Adsorption Equilibrium, Kinetics, and Thermodynamic Studies of

Article Cite This: J. Chem. Eng. Data XXXX, XXX, XXX−XXX

pubs.acs.org/jced

Adsorption Equilibrium, Kinetics, and Thermodynamic Studies of Fluoride Adsorbed by Tetrametallic Oxide Adsorbent Sapna Raghav† and Dinesh Kumar*,‡ †

Department of Chemistry, Banasthali University, Rajasthan, 304022, India School of Chemical Science, Central University of Gujarat, Gandhinagar, 382030, India

‡

S Supporting Information *

ABSTRACT: This study investigated the performance of fluoride adsorption onto a specific tetrametallic oxide adsorbent Fe−Al−Ce-Ni (FACN) and the effect of temperature on adsorption performance. The adsorption performance was determined by adsorption equilibrium, kinetics, and thermodynamic parameters. The adsorption, kinetic, and thermodynamic parameters were compared alternatively. The fluoride adsorption capacity was obtained from four different adsorption isotherm models, namely, Langmuir, Freundlich, Temkin, and Dubinin−Radushkevich (D−R), and Freundlich was found to best fit model. Fluoride removal rate using adsorption (0.27 min−1) was obtained faster than reactive adsorption (0.04 min−1). Several thermodynamic parameters such as enthalpy, Gibbs free energy, entropy (ΔS > 0), and adsorption activation energy were calculated which demonstrated the feasibility and spontaneity (ΔG < 0) and exothermic nature of (ΔH < 0) the fluoride adsorption process. The adsorption process was controlled by a physical mechanism, and the maximum adsorption capacity was found to be 250 mg/g. To our knowledge, this is the first report on the synthesis of tetrametallic oxide adsorbent for fluoride adsorption, and the feasibility of the adsorption process was ratified by three van’t Hoff plots. especially in children.3,4 High fluoride concentration in drinking water shows a negative effect and causes harmful diseases such as dental fluorosis, skeletal fluorosis, osteoporosis, arthritis, brittle bones, infertility, cancer, Alzheimer, brain damage, and thyroid disorder, etc. Worldwide, about 200 million people are affected due to high F¯ concentration in drinking water, which is more than the value of 1.5 mg/L, a guideline line set by WHO.5 Thus, choosing a technology for the removal of F¯ from drinking water represents a challenge. There are numerous methods for F¯ removal such as reverse osmosis,6 electrocoagulation,7 nanofiltration,8 ion exchange,9 and adsorption.10 Among these, adsorption is one of the best techniques, because of its low cost, easy operation, minimal waste disposal, and environment friendliness, and it has the added advantage that it can be applied to a decentralized water supply system.11−16 Many adsorbents from natural to synthetic have been utilized for F¯ removal from water. To date, many synthetic single and multicomponent metal oxides have been reported, such as aluminum oxide,17 magnesium oxide,18 Mn−Ce,19 Mn−Zr,20 Fe−Al,21 Fe−Ti,22 Fe−Ca−Ce,23 and Fe−Al−Ce.24 Most of the reported multicomponent composites were iron or aluminum-based metallic oxides and their improved adsorption properties

1. INTRODUCTION The general term “fluorosis” is closely related to higher levels of fluoride (F¯) concentration in water. There are many ores, minerals, and rocks present inside the earth’s crust, which are the natural resources of fluoride. These natural resources have been in contact with groundwater for many years. So, because of the long contact time of ores, minerals, and rocks with groundwater, there is a constant leaching of fluoride ions that is responsible for the high fluoride concentration in groundwater as well as oceanic water. In addition, there are many active volcanoes present in the ocean, and their volcanic eruption is also responsible for the increase in salinity of ocean water as lava contains many salts, ores, minerals, and rocks, etc. Hence the coastal areas of many countries have a high fluoride concentration in their waters. In India, fluoride was initially identified in drinking water at the Nellore district of Andhra Pradesh, in 1937. In 2008, Scientific American magazine reported on the severity of F¯ poisoning in a report called “Second thoughts about fluoride.”1 Increasing F¯ concentration in the water is recognized as a major threat to human health by the World Health Organization (WHO), European Union (EU), and Environmental Protection Agency (EPA).2 Long-term intake of fluoride in high concentration through drinking water can lead to health problems. In fact, the effect of fluoride depends upon the fluoride concentration in water. A smaller concentration of fluoride in drinking water is essential for enamel development, © XXXX American Chemical Society

Received: January 8, 2018 Accepted: March 7, 2018

A

DOI: 10.1021/acs.jced.8b00024 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

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obtained then 6 M NaOH solution was added dropwise until a pH of 9−10 was obtained. A brown precipitate was observed upon stirring the solution for 4−5 h at room temperature. The resulting precipitate was allowed to stand for 24 h at room temperature. It was centrifuged and washed with double distilled water until the pH of precipitates reached a value of 6.0−6.5. After that the precipitates were dried overnight at 100 °C and cooled at room temperature to obtain the solid powder. As a crux, the solid powder was dried in a muffle furnace by heating from room temperature to 600 °C @ 2 °C/min, to remove all volatile gases. It was then cooled at room temperature and crushed into fine powder to obtain the activated solid powder of FACN. 2.3. Batch Reactive Adsorption Study. The batch reactive adsorption was evaluated in terms of the effect of adsorbent dose, pH, temperature, initial F¯ concentration, time, co-ion interference, and regeneration studies. In all the experiments, 100 mL of F¯ ion solution of various concentrations from 5 to 100 mg/L, the adsorbent dose of 0.4 g/L, pH 5.0, and agitation speed of 400 rpm were maintained consistently. One milliliter total ionic strength adjuster buffer (TISAB) solution was added in all the sample solutions to preserve pH 5.0 and to eliminate the interference effect of complexing ions. The pH studies were done by maintaining the pH of a solution using 1.0 M HNO3 or 1.0 M NaOH. To explore the competitive effects of various coexisting anions (chloride, nitrate, sulfate, bicarbonate, and phosphate) on F¯ removal, 100 mL of 10 mg/L F¯ solution was added to each solution, that is, containing five concentration levels (1, 10, 20, 50, and 100 mg/L) of the coexisting ions. The residual F¯ concentration in the aqueous phase was analyzed by a F¯ ion meter equipped with an Orion F¯ ion selective electrode (Thermo Scientific Orion, USA). 2.4. Batch Adsorption and Kinetic Study. The adsorption isotherms were studied by using the initial fluoride concentration of 5−100 mg/L. For the experiment, 100 mL of each concentration was taken in a 250 mL conical flask with adsorbent dose 0.4 g/L, pH 5.0, contact time 45 min, stirring speed 400 rpm and at three different temperatures (303, 313, and 323 K). Kinetics study experiment was carried out by adding 0.4 g of adsorbent into 250 mL conical flasks, containing 100 mL of 10 and 20 mg/L F¯ solution at three different temperatures 303, 313, and 323 K. The sample was stirred for 50 min; after the regular interval of 5 min, 5 mL suspension was taken out and analyzed for residual F¯. The % adsorption and adsorption capacities (qe and qt) of F¯ were calculated using equations 1−3:

resulted from the synergistic reactions among the component species. But tetrametallic mixed oxide has not been used to date for fluoride adsorption. From the literature survey, we concluded that bimetallic oxide is a better adsorbent for F¯ than monometallic oxide; likewise, the trimetallic oxide is better than bimetallic oxide. In this context, we have successfully developed a tetrametallic oxide adsorbent for F¯ adsorption. Synthesis of tetrametallic oxide adsorbent is a unique kind of synthesis in the world of multimetallic adsorbents. The tetrametallic oxide has better adsorption capacity (except few cases), high thermal stability, good regenerability, and wide pH applicability than other multimetallic adsorbents, because of the improved synergistic reactions among the components. Adsorption capacity depends on experimental conditions also, that is, the volume of solution, dose, and initial fluoride concentration. On varying experimental conditions, adsorption capacity varies up to a certain level. Another unique feature is its high fluoride removal rate (0.27 min−1), that is, the adsorption of maximum (98.5%) fluoride ion from fluorinated water occurs within 45 min in permissible limit. The objective of this study is to determine and compare the adsorption performances of Fe−Al−Ce−Ni (FACN) in fluoride removal from aqueous solutions at different temperatures in view of adsorption equilibrium, kinetic, and thermodynamic studies. Equilibrium data were fitted to Langmuir, Freundlich, Temkin, and Dubinin−Radushkevich (D−R) equations to determine the best-fitted isotherm. The kinetics and thermodynamics of fluoride adsorption onto FACN in aqueous systems were investigated by considering the effects of fluoride concentration, adsorbent dose, temperature, and pH. The kinetic adsorption results were analyzed using pseudo-first-order and pseudo-second-order reactions, intraparticle diffusion and Bangham’s model as well as column study. The thermodynamic parameters were calculated by using Langmuir constant (KL), solute distribution coefficient (Kd), and equilibrium fluoride concentration (isosteric adsorption enthalpy), respectively, to compare the different thermodynamic calculation methods. Furthermore, the reusability of FACN for fluoride removal was examined. As adsorption relies on temperature, the appraisal of thermodynamic parameters portrays the feasibility of adsorption process and spontaneity of the system. The improved adsorption capacity of developed adsorbent has been checked by using three van’t Hoff plots, to ensure the feasibility of the process. This work provides elaborate thermodynamic study by using three van’t Hoff plots for the calculation of the thermodynamic parameter, whereas no literature has been reported which contains more than two van’t Hoff plots.

%adsorption = (Co − Ce)100/Co

2. MATERIALS AND METHODS 2.1. Materials and Reagents. The chemicals used were iron(III) nitrate [Fe(NO3)3·9H2O], aluminum nitrate [Al(NO 3 ) 3 ·9H 2 O], ceric ammonium nitrate ((NH 4 ) 2 [Ce(NO3)6]), nickel(II) nitrate [Ni(NO3)2·6H2O], sodium fluoride (NaF), sodium hydroxide (NaOH), hydrochloric acid (HCl), and sulfuric acid (H2SO4). All chemicals used were of analytical reagent (A.R.) grade and without further purified. 2.2. Synthesis of Tetrametallic Oxide Adsorbent. The FACN adsorbent was synthesized by the coprecipitation method.5 Briefly, 37.051 g of Al, 29.08 g of Ni, 40.4 g of Fe, and 86.4 g of of Ce salts were mixed in 1000 mL of double distilled water in the same molar ratio of 0.1. The mixed metal salt solution was stirred until a homogeneous solution was

(1)

adsorption capacity at initial concentration (qe) = (Co − Ce)V /m

(2)

adsorption capacity at time (qt) = (Co − C t)V / m

(3)

2.5. Column Study. A glass column of 2 cm diameter and 50 cm length was used for the evaluation of column performance of FACN. The bottom of the column was filled with glass wool to provide support to the adsorbent then 2 g of FACN was filled in the column with a bed height of 2 cm. A known concentration of 10 mg/L F¯ solution was passed through the column with a constant flow rate of 0.5, 1.0, 1.5 mL/min. The treated water was collected in a conical flask and B



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adsorbent surface favors the adsorption of F¯ and negative surface disfavors F¯ adsorption. The pHpzc of FACN was found greater than the pH, so it shows electrostatic adsorption.25 At low pH protonation of the adsorbent surface occurs (pH < pHpzc)

analyzed for the residual F¯ concentration at different intervals of time and with different flow rates. 2.6. Characterization and Analysis. The adsorbent was characterized by many techniques to prove their structural features viz. Brunauer−Emmett−Teller (BET), field emission scanning electron microscopy (FESEM), energy dispersive Xray (EDS), X-ray diffraction (XRD), atomic force microscopy (AFM), Fourier transform infrared spectroscopy (FTIR), X-ray photoelectron spectroscopy (XPS), thermal gravimetric analysis (TGA), zeta (zeta sizer nano series-ZS90), and differential scanning calorimetry (DSC). The total surface area and pore size distribution were determined by N2 adsorption−desorption isotherm using BET method; a 25 mg sample was degasified by 3 h at 300 °C. The surface morphology of the adsorbent was characterized by FESEM at both stages before fluoride adsorption and after fluoride adsorption. About 5 mg of the sample was first placed on the carbon tape and sputter-coated with gold to make it conducting and then placed in the specimen holder. EDS analysis of the adsorbent was carried out to determine the mean element proportions of the adsorbent. The composition of the adsorbent was investigated by XRD. The XRD analysis was performed with Cu Kα radiation (40 kV, 40 mA) at a scanning speed of 2°/min and a scanning step of 0.02°. For AFM analysis, 20 mg of adsorbent was taken into 1 mL DMSO and then sonicated for 30 min. A small drop of solution was transferred immediately on a 1 × 1 cm2 glass slide to form a thin layer and dried at room temperature, after which it was taken for AFM analysis. FTIR was used for the characterization of the functional group involved in the structure and changes in bond length. The adsorbent was dried at room temperature and mixed with KBr to form a pellet and then, the FTIR spectra were recorded in the range of 4000−400 cm−1. The zeta sizer was used for zeta potential measurement. Additionally, the surface chemistry and chemical state of the element before and after adsorption were determined by XPS. TGA and DSC studies were carried out to monitor the characteristic physical and chemical changes in the FACN adsorbent.

M−OH + H+ → M−OH 2+

At high pH deprotonation of the adsorbent surface occurs (pH > pHpzc) M−OH + OH− → M−O− + H 2O

3.2. Batch Reactive Adsorption Studies. 3.2.1. Variation of pH. The variation of pH on F¯ % adsorption and qe were studied by using 10 different samples having a pH range of 2− 11. From Figure 1A, it was found that on varying pH of the

3. RESULTS AND DISCUSSION 3.1. Point of Zero Charge and Zeta (ζ) Potential. Point of zero charge (pHpzc) is a good tool for the analysis of the surface charge of adsorbent. The point of zero charge is the stage of the material when the electrical charge density on a surface is zero and, in that condition, adsorption takes place through the ion−exchange process. For the study of pHpzc eight samples of 0.1 M NaCl solution were prepared, and the pH of solution was maintained from 2 to 9 by using dilute H2SO4 and NaOH. A 0.1 g sample of adsorbent was added to each solution. After 24 h, each solution was separated into two different tubes (batch A and B). Batch A tubes were used for calculation of the ζ potential, and batch B tubes were used for measurement of change in pH. Each sample was centrifuged, and the supernatant was utilized for the ζ potential measurement. The ζ potential for pH 2−7 was positive while that for pH 8 and 9 was negative. The curve was plotted as ζ potential vs pH. From Figure S1A, the pHpzc was found to be 6.2. For batch B, pHpzc was also calculated by recording the final pH of each solution. The curve plotted against ΔpH [pHinitial− pHfinal] vs pHinitial and pHpzc was found to be 6.2 (Figure S1B). If pH < pHpzc, the surface charge would be positive and if pH > pHpzc, it would be negative. Therefore, the positive

Figure 1. (A) Effect of pH and (B) adsorbent dose on the adsorption of F¯ ion.

solution, % adsorption (Figure 1Aa) was increased up to pH 5.0, that is, 99.84%, after that it was decreased slightly up to 96.88 at pH 11, and the same trend was found in the qe26 (2.496−2.42 mg/g) (Figure 1Ab). The decrease in % adsorption and adsorption capacity were due to the negatively charged adsorbent surface, which tends to repel the negatively charged F¯ ion electrostatically. Additionally, the concentration of OH− ion also increases, which competes with the F¯ for the active binding sites of the adsorbent. The mechanism by which F¯ tends to adsorb by FACN can be explained mainly by the hard and soft (Lewis) acids and bases principle. 3.2.2. Variation of Adsorbent Dose and Distribution Coefficient. The effect of varying adsorbent dose from 0.1 to 0.5 g/L on % adsorption was studied. Figure 1Ba shows that on increasing the adsorbent dose up to 0.4 g/L, % adsorption increased up to 98.08, after that on increasing adsorbent dose % C



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Figure 2. (A) Effect of temperature and (B) coexisting anions on F¯ adsorption.

thickness of the limit layer decreases because of the high propensity of the F¯ ions to escape from the adsorbent surface to the arrangement stage which results in a decrease in adsorption as temperature increases. The increase in temperature increases the rate of diffusion of the fluoride ions toward the adsorbent surface and at the same time, some adsorbate ions acquire enough thermal energy and diffuse away from the monolayer surface. This implies that the rate of desorption was higher than the rate of adsorption and proved that F¯ adsorption on FACN was exothermic. The temperature of the adsorption medium could be important for an energydependent mechanism in the F¯ ions binding process.27 3.2.4. Influence of Inorganic Anion. Anions present in water can hinder the sorption processes. The effect of coexisting anions, chloride, sulfate, nitrate, phosphate, and bicarbonate on the F¯ removal on FACN adsorbent, was studied by the batch method (Figure 2B). Chloride does not have much effect on F¯ adsorption, but phosphate, nitrate, carbonate, and sulfate show a significant effect on F¯ removal. Carbonate shows a significantly higher reduction in F¯ adsorption efficiency due to change in pH as well as the competing effect of this co-ion for the active site of the adsorbent. The affinity sequence for the anion adsorption on this adsorbent is carbonate > phosphate > sulfate > nitrate > chloride.4 3.2.5. Variation of Equilibrium Concentration. The variation in the initial F¯ concentration was investigated by using different fluoride concentration solutions in the range of 5−100 mg/L with adsorbent dose 0.4 g/L, time 45 min, pH 5.0 (Figure 3A). The % adsorption decreased significantly with increasing initial F¯ concentration. This can be explained by the

adsorption was found constant up to 0.5 g/L (98.09%), so 0.4 g/L adsorbent dose was taken as standard. At lower adsorbent dose, the number of active adsorption sites was fewer for the adsorption of F¯. An increase in the active sites for F¯ adsorption is possible by increasing the adsorbent dose.11 To clarify the nature of the adsorbent surface, the distribution coefficient (Kd) was used, which describes the binding ability of an adsorbent surface to adsorbate. The Kd values expressed as Kd = (C0 − Ce) × V /(Ce × m)

(4)

Tian et al. reported that the Kd value for an element showed two different situations at a given pH value; if the surface was homogeneous, the Kd value stayed the same with adsorbent dose, whereas if the surface was heterogeneous, the Kd value first would decrease with increasing the adsorbent dose and then increase with increase in dose. As shown in Figure 1Bb, the Kd value decreased with increasing the adsorbent dose and then increased with dose, indicating that the surface of the FACN is heterogeneous. 3.2.3. Variation of Temperature. Temperature variation of the sample affects the adsorption process by changing the molecular connections and the solvency of the adsorbate. Adsorption studies were done at three different temperatures (303, 313, and 323 K) using FACN for initial F¯ ion concentration of 5−50 mg/L. The % adsorption was decreased from 95.8 to 72.88%, 95.1 to 71.933%, and 95.1 to 67.64% at 303, 313, and 323 K, respectively (Figure 2A). These results indicate that high temperature works against the fluoride adsorption by FACN. At high temperature, the D



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3.2.6. Variation of Contact Time. To study the effect of contact time on the adsorption process, contact time varied from 5 to 50 min, and the curve plotted against contact time vs % adsorption and adsorption capacity. Figure 3B elucidates that on increasing the contact time up to initial 45 min the adsorption capacity and % adsorption (80.4−98.5%) increase up to maximum, and after 45 min of contact time the % adsorption and adsorption capacity was constant. This was probably because of attainment of equilibrium on the adsorption process, decreased number of available adsorbent site and very low concentration of fluoride ion in solution.28 3.3. Equilibrium Adsorption Isotherms. Adsorption isotherms are represented as the number of adsorbate molecules per unit mass of adsorbent as a function of equilibrium concentration in bulk solution at a constant temperature. In this study, adsorption of fluoride onto FACN was carried out, and four isotherm models were tested, which are tabulated in Table S1, namely, Langmuir, Freundlich, Temkin, and D−R isotherm.29−33 Langmuir model assumes that a fixed number of adsorption sites are available on the surface of the adsorbent; each site can take up one molecule only, that is, monolayer adsorption, and the energy of the adsorption is constant. 1/qe = 1/(qm × KL × Ce) + 1/qm

(5)

or Ce/qe = 1/(KL × qm) + Ce/qm Figure 3. Effect of (A) initial F¯ concentration and (B) contact time on F¯ adsorption.

The maximum adsorption capacity (qm) is calculated by plotting the curve 1/qe vs 1/Ce (Figure 4A) which gives the maximum adsorption capacity as well as the Langmuir constant which relates to the energy of adsorption. The maximum adsorption capacity is 250 mg/g at 303 K (Table 1). A similar observation was reported for many adsorbents (Table 2). The

presence of more F¯ ions per unit number of adsorbent sites, which leads to a saturation of the coordination sites.25

Figure 4. Adsorption curves: (A) Langmuir, (B) Freundlich, (C) Temkin, and (D) D−R at (a = 303 K, b = 313 K, c = 323 K). E



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Kf and n are Freundlich constants which denote the approximate indicators of adsorption capacity and intensity of adsorption in the adsorption process, respectively. The curve plotted against ln qe vs ln Ce which gives the value of Kf and n, as 3.22−2.84 and 1.92−1.88, respectively (Table 1). The Freundlich curve plotted at three different temperatures (Figure 4B). The smallest value of 1/n, that is, 0.52 and higher value of n signifies an active interaction between FACN and fluoride. The value of Kf decreases with an increase in temperature indicating the exothermic nature of the adsorption. The Temkin model was tested for equilibrium description at three different temperatures. This model considers the adsorbent−adsorbate interactions.

Table 1. Adsorption Parameters of F¯ Adsorption curve

303 K Langmuir 0.0044 0.97 2.14 250 0.047 Freundlich 0.52 0.99 1.92 1.17 3.22 Temkin 552.44 0.012 1.012 0.93 D−R 161.8 8.76 0.94 238.9

1/qm R2 KL (L/mg) qm (mg/g) RL 1/n R2 N Ln Kf Kf (mg1−1/n/gL) BT ln AT AT R2 qm (mg/g) K (mol2/J2) ×10−3 R2 Em (J/mol)

313 K

323 K

0.0050 0.96 2.12 200 0.045

0.008 0.96 2.00 125 0.044

0.53 0.992 1.88 1.085 2.96

0.52 0.98 1.92 1.04 2.84

543.27 0.020 1.02 0.92

562.98 0.016 1.016 0.93

170.45 8.17 0.91 247.3

172.8 6.5 0.93 277.3

adsorbent

qe (mg/g) 8.21 10.47 70 3.56 198 178 384.6 45.45 23.99 250 (present study)

1 ln Ce n

(9)

ln qe = ln qm − K ϵ2

(10)

ϵ = RT × ln[1 + (1/Ce)]

(11)

1 2K

(12)

For the D−R model, the curve plotted between ln qe and ϵ2 is shown in Figure 4D. The qm was 172.8 mg/g, and this indicates the maximum monolayer value of fluoride adsorption per unit weight of membrane (Table 1). From the D−R equation and according to the value of mean free energy Em, an adsorption process may be one of three types: (i) physisorption if Em < 8.0 kJ/mol, (ii) ion−exchange if Em = 8.0−16.0 kJ/mol, and (iii) chemisorption if Em > 16.0−400 kJ/mol. The Em value in the present system is 0.23, 0.24, and 0.27 kJ/mol at 303, 313, and 323 K, respectively. It confirms that as the adsorption process and concentrations of both adsorbate and adsorbent are involved in the rate-determining step; the adsorption may be a physisorption. The experimental results of the adsorption models were compared, and the best fit correlation value of Freundlich model was found to be greater than others. The order of adsorption model to fit in the adsorption data according to R2 value is Freundlich > Langmuir > D−R > Temkin model; other details are given in Table 1. 3.4. Kinetics of Reactive Adsorption. The kinetics of fluoride removal process was analyzed by first-order and second-order kinetics, which can be expressed as

(6)

The RL value indicates the shape of isotherm and the nature of the adsorption process, unfavorable for RL > 1, linear for (RL = 1), favorable for (0 < RL < 1), and irreversible (RL = 0). The value of RL for FACN was found in the range of 0.047−0.044, which is less than unity, this shows the effective interaction between the FACN and fluoride. The adsorption data were also analyzed by the Freundlich isotherm model, which is suitable for heterogeneous surfaces. This model explains both mono- as well as a multilayer adsorption process, it also explains that the adsorbent has surfaces of varied affinities or adsorption on heterogeneous surfaces.34 ln qe = ln K f +

qe = (RT × ln A T)/BT + (RT × ln Ce)/BT

Em =

feasibility of Langmuir isotherm can be described by a separation factor (RL), which can be determined by the following equation: RL = 1/(1 + KL × Co)

(8)

AT is the equilibrium binding constant (L/g) and BT (J/mol) is Temkin constant related to the heat adsorption (J/mol). Temkin curve is plotted over a range of temperature as shown in Figure 4C. The Temkin constant related to the heat of adsorption was in the range of 0.54−0.56 kJ/mol. It has been reported that the heat of sorption value, intraparticle diffusion model > Bangham model > pseudo-first-order; other details are given in Table 3. 3.6. Adsorption Thermodynamics. Thermodynamic parameters are very important to check the spontaneity and feasibility of the adsorption process. They provide necessary information to design an adsorption process. Usually, thermodynamic parameters, that is, heat of enthalpy (ΔH), Gibbs free energy (ΔG), and entropy (ΔS) are major parameters which govern the feasibility and spontaneity. For calculating these parameters, KL (Langmuir constant), Kd (solute coefficient distribution), and Ce (equilibrium fluoride concentration) which changes with temperature are given in Tables 4−7. In this study, thermodynamic parameters were calculated by eqs 23−30.45−47

Langmuir constant

(23)

ΔG = ΔH − T ΔS

(24)

ΔG = −RT × ln KL

(27)

(28)

The ΔS could be calculated by the Gibbs− Helmholtz equation ΔS = (ΔH − ΔG)/T

(29)

ΔG = −nRT

(30)

where n is Freundlich constant. Using the above equation, the plot of ln Kd vs −1/T, ln Ce vs 1/T, and ln KL vs −1/T is shown in Figure 8(A,B) and Figure 9A. The values of thermodynamic parameters can be directly obtained from the slope and intercept of the van’t Hoff plots (Tables 4−7). The adsorption of fluoride was more favorable at 303 K and the adsorption gradually decreased upon increasing the temperature, which reveals that at higher temperature the adsorbate−adsorbent interaction weakened. The high temperature is not favorable for F¯ adsorption. In the present study, a lower temperature is found to be suitable for adsorption process, which is evident by ΔG values obtained at three different temperatures. The negative ΔG value indicated the feasibility of the reaction and spontaneous nature of the adsorption at a given temperature. The increasing value of ΔG with an increase in temperature indicates a decreasing trend in the degree of feasibility of F¯ adsorption. The ΔH values were observed in the range of −20 to 40 kJ/mol for physical adsorption and −80 to 400 kJ/mol for chemical adsorption. In the present study, the negative value of ΔH indicates that the type of adsorption is physisorption, and it is exothermic in nature. The positive ΔS suggests the increased randomness at

van’t Hoff Plot

ΔH (kJ/mol) ΔS (kJ/mol) R2

(26)

ln Ce = −ln Ko + ΔH /RT

Table 4. Thermodynamic Parameters for F¯ Adsorption ΔG (kJ/mol)

ln KL = ΔS /R − ΔH /RT

Equilibrium fluoride concentration

Distribution coefficient ln Kd = ΔS /R − ΔH /RT

(25)

−8.3 (303 K) −7.9 (313 K) −7.4 (323 K) −6.45 (303 K) 0.11 (303 K) 0.95 (303 K) I



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Table 5. Thermodynamic Parameters for F¯ Adsorption Kd T/K

ΔH

ΔS

ΔG = ΔH − TΔS (kJ/mol)

ΔG = −8.314 × RT × ln Kd (kJ/mol)

T/K

T/K

Co

303

313

323

(kJ/mol)

(kJ/mol)

303

313

323

303

313

323

R2

10 20 30 40 50

21.9 9.3 7.3 4.7 4.5

20.9 8.3 5.8 4.3 4.0

19.1 6.2 5.1 4.1 3.9

−5.6 −16.1 −14.2 −6.1 −7.7

0.006 0.034 0.030 0.007 0.012

−3.5 −2.6 −2.3 −8.3 −11.6

−3.4 −2.7 −2.38 −8.3 −11.8

−3.4 −2.7 −2.4 −8.4 −11.9

−7.7 −5.6 −5.3 −4.7 −3.8

−7.9 −5.5 −4.6 −4.3 −3.6

−7.9 −4.9 −4.4 −4.1 −3.5

0.97 0.93 0.91 0.90 0.85

has lower Ea values (5−40 kJ/mol) while the chemical adsorption has higher Ea values (40−800 kJ/mol). The activation energy of fluoride adsorption onto FACN adsorbent was found lower than the range of physical adsorption. The significantly lower value of the activation energy supports the high F¯ affinity of FACN and the feasibility of the reaction.48 3.8. Column Studies. Industrial scale operations aim to achieve the maximum flow rate possible for operating at higher throughput. Hence, the effect of flow rate was studied by taking 2 g of adsorbent in a 2.0 cm inner diameter column, with bed height 2 cm and passing water containing 10 mg/L F¯ at different flow rates (0.5, 1.0, and 1.5 mL/min). The breakthrough curves obtained are shown in Figure S2. Higher flow rates resulted in earlier breakthrough due to insufficient contact time. The adsorption capacity decreases with an increase in flow rate from 0.5 to 1.5 mL/min, due to insufficient contact time, whereas it increases at a low flow rate up to 0.5 mL/min. To describe the fixed bed column behavior and scale it up for industrial applications, various breakthrough models shown in Table S5 viz., Thomas, Yoon− Nelson, Bohart−Adam, and Clark models were used to fit the experimental data of the column studies. Each one of these models is different from the other in terms of the type of adsorption isotherm, inclusion or exclusion of chemical reaction, significance, and the type of rate law used.49,50 The Thomas model is a general model that is widely used to describe column performance where external and internal diffusion resistances are extremely small. This model assumes that the adsorption behavior follows second-order reversible reaction kinetics, Langmuir isotherm without axial dispersion.

Table 6. Isosteric Enthalpy for F¯ Adsorption Co

ΔH (kJ/mol)

−ln Ko

R2

20 30 40 50

−14.12 −12.38 −4.93 −6.20

0.001 0.002 0.020 0.009

0.87 0.95 0.91 0.92

Table 7. Thermodynamic Parameters for F¯ Adsorption T/K

303

313

323

ΔG = −nRT ΔS = (ΔH − ΔG)/T

−4.81 0.576

−4.86 0.06

−5.15 0.059

the solid−solution interface during the adsorption of F¯ onto FACN. It is generally known that a large positive ΔS or large negative ΔH will be helpful for spontaneous adsorption. From all the three van’t Hoff plots the result comes out to be the same, that is, ΔG < 0, ΔH < 0, ΔS > 0. 3.7. Activation Energy (Ea). To calculate the activation energy of the process utilizing the Arrhenius equation we use

k 2 = A e−Ea /(RT )

(31)

The logarithm form of eq 31 illustrates a linear relationship of the Arrhenius equation and can be drawn as eq 32. log k 2 = log A − Ea /(2.303RT )

(32)

The plot of log k2 vs 1/T generates a straight line (Figure 9B); the values of Ea (1.1 kJ/mol) and A (0.055 g/mg/min) were calculated from the slope and intercept of the plot, respectively. The magnitude of Ea indicates the type of adsorption, which can be physical or chemical. The physical adsorption process

Figure 8. (A) van’t Hoff plot w.r.t. Kd for F¯ adsorption and (B) van’t Hoff plot w.r.t. Ce for F¯ adsorption. J



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Figure 9. (A) van’t Hoff plots w.r.t. KL for F¯ adsorption and (B) activation energy plot for F¯ adsorption.

Figure 10. Breakthrough curves: (A) Thomas model, (B) Yoon−Nelson model, (C) Bohart−Thomas model, and (D) Clark model at (a) 0.5, (b) 1.0, and (c) 1.5 mL/min.

The Clark model is based on the use of a mass-transfer concept in combination with the Freundlich isotherm. The Clark model assumes that the shape of the mass transfer zone is constant, and all the adsorbate is removed at the end of the column.

Linear form: ln(Ce/Co − 1) = k THqom /Q − k THCot

(33)

Yoon and Nelson developed a less complicated model, which requires no detailed data concerning the characteristic of the adsorbate, properties of the adsorption bed, and the type of adsorbent. In this model, it is assumed that the rate of decrease in the probability of adsorbate molecule is proportional to the probability of the adsorbate adsorption and the probability of an adsorbate breakthrough on the adsorbent. Linear form: ln[Ce/(Co − Ce)] = k YNt − k YNτ

Linear form: (36)

A = exp(K CNoZ /v)

(37)

r = K CCo

(38)

Breakthrough curves for different flow rates were fitted to the models as shown in Figure 10. All four models have well fitted the experimental data with very good R2. The data of Table S6 show that kTh and KC increase with an increase in flow rate, and kBA and KYN decrease with an increase in flow rate. This performance due to F¯ adsorption by FACN follows secondorder reversible reaction kinetics and Freundlich adsorption isotherm.51 3.9. Regeneration and Desorption Studies. The exhausted FACN were regenerated with eluents such as dilute HCl and dilute NaOH. All the regeneration experiments were

(34)

Bohart-Adams assumed that equilibrium is not instantaneous. The adsorption capacity remains on the adsorbent and the concentrations of the adsorbing species are proportional to the rate of adsorption. Linear form: ⎡ k NoH ⎤ ln(Ce/Co) = kABCot − ⎢ ⎣ v ⎥⎦

ln(Co/Ce)n − 1 − 1 = A e−rt

(35) K



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carried out at room temperature. Mineral acids elute F¯ from the exhausted adsorbent as HF and NaOH as NaF. The desorption and reuse of FACN at Co = 10 mg/L were evaluated for six regeneration cycles and adsorption capacity decreased from 98.9 to 50%, as presented in Figure 11A and

Table 9. Calculation Results of Surface Area, Pore Size, and Pore Volume Using BJH Analysis parameter BJH pore volume (cc/g) BJH pore radius (Å) BJH surface area (m2/g)

Table S7, and desorption data with different concentrations of NaOH and HCl (0.2−1.0 M) were presented in Figure 11B and Table S8. The maximum removal efficiency for reuse occurred after desorption using 1.0 M NaOH solution and an overly high concentration (1.5 M) was harmful to reuse which may be because of the residual OH− on the FACN or the destruction of the structure of the adsorbent. Therefore, the optimal concentration of the NaOH solution for regeneration was 1.0 M. 3.10. Adsorbent Characterization. BET analysis was carried out to determine the surface area, pore volume, and average pore radius of adsorbent by using different models. For the surface area analysis, 25 mg of adsorbent was degassed at 300 °C for 3 h and then analyzed. The whole process was completed in 24 h. In addition, to the multipoint BET analysis, the pore size distribution curve determined from Barrett− Joyner−Halenda (BJH) analysis shows (Figure S3) that the main pore size distribution of the adsorbent and the complete data of adsorption and desorption shown in Tables 8 and 9. Adsorption−desorption isotherm exhibits a type I isotherm. Table 8. Calculation Results of Surface Area, Pore Size, and Pore Volume Using BET Analysis parameter

FACN

FACN-F

184.47 29.41 51.43

197.339 19.57 31.87

FACN

FACN-F

0.252 31.880 21.546 342.988 102.03 1858.961

19.548 21.477 24.667 297.206 102.210 1444.314

The surface area of the adsorbent was 184.47 and 197.339 m2/g before and after treatment, respectively. The surface area was calculated using the multipoint BET equation in the p/po range of 0.05−0.30. It confirms that the surface area of the adsorbent increases after the F¯ treatment which is a clear demonstration of a decrease in pore size by the accumulation of F¯ in the pores of the adsorbent.10 This considerable change in the structure of the adsorbent increases the surface area and decreases pore size of the adsorbent because of the fluoride adsorption. Improved surface area favors the structural change occurring during the adsorption process.10 The pore size distribution mainly concentrates in a range less than 51 nm, revealing that the composite was a mesoporous material with a significantly large specific surface area. The surface characteristics of the adsorbent were determined by FESEM method, the adsorbent was analyzed before and after the treatment of fluoride. The surface morphology of the adsorbent presents a cotton-like structure having many pores and a rough surface having copious bumps on the surface of the adsorbent. The pores may be used to adsorb the F¯ on its upper surface. Figure S4(A1, A2, B1, B2) shows the FESEM image of adsorbent before and after fluoride adsorption at 1 and 2 μm, respectively. The elemental composition of the adsorbent before and after F¯ adsorbed was determined by EDX. Images in Figure S5(A, B) designate the main components of adsorbent before the F¯ is adsorbed, that is, Ce, Al, Fe, Ni, and O, and after fluoride is adsorbed, they are Fe, Al, Ce, Ni, O, and F. No other peak was detected in the EDX pattern, indicating the purity of the FACN adsorbent. The presence of a fluoride ion peak in the EDX spectrum after fluoride adsorption confirmed adsorption. The mapping images further support the elemental composition of the adsorbent after and before F¯ adsorb (Figure 12). No other peak was detected in the EDX pattern, indicating the purity of the FACN adsorbent. AFM images were used to characterize the surface morphology of adsorbent: the uniformity and grain size is shown in Figure S6. The result of the FTIR study of FACN mixed oxide showed sharp absorption peaks located 3742, 3599, 1620, 1519, 1326, 689, 588, and 520 cm−1 (Figure 13). After fluoride adsorption, the peaks decrease in intensity, but there is a slight increase in frequency due to stronger binding of fluoride with metal ions. The slight shift of peaks was at 3749, 3597, 1632, 1531, 1335, 695, 602, 527, 430 cm−1. The extra peak appearing at 430 cm−1 corresponds to metal F¯ vibrations. A general slight shift of frequency and a decrease in intensity of peaks are attributed to F− ions adsorbed to FACN. The FTIR spectroscopic study after adsorption confirmed that the F¯ ion has a strong binding ability with the surface of FACN mixed oxide adsorbent. The XPS technique was used to analyze the composition and chemical state of FANC adsorbent before and after fluoride adsorption. The XPS spectra were studied to understand the adsorption mechanism (Figure 14 and Figures S7−S9). The

Figure 11. (A) Cyclic adsorption−regeneration run; (B) NaOH and HCl for regeneration of adsorbents.

surface area (m2/g) pore volume (cc/g) average pore size (nm)

A D A D A D

L



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Figure 12. Mapping images of FACN and FACN-F.

adsorption, a new F 1s peak was found at 684.56 eV, indicating the existence of fluoride on the adsorbent surface. The occurrence of a new peak at 684.56 eV in high resolution after fluoride removal was assigned to the F 1s spectrum, indicating that fluoride was bound to the adsorbents. These results confirmed that the adsorbent was effective for fluoride removal. On the other hand, the binding energy (BE) values of O 1s, Fe 3d, Al 2p, Ce 3d5/2,3/2, and Ni 3d3/2 of the assembly were remarkably shifted as summarized in Table 10, indicating Table 10. Binding Energy of Respective Elements of Adsorbent before and after Fluoride Adsorbed elements Fe Fe−F Al Al−F Ce Ce−F Ni Ni−F O O−F F

Figure 13. FTIR Spectra of FACN and FACN-F.

binding energy (eV) 710.58, 719.69, 710.96, 720.36, 72.17 73.17 881.61, 888.50, 881.66, 888.92, 854.12, 860.28, 854.55, 860.92, 529, 531 528, 530 684.56

730.13 731.08

897.42, 897.47, 872.12, 872.62,

899.90 900.60 881.80 882.03

that an alteration of the local bonding environments occurred. Likely, it is due to the specific adsorption mechanism (Figure 15) of fluoride onto the surface of the assembly.30,53 XRD analysis of FACN before and after fluoride adsorbed was conducted using the X-ray diffractometer in the range of the diffraction angle 2θ = 20−90°. A comparison with the Joint Committee on Powder Diffraction Standards (JCPDS) file confirms the that the peaks present corresponded mainly to CeO2, Fe2O3, NiO, Al2O3 (JCPDS no. 78-0694, 86-0550, 78-

Figure 14. XPS spectra of FACN and FACN-F.

wide scan XPS spectra showed that the main elements in the shell of the adsorbent were O, Fe, Al, Ce, and Ni. After M



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Figure 15. Mechanism for fluoride adsorption on FACN adsorbent.

method. The percent adsorption and maximum adsorption capacity of the adsorbent obtained was 98.5% and 250 mg/g, respectively, at pH 5.0, the contact time was 45 min, and adsorbent dose was 0.4 g. The adsorption equilibrium, kinetics, and thermodynamics of F¯ adsorption have been studied in the present work. The equilibrium data have been analyzed by Langmuir, Freundlich, Temkin, and D−R isotherm models. Freundlich is a best-fit model among the tested models. This shows that the adsorption process follows the multilayer adsorption. The calculated values of Ea and Em are 1.1 and 0.23 kJ/mol, respectively. The D−R model indicated that the adsorption process was physisorption. Kinetic studies showed that the adsorption adhered to the pseudo-second-order model because the theoretical and experimental sorption capacities were in excellent agreement (R2 ≥ 0.997). F¯ removal rate using adsorption (0.27 min−1) was much faster than reactive adsorption (0.04 min−1). Column studies results are showed that the adsorption behavior follows second-order reversible reaction kinetics which is also proved by pseudo-second-order kinetic model. The change in Gibb free energy was negative at all temperatures, indicating the feasibility and spontaneity of the adsorption process. The value of ΔH and ΔS indicated the exothermic nature of adsorption and the randomness at the solution-solid interface, respectively.

0429, 83-2080) in the FACN (Figure S10) XRD pattern, while in FACN-F including the metal oxide peaks, AlF3, CeF2, CeF3, NiF2, and FeF3 (JCPDS no 77-0252, 30-0452, 81-2268, 850481) were also observed which clearly demonstrates the fluoride adsorption by each element of the mixed oxide. Thermal gravimetric analysis (TGA) of synthesized FACN adsorbent showed the high thermal stability of the compound. The result of the TGA (Figure S11(A)) of FACN shows that the first degradation (initial weight loss of 7.5%) occurred at 100 °C which might be due to the evaporation of physically adsorbed water molecules, thereafter the weight loss of 11.5% up to 400 °C is mainly due to the loss of water from the hydroxyl groups, and later above 400 °C, the curve eventually flattened. Further, the compound has no degradation up to 800 °C; that is, the compound is stable up to 800 °C. To obtain DSC results of FACN, the DSC experiment was analyzed up to 600 °C, and the DSC thermogram showed an exothermic peak at 100 °C corresponding to water evaporation, and no further peak was found in the thermogram (Figure S11(B)).

4. FIELD STUDY The fluoride removal property of FACN adsorbent was also ratified with the field sample taken from nearby Banasthali village. Almost 0.4 g of adsorbent was taken into 100 mL of fluorinated water and shaken for 40 min at room temperature. Table S9 represents the results of the field trial. The response of FACN adsorbent was promising because it removed the fluoride ion to below the permissible limit and other water quality parameters from the tested water (Table S9). Hence FACN adsorbent can be used effectively for fluoride removal from water.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jced.8b00024. Additional figures and tables as described in the text (PDF)

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5. CONCLUSION The present study investigated the adsorption performance onto FACN. The FACN was synthesized by the coprecipitation

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. N



Article

ORCID

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Dinesh Kumar: 0000-0001-5488-951X Author Contributions

Dinesh Kumar designed the project and organized the research. Rekha Sharma and Ankita Dhillon carried out the experiments, analyzed the data and wrote the manuscript. All authors reviewed the manuscript. Funding

We gratefully acknowledge support from the Ministry of Science and Technology and Department of Science and Technology, Government of India under the Scheme of Establishment of Women Technology Park, for providing the necessary financial support to carry out this study vide letter No, F. No SEED/WTP/063/2014 Notes

The authors declare no competing financial interest.

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ABBREVIATIONS

Symbol Unit Abbreviation

Co = mg/L initial fluoride concentration Ce = mg/L equilibrium fluoride concentration qe = mg/g adsorption capacity at equilibrium concentration qt = mg/g adsorption capacity at time V = mL volume of solution m = g mass of adsorbent Kd = L/g distribution coefficient KL = L/mg equilibrium adsorption constant related to the affinity of binding sites RL = - separation factor qo = mg/g maximum monolayer adsorption capacity Kf = (mg/g) (L/mg)1/n adsorption capacity of the adsorbent n = - intensity of adsorption AT = L/g binding constant BT = J/mol temkin constant R = J/mol universal gas constant T = °C Temperature Em = kJ/mol mean energy kap = min−1 apparent rate constant k1 = min−1 first order rate constant k2 = g/mg/min second order rate constant ki = mg/g/min intraparticle diffusion rate constant Xi = - Constant ΔG = kJ/mol Gibbs energy of adsorption ΔG⧧ = kJ/mol Gibbs energy of reactive adsorption ΔH = kJ/mol enthalpy of adsorption ΔH⧧ = kJ/mol enthalpy of reactive adsorption ΔS = kJ/mol entropy of adsorption ΔS⧧ = kJ/mol entropy of reactive adsorption Ea = kJ/mol activation energy of reaction A = - pre-exponential factor

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