and the adsorption energy, being the latter a measure of the heterogeneity of surface adsorption sites. All these parameters also appear in the Sips equation maintaining their meaning. Results of the mathematical modelling are summarized in Table 4. As representative, Fig. 6 displays the isotherms obtained at pH 7 for each adsorbent as the amount of mercury adsorbed at equilibrium, qe (mg g−1), against the equilibrium mercury concentration in the liquid phase, Ce (mg L−1) along with the fitting results.

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The comparison of regression coefficients for non-linear plots (Table 4) of Langmuir and Freundlich isotherms indicates that the experimental adsorption of Hg(II) on titania and TNT materials is better fitted by the Langmuir equation (most R2 values > 0.97) rather than the Freundlich model. However, the comparison with the Sips model reveals that experimental data are even better reproduced by the latter isotherm since the corresponding R2 parameter is always higher compared with Freundlich or Langmuir equations. Moreover, as Fig. 6 demonstrates, the agreement between experimental and Sips calculated isotherms is very significant within the whole concentration range for every adsorbent. Therefore, it can be concluded that the Hg(II) adsorption on TiO2 and TNT mainly entails the formation of a superficial monolayer but there is a certain degree of heterogeneity. The influence of pH on Hg(II) adsorption is well illustrated by Fig. 7, which shows the Hg(II) isotherms for TiO2-P25, TiO2-SG and TNT materials at the different pH values over the range 2 to 10, along with the corresponding fit to the Sips model. In general, isotherms are type I or Ltype according to IUPAC [49] and Giles classifications [50] respectively, with an initial moderate slope until reaching a plateau value for the higher Hg(II) concentration. The isotherms determined for TNT at acidic pH could be classified as S-type of Giles classification what indicates that adsorption is more difficult at very low concentration. This behaviour is ascribed to monofunctional adsorbate molecules, which have moderate intermolecular attraction and meets strong competition for substrate sites with solvent or another adsorbed species [Error! Bookmark not defined.]. For each adsorbent the concentration of aqueous mercury needed to reach saturation increased with the increment of pH. As seen in Table 4, for titania samples, mercury adsorption capacities for the monolayer, Q0, rised for TiO2-P25 from 4.3 mg g−1at pH 2 to 100 mg g−1 at pH 10, and were found to be higher for TiO2-SG yielding values of 6.3 and 121 mg g−1 at pH 2 and pH 10, respectively. Two experimental works have been reported regarding the adsorption of Hg(II) on TiO2 nanoparticles. Anatase sample synthesized by Dou et al. [Error! Bookmark not defined.] yielded an adsorption maximum capacity of 101 mg/g at pH=6, while Ghassemi et al. [Error!

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Bookmark not defined.] also obtained significant Hg(II) values for the maximum adsorption of 167 mg/g at pH = 8 by TiO2 nanopaticles consisting of 87% rutile and 13% anatase. The differences with the results presented in this study could be mainly attributed to the higher surface area of the mentioned samples, 210 and 99 m2/g, respectively, which clearly display more active sites, achieving a higher mercury adsorption capacity. Regarding the titanate nanotubes, they showed an adsorption capacity higher than both titania samples at each pH, with Q0 values ranging from 14 mg g−1 at pH 2 to 140 mg g−1 at pH 10. It is worthy to notice, however, that the differences in the adsorption capacity between TNT and titania samples were found to be less pronounced as the pH was increased.

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The increasing values of the estimated Sips constants for both titania materials as the pH was rised, points to a gradually more intense adsorption of Hg(II) over the solid surface, especially at basic conditions. A more intense adsorption was also found at high pH for the TNT sample, since a significant increment in the Sips constant is observed when solution pH was modified from acidic to neutral and basic pH. The n parameter was found to decrease as a function of pH indicating that the adsorption interaction is generally more homogeneous when the pH increases.

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The comparison with bibliography evidences the major adsorption capacities found in the present study and those previously reported for other inorganic materials such as silica, alumina, iron oxides or derivatives hydroxides [Error! Bookmark not defined., Error! Bookmark not defined.-Error! Bookmark not defined.]. Moreover, it is worthy to remark the significant ability of titania and titanate nanotubes to remove mercury(II) species from water, being possible to achieve an adsorption percentage higher than 99% for pH 7 and pH 10 when the initial mercury concentration is low (10 mg·L-1). This result is especially relevant when considering that both TiO2 materials studied in the present work are scarcely porous.

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3.2.3 Comparison between adsorbents as a function of pH The pH of the solution has been recognized to be one of the factors with significant impact on the adsorption of metals onto oxides and titanates nanostructures. The experimental variations in the extent of adsorption might be related to the induced changes in both the physicochemical properties of the adsorbent surface and the chemical speciation of the metal [Error! Bookmark not defined., 18,51]. According to the speciation diagram (Fig. 8), molecular HgCl2 are the predominant species at pH 2 and pH 4.5, with 10% of HgClOH being also present at the latter pH. At pH 7 the hydroxochloro complex is the main form (50%) in the aqueous solution along with HgCl2 (17%) and Hg(OH)2 (33%), the latter found as the unique compound at pH 10. The lower Hg adsorption capacities found at pH 2 for titania and TNT samples can be therefore attributed to HgCl2 species which are characterized by their low tendency to be adsorbed on oxide surfaces [7,9]. The increase of pH to 4.5 results in a significant enhancement of Hg adsorption capacity (by a factor between 3 and 4) likely due to the presence of HgClOH whose charge asymmetry has been considered to account for its affinity for hydroxylated surfaces [18]. In order to get further information on the adsorption of HgClOH species an additional set of experiments were done with TiO2-P25 at low Hg(II) aqueous concentration, namely in the initial part of the isotherm. In this region, the amount of adsorbate loaded strongly depends on the equilibrium concentration and hence on the initial amount of the adsorbate, being the observed changes very significant. The experiments were carried out at 10 mg L-1 of Hg(II) and natural pH (4.5) modifying the TiO2 amount dispersed in solution. Unexpectedly, the percentage of mercury removal decreased from 19% to 6.5% when the titania concentration varied in the range 0.2 to 4 g L−1, despite the increasing surface area exposed. This result can be explained by means of the detected decrease in the pH value from 4.5 at the lower TiO2 concentration, i.e. 0.2 g L-1, to 3.7 for the higher one, i.e. 4 g L-1. As can be read in Fig. 8, this pH change induces a significant variation, from 10% to 4%, in the percentage of HgClOH present in solution, which is in agreement with the observed reduction on the mercury removal. The results obtained demonstrate on one side, the crucial influence of pH on the overall mercury adsorption capacity as it determines the nature of Hg

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species available to be adsorbed. On the other side, it clearly proves the much higher affinity of HgClOH in comparison with HgCl2 for titania and TNT surface. As for Hg(OH)2 species, the significant enhancement in Hg sorption found at pH 7 and 10 evidences they are also prone to be adsorbed onto both, TiO2 and titanates materials. The other important factor to explain the influence of pH on the adsorption of mercury is the nature of the adsorbent surface. It is generally accepted that Hg sorption takes place on hydroxylated sites [7, 10, 18, 52, 53] hence the hydroxylation degree of the surface would determine the overall adsorption for a given adsorbate species. According to the zero point charge of TiO2 (pH = 6-7) [64-65] at natural pH only a small fraction of Ti−OH species must be present on the titania surface. By contrast, Ti-OH and negatively charged TiO− moieties would predominate respectively, at pH 7 and 10 on the surface of both TiO2 samples, what is in accordance with the significant rise in Hg(II) uptake with increasing pH. For the different pH values, the adsorption of mercury could be described by the formation of a surface complex between the existing metal species, HgClOH and/or Hg(OH)2, and the hydroxyl groups of the titania surface according to the following reactions [18]: ≡Ti-OH + OHHgCl Ti-O(HOHgCl)− + H+ (8) ≡TiOH + Hg(OH)2 ≡TiOHg(OH)2– + H+ (9) ≡TiO– + Hg(OH)2 ≡TiOHg(OH)2– (10)

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Therefore, both the increment on the amount of mercury hydroxylated species able to be adsorbed and the available active sites on the titania surface explain the higher adsorption capacities found at high pH values. Moreover, according to the values estimated for the b and n parameters of the Sips isotherm, the interaction becomes more intense and more homogeneous as pH is increased. In agreement with the present results, a similar trend of increasing mercury sorption capacity upon enhancement of the surface hydroxylation has been also reported for other metallic oxides and hydroxides such as quartz [Error! Bookmark not defined.], gibbsite [Error! Bookmark not defined.], goethite [52], α-SiO2 (quartz) and amorphous hydrous ferric oxide [Error! Bookmark not defined.]; amorphous silica and goethite [53]; and kaolinite [Error! Bookmark not defined.].

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The comparison between TiO2-P25 and TiO2-SG revealed a higher adsorption capacity for the latter sample at any pH value (Table 5). Both TiO2 materials exhibit similar superficial area and pore volume (Table 2) so the observed differences cannot be related to their textural properties. As seen in Fig. 9, when mercury adsorption was normalized to superficial area, TiO2-SG was also found to be better adsorbent than TiO2-P25. A plausible explanation can be found on their dissimilar crystalline phase composition, as TiO2P25 sample contains 20% of rutile phase and 80% anatase, while anatase was the unique crystalline phase detected for TiO2-SG sample. Some previous works reported a relationship between crystallite phases and the amount of hydroxyl groups and water molecules adsorbed on the surface, which was reported to be smaller for rutile samples [54,55]. This factor could account for the differences observed between TiO2 P25 and TiO2-SG, as the higher content of anatase of the latter might induce a higher surface hydroxylation and then the presence of more available adsorption sites for Hg(II). Another aspect that could be related to the crystalline phase composition is the chemical nature of the dissimilar surface planes exposed for adsorption on the TiO2 surface. As an example, G. Martra studied by FT-IR the adsorption of probe molecules on two titania samples (TiO2 P25 and pure TiO2 anatase Merck) demonstrating that the nature and strength of acidic and basic centers were highly dependent on the exposed faces.[56]. On the other hand, it has been also pointed that the mixing of anatase and rutile phases might also cause crystalline defects and disorder that could inhibit adsorption [57] what is in

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agreement with the lower adsorption capacity found for TiO2-P25. Moreover, two other aspects related to the synthesis should also be taken into account: the probable presence of amorphous phase in the TiO2-SG that would increase the adsorption sites, yielding an enhancement of the overall adsorption capacity, and the possible contribution of organic traces remaining in the surface due to the preparation procedure of the TiO2-SG sample.

an

us

cr

ip t

Regarding the behavior of titanate nanotubes for adsorption, it has been argued that the removal of metallic cations such as Pb2+, Cd2+, Cu2+, and Cr3+ is due to the ion exchange mechanism, with Na+ and even H+ as ions involved [Error! Bookmark not defined., Error! Bookmark not defined., 58,59,60]. However, nanotubes are also able to adsorb neutral and anionic molecular species such as arsenite or arsenate compounds, thus pointing to a surface complexation mechanism instead of an electrostatic interaction [61]. Trititanate structures consist of disorder TiO6 octahedra, with unshared oxygen atoms forming negative layers, among which the cationic species Na+ or H+ are located, so that both chemical sites –Ti-ONa and -TiOH can be found [Error! Bookmark not defined.]. Besides, due to the open nanotubes morphology, the TNT sample exhibits a high surface area thus displaying a high number of active sites. Taking into account the speciation of Hg(II) in solution (Fig. 8), the adsorption of mercury on the nanotubes material is expected to take place by a surface complexation mechanism through reactions analogue to those above described for titanium dioxide (Eq. 8-10).

d

M

Comparing the adsorption performance between TNT and TiO2 samples, the results revealed a greater adsorption capacity for the former material explained by its higher surface area which provides more active sites (Table 2). The net differences become gradually smaller as the pH increased, indicating that the nature and/or total amount of active groups available on TNT and TiO2 surfaces were similar at basic pH. Actually, when adsorption capacity was normalized to the surface area (Fig. 9), mercury uptake increased in the order TNT< TiO2-P25 < TiO2-SG demonstrating that the number of sorption sites available per area unit is smaller for TNT, thus being the TiO2-SG the more active material.

Ac ce p

te

The point of zero charge (PZC) of titanate nanotubes has been reported to be in the pH range 3 to 5 [Error! Bookmark not defined., Error! Bookmark not defined., Error! Bookmark not defined.], value significantly lower than that of TiO2 (pH = 6-7) [62,63]. Therefore, the pH range where TiOH and TiO−Na+ groups are present in the surface might be larger for TNT. Nevertheless, TiO−Na+ sites might be expected to be predominant over TiOH as the pH increases which could be the reason for the smaller adsorption capacity per area unit exhibited by TNT sample in comparison to TiO2 materials. Accordingly, adsorption of Hg(II) over TNT surface could take place mainly on TiOH active sites, while ionic exchange with sodium cations would not be the principal mechanism of adsorption. As a proof, the X-ray diffraction pattern of titanate nanotubes (not shown) remained unaltered after adsorption of Hg(II). This is in contrast with the intensity reduction of interlayer distance reflection previously reported when ionic metals such as Cu2+ or Pb2+ were adsorbed by ion exchange of Na+ of titanate nanotubes [Error! Bookmark not defined., Error! Bookmark not defined.].

1. Conclusions The adsorption of Hg(II) from aqueous solution over TiO2 (TiO2-P25 and TiO2-SG) and titanate nanotubes was investigated in this work. The kinetic and equilibrium results revealed that the adsorption of mercury species was significantly dependent on the solution pH as well as on the type of adsorbent. The kinetics of mercury adsorption was well fitted by a pseudo-second order model in the studied conditions. The adsorption rates followed the order TiO2-SG << TNT < TiO2-P25, indicating that porosity is not a critical variable in the global kinetics since on one side, both titania samples displayed appreciable differences although exhibited similar textural properties. On the other side, the adsorption rate for TNT sample was just slightly smaller than TiO2 -P25 despite the large difference between the superficial areas of both samples.

11 Page 11 of 25

ip t

Equilibrium isotherms of Hg(II) were best described by the Sips model for every adsorbent and pH. It was found that maximum adsorption capacities of Hg(II) strongly increased as pH rised, with higher values for titanate nanotubes. As for titania samples, TiO2-SG exhibited the highest mercury uptake. At pH 10, a maximum Hg(II) adsorption capacity of 100, 121, and 140 mg g−1 was achieved for TiO2-P25, TiO2-SG, and TNT, respectively. However, when maximum capacities were normalized to the surface area, titanate nanotubes showed smaller adsorption capacities than titania samples. Remarkably, titania and titanate nanotubes showed a significant ability to remove Hg(II) species from water, achieving adsorption percentages higher than 99% for pH 7 and pH 10 at low mercury concentration, despite the scarce porosity of both TiO2 materials.

us

cr

Adsorption was explained taking into account that hydroxylated mercury species (HgClOH and Hg(OH)2) are effectively adsorbed on the materials probably through the interaction with – TiOH and –TiO− active sites available on the TiO2 or TNT surface. For titanate nanotubes it was proposed that Hg(II) could be adsorbed on TiOH active sites mainly through surface complexation rather than by an ionic mechanism.

M

an

Acknowledgements Authors thank the Spanish Government for the financial support through the projects CTM2009-08649 and CTM2012-34988 and also thank Regional Government of Madrid that also financed this research through the Project 2013-MAE2716 “REMTAVARES”.

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Fig. 1 X-Ray diffraction patterns of a) TiO2-P25; b) TiO2-SG; and c) TNT samples. Fig. 2 SEM images of (a) commercial TiO2-P25, and (b) synthesized TiO2-SG. (c) TEM micrograph of synthesized titanate nanotubes. Fig. 3 Nitrogen adsorption-desorption isotherms at 77 K corresponding to a) TiO2-P25; b) TiO2 sol-gel; and c) titanate nanotubes. Fig. 4 Hg(II) concentration kinetic profiles found for (a) TiO2-P25 at different pH and (b) titania and titanate adsorbents at pH 7. Lines are the estimated values from the pseudo-second order fits.

15 Page 15 of 25

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Fig. 5 Linear fit of the kinetic data: (a) pseudo-first-order model and (b) pseudo-second-order model for TiO2-P25 at different pH (2, 4.5 and 7) and (c) pseudo-second-order model for TiO2 P25, TiO2-SG, and TNT at pH 7. Fig. 6 Experimental adsorption isotherms obtained at pH 7 for (a) TiO2-P25, (b) TiO2-SG, and (c) TNT along with the fitting results using the Langmuir, Freundlich, and Sips models. Fig. 7 Hg(II) equilibrium adsorption isotherms obtained for (a) TiO2-P25, (b) TiO2-SG and (c) TNT. Solid lines correspond to the estimated Sips model. Fig. 8 Speciation diagram for aqueous mercury chloride at 100 mg L−1as a function of pH. Fig. 9 Maximum adsorption capacity of Hg(II) per area unit for TiO2 -P25, TiO2-SG, and TNT at different pH.

Reaction

Log K H2O + HgO ↓

Hg + OH

−

HgOH

3.6

+

10.9

Hg2+ + 2 OH−

Hg(OH)2

22.3

Hg2+ + 3 OH−

Hg(OH)3−

21.46

−

Hg + Cl + OH Hg2+ + Cl−

−

10.44

HgClOH

HgCl+ HgCl2

Hg2+ + 3 Cl−

HgCl3−

Hg2+ + 4 Cl−

HgCl42−

14.0

14.2

15.5

te

d

Hg2+ + 2 Cl−

7.3

M

2+

us

2+

an

Hg(OH)2

cr

Table 1 Stability constants of mercury species used in the MINEQL+ speciation calculations

Table 2 Textural properties of TNT and TiO2 materials

Ac ce p

Sample TiO2-P25 TiO2-SG TNT

SBET (m2g-1) 47 38 154

VP (cm3g-1) 0.10 0.11 0.38

Table 3 Estimated parameters from pseudo-second order modelling of kinetic adsorption data for TiO2-P25 (pH values: 2, 4.5, 7), TiO2-SG and TNT at pH 7. Adsorbent

pH

k2 (g mg-1 min-1) qe (mg g-1)

qeexp (mg g-1)

v0 (mg g-1·min-1)

R2

SD

TiO2-P25

2

0.015

4.26

3.82

0.27

0.9942

1.410

TiO2-P25

4.5

0.031

12.19

12.53

4.6

0.9990

0.217

TiO2-P25

7

0.085

31.25

31.96

83.3

0.9990

0.034

TiO2-SG

7

0.013

40.00

39.25

21.3

0.9999

0.018

TNT

7

0.024

52.63

52.86

66.7

0.9999

0.010

Table 4 Parameters and regression coefficients obtained from non-linear modelling of equilibrium experimental results. pH Freundlich

Langmuir

Sips

16 Page 16 of 25

KF

R2

n

(mg1-n Ln g-1)

b

(mg g-1)

(L mg-1)

n

R2

0.8

0.33

0.813

5.3

0.03

0.910

4.3

0.0002

2.5

0.997

4.5 1.9

0.39

0.862

15

0.04

0.953

13

0.003

2.0

0.981

7

4.4

0.51

0.924

71

0.02

0.976

55

0.005

1.5

0.997

10

22

0.31

0.951

86

0.17

0.972

100

0.196

0.73

0.980

2

0.59

0.47

0.854

8.5

0.02

0.931

6.3

0.00093

2.0

0.971

4.5 1.51

0.56

0.952

33.6

0.02

0.977

25

0.0036

1.6

0.980

7

5.7

0.58

0.950

96.3

0.03

0.979

72

0.013

1.5

0.987

10

20

0.40

0.971

98.6

0.13

0.982

121

0.14

0.8

0.987

2

2.4

0.34

0.861

16

0.03

0.950

13

0.0035

1.8

0.985

4.5 1.6

0.69

0.912

74

0.01

0.940

39

0.00017

2.5

0.994

7

68

0.44

0.886

109

2.4

0.973

92

8.4

1.6

0.986

10

77

0.49

0.960

126

2.0

0.978

140

1.5

0.89

0.984

M

an

TNT

(L mg-1)

Q0

ip t

TiO2-SG

(mg g-1)

R2

cr

TiO2-P25

b

us

2

Q0

Figure 1 a)

a

a a

aa

ra a r

r

b) a

Ac ce p

a

a

a

te

r

Intensity (a.u.)

ar a

a

d

a: anatase r: rutile

aa

a

a

a a

a aa

a

c)

10

20

30

40

50

60

70

80

90

2θ (degrees)

17 Page 17 of 25

Figure 2

us

cr

ip t

(a)

te

Ac ce p

(c )

d

M

an

(b)

18 Page 18 of 25

Figure 3 300 TNT TiO2-SG TiO2-P25

200

ip t

V ads (cm3/g STP)

250

150 100

0 0.0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

cr

50

1.0

Ac ce p

te

d

M

an

us

P / Po

19 Page 19 of 25

Figure 4 100

(a)

(b)

90

90

80

80

TiO2-P25

70

70

TiO2-SG

60

TNTs

-1

C (mg Hg(II) L )

50 40 30 pH 2 pH 4.5 pH 7

20 10

50

ip t

60

40 30 20

cr

-1

C (mg Hg(II) L )

100

10 0

0 20

40

60

80 100 120 140 160 180

0

20

40

60

80 100 120 140 160 180

us

0

Adsorption time (min)

Ac ce p

te

d

M

an

Adsorption time (min)

20 Page 20 of 25

Figure 5 4 (a)

3

pH 2 pH 4.5 pH 7

2

0

ip t

Ln (qe-qt)

1

-1 -2

cr

-3 -4 0

25

50

75

100 125 150 175 200 225 250

us

Adsorption time (min)

50 (b)

30

an

pH 2 pH 4.5 pH 7

20

M

t/qt (min·g/ mg)

40

10

0 25

50

75

100 125 150 175 200 225 250

d

0

Adsorption time (min)

(c)

TiO2-P25

7

TiO2-SG TNT

6 5

Ac ce p

t/qt (min·g/ mg)

te

8

4 3 2 1 0

0

25

50

75

100 125 150 175 200 225 250

Adsorption time (min)

21 Page 21 of 25

Figure 6. 60 (a) 50

20

ip t

30

Experimental Freundlich Langmuir Sips

10

cr

qe (mg/g)

40

0 0

20

40

60

80

100

120

140

160

us

Ce (mg/L)

70 (b) 60

an

40 30 Experimental Freundlich Langmuir Sips

M

qe (mg/g)

50

20 10 0 10

20

30

40

Ce (mg/L)

90

60

70

te

100

50

d

0

(c)

80

Ac ce p

70

qe (mg/g)

60 50 40 30 20

Experimental Freundlich Langmuir Sips

10

0 0.0

0.5

1.0

1.5

2.0

2.5

Ce (mg/L)

22 Page 22 of 25

Figure 7 90

90 pH 2 pH 4.5 pH 7 pH 10

pH 2 pH 4.5 pH 7 pH 10

70

60

60

50

50

qe (mg/g)

qe (mg/g)

70

(b)

80

40 30

40 30

20

20

10

10

0

ip t

80

0 0

20

40

60

80

100 120

140 160 180 200

0

20

40

60

120

(c.1)

(c.2) 100

30

140 160 180 200

qe (mg/g)

20 15

pH 2 pH 4.5

5 0

60 40

M

10

an

80

25

qe (mg/g)

100 120

Ce (mg/L)

40 35

80

us

Ce (mg/L)

cr

(a)

pH 7 pH 10

20 0

0

20

40

60

80

100

120

140

160

180

0.0

Ce (mg/L)

0.5

1.0

1.5

2.0

2.5

3.0

Ac ce p

te

d

Ce (mg/L)

23 Page 23 of 25

Figure 8 100 90 80 HgCl2 (aq.)

70 50

ip t

% Total

+

HgCl (aq.) HgClOH (aq.) Hg(OH)2 (aq.)

60

−

Hg(OH)3 (aq.)

40 30

cr

20 10 0 2

4

6

8

10

12

14

us

0

Ac ce p

te

d

M

an

pH

24 Page 24 of 25

Figure 9. 4 TiO2-P25 TiO2-SG

3

ip t

Q0/area (mg / m2)

TNT

2

0

2

4

6

8

10

Ac ce p

te

d

M

an

pH

us

cr

1

25 Page 25 of 25