Journal of Colloid and Interface Science 440 (2015) 84–90
Contents lists available at ScienceDirect
Journal of Colloid and Interface Science www.elsevier.com/locate/jcis
Adsorption, kinetic and thermodynamic studies for manganese extraction from aqueous medium using mesoporous silica Salah Ali Mahgoub Idris Chemistry Department, Faculty of Science, University of Omar Al-Mukhtar, Tobruk, Libya
a r t i c l e
i n f o
Article history: Received 26 July 2014 Accepted 15 October 2014 Available online 29 October 2014 Keywords: Modified silica Water Manganese Adsorption Kinetic Thermodynamic
a b s t r a c t This paper describes studies of functionalized mesoporous silica employed as adsorbent for Mn(II) from aqueous solutions. The surface area of MCM-41 and diethylenetriamine functionalized-MCM-41 used in this study were 760 and 318 m2 g1 (N2 adsorption). A strong dependence on pH in the Mn(II) adsorption capacity and best results were obtained at pH 6.5–7. The adsorption onto the diethylenetriamine functionalized-MCM-41 followed the pseudo-second-order kinetic model and the highest reaction rate 0.324 min1 was observed at low initial concentration 10 ppm. The equilibrium data showed excellent correlation with the Langmuir isotherm model and the maximum adsorption capacity of Mn(II) reached 88.9 mg/g for DETA-MCM-41 indicating that the adsorption occurs on a homogeneous surface by monolayer sorption without interaction between the adsorbed ions. These data contribute to the understanding of mechanisms involved in mesoporous silica and provide some practical clues to improve the adsorption efficiency (uptake capacity and kinetics) of Mn(II) ions. Ó 2014 Elsevier Inc. All rights reserved.
1. Introduction Manganese (Mn) is an element that is present in the environment. It is commonly found in drinking water and is essential for human health at low concentrations [1]. However, excessive concentrations of Mn(II) could result in metallic tasting water and many health problems. Manganese is present in ground water as a divalent ion Mn++ and is considered a pollutant mainly because it may causes nervous system damage, leading to Parkinson’s disease, and injures arterial breastwork and the myocardium. The World Health Organization has set a guideline of 0.4 mg/L for Mn(II) intake control from drinking water [2]; and the limit value in the Chinese Standard for Drinking Water Quality is 0.1 mg/L [3]. In neutral and alkaline environments, Mn(II) is difficult to be oxidized [4]. Conventional treatment for Mn(II) removal generally requires the use of strong oxidizing agents such as potassium permanganate, chlorine, hypochlorite, chlorine dioxide or ozone [5]. The generation of other pollutants or toxins is the major disadvantage of these oxidizers. Although chlorine or KMnO4 oxidation combined with sand filtration is usually applied for Mn(II) removal, its concentration in the treated water is still difficult to meet the accepted levels requirements for drinking water [6]. In fact, as the effects of pollutants accumulation and sedimentation, the concentration of Mn(II)
E-mail address: [email protected] http://dx.doi.org/10.1016/j.jcis.2014.10.022 0021-9797/Ó 2014 Elsevier Inc. All rights reserved.
is variable, and the variation of water quality such as solution pH may cause a sudden release of Mn(II) [7]. To develop extraction method for removal of Mn(II) from solution, in the presence of other heavy metals, functionalized mesoporous silica nanoparticles have been explored due to their high metal uptake capacities, good physical and chemical stabilities, uniform structures made of mesopore channels regularly organized in space [8–10]. Summaries of some reported studies of Mn(II) ions adsorption for different materials are presented in Table 1. Although these adsorption capacities correspond to different experimental conditions, they are representative of the tendency to uptake Mn(II) ions. Silica based adsorbents are of particular interest as they exhibit enhanced accessibility to active centers [13,14], and fast mass transfer rates in the regular porous environment [15,16]. These combination of factors have led to high capacity adsorbents [16]. To this end no papers reported the results of kinetic studies and thermodynamic results on Mn(II) adsorption onto mesoporous silica. The aim of this work was to study the adsorption of Mn(II) ions onto amino-functionalized mesoporous silica. The linearized Langmuir and Freundlich equations were used to fit the equilibrium isotherms. Kinetic data were analyzed by pseudo-first, pseudosecond and intra-particle diffusion models to obtain the kinetic constants and thermodynamic study of Mn(II) removal and postulate the mechanism of Mn(II) removal by porous materials. This
S.A.M. Idris / Journal of Colloid and Interface Science 440 (2015) 84–90
work is part of a series of adsorption studies in well characterized mesoporous silica (MCM-41) [17,18]. 2. Experimental 2.1. Materials and reagent Cetyltrimethylammonium bromide (CTAB), 98%, was supplied by BDH. Aqueous ammonia (30% NH3), tetraethoxysilane (TEOS) 98%, hydrogen peroxide solution (30 wt.%), N-(3-trimethoxysilylpropyl) diethylenetriamine (DETA-TMS), ethanol absolute, 1000 lg mL1 of Mn(II) standard solution and toluene (+99%) were purchased from Sigma Aldrich. Nitric acid (65 wt.%) was provided by Fisher Scientific. Glassware was soaked in 5% HNO3 overnight and cleaned with deionised water before use. All products were used as supplied and deionised water was used throughout this work.
85
2.3. Characterization and Analysis The surface areas of the MCM-41 were measured using nitrogen physisorption isotherms. The Brumauer–Emmett–Teller (BET) surface areas were calculated using experimental points at a relative pressure (P/P0) of 0.05–0.25. The total pore volume was calculated from the N2 amount adsorbed at the P/P0 of 0.99 for each sample. Both samples exhibited a Type IV adsorption isotherm typical of mesoporous solids. Desorption isotherms were used to calculate the pore diameters. Elemental analysis (EA) was carried out using an Exeter Analytical CE440 elemental function. Total concentration of Mn(II) in water samples was determined using a PerkinElmer AAnalyst200 flame atomic absorption spectrometry (FAAS) instrument. Calibration standards (0.0, 0.1, 0.5, 1, 1.5 and 2 lg cm3) were prepared in 5% HNO3. Method parameters used are shown in Table 2. The method detection limit (MDL) for Mn(II) was 0.015 lg cm3. 2.4. Adsorption process and effect of pH
2.2. Mesoporous silica preparations and functionalization The synthesis method and functionalization MCM-41 was prepared according to the method reported in Ref. [17]. Approximately 9.0 g of CTAB was dissolved under slight warming (35 °C) in a mixture of 210 mL of distilled H2O and 100 mL of aqueous NH3. To this clear solution, 50 mL of TEOS was slowly added under stirring. After further stirring for 4 h, the gel was aged at room temperature for 48 h in a closed container. The product was obtained by filtration, washed with 800 mL of distilled H2O, and dried in air at room temperature. To remove the surfactant microwave digestion (MWD) was performed by using a MARS 5 microwave digestion system (CEM Corporation, Buckingham, UK) was used at an operating power of approximately 1600 W. The pressure and temperature inside the microwave were controlled to be lower than 1.3 MPa and 200 °C, respectively. Samples (approx. 0.1 g) were added to multiple Teflon vessels to which 1.5 mL of HNO3 and 0.70 mL of H2O2 were added. Microwave digestion was operated at a working frequency of 2450 MHz and 220 V for 15 min. The product (MCM-41) was filtered, washed with copious amounts of distilled H2O and dried at 100 °C for 2 h. Surface modification of MCM-41 was carried out by condensation using the N-(3-trimethoxysilylpropyl) diethylenetriamine with MCM-41 (Scheme 1). Briefly, approximately 5 g of MCM-41 was pre-treated at 140 °C for 2 h before being immersed in 50 cm3 of toluene and 10 cm3 of DETA-TMS, in a 250 cm3 flask. The mixture was refluxed for 4 h and the solid produced was filtered, washed with 100 cm3 ethanol, and oven-dried at 80 °C for 2 h to produce a DETA-MCM-41 sorbent. Table 1 Manganese ions adsorption capacity for different materials. Material
Mn(II) uptake (mg g1)
References
Clinoptilolite–Fe system Clinoptilolite from Turkey Na–montmorillonite Granular activated carbon Nigerian kaolinite clay Manganese oxide coated zeolite
27.12 4.22 3.22 2.54 111.11 27.5
[11] [12] [2] [4] [6] [7]
Approximately 10 mg of functionalized MCM-41 was suspended in 20 mL of solution containing 20 lg mL1 Mn(II) ions in 50 mL beaker and the solution was stirred (250 rpm) using MyLab (Magnetic Stirrer with Hot Plate-SLMSH300) for approximately 2 h. After this time the solution was removed and analyzed for Mn(II) by flame atomic absorption spectrometry (FAAS). This process was repeated until saturation of the sorbent as indicated by a measurement of Mn(II) in the solution aliquot. The extraction was examined at various pH values (between 1 and 9) with solution modification achieved via the addition of small amounts of 1 M ammonium hydroxide. 2.5. Adsorption isotherms The Langmuir [19] or Freundlich [20] models were applied to measured data to study adsorption isotherms. Solutions containing initial concentrations of Mn(II) at 10, 50, 100 or 200 lg cm3 were prepared. To each solution 0.05 g of sorbent was added and the solution was stirred at 250 rpm for 120 min at room temperature, solutions were adjusted to provide a pH of 7. The amounts of Mn(II) extracted at equilibrium, qe (mg/g) were calculated according to Eq. (1):
qe ¼
C0 Ce V W
ð1Þ
where C0 and Ce (mg/g) were the liquid phase initial and equilibrium concentrations of the Mn(II) respectively. V was the volume of the solution (cm3), and W was the mass of sorbent (g) used [21]. The sorption equilibrium data were analyzed according to Langmuir Eq. (2) and Freundlich Eq. (3) isotherm models [22].
Ce 1 1 þ Ce ¼ qm b qm qe ln qe ¼ ln K f þ
1 ln C e n
ð2Þ
ð3Þ
where qe and Ce were the equilibrium concentrations of the Mn(II) ions in the adsorbed and liquid phases in mg/g and mg/L, respec-
Scheme 1. Modification of MCM-41 using N-(3-trimethoxysilylpropyl) diethylenetriamine.
86
S.A.M. Idris / Journal of Colloid and Interface Science 440 (2015) 84–90
The intraparticle diffusion model: To study the mechanism of the particle diffusion the Morris–Weber equation was applied:
Table 2 Conditions used in FAAS. Metal
Mn
Wave length (k) nm Slit (mm) Sample flow rate (mL/min) Current reading value (mA) Lamp current used (mA) Optimum fuel flow rate (mL/min)
279.48 1.8/0.6 7.2 30 10/12 3
pffiffi qt ¼ kid t
ð7Þ
where qt is the amount of metal ion sorbed (mg L1) at time t and kid is the intraparticle diffusion rate coefficient (mg L1 min1/2). 2.7. Thermodynamic study
tively. qm (mg/g) and b (L/mg) were the Langmuir constants. Whereas the qm is the maximum monolayer capacity and b was the adsorption affinity onto the adsorption. Kf (mg/g) and n (L/ mg) were the Freundlich constants which are related to the sorption capacity and intensity, respectively. The Langmuir and Freundlich constants were calculated from the slope and intercept of the linear plot obtained from Eqs. (2) and (3) respectively. For predicting the favorability of an adsorption system, the Langmuir equation can also be expressed in terms of a dimensionless separation factor (RL) by using the Langmuir constant b and the initial concentrations of the Mn(II) (Eq. (4)).
RL ¼
1 1 þ C0 b
ð4Þ
When, RL > 1, RL = 1, 0 < RL < 1 and RL = 0, indicates unfavorable, linear, favorable and irreversible, adsorption isotherms, respectively [23,24].
Kinetic studies to determine the rate of Mn(II) removal from water samples were conducted for DETA-MCM-41. Solutions were prepared with the same initial Mn(II) concentration of 10 lg cm3 and stirred with 0.05 g of each sorbent at 250 rpm. 25 cm3 aliquots of each solution were stirred at 25 °C for 1, 5, 10, 20, 30 or 40 min. After each time period, solutions were filtered and analyzed by FAAS to determine the concentration of Mn(II) in the final solution. The kinetics of Mn(II) adsorption onto the surface of the silica nanoparticles were analyzed using pseudo first-order [25], pseudo second-order [26] and intraparticle diffusion [27,28] kinetic models. The conformity between experimental data and the model predicted values was expressed by the correlation coefficients (R2). A relatively high R2 value (close or equal to 1) was used to indicate best fit to the kinetic model. The pseudo first-order equation: The pseudo first-order equation is generally presented as follows:
k1 t 2:303
ð5Þ
where qe and qt are the adsorption capacity at equilibrium and at time t, respectively (mg g1), k1 is the rate coefficient of pseudo first-order adsorption (L min1). The plot of log(qe qt) vs. t should give a linear relationship from which k1 and qe can be determined from the slope and intercept of the plot, respectively. The pseudo second-order equation: The pseudo second-order adsorption kinetic rate equation is expressed as:
t 1 1 ¼ þ t qt k2 q2e qe
DG0 ¼ RT ln K c
ð8Þ
where Kc is the adsorption equilibrium constant, R is the gas constant and T is the absolute temperature (K). The adsorption equilibrium constant (Kc) can be calculated from:
Kc ¼
Fe 1 Fe
ð9Þ
where Fe is the fraction attainment of Mn(II) ion adsorbed at equilibrium time, and is obtained by the expression
2.6. Adsorption kinetics study
logðqe qt Þ ¼ logðqe Þ
Thermodynamic studies for Mn(II) removal from water samples were conducted for DETA-MCM-41. Solutions were prepared with the same initial Mn(II) concentration of 50 lg cm3 and stirred with 0.05 g DETA-MCM-41 at 250 rpm. 25 cm3 aliquots of each solution were stirred at 25, 35, 45 and 55 °C for 1, 5, 10, 20, 30 or 40 min. After each time period, solutions were filtered and analyzed by FAAS to determine the concentration of Mn(II) in the final solution. The free energy (DG0) of the adsorption reaction is given by the following equation:
ð6Þ
where k2 is the rate coefficient of pseudo second-order adsorption t and t, should give a linear relationqt
(g mg1 min1). The plot of
ship from which qe and k2 can be determined from the slope and intercept of the plot, respectively.
Fe ¼
C0 Ce C0
ð10Þ
where C0 and Ce are the initial and equilibrium concentrations of Mn(II) ion in solution (mg/L). The value of the adsorption equilibrium constant (Kc) for the adsorption of Mn(II) ion on the adsorbent were calculated at different temperature and at equilibrium time using Eqs. (9) and (10). The Gibbs free energy can be represented as follows:
DG0 ¼ DH0 T DS0
ð11Þ
The values of enthalpy change (DH0) and entropy change (DS0) calculated from the intercept and slope of the plot of DG0 vs. T [29]. The activation energy for Mn(II) adsorption was calculated by the Arrhenius equation DE k2 ¼ AeðRT Þ
ð12Þ
where DE is the activation energy (kJ/mol), A is the frequency factor, T is the absolute temperature (K), and R is the gas constant. From the plot of ln(k2) vs. 1/T, the activation energy DE for the adsorption of Mn(II) can be calculated. 3. Results and discussion 3.1. Materials characterization The MCM-41 and functionalized-MCM-41 materials were characterized using BET and the physicochemical properties of both are summarized in Table 3. The N2 sorption isotherms (Fig. 1) were type IV for both samples confirming their mesoporous nature, however the quantities of nitrogen gas adsorbed with MCM-41 was higher compared with DETA-MCM-41 due to the surface of DETA-MCM-41 was occupied with the functional groups. Elemental analysis was used to estimate the amount of molecules (L0) attached to functionalized samples from the percentage
S.A.M. Idris / Journal of Colloid and Interface Science 440 (2015) 84–90
87
Table 3 Physicochemical properties of MCM-41 and DETA- MCM-41. Sample Name
BET surface area (m2 g1)a
Pore size (nm)b
Pore volume (cm3 g1)c
MCM-41 DETA- MCM-41
760 318
6.74 6.18
0.99 0.63
a Calculated by the linear BET equation from sorption data in a relative pressure range from 0.05 to 0.25. b Calculated by the BJH model from the adsorption branches of isotherms. c Calculated from N2 amount adsorbed at a relative pressure P/P0 of 0.99.
Fig. 2. Adsorption capacities of functionalized MCM-41 for Mn(II) extraction over different pH values.
Fig. 1. Nitrogen adsorption isotherms of MCM-41 and DETA-MCM-41.
Table 4 Elemental analysis data recorded for the MCM-41. Silica
%C
%H
%N
L0 (mmol/g)a
MCM-41 DETA-MCM-41
Trace/Nil 16.8
0.65 4.72
Trace/Nil 10.8
– 2.57
Fig. 3. Chemical diagram species for Mn(II) aqueous solution. C0 = 1.82 mM Mn(II).
a Functionalisation degree (L0 = millimoles of ligand per gram of functionalised silica).
of nitrogen, in the functionalized mesoporous silica [30], using Eq. (13):
L0 ¼
%N 10 nitrogen atomic weight
ð13Þ
The calculated L0 (Table 4) values for DETA-MCM-41 were 2.57 mmol of ligand per gram of functionalized silica indicating the successfully functionalization for diethylenetriamine functional groups. 3.2. Effect of pH on Mn(II) adsorption The pH of the aqueous solution is an important controlling parameter in the adsorption processes and metal ion removal usually increases with increasing pH values [31]. This Fig. 2 shows that the amount of Mn(II) ions adsorbed on DETA-MCM-41 increases with the increase of the pH. This phenomenon appears to be due to the fact that functional group and the surface of mesoporous silica were highly selective for H3O+ ions when the H3O+ ions concentration was high. Thus, at
lower pH values the H3O+ ions compete with Mn(II) ions for the exchange sites in DETA-MCM-41 leading to a low removal of the metal ion. Furthermore, the extensive repulsion of metal ions due to protonation of active sites of DETA-MCM-41 surface at lower pH may be another reason for decrease in Mn(II) ions adsorption at pH < 5. The increase in Mn(II) ion removal with the increases in pH can be explained on the basis of a decrease in competition between protons and Mn(II) cations for the same functional groups and by the decrease in positive surface charge, which results in a lower electrostatic repulsion between the surface of DETA-MCM-41 and Mn(II) ions before ion-exchange. Adsorption results showed an optimum pH value of about 7 (equilibrium pH fluctuated between 6.5 and 7.0) which is also important in the disposal of the treated effluents (required for neutrality of the bodies ‘‘receivers’’ water or water reuse). The decrease in adsorption at higher pH (above pH 7), is probably due to the formation of insoluble Mn(II) hydroxide Mn(OH)2 (Fig. 3) or the pH may affect the ionization degree (species formation) of the adsorbate and the surface property of the adsorbent [31]. For example, the heavy metal ions may form a complex with inorganic ligands such as OH, changing its adsorption degree. The extent of these compounds formation varies with pH, ionic composition and the adsorbate ion. Results obtained showing that the amount of Mn(II) extracted from water using amino functionalized MCM-41 can be
88
S.A.M. Idris / Journal of Colloid and Interface Science 440 (2015) 84–90
Table 5 Isotherm parameters for Mn(II) sorption on functionalised mesoporous silica. Adsorbents
Langmuir
DETA-MCM-41
Freundlich 2
qm (mg/g)
b (L/mg)
RL
R
Kf (mg/g)
n (L/mg)
R2
88.9
8.34
0.55
0.9998
26.4
4.2
0.7889
Fig. 5. The intraparticle diffusion model for Mn(II) adsorption onto DETA-MCM-41.
isotherms are the most commonly used isotherms for different adsorbent/adsorbate systems to explain solid–liquid adsorption systems and to predict their equilibrium parameters [32–34]. The relevant parameters for these isotherms are presented in Table 5. As seen from Table 5, the R2 values obtained from the Langmuir model are much closer to one than are those from the Freundlich model, suggesting that the Langmuir model is better than the Freundlich isotherm. Thus the adsorption can be described by the Langmuir isotherm and the metal ion adsorption occurs on a homogeneous surface by monolayer sorption without interaction between the adsorbed ion [35] and the adsorption capacity is 88.9 mg/g. 3.4. Adsorption kinetics Fig. 4. Pseudo-first order (a) and Pseudo-second order (b) kinetics models of Mn(II) onto DETA-MCM-41 sorbent.
influenced by changing the pH of aqueous medium due to ionization degree (species formation) of the Mn(II) and the functional groups property attached to the surface of mesoporous silica. 3.3. Adsorption isotherms The adsorption isotherms were studied to find the relationship between equilibrium adsorption capacity and equilibrium concentration at a certain temperature. Langmuir and Freundlich
The experimental kinetic data were fitted using a pseudofirst-order kinetic model and pseudo-second-order kinetic model (Fig. 4). The results are shown in Table 6. It can be seen that the obtained R2 values of the pseudo-second-order model (>0.996) were better than R2 of the pseudo-first – order model (0.949–0.993), suggesting that the adsorption process is second-order. Moreover, the calculated qe values were much closer to the experimental values in the pseudo-second-order kinetic model than the pseudofirst-order kinetic model indicating that the adsorption process is second-order. As seen in Table 6, when the initial ion concentration increases from 10 to 200 lg/ml, the pseudo-second-order constants
Table 6 Kinetic parameters for the adsorption of manganese on the adsorbent. Adsorbents
DETA-MCM-41 DETA-MCM-41 DETA-MCM-41 DETA-MCM-41
C0 (lg/mL)
10 50 100 200
qe (exp) (mg/g)
5.00 25.00 49.78 88.45
The pseudo first-order
The pseudo second-order
k1 (min1)
qe (cal) (mg/g)
R2
k2 (min1)
qe (cal) (mg/g)
R2
0.164 0.219 0.130 0.156
2.07 11.49 29.90 58.26
0.9803 0.9489 0.9930 0.9769
0.324 0.052 0.014 0.008
5.07 25.58 51.29 90.98
0.9996 0.9986 0.9976 0.9969
89
S.A.M. Idris / Journal of Colloid and Interface Science 440 (2015) 84–90 Table 7 Thermodynamic parameters for the adsorption of Mn(II) on the adsorbent. Adsorbents
Reaction temp. (K)
Kc
DG0 (kJ/mol)
DH0 (kJ/mol)
DS0 (J/mol K)
DE0 (kJ/mol)
DETA-MCM-41 DETA-MCM-41 DETA-MCM-41 DETA-MCM-41
298 308 318 328
12,499 16,666 24,999 49,999
23.37 24.89 26.77 29.51
37.34
202.8
4.07
Fig. 6. (a) Relationship between Gibbs free energy change and temperature of adsorption of Mn(II) onto the adsorbent. (b) Arrhenius plot for Mn(II) adsorption on the adsorbent.
(k2) decrease from 0.324 to 0.052, 0.014 to 0.008 for DETA-MCM-41 which indicate that the active sites on the surface of mesoporous silica were occupied by Mn(II). Adsorption capacity is 90.98 mg/g which is high and that maybe favorability of Mn(II) to bonded to hard ligand as Mn(II) is hard metal and also may be due to the adsorption of Mn(II) inside pores in addition to the surface adsorption. Fig. 5 shows the intraparticle diffusion model for Mn(II) adsorption onto DETA-MCM-41. The amount adsorbed of Mn(II) ions from aqueous solutions on the adsorbent at different concentrations as a function of reaction time. It was a clear that a higher initial rate of removal within the first 10 min followed by a slower subsequent removal rate till reaching equilibrium. With increasing the initial concentration the slower adsorption process was observed but it reached in 20 min in all cases indicating that more adsorption occurred with more active sites (functional groups) available. The performance of DETA-MCM-41 toward Mn(II) seems to be faster at low Mn(II) initial concentrations as more active sites on the surface will be available and faster adsorption rate observed. 3.5. Adsorption thermodynamics The thermodynamic parameters of the adsorption process, such as free energy (DG0) of the adsorption, enthalpy change (DH0) and entropy change (DS0) and activation energy (DE) were determined and are listed in Table 7. The variation of equilibrium constant (Kc) with temperature, as summarized in Table 7, showed that Kc values were increased slightly by increasing in adsorption temperature, thus implying a strengthening of adsorbate–adsorbent interactions at higher temperature. Negative values of DG0 confirm that the process is spontaneous, or it occurs on its own without any outside input at the specified temperature. The value of DH0 was found to be positive confirming the endothermic nature of the adsorption
process. The positive values of DS0 show the increased randomness at the solid/solution interface with some structural changes in the adsorbate and adsorbent (Fig. 6(a)). The activation energy values for the adsorption of Mn(II) were found to be 4.07 kJ/mol which is low and means just small energy required for this type of adsorption (Fig. 6(b)). These thermodynamic results confirmed that the DETA-MCM-41 can adsorb the Mn(II) ions from water samples even at high temperature (up to 55 °C) with good performance. 4. Conclusions In this study, DETA-MCM-41 was prepared and the adsorption of Mn(II) ions by this adsorbent was investigated. A strong dependence on pH in the Mn(II) ions adsorption capacity and best results were obtained at pH 6.5–7. The equilibrium data well fitted the Langmuir sorption isotherms (R2 values are closer to one), and the maximum adsorption capacity of Mn(II) reached 88.9 mg/g for DETA-MCM-41 indicating that the adsorption occurs on a homogeneous surface by monolayer sorption without interaction between the adsorbed ion. The adsorption process is fitted by pseudosecond-order kinetic model and when the initial ion concentration increases from 10 to 200 lg/ml, the pseudo-second-order constants (k2) decreased indicates that the available active sites on the adsorbents are saturated rapidly by Mn(II) ion. The adsorption thermodynamic parameters revealed that the uptake reactions of Mn(II) to adsorbents are spontaneous and endothermic. References [1] F. Raji, M. Pakizeh, Appl. Surf. Sci. 301 (2014) 568–575. [2] O. Abollino, M. Aceto, M. Malandrino, C. Sarzanini, E. Mentasti, Water Res. 37 (2003) 1619–1627. [3] O. Hakami, Y. Zhang, C.J. Banks, Water Res. 46 (2012) 3913–3922. [4] A. Jusoh, W. Cheng, W. Low, A. Aini, M. Noor, Desalination 182 (2005).
90 [5] [6] [7] [8] [9]
[10] [11] [12] [13] [14] [15] [16] [17] [18] [19] [20]
S.A.M. Idris / Journal of Colloid and Interface Science 440 (2015) 84–90 Z. Teng, J.Y. Huang, K. Fujita, S. Takizawa, Desalination 139 (2001) 411–418. F.A. Dawodua, K.G. Akpomie, J. Mater. Res. Technol. 3 (2014) 129–141. S.R. Taffarel, J. Rubio, Miner. Eng. 23 (2010) 1131–1138. Y. Wan, D. Zhao, ChemInform 38 (2007) 2821–2860. F. Hoffmann, M. Fröba, Silica-based mesoporous organic–inorganic hybrid materials, in: The Supramolecular Chemistry of Organic–Inorganic Hybrid Materials, John Wiley & Sons, Inc., 2010, pp. 37–111. N. Shevchenko, V. Zaitsev, A. Walcarius, Environ. Sci. Technol. 42 (2008) 6922– 6928. M.K. Doula, Water Res. 40 (2006) 3167–3176. E. Erdem, N. Karapinar, R. Donat, J. Colloid Interface Sci. 280 (2004) 309–314. L. Mercier, T.J. Pinnavaia, Adv. Mater. (Weinheim, Ger.) 9 (1997) 500–503. J. Brown, R. Richer, L. Mercier, Micropor. Mesopor. Mater. 37 (2000) 41–48. A. Walcarius, M. Etienne, J. Bessière, Chem. Mater. 14 (2002) 2757–2766. A. Walcarius, C. Delacôte, Chem. Mater. 15 (2003) 4181–4192. S.A. Idris, C.M. Davidson, C. McManamon, M.A. Morris, P. Anderson, L.T. Gibson, J. Hazard. Mater. 185 (2011) 898–904. S.A. Idris, S.R. Harvey, L.T. Gibson, J. Hazard. Mater. 193 (2011) 171–176. I. Langmuir, J. Am. Chem. Soc. 40 (1918) 1361–1403. J.-U.-R. Memon, S.Q. Memon, M.I. Bhanger, M.Y. Khuhawar, J. Hazard. Mater. 163 (2009) 511–516.
[21] G. Purna Chandra Rao, S. Satyaveni, A. Ramesh, K. Seshaiah, K.S.N. Murthy, N.V. Choudary, J. Environ. Manage. 81 (2006) 265–272. [22] G. Zolfaghari, A. Esmaili-Sari, M. Anbia, H. Younesi, S. Amirmahmoodi, A. Ghafari-Nazari, J. Hazard. Mater. 192 (2011) 1046–1055. [23] T.Y. Guo, Y.Q. Xia, G.J. Hao, M.D. Song, B.H. Zhang, Biomaterials 25 (2004) 5905–5912. [24] J. Pan, X. Zou, X. Wang, W. Guan, Y. Yan, J. Han, Chem. Eng. J. (Lausanne) 162 (2010) 910–918. [25] S. Lagergren, Handlingar 24 (1898) 1–39. [26] Y. Ho, G. McKay, D. Wase, C. Foster, Adsorpt. Sci. Technol. 18 (2000) 639–650. [27] S. Srivastava, R. Tygir, N. Pant, Water Res. 23 (1989) 1161–1165. [28] W. Weber, J. Morris, Am. Soc. Civ. Eng. 89 (1963) 31–60. [29] T. Wajima, K. Sugawara, Fuel Process. Technol. 92 (2011) 1322–1327. [30] D. Pérez-Quintanilla, I.d. Hierro, M. Fajardo, I. Sierra, J. Hazard. Mater. 134 (2006) 245–256. [31] S.R. Taffarel, J. Rubio, Miner. Eng. 22 (2009) 336–343. [32] Y.-S. Ho, W.-T. Chiu, C.-S. Hsu, C.-T. Huang, Hydrometallurgy 73 (2004) 55–61. [33] Y.-S. Ho, Water Res. 37 (2003) 2323–2330. [34] S. Hong, C. Wen, J. He, F. Gan, Y.-S. Ho, J. Hazard. Mater. 167 (2009) 630–633. [35] Y. Sag˘, Y. Aktay, Process Biochem. 36 (2001) 1187–1197.
Contents lists available at ScienceDirect
Journal of Colloid and Interface Science www.elsevier.com/locate/jcis
Adsorption, kinetic and thermodynamic studies for manganese extraction from aqueous medium using mesoporous silica Salah Ali Mahgoub Idris Chemistry Department, Faculty of Science, University of Omar Al-Mukhtar, Tobruk, Libya
a r t i c l e
i n f o
Article history: Received 26 July 2014 Accepted 15 October 2014 Available online 29 October 2014 Keywords: Modified silica Water Manganese Adsorption Kinetic Thermodynamic
a b s t r a c t This paper describes studies of functionalized mesoporous silica employed as adsorbent for Mn(II) from aqueous solutions. The surface area of MCM-41 and diethylenetriamine functionalized-MCM-41 used in this study were 760 and 318 m2 g1 (N2 adsorption). A strong dependence on pH in the Mn(II) adsorption capacity and best results were obtained at pH 6.5–7. The adsorption onto the diethylenetriamine functionalized-MCM-41 followed the pseudo-second-order kinetic model and the highest reaction rate 0.324 min1 was observed at low initial concentration 10 ppm. The equilibrium data showed excellent correlation with the Langmuir isotherm model and the maximum adsorption capacity of Mn(II) reached 88.9 mg/g for DETA-MCM-41 indicating that the adsorption occurs on a homogeneous surface by monolayer sorption without interaction between the adsorbed ions. These data contribute to the understanding of mechanisms involved in mesoporous silica and provide some practical clues to improve the adsorption efficiency (uptake capacity and kinetics) of Mn(II) ions. Ó 2014 Elsevier Inc. All rights reserved.
1. Introduction Manganese (Mn) is an element that is present in the environment. It is commonly found in drinking water and is essential for human health at low concentrations [1]. However, excessive concentrations of Mn(II) could result in metallic tasting water and many health problems. Manganese is present in ground water as a divalent ion Mn++ and is considered a pollutant mainly because it may causes nervous system damage, leading to Parkinson’s disease, and injures arterial breastwork and the myocardium. The World Health Organization has set a guideline of 0.4 mg/L for Mn(II) intake control from drinking water [2]; and the limit value in the Chinese Standard for Drinking Water Quality is 0.1 mg/L [3]. In neutral and alkaline environments, Mn(II) is difficult to be oxidized [4]. Conventional treatment for Mn(II) removal generally requires the use of strong oxidizing agents such as potassium permanganate, chlorine, hypochlorite, chlorine dioxide or ozone [5]. The generation of other pollutants or toxins is the major disadvantage of these oxidizers. Although chlorine or KMnO4 oxidation combined with sand filtration is usually applied for Mn(II) removal, its concentration in the treated water is still difficult to meet the accepted levels requirements for drinking water [6]. In fact, as the effects of pollutants accumulation and sedimentation, the concentration of Mn(II)
E-mail address: [email protected] http://dx.doi.org/10.1016/j.jcis.2014.10.022 0021-9797/Ó 2014 Elsevier Inc. All rights reserved.
is variable, and the variation of water quality such as solution pH may cause a sudden release of Mn(II) [7]. To develop extraction method for removal of Mn(II) from solution, in the presence of other heavy metals, functionalized mesoporous silica nanoparticles have been explored due to their high metal uptake capacities, good physical and chemical stabilities, uniform structures made of mesopore channels regularly organized in space [8–10]. Summaries of some reported studies of Mn(II) ions adsorption for different materials are presented in Table 1. Although these adsorption capacities correspond to different experimental conditions, they are representative of the tendency to uptake Mn(II) ions. Silica based adsorbents are of particular interest as they exhibit enhanced accessibility to active centers [13,14], and fast mass transfer rates in the regular porous environment [15,16]. These combination of factors have led to high capacity adsorbents [16]. To this end no papers reported the results of kinetic studies and thermodynamic results on Mn(II) adsorption onto mesoporous silica. The aim of this work was to study the adsorption of Mn(II) ions onto amino-functionalized mesoporous silica. The linearized Langmuir and Freundlich equations were used to fit the equilibrium isotherms. Kinetic data were analyzed by pseudo-first, pseudosecond and intra-particle diffusion models to obtain the kinetic constants and thermodynamic study of Mn(II) removal and postulate the mechanism of Mn(II) removal by porous materials. This
S.A.M. Idris / Journal of Colloid and Interface Science 440 (2015) 84–90
work is part of a series of adsorption studies in well characterized mesoporous silica (MCM-41) [17,18]. 2. Experimental 2.1. Materials and reagent Cetyltrimethylammonium bromide (CTAB), 98%, was supplied by BDH. Aqueous ammonia (30% NH3), tetraethoxysilane (TEOS) 98%, hydrogen peroxide solution (30 wt.%), N-(3-trimethoxysilylpropyl) diethylenetriamine (DETA-TMS), ethanol absolute, 1000 lg mL1 of Mn(II) standard solution and toluene (+99%) were purchased from Sigma Aldrich. Nitric acid (65 wt.%) was provided by Fisher Scientific. Glassware was soaked in 5% HNO3 overnight and cleaned with deionised water before use. All products were used as supplied and deionised water was used throughout this work.
85
2.3. Characterization and Analysis The surface areas of the MCM-41 were measured using nitrogen physisorption isotherms. The Brumauer–Emmett–Teller (BET) surface areas were calculated using experimental points at a relative pressure (P/P0) of 0.05–0.25. The total pore volume was calculated from the N2 amount adsorbed at the P/P0 of 0.99 for each sample. Both samples exhibited a Type IV adsorption isotherm typical of mesoporous solids. Desorption isotherms were used to calculate the pore diameters. Elemental analysis (EA) was carried out using an Exeter Analytical CE440 elemental function. Total concentration of Mn(II) in water samples was determined using a PerkinElmer AAnalyst200 flame atomic absorption spectrometry (FAAS) instrument. Calibration standards (0.0, 0.1, 0.5, 1, 1.5 and 2 lg cm3) were prepared in 5% HNO3. Method parameters used are shown in Table 2. The method detection limit (MDL) for Mn(II) was 0.015 lg cm3. 2.4. Adsorption process and effect of pH
2.2. Mesoporous silica preparations and functionalization The synthesis method and functionalization MCM-41 was prepared according to the method reported in Ref. [17]. Approximately 9.0 g of CTAB was dissolved under slight warming (35 °C) in a mixture of 210 mL of distilled H2O and 100 mL of aqueous NH3. To this clear solution, 50 mL of TEOS was slowly added under stirring. After further stirring for 4 h, the gel was aged at room temperature for 48 h in a closed container. The product was obtained by filtration, washed with 800 mL of distilled H2O, and dried in air at room temperature. To remove the surfactant microwave digestion (MWD) was performed by using a MARS 5 microwave digestion system (CEM Corporation, Buckingham, UK) was used at an operating power of approximately 1600 W. The pressure and temperature inside the microwave were controlled to be lower than 1.3 MPa and 200 °C, respectively. Samples (approx. 0.1 g) were added to multiple Teflon vessels to which 1.5 mL of HNO3 and 0.70 mL of H2O2 were added. Microwave digestion was operated at a working frequency of 2450 MHz and 220 V for 15 min. The product (MCM-41) was filtered, washed with copious amounts of distilled H2O and dried at 100 °C for 2 h. Surface modification of MCM-41 was carried out by condensation using the N-(3-trimethoxysilylpropyl) diethylenetriamine with MCM-41 (Scheme 1). Briefly, approximately 5 g of MCM-41 was pre-treated at 140 °C for 2 h before being immersed in 50 cm3 of toluene and 10 cm3 of DETA-TMS, in a 250 cm3 flask. The mixture was refluxed for 4 h and the solid produced was filtered, washed with 100 cm3 ethanol, and oven-dried at 80 °C for 2 h to produce a DETA-MCM-41 sorbent. Table 1 Manganese ions adsorption capacity for different materials. Material
Mn(II) uptake (mg g1)
References
Clinoptilolite–Fe system Clinoptilolite from Turkey Na–montmorillonite Granular activated carbon Nigerian kaolinite clay Manganese oxide coated zeolite
27.12 4.22 3.22 2.54 111.11 27.5
[11] [12] [2] [4] [6] [7]
Approximately 10 mg of functionalized MCM-41 was suspended in 20 mL of solution containing 20 lg mL1 Mn(II) ions in 50 mL beaker and the solution was stirred (250 rpm) using MyLab (Magnetic Stirrer with Hot Plate-SLMSH300) for approximately 2 h. After this time the solution was removed and analyzed for Mn(II) by flame atomic absorption spectrometry (FAAS). This process was repeated until saturation of the sorbent as indicated by a measurement of Mn(II) in the solution aliquot. The extraction was examined at various pH values (between 1 and 9) with solution modification achieved via the addition of small amounts of 1 M ammonium hydroxide. 2.5. Adsorption isotherms The Langmuir [19] or Freundlich [20] models were applied to measured data to study adsorption isotherms. Solutions containing initial concentrations of Mn(II) at 10, 50, 100 or 200 lg cm3 were prepared. To each solution 0.05 g of sorbent was added and the solution was stirred at 250 rpm for 120 min at room temperature, solutions were adjusted to provide a pH of 7. The amounts of Mn(II) extracted at equilibrium, qe (mg/g) were calculated according to Eq. (1):
qe ¼
C0 Ce V W
ð1Þ
where C0 and Ce (mg/g) were the liquid phase initial and equilibrium concentrations of the Mn(II) respectively. V was the volume of the solution (cm3), and W was the mass of sorbent (g) used [21]. The sorption equilibrium data were analyzed according to Langmuir Eq. (2) and Freundlich Eq. (3) isotherm models [22].
Ce 1 1 þ Ce ¼ qm b qm qe ln qe ¼ ln K f þ
1 ln C e n
ð2Þ
ð3Þ
where qe and Ce were the equilibrium concentrations of the Mn(II) ions in the adsorbed and liquid phases in mg/g and mg/L, respec-
Scheme 1. Modification of MCM-41 using N-(3-trimethoxysilylpropyl) diethylenetriamine.
86
S.A.M. Idris / Journal of Colloid and Interface Science 440 (2015) 84–90
The intraparticle diffusion model: To study the mechanism of the particle diffusion the Morris–Weber equation was applied:
Table 2 Conditions used in FAAS. Metal
Mn
Wave length (k) nm Slit (mm) Sample flow rate (mL/min) Current reading value (mA) Lamp current used (mA) Optimum fuel flow rate (mL/min)
279.48 1.8/0.6 7.2 30 10/12 3
pffiffi qt ¼ kid t
ð7Þ
where qt is the amount of metal ion sorbed (mg L1) at time t and kid is the intraparticle diffusion rate coefficient (mg L1 min1/2). 2.7. Thermodynamic study
tively. qm (mg/g) and b (L/mg) were the Langmuir constants. Whereas the qm is the maximum monolayer capacity and b was the adsorption affinity onto the adsorption. Kf (mg/g) and n (L/ mg) were the Freundlich constants which are related to the sorption capacity and intensity, respectively. The Langmuir and Freundlich constants were calculated from the slope and intercept of the linear plot obtained from Eqs. (2) and (3) respectively. For predicting the favorability of an adsorption system, the Langmuir equation can also be expressed in terms of a dimensionless separation factor (RL) by using the Langmuir constant b and the initial concentrations of the Mn(II) (Eq. (4)).
RL ¼
1 1 þ C0 b
ð4Þ
When, RL > 1, RL = 1, 0 < RL < 1 and RL = 0, indicates unfavorable, linear, favorable and irreversible, adsorption isotherms, respectively [23,24].
Kinetic studies to determine the rate of Mn(II) removal from water samples were conducted for DETA-MCM-41. Solutions were prepared with the same initial Mn(II) concentration of 10 lg cm3 and stirred with 0.05 g of each sorbent at 250 rpm. 25 cm3 aliquots of each solution were stirred at 25 °C for 1, 5, 10, 20, 30 or 40 min. After each time period, solutions were filtered and analyzed by FAAS to determine the concentration of Mn(II) in the final solution. The kinetics of Mn(II) adsorption onto the surface of the silica nanoparticles were analyzed using pseudo first-order [25], pseudo second-order [26] and intraparticle diffusion [27,28] kinetic models. The conformity between experimental data and the model predicted values was expressed by the correlation coefficients (R2). A relatively high R2 value (close or equal to 1) was used to indicate best fit to the kinetic model. The pseudo first-order equation: The pseudo first-order equation is generally presented as follows:
k1 t 2:303
ð5Þ
where qe and qt are the adsorption capacity at equilibrium and at time t, respectively (mg g1), k1 is the rate coefficient of pseudo first-order adsorption (L min1). The plot of log(qe qt) vs. t should give a linear relationship from which k1 and qe can be determined from the slope and intercept of the plot, respectively. The pseudo second-order equation: The pseudo second-order adsorption kinetic rate equation is expressed as:
t 1 1 ¼ þ t qt k2 q2e qe
DG0 ¼ RT ln K c
ð8Þ
where Kc is the adsorption equilibrium constant, R is the gas constant and T is the absolute temperature (K). The adsorption equilibrium constant (Kc) can be calculated from:
Kc ¼
Fe 1 Fe
ð9Þ
where Fe is the fraction attainment of Mn(II) ion adsorbed at equilibrium time, and is obtained by the expression
2.6. Adsorption kinetics study
logðqe qt Þ ¼ logðqe Þ
Thermodynamic studies for Mn(II) removal from water samples were conducted for DETA-MCM-41. Solutions were prepared with the same initial Mn(II) concentration of 50 lg cm3 and stirred with 0.05 g DETA-MCM-41 at 250 rpm. 25 cm3 aliquots of each solution were stirred at 25, 35, 45 and 55 °C for 1, 5, 10, 20, 30 or 40 min. After each time period, solutions were filtered and analyzed by FAAS to determine the concentration of Mn(II) in the final solution. The free energy (DG0) of the adsorption reaction is given by the following equation:
ð6Þ
where k2 is the rate coefficient of pseudo second-order adsorption t and t, should give a linear relationqt
(g mg1 min1). The plot of
ship from which qe and k2 can be determined from the slope and intercept of the plot, respectively.
Fe ¼
C0 Ce C0
ð10Þ
where C0 and Ce are the initial and equilibrium concentrations of Mn(II) ion in solution (mg/L). The value of the adsorption equilibrium constant (Kc) for the adsorption of Mn(II) ion on the adsorbent were calculated at different temperature and at equilibrium time using Eqs. (9) and (10). The Gibbs free energy can be represented as follows:
DG0 ¼ DH0 T DS0
ð11Þ
The values of enthalpy change (DH0) and entropy change (DS0) calculated from the intercept and slope of the plot of DG0 vs. T [29]. The activation energy for Mn(II) adsorption was calculated by the Arrhenius equation DE k2 ¼ AeðRT Þ
ð12Þ
where DE is the activation energy (kJ/mol), A is the frequency factor, T is the absolute temperature (K), and R is the gas constant. From the plot of ln(k2) vs. 1/T, the activation energy DE for the adsorption of Mn(II) can be calculated. 3. Results and discussion 3.1. Materials characterization The MCM-41 and functionalized-MCM-41 materials were characterized using BET and the physicochemical properties of both are summarized in Table 3. The N2 sorption isotherms (Fig. 1) were type IV for both samples confirming their mesoporous nature, however the quantities of nitrogen gas adsorbed with MCM-41 was higher compared with DETA-MCM-41 due to the surface of DETA-MCM-41 was occupied with the functional groups. Elemental analysis was used to estimate the amount of molecules (L0) attached to functionalized samples from the percentage
S.A.M. Idris / Journal of Colloid and Interface Science 440 (2015) 84–90
87
Table 3 Physicochemical properties of MCM-41 and DETA- MCM-41. Sample Name
BET surface area (m2 g1)a
Pore size (nm)b
Pore volume (cm3 g1)c
MCM-41 DETA- MCM-41
760 318
6.74 6.18
0.99 0.63
a Calculated by the linear BET equation from sorption data in a relative pressure range from 0.05 to 0.25. b Calculated by the BJH model from the adsorption branches of isotherms. c Calculated from N2 amount adsorbed at a relative pressure P/P0 of 0.99.
Fig. 2. Adsorption capacities of functionalized MCM-41 for Mn(II) extraction over different pH values.
Fig. 1. Nitrogen adsorption isotherms of MCM-41 and DETA-MCM-41.
Table 4 Elemental analysis data recorded for the MCM-41. Silica
%C
%H
%N
L0 (mmol/g)a
MCM-41 DETA-MCM-41
Trace/Nil 16.8
0.65 4.72
Trace/Nil 10.8
– 2.57
Fig. 3. Chemical diagram species for Mn(II) aqueous solution. C0 = 1.82 mM Mn(II).
a Functionalisation degree (L0 = millimoles of ligand per gram of functionalised silica).
of nitrogen, in the functionalized mesoporous silica [30], using Eq. (13):
L0 ¼
%N 10 nitrogen atomic weight
ð13Þ
The calculated L0 (Table 4) values for DETA-MCM-41 were 2.57 mmol of ligand per gram of functionalized silica indicating the successfully functionalization for diethylenetriamine functional groups. 3.2. Effect of pH on Mn(II) adsorption The pH of the aqueous solution is an important controlling parameter in the adsorption processes and metal ion removal usually increases with increasing pH values [31]. This Fig. 2 shows that the amount of Mn(II) ions adsorbed on DETA-MCM-41 increases with the increase of the pH. This phenomenon appears to be due to the fact that functional group and the surface of mesoporous silica were highly selective for H3O+ ions when the H3O+ ions concentration was high. Thus, at
lower pH values the H3O+ ions compete with Mn(II) ions for the exchange sites in DETA-MCM-41 leading to a low removal of the metal ion. Furthermore, the extensive repulsion of metal ions due to protonation of active sites of DETA-MCM-41 surface at lower pH may be another reason for decrease in Mn(II) ions adsorption at pH < 5. The increase in Mn(II) ion removal with the increases in pH can be explained on the basis of a decrease in competition between protons and Mn(II) cations for the same functional groups and by the decrease in positive surface charge, which results in a lower electrostatic repulsion between the surface of DETA-MCM-41 and Mn(II) ions before ion-exchange. Adsorption results showed an optimum pH value of about 7 (equilibrium pH fluctuated between 6.5 and 7.0) which is also important in the disposal of the treated effluents (required for neutrality of the bodies ‘‘receivers’’ water or water reuse). The decrease in adsorption at higher pH (above pH 7), is probably due to the formation of insoluble Mn(II) hydroxide Mn(OH)2 (Fig. 3) or the pH may affect the ionization degree (species formation) of the adsorbate and the surface property of the adsorbent [31]. For example, the heavy metal ions may form a complex with inorganic ligands such as OH, changing its adsorption degree. The extent of these compounds formation varies with pH, ionic composition and the adsorbate ion. Results obtained showing that the amount of Mn(II) extracted from water using amino functionalized MCM-41 can be
88
S.A.M. Idris / Journal of Colloid and Interface Science 440 (2015) 84–90
Table 5 Isotherm parameters for Mn(II) sorption on functionalised mesoporous silica. Adsorbents
Langmuir
DETA-MCM-41
Freundlich 2
qm (mg/g)
b (L/mg)
RL
R
Kf (mg/g)
n (L/mg)
R2
88.9
8.34
0.55
0.9998
26.4
4.2
0.7889
Fig. 5. The intraparticle diffusion model for Mn(II) adsorption onto DETA-MCM-41.
isotherms are the most commonly used isotherms for different adsorbent/adsorbate systems to explain solid–liquid adsorption systems and to predict their equilibrium parameters [32–34]. The relevant parameters for these isotherms are presented in Table 5. As seen from Table 5, the R2 values obtained from the Langmuir model are much closer to one than are those from the Freundlich model, suggesting that the Langmuir model is better than the Freundlich isotherm. Thus the adsorption can be described by the Langmuir isotherm and the metal ion adsorption occurs on a homogeneous surface by monolayer sorption without interaction between the adsorbed ion [35] and the adsorption capacity is 88.9 mg/g. 3.4. Adsorption kinetics Fig. 4. Pseudo-first order (a) and Pseudo-second order (b) kinetics models of Mn(II) onto DETA-MCM-41 sorbent.
influenced by changing the pH of aqueous medium due to ionization degree (species formation) of the Mn(II) and the functional groups property attached to the surface of mesoporous silica. 3.3. Adsorption isotherms The adsorption isotherms were studied to find the relationship between equilibrium adsorption capacity and equilibrium concentration at a certain temperature. Langmuir and Freundlich
The experimental kinetic data were fitted using a pseudofirst-order kinetic model and pseudo-second-order kinetic model (Fig. 4). The results are shown in Table 6. It can be seen that the obtained R2 values of the pseudo-second-order model (>0.996) were better than R2 of the pseudo-first – order model (0.949–0.993), suggesting that the adsorption process is second-order. Moreover, the calculated qe values were much closer to the experimental values in the pseudo-second-order kinetic model than the pseudofirst-order kinetic model indicating that the adsorption process is second-order. As seen in Table 6, when the initial ion concentration increases from 10 to 200 lg/ml, the pseudo-second-order constants
Table 6 Kinetic parameters for the adsorption of manganese on the adsorbent. Adsorbents
DETA-MCM-41 DETA-MCM-41 DETA-MCM-41 DETA-MCM-41
C0 (lg/mL)
10 50 100 200
qe (exp) (mg/g)
5.00 25.00 49.78 88.45
The pseudo first-order
The pseudo second-order
k1 (min1)
qe (cal) (mg/g)
R2
k2 (min1)
qe (cal) (mg/g)
R2
0.164 0.219 0.130 0.156
2.07 11.49 29.90 58.26
0.9803 0.9489 0.9930 0.9769
0.324 0.052 0.014 0.008
5.07 25.58 51.29 90.98
0.9996 0.9986 0.9976 0.9969
89
S.A.M. Idris / Journal of Colloid and Interface Science 440 (2015) 84–90 Table 7 Thermodynamic parameters for the adsorption of Mn(II) on the adsorbent. Adsorbents
Reaction temp. (K)
Kc
DG0 (kJ/mol)
DH0 (kJ/mol)
DS0 (J/mol K)
DE0 (kJ/mol)
DETA-MCM-41 DETA-MCM-41 DETA-MCM-41 DETA-MCM-41
298 308 318 328
12,499 16,666 24,999 49,999
23.37 24.89 26.77 29.51
37.34
202.8
4.07
Fig. 6. (a) Relationship between Gibbs free energy change and temperature of adsorption of Mn(II) onto the adsorbent. (b) Arrhenius plot for Mn(II) adsorption on the adsorbent.
(k2) decrease from 0.324 to 0.052, 0.014 to 0.008 for DETA-MCM-41 which indicate that the active sites on the surface of mesoporous silica were occupied by Mn(II). Adsorption capacity is 90.98 mg/g which is high and that maybe favorability of Mn(II) to bonded to hard ligand as Mn(II) is hard metal and also may be due to the adsorption of Mn(II) inside pores in addition to the surface adsorption. Fig. 5 shows the intraparticle diffusion model for Mn(II) adsorption onto DETA-MCM-41. The amount adsorbed of Mn(II) ions from aqueous solutions on the adsorbent at different concentrations as a function of reaction time. It was a clear that a higher initial rate of removal within the first 10 min followed by a slower subsequent removal rate till reaching equilibrium. With increasing the initial concentration the slower adsorption process was observed but it reached in 20 min in all cases indicating that more adsorption occurred with more active sites (functional groups) available. The performance of DETA-MCM-41 toward Mn(II) seems to be faster at low Mn(II) initial concentrations as more active sites on the surface will be available and faster adsorption rate observed. 3.5. Adsorption thermodynamics The thermodynamic parameters of the adsorption process, such as free energy (DG0) of the adsorption, enthalpy change (DH0) and entropy change (DS0) and activation energy (DE) were determined and are listed in Table 7. The variation of equilibrium constant (Kc) with temperature, as summarized in Table 7, showed that Kc values were increased slightly by increasing in adsorption temperature, thus implying a strengthening of adsorbate–adsorbent interactions at higher temperature. Negative values of DG0 confirm that the process is spontaneous, or it occurs on its own without any outside input at the specified temperature. The value of DH0 was found to be positive confirming the endothermic nature of the adsorption
process. The positive values of DS0 show the increased randomness at the solid/solution interface with some structural changes in the adsorbate and adsorbent (Fig. 6(a)). The activation energy values for the adsorption of Mn(II) were found to be 4.07 kJ/mol which is low and means just small energy required for this type of adsorption (Fig. 6(b)). These thermodynamic results confirmed that the DETA-MCM-41 can adsorb the Mn(II) ions from water samples even at high temperature (up to 55 °C) with good performance. 4. Conclusions In this study, DETA-MCM-41 was prepared and the adsorption of Mn(II) ions by this adsorbent was investigated. A strong dependence on pH in the Mn(II) ions adsorption capacity and best results were obtained at pH 6.5–7. The equilibrium data well fitted the Langmuir sorption isotherms (R2 values are closer to one), and the maximum adsorption capacity of Mn(II) reached 88.9 mg/g for DETA-MCM-41 indicating that the adsorption occurs on a homogeneous surface by monolayer sorption without interaction between the adsorbed ion. The adsorption process is fitted by pseudosecond-order kinetic model and when the initial ion concentration increases from 10 to 200 lg/ml, the pseudo-second-order constants (k2) decreased indicates that the available active sites on the adsorbents are saturated rapidly by Mn(II) ion. The adsorption thermodynamic parameters revealed that the uptake reactions of Mn(II) to adsorbents are spontaneous and endothermic. References [1] F. Raji, M. Pakizeh, Appl. Surf. Sci. 301 (2014) 568–575. [2] O. Abollino, M. Aceto, M. Malandrino, C. Sarzanini, E. Mentasti, Water Res. 37 (2003) 1619–1627. [3] O. Hakami, Y. Zhang, C.J. Banks, Water Res. 46 (2012) 3913–3922. [4] A. Jusoh, W. Cheng, W. Low, A. Aini, M. Noor, Desalination 182 (2005).
90 [5] [6] [7] [8] [9]
[10] [11] [12] [13] [14] [15] [16] [17] [18] [19] [20]
S.A.M. Idris / Journal of Colloid and Interface Science 440 (2015) 84–90 Z. Teng, J.Y. Huang, K. Fujita, S. Takizawa, Desalination 139 (2001) 411–418. F.A. Dawodua, K.G. Akpomie, J. Mater. Res. Technol. 3 (2014) 129–141. S.R. Taffarel, J. Rubio, Miner. Eng. 23 (2010) 1131–1138. Y. Wan, D. Zhao, ChemInform 38 (2007) 2821–2860. F. Hoffmann, M. Fröba, Silica-based mesoporous organic–inorganic hybrid materials, in: The Supramolecular Chemistry of Organic–Inorganic Hybrid Materials, John Wiley & Sons, Inc., 2010, pp. 37–111. N. Shevchenko, V. Zaitsev, A. Walcarius, Environ. Sci. Technol. 42 (2008) 6922– 6928. M.K. Doula, Water Res. 40 (2006) 3167–3176. E. Erdem, N. Karapinar, R. Donat, J. Colloid Interface Sci. 280 (2004) 309–314. L. Mercier, T.J. Pinnavaia, Adv. Mater. (Weinheim, Ger.) 9 (1997) 500–503. J. Brown, R. Richer, L. Mercier, Micropor. Mesopor. Mater. 37 (2000) 41–48. A. Walcarius, M. Etienne, J. Bessière, Chem. Mater. 14 (2002) 2757–2766. A. Walcarius, C. Delacôte, Chem. Mater. 15 (2003) 4181–4192. S.A. Idris, C.M. Davidson, C. McManamon, M.A. Morris, P. Anderson, L.T. Gibson, J. Hazard. Mater. 185 (2011) 898–904. S.A. Idris, S.R. Harvey, L.T. Gibson, J. Hazard. Mater. 193 (2011) 171–176. I. Langmuir, J. Am. Chem. Soc. 40 (1918) 1361–1403. J.-U.-R. Memon, S.Q. Memon, M.I. Bhanger, M.Y. Khuhawar, J. Hazard. Mater. 163 (2009) 511–516.
[21] G. Purna Chandra Rao, S. Satyaveni, A. Ramesh, K. Seshaiah, K.S.N. Murthy, N.V. Choudary, J. Environ. Manage. 81 (2006) 265–272. [22] G. Zolfaghari, A. Esmaili-Sari, M. Anbia, H. Younesi, S. Amirmahmoodi, A. Ghafari-Nazari, J. Hazard. Mater. 192 (2011) 1046–1055. [23] T.Y. Guo, Y.Q. Xia, G.J. Hao, M.D. Song, B.H. Zhang, Biomaterials 25 (2004) 5905–5912. [24] J. Pan, X. Zou, X. Wang, W. Guan, Y. Yan, J. Han, Chem. Eng. J. (Lausanne) 162 (2010) 910–918. [25] S. Lagergren, Handlingar 24 (1898) 1–39. [26] Y. Ho, G. McKay, D. Wase, C. Foster, Adsorpt. Sci. Technol. 18 (2000) 639–650. [27] S. Srivastava, R. Tygir, N. Pant, Water Res. 23 (1989) 1161–1165. [28] W. Weber, J. Morris, Am. Soc. Civ. Eng. 89 (1963) 31–60. [29] T. Wajima, K. Sugawara, Fuel Process. Technol. 92 (2011) 1322–1327. [30] D. Pérez-Quintanilla, I.d. Hierro, M. Fajardo, I. Sierra, J. Hazard. Mater. 134 (2006) 245–256. [31] S.R. Taffarel, J. Rubio, Miner. Eng. 22 (2009) 336–343. [32] Y.-S. Ho, W.-T. Chiu, C.-S. Hsu, C.-T. Huang, Hydrometallurgy 73 (2004) 55–61. [33] Y.-S. Ho, Water Res. 37 (2003) 2323–2330. [34] S. Hong, C. Wen, J. He, F. Gan, Y.-S. Ho, J. Hazard. Mater. 167 (2009) 630–633. [35] Y. Sag˘, Y. Aktay, Process Biochem. 36 (2001) 1187–1197.
