The removal of Mn2+ ions in aqueous solution using hydrotalcite material

JOURNAL OF SCIENCE OF HNUE DOI: 10.18173/2354-1059.2015-00078 Chemical and Biological Sci. 2015, Vol. 60, No. 9, pp. 51-58 This paper is available online at Received November 3, 2015. Accepted December 2, 2015. Contact Ly Bich Thuy, e-mail address: thuy.lybich@hust.edu.vn 51 THE REMOVAL OF Mn 2+ IONS IN AQUEOUS SOLUTION USING HYDROTALCITE MATERIAL Pham Thi Hanh and Ly Bich Thuy School of Environmental Science and Technology, Hanoi University of Science and Technology Abstract. This study evaluated the adsorption capacity of hydrotalcite material (HT) for the removal of Mn2+i an aqueous solution. Factors affecting the adsorption capacity of HT for the removal of Mn2+ ions from aqueous solutions, such as initial pH, contact time and initial Mn2+ concentration, were investigated using batch experiments. This study made use of pseudo-first-order and pseudo-second-order kinetic models. Langmuir and Freundlich adsorption isotherms were investigated. The obtained results show that HT removes Mn2+ effectively (> 97%) at a temperature of 25 oC when the solid/liquid ratio is 0.5% and the initial Mn2+ concentration is 5 mg/L. The pseudo-second-order model fit the observed ata better than the pseudo-first-order model. The equilibrium data correlated very well with the Freundlich isotherm model which implies multilayer adsorption. The obtained results suggested that HT is a good adsorbent for removal of Mn2+ ions in aqueous solution. Key words: Hydrotalcite material, Mn2+ removal. 1. Introduction Manganese is usually present in groundwater as a divalent ion (Mn2+) and it is considered to be a groundwater contaminant in several areas of the Red and Mekong River deltas [1, 2]. Manganese removal is a matter of concern because groundwater is the main water source of drinking water in Vietnam, especially in rural areas. Almost every rural household in Vietnam uses sand filtration which removes many impurities from the groundwater. However, many studies have shown that after sand filtration the manganese level in the water exceeds the maximum amount for drinking water as recommended by both the World Health Organization (0.4 mg/L), and Vietnamese law QCVN 01:2009/BYT (0.3 mg/L) [3]. Adsorption could provide additional treatment after the primary sand filtration. Previously, research has been done to determine the efficacy of clinoptilolite (Clin), clinoptilolite-iron (Clin-Fe), manganese oxide coated zeolite (MOCZ) and carbon nanotubes to serve as adsorbants to remove Mn2+ [4-6]. It was found that the adsorption capacity of Clin was 7.69 mg/g while that of Clin-Fe was 27.12 mg/g. However, after using Clin and Clin-Fe Pham Thi Hanh and Ly Bich Thuy 52 as an adsorbent, an undesirable amount of Ca2+, Na+, K+ ions were introduced in the aqueous solution [4]. MOCZ has low adsorption capacity, about 1.1 mg/g [5], and so there are still a need for a manganese adsorption material that is both efficientand economical. Hydrotalcites (HTs) are layered double hydroxides, also known as anionic clays, with a general formula: [M2+1-xM 3+ x(OH)2] x+[(An-)x/n.mH2O] x- (x = 0.2 - 0.33), where M2+ and M3+ are divalent and trivalent metal ions, respectively, and An- is an n-valent anion [7]. The structure of HT is that of a positive charged brucite-like octahedral layer and a negatively charged interlayer containing anions and water molecules [8]. This po itive charged layer is formed by a partial substitution of a trivalent metal for a divalent one. The layers can be stacked, and the balancing interlayer anions can be exch ged with other anions. HTs have received increasing attention in recent years as potentialion-exchangers and catalysts [9-11]. This study evaluates the feasibility of using HT to remove Mn2+ ions from an aqueous solution. In this study, batch equilibrium and kinetic tests were carried out using laboratory solutions containing Mn2+ ions. 2. Content 2.1. Material and methods * Materials [Mg4Al2(OH)12] 2+[NO3.CO3.zH2O] 2- (MA) were prepared using a coprecipitation method that is presented elsewhere [12]. With this method, a solution of Mg(NO3)2 and Al(NO3)3 (Mg2+: Al3+ = 2:1) are gradually added to a solution of Na2CO3 under constant mixing to create a precipitate. The amount of precipitate was increased by continuously mixing of obtained products under thermal controlled conditions. * Batch experiments to invest the affected factors on adsorption capacity Batch experiments were carried out in glass conical flasks (50 mL) using 20 mL of Mn2+ ion solution. Except for one experiment done to determine the effect of the initial Mn2+, Co of 5 mg/L was kept constant in all experiments. HT with a solid o liquid ratio of 0.5% was applied in all experiments except for one experiment done to determine the effect of the solid to liquid ratio. An initial pH of 6.8 was maintained in all experiments except for one experiment done to determine effect of initial pH The mixture was then mixed in a shaker (Jeiotech BS-31, Korea) at a speed of 150 rpm and a temperature of 25 ± 2 oC. The supernatant was then filtered through a 0.45 μm filter membrane and the Mn2+ concentration was determined. The Mn2+ ion concentration was estimated after catalytic oxidization of Mn2+ to Mn7+ in a H2SO4 solution using K2S2O8 as the oxidazing agent and AgNO3 as the catalyst. The Mn7+ was then an lyzed using a spectrophotometer (UV-VIS 1201) at  = 520 nm. The pH of the solution was adjusted using HNO3 or a NaOH solution. The removal efficiency was calculated using equation (1): 100 )( (%)    o to C CC Efficiency (1) where Co and Ct are the initial conce tration of Mn 2+ and the concentration of Mn2+ at time t, respectively. * Determination of adsorption isotherm model Adsorption isotherm models were used to characterize the interactions of ions with sorbents. The Langmuir and Freundlich models are commonly used to describe the adsorption isotherms at constant temperature [13]. The Langmuir model is based on the assumption that maximum adsorption occurs when a saturated monolayer of solute molecules is present on the The removal of Mn 2+ ions in aqueous solution using hydrotalcite material 53 adsorbent surface, the energy of adsorption is constant and there is no migration of adsorbate molecules in the surface plane [14]. This is represented by the following equation: max . . 1 K . L e L e K C q q C   (2) Equation (2) can be rearranged in linear form: max max 1 .K e e e L C C q q q   (3) where qe and qmax are the equilibrium and maximum uptake capacities (mg/g); Ce is the equilibrium concentration (mg/L), KL is the equilibrium constant (L/mg). The Freundlich equation is the empirical relationship whereby it is assumed that the adsorption energy of an ion binding to a site on an adsorbent depends on whether or not the adjacent sites are already occupied [14]. It is represented below: 1 . ne F eq K C (4) The linear form of equation (4) is: 1 ln ln lne F eq K C n   (5) where KF and n are the Freundlich constants which are characteristic of the system. In order to determine the adsorption isotherm model, batch mode experiments were carried out with initial concentrations of Mn2+ varyi g from 1 to 50 mg/L. All other parameters: mixing speed of 150 rpm, temperature of 25 ± 2 oC, volume of the Mn2+ solution (20 mL), 0.5% solid/liquid ratio, pH of 6.8 and contact time of 60 min, were kept constant. * Determination of sorption kinetic In order to analyze the sorption kinetics of ions, the pseudo-first-order kinetic and pseudo-second-order kinetic models were applied to the data obtained from the investigation of the effect of contact time. The first-orde rate equation is one of the most widely used to determine the sorption of a solute from a liquid solution [15] and is represented as 1log( ) log 2.303e t e k q q q t   (6) where, k1 is the rate constant of first-order sorption (min −1) and qe and qt denote the amounts of sorption at equilibrium and time t (mg/g), respectively. A plot of log (qe−qt) versus t should give a straight line to confirm the applicability of the first-order kinetic model. The pseudo-second-order equation based on adsorption equilibrium capacity may be expressed in the form: 2 2 1 1 . t e e t t q k q q   (7) where k2 (g mg −1min−1) is the rate equilibrium constant of the pseudo-second-order reaction. A plot of t/qt versus t should give a linear relationship to confirm the applicability of the second-order kinetic model. 2.2. Results and discussion 2.2.1. Effect of contact time The effect of contact time from 5 to 180 minute was investigated. Figure 1 shows that when contact time increased from 5 to 30 min, Mn2+ removal increased from 78% to 88% and Mn2+concentrations decreased from 5 mg/L to 0.61 mg/L. The initial rapid adsorption phase may be due to a greater number of adsorption sites available for adsorption of Mn2+ initially, Pham Thi Hanh and Ly Bich Thuy 54 and the adsorption rate decreased gradually s the number of available sites fell. The fast uptake of Mn2+ by HT could be attributed to the highly porous and mesh structure of HT which provides ready access and a large surface area for the sorption of metals on the binding sites. The equilibrium time for Mn2+ adsorption onto HT is 60 minutes. Th refore a contact time of 60 minutes was selected for the next adsorption experiment. 2.2.2. Effect of initial pH The pH of the solution is the most significant factor in adsorption of metal ions as it has a major effect on the protonation and deprotonation of the adsorbent and adsorbate functional groups [4]. Figure 2 and Table 1 show the effect of pH on Mn2+ ion adsorption. When the initial pH increased from 4.05 to 6.11 the percentage of the removal increased from 86.8 to 94.7%. As evident from Figure 2, Mn2+ ions were adsorbed onto HT better when the pH was 6.11 - 8.53. A maximum Mn2+ adsorption of 95.7% was observed at an initial pH of 7.5. Table 1. Final pH and Mn 2+ removal efficiency at different initial pH Initial pH 4.05 5.12 6.11 7.05 7.51 8.03 8.53 9.00 Final pH 7.42 8.07 8.25 8.13 8.25 8.05 8.53 9.01 Removal efficiency, % 86.8 88.8 94.7 95.2 95.7 95.5 95.0 94.8 At low initial pH, a positively charged density on surface sites led to an incr ase in electrostatic repulsion between the surface positive charges and the positively charged metal ions and therefore caused a decrease in the adsorption capacity of the adsorbent. Moreover, at low pH, the concentration of H+i ns is high and removal of Mn2+ ions was inhibited, possibly as a result of competitiv adsorption between H+ and Mn2+ on the exchangeable sites of the surface which had an apparent preponderance of H+ ions. As the initial pH increased, the negative charge density on the HT surface increases due to th deprotonation of the metal binding sites and thus the adsorption of Mn2+ ions increased. A maximum Mn2+ adsorption of 95.7% was observed at an initial pH = 7.5. Therefore, pH = 7.5 was selected for next experiment. Figure 1. Effect of contact time on the adsorption of Mn 2+ on HT Figure 2. Effect of pH on the adsorption of Mn 2+ on HT The removal of Mn 2+ ions in aqueous solution using hydrotalcite material 55 2.2.3. Effect of initial solution concentration Figure 3. Effect of initial concentration of Mn 2+ sorption by HT The adsorption capacity and the adsorption efficiency of Mn2+ at different initial Mn2+ concentrations are shown in Figure 3. It was observed that the adsorption capacity of Mn2+ on HT increased from 0.2 to 19 mg/g when the initial Mn2+ concentration increased from 1 to 500 mg/L, with the removal of Mn2+ decreasing from 100% to 18.6%. It is readily understood that the number of available adsorption sites increases at high adsorbent concentration which results in an increased amount of adsorbed Mn2+. The plot of amount adsorbed versus adsorbent concentration reveals that the uptake capacity was higher at lower concentrations and reduced at higher concentrations. 2.2.4. Adsorption isotherm The linearized Langmuir and Freundlich isotherms obtained for Mn2+ are p esented in Figures 4 and 5. The Langmuir and Freundlich adsorption constants calculated from the corresponding isotherms and their respective coefficients are presented in Table 2. Both isotherms are considered functions of Ce which correspond to the equilibrium distribution of ions between aqueous and solid phases as the initial Mn2+ ion concentration increases. The values of R2 are regarded as a measure of goodness- f-fit of experimental data in the isotherm models [15]. As seen in Table 2, Freundlich modelling exhibited a better fit to the data for Mn2+ ions. Therefore, the sorption process described in this study may be interpreted as being multilayer adsorption. According to the Freundlich model, the Freundlich constant, KF (adsorption capacity) value of 1.95 mg.g-1.L-1 and the n values at equilibrium of greater than 1 (n = 2.82) reflects favorable adsorption. The results indicate that the HT has a strong adsorption capacity for Mn2+ ions in solution. Table 2. Adsorption isotherm parameters of Mn 2+ ions on HT Langmuir model Freundlich model Parameters qmax (mg/g) KL (L/mg) R 2 n KF (mg/L.g) R 2 3.5 0.3 0.926 2.82 1.95 0.991 Pham Thi Hanh and Ly Bich Thuy 56 Figure 4. The Langmuir adsorption isotherm for Mn 2+ by HT Figure 5. The Freundlich adsorption isotherm for Mn 2+ by HT 2.2.5. Kinetic models First and second adsorption models were applied to interpret the experimental data and the results are shown in Figures 6 and 7 and Table 3. The results indicat that the adsorption process of Mn2+ ions on the HT does not follow a pseudo first-order reaction. The correlation coefficient, obtained from the pseudo-first-order kinetic model, was low (0.69) and the calculated qe differs from the experimental results (qe(exp) = 4.75 mg/g). The second-order equation fitted well with the experimental data (Figure 7). The correlation coefficient for the pseudo-second-order kinetic model was higher than that for the pseudo-first-order kinetic model (0.999 > 0.692) and the theoretical qe values (4.8 mg/g) agreed well with the experimental value (4.75 mg/g). Figure 6. The pseudo-first-order kinetic model for the adsorption of Mn 2+ ions onto HT Figure 7. The pseudo-second-order kinetic model for the adsorption of Mn 2+ ions onto HT Table 3. Kinetic constants for Mn 2+ ion adsorption onto HT Pseudo-first-order kinetic model Pseudo-second-order kinetic model Experimental qe(mg/g) k1×10 3, min-1 qe(theor.), mg/g R 2 k2, g.mg1.min-1 qe(theor.), mg/g R 2 6.909 0.76 0.692 0.132 4.8 0.999 4.75 The removal of Mn 2+ ions in aqueous solution using hydrotalcite material 57 3. Conclusion The results obtained in this study clearly demonstrate the potential use of HT for the removal of Mn2+ from aqueous solutions. 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