Synthesis, Characterization, and Application as a Mercury(II) Sorbent

Sep 28, 2002 - Acrylamide was graft polymerized onto banana stalk, BS (Musa .... from the main fruit market of the Municipal Corporation, Trivandrum, ...
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Ind. Eng. Chem. Res. 2002, 41, 5341-5352

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Synthesis, Characterization, and Application as a Mercury(II) Sorbent of Banana Stalk (Musa paradisiaca)-Polyacrylamide Grafted Copolymer Bearing Carboxyl Groups I. G. Shibi and T. S. Anirudhan* Department of Chemistry, University of Kerala, Kariavattom, Trivandrum 695581, India

Acrylamide was graft polymerized onto banana stalk, BS (Musa paradisiaca), using the ferrous ammonium sulfate/H2O2 redox initiator system in an aqueous medium. The effects of reaction variables such as time, temperature, and monomer and initiator concentrations on the percentage grafting were studied. A new adsorbent carrying a carboxylate functional group at the chain end was synthesized by surface modification of polyacrylamide-grafted BS (PGBS-COOH). Infrared spectroscopy and acid-base titration were used to confirm graft copolymer formation and carboxylate functionalization. X-ray diffraction and SEM studies were carried out to investigate the crystallinity and morphology of the adsorbents. XRD studies indicated that the grafting of polyacrylamide resulted in a consequent decrease in crystallinity. SEM studies of PGBS-COOH clearly indicated that polyacrylamide grafts deposited more on the surface of the unit cell than in the intercellular gaps. A probable mechanism for graft copolymerization and surface functionalization is also suggested. The use of this adsorbent material for the removal of Hg(II) from water and wastewater was investigated using the batch adsorption technique. The adsorption of Hg(II) on the adsorbent was found to be pH-, time-, concentration-, and temperature-dependent. The optimum pH range for the process was found to be 6.0-9.0. Maximum removals of 99.3 and 84.1% were observed at Hg(II) concentrations of 50 and 100 mg/L, respectively, in this pH range. Removal of Hg(II) is adversely affected by increasing initial Hg(II) concentration. The adsorption process follows pseudo-second-order kinetics. Kinetic parameters as functions of initial concentration and temperature were calculated. Hg(II) adsorption was found to decrease with increasing ionic strength. The L-type adsorption isotherm obtained for the adsorbent indicated a favorable process and fitted the Langmuir isotherm model well. The adsorption capacity for Hg(II) calculated using the Langmuir isotherm equation was 137.89 mg/g at 30 °C, which increased to 210.50 mg/g at 60 °C. Thermodynamic parameters such as the changes in free energy, enthalpy, and entropy were calculated to predict the nature of adsorption. The isosteric heat of adsorption was found to be 49.89 ( 1.31 kJ/mol and was independent of surface coverage. Chlor-alkali industrial wastewater samples were treated with this adsorbent to demonstrate its efficiency in removing Hg(II) from industrial wastewater. Recovery of Hg(II) after adsorption and regeneration of the adsorbent for several cycles can be carried out by treatment of the loaded adsorbent with 0.2 M HCl. Introduction The removal of heavy metals from natural waters and industrial wastewaters is a technological challenge with respect to environmental and industrial applications. The traditional precipitation process does not always achieve sufficient metal removal to meet pollution control limits, and the use of activated carbons and synthetic ion-exchange resins is often quite expensive. Adsorbent materials derived from suitable waste biomass and agricultural byproducts can be used for the effective removal and recovery of heavy metals from wastewaters. The advantages of using these materials are their effectiveness in reducing the concentrations of heavy metals to very low levels and the use of inexpensive materials. Various types of waste biomass have been studied for their heavy metal uptake capacities and suitability to be used as a basis for adsorbent development.1-4 However, the application of native biomass was found to be limited by leaching of organic * To whom correspondence should be addressed. E-mail: [email protected].

substances such as lignin, pectin, tannin, etc., into the solution. New research in this direction indicates that surface modification of the adsorbent materials prevents the leaching of organic substances and also enhances the adsorption efficiency of the adsorbent.5 Graft polymerization on solids followed by functionalization has been used as an important technique for modifying physical and chemical properties of the adsorbent and improving its adsorption capacity. Graft polymerization of waste biomass can also prevent organic substances from leaching out of the biomass. This modification can also increase the stability of the adsorbent materials, which is an important aspect of the commercial development of biosorbent materials. Materials such as tea leaves,6 coconut coir,7,8 bagasse pith,9 saw dust,10 and cellulose11 have been used as polymer supports for the preparation of adsorbent materials having different functional groups. Previous investigations12,13 have demonstrated that carboxylate functional groups substituted into the backbone of a polymerized material can act as adsorption sites for metal ions from aqueous solutions. From the literature,14 it is clear that poly-

10.1021/ie020245f CCC: $22.00 © 2002 American Chemical Society Published on Web 09/28/2002

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acrylamide-grafted lignocellulosic material with carboxylate functionalities exhibits a high potential for the removal of heavy metal ions from aqueous solutions. The suitability of lignocellulosic materials as precursors for graft polymerization is determined by their local and bulk availability. Banana stalk (Musa paradisiaca, axis of inflorescence) another lignocellulosic material is a widely available natural material in the tropical developing world. After the banana fruit has been used, the banana stalk has no use and is usually discarded, sometimes causing disposal problems. If banana stalk could find a suitable application in controlling environmental pollution, it would be another milestone in our march toward economic development. Banana stalk basically contains cellulose, hemicellulose, pectin, tannin, and lignin.15 The present paper reports on investigations on the synthesis, characterization, and sorption properties of polyacrylamide-grafted banana stalk bearing carboxylate functional groups and its application for the removal of Hg(II) ions from water and industrial wastewater. Experimental Section Preparation of Adsorbent. Banana stalk (BS) used as the starting material was obtained from the main fruit market of the Municipal Corporation, Trivandrum, India. The raw material was washed with distilled water to remove dirt and grime and air-dried. Before the initiation of graft polymerization, the BS was defatted with an ethanol/benzene (1:2 v/v) mixture in a Soxhlet apparatus for 8 h. After refluxing, the material was washed with alcohol, dried, and sieved to -80+230 mesh size (average size ) 0.096 mm). Graft polymerization of acrylamide (AAm) onto BS was conveniently carried out in water using a ferrous ammonium sulfate (FeAmSO4)/H2O2 initiator system.16 A sample-to-liquid ratio of 1:200 was used. One gram of BS was transferred into a 500-mL reaction flask equipped with a mechanical stirrer, and then 20 mL of FeAmSO4 solution (8.0 × 10-3 M) was added. Then, 20 mL of H2O2 (0.12 M) and 160 mL of water were added to the reaction vessel. Purified nitrogen was passed through the vessel (5 min) for deaeration. The polymerization was started by adding 5.0 g of monomer (AAm). The vessel was kept in a water bath maintained at 70 °C. After specified time intervals, polymerization was terminated by quenching the vessel in ice-cold water. The homopolymer was separated by water extraction with a Soxhlet apparatus and dried at 70 °C. The effects of the concentrations of FeAmSO4 [(1.0-12.0) × 10-4 M]; H2O2 [(0.2-1.4) × 10-2 M]; and monomer, AAm (5.0-40.0 g/L), at different temperatures were studied to determine the effects of these parameters on the grafting yield. The graft yield (%) was calculated as

graft yield (%) ) (dry weight of grafted BS dry weight of original BS) × 100/dry weight of original BS (1) From the above experiments, it was possible to design operating conditions for preparing bulk quantity of polyacrylamide-grafted BS (PGBS). The general scheme for the preparation of carboxylate-functionalized PGBS is shown in Scheme 1. The dried PGBS (50 g) was refluxed with a known volume (500 mL) of ethylenediamine [(en)2] in toluene for 8 h and then washed with

Scheme 1. Preparation of PGBS-COOH

toluene and dried. One part by weight of the material was then refluxed with an equal part by weight of succinic anhydride in 1,4-dioxane at pH 4.0 for 6 h. The excess succinic anhydride was washed out with 1,4dioxane and dried. The carboxylic acid bound PGBS (PGBS-COOH) was sieved, and particles having an average diameter of 0.096 mm were used for the experiments. Adsorbents were examined by scanning electron microscopy (SEM). The samples were mounted on aluminum micro-studs, gold coated, and analyzed with an S-2400 Hitachi scanning electron microscope based at the Sree Chitra Thirunal Medical Science & Technology, Trivandrum, India. X-ray diffraction patterns were obtained with a Siemens D 5005 X-ray unit using Nifiltered Cu KR radiation. FTIR spectra of the adsorbents were recorded between 4000 and 400 cm-1 using the KBr method with a Perkin-Elmer IR-180 spectrophotometer. The BET surface areas of the adsorbents were calculated from the N2 adsorption isotherms using a Quantasorb surface area analyzer (model QS/7). TG/ DTA plots of the samples were obtained using a Metler Toledo Thermoflex instrument. A sample portion of 10 mg was heated from room temperature to 600 °C at 10 °C/min in flowing N2. The Boehm titration17 method using NaOH and NaHCO3 was used to estimate the total acidic group and carboxylic acid group contents, respectively, in the adsorbent. Potentiometric titration18 was carried out to determine the surface charge density as a function of pH of the adsorbent. The cationexchange capacity was determined by the column process using 1 M NaNO3 as the eluent at a flow rate of 0.5 mL/min. The effluent was titrated against a standard alkali solution to determine the total amount of H+ released. The apparent density of the adsorbent was also calculated from the displacement of nitrobenzene according to the pycnometric method. Adsorption Experiments. Batch adsorption experiments were carried out to determine the optimum pH for Hg(II) adsorption. Tests were performed by shaking 0.1 g of the adsorbent with 50 mL of aqueous Hg(II) solution of the desired concentration at various pH’s (2.0-9.0) and at room temperature (30 °C) in several 100-mL stoppered bottles for 4 h using a temperaturecontrolled water bath shaker. The initial pH’s of these solutions were adjusted using 0.1 M HCl and NaOH, and all of the pH measurements were carried out with a Systronics pH meter (model µ-360). The solutions were stirred at a constant speed of 200 rpm. At the end of the reaction period, each reaction mixture was filtered (with Whatmann no. 42 filter paper) to separate the supernatant and PGBS-COOH. The pH of the supernatant was measured at the end of the experiments. The equilibrium concentration of Hg(II) in supernatant was

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determined using a mercury analyzer (Perkin-Elmer model MAS-50A). Several batch experiments were carried out to study the effect of a number of variables such as the initial Hg(II) concentration (50-150 mg/L), temperature (30-60 °C), pH (2.0-9.0), and adsorbent dose. Kinetic batch experiments were performed by agitating 50 mL of Hg(II) solution of 50-150 mg/L concentration at pH 6.0 with 0.1 g of adsorbent for retention times varying from 1 to 450 min using water bath shaker. At the end of each predetermined time interval, the sorbate was separated by filtration, and the supernatant was analyzed for Hg(II) concentration. Adsorption isotherms were obtained by the batch method, employing 0.1 g of adsorbent and 50 mL of Hg(II) solution at different concentrations (75-750 mg/L). The pH’s of the solutions were adjusted to 6.0. The samples were shaken in a temperature-controlled water bath shaker at different temperatures (30-60 °C) for 4 h. The adsorbent was separated by filtration, and the quantity of Hg(II) adsorbed was determined as described earlier. The amount of Hg(II) adsorbed on the adsorbent was calculated by subtracting the final concentration in solution from the initial concentration. Batch adsorption experiments were also conducted to determine the adsorption capacity of commercial synthetic polymer-based cation-exchange resin. For this purpose, the commercial carboxylate-functionalized cation exchanger Ceralite IRC-50 (similar to Amberlite IRC-50) was obtained from Central Drug House, Bombay, India. No pretreatment was applied to the Ceralite IRC-50, which was used as obtained. Desorption and Regeneration Studies. Desorption studies were carried out by the batch technique under similar conditions. About 50 mL of industrial wastewater containing 50 mg/L of Hg(II) was treated with 0.1 g of adsorbent for 4 h at pH 6.0. The solution was then filtered, and the filtrate was analyzed for Hg(II). The Hg(II)-laden adsorbent was washed with deionized water to remove any unadsorbed Hg(II). The spent adsorbent was then placed in a 100-mL stoppered bottle containing 50 mL of 0.2 M HCl. The suspension was shaken for 6 h and then filtered. The concentration of desorbed Hg(II) in the filtrate was determined. The above adsorption and desorption procedures were repeated for four cycles using the same adsorbent to assess the effect of desorption on the ability of the adsorbent to take up Hg(II) again. Following each desorption step, the adsorbent was washed with deionized water, dried to constant weight, and reloaded with Hg(II). All of the experiments were done in duplicate or triplicate and were also reproduced on different days. The maximum variation in the batch adsorption data between duplicate/ triplicate runs was 3.7%, and that between experiments performed on two different days was 4.8%. The data presented both in the tables and in the figures are average values. Results and Discussion Mechanism of Polymer Grafting. The graft copolymerization of AAm onto BS was carried out at different temperatures ranging from 30 to 70 °C (Figure 1). The graft yield increased with increasing temperature. This result can be attributed to the following factors: (i) enhanced swellability, (ii) increased mobility of the monomer and initiator, (iii) higher rate of diffusion of

Figure 1. Effects of reaction time and temperature on the graft copolymerization of AAm onto BS.

Figure 2. Effect of ferrous ammonium sulfate concentration on the graft coplymerization of AAm onto BS.

the monomer and initiator, and (iv) higher rate of initiation and propagation of the graft. The results also show that the graft yield increases at the beginning and subsequently levels off. The percentage graft yield increases with increasing concentration of FeAmSO4 up to 8.0 × 10-4 M in the case of higher temperatures (above 40 °C), and thereafter, the graft yield decreases slightly (Figure 2). In the case of 30° C, the percentage graft yield increases up to 8.0 × 10-4 M and then assumes steady values. A possible reason for the above trend is the formation of a greater number of active sites as a result of the increase in FeAmSO4 concentration up to a limiting value. In addition, the concentration of the initiating free radical might also decrease because of the mutual combination. Therefore, an 8.0 × 10-4 M FeAmSO4 concentration was used for the study of all other effects. Figure 3 shows that the graft yield increased with increasing H2O2 concentration up to 1.2 × 10-2 M and remained constant thereafter. This might be due to the enhanced rate of termination, which balances the rate of propagation.19 The effect of monomer was studied by

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Figure 3. Effect of H2O2 concentration on the graft coplymerization of AAm onto BS.

Figure 5. FTIR spectra of BS and PGBS-COOH.

where ∼BS- is the phenolate ion of the lignin, ∼BS• is the phenolic radical, and M is the monomer (AAm).

propagation

∼BS-M• + M f ∼BS-M2• ∼BS-Mn-1• + M f ∼BS-Mn•

(5)

termination ∼BS-Mn• + ∼BS-Mn• f PGBS (graft copolymer) (6) ∼BS-Mn• + Fe3+ f Fe2+ + ∼BS-Mn (PGBS/graft copolymer) (7)

Figure 4. Effect of monomer concentration (AAm) on the graft copolymerization of AAm onto BS.

varying the concentration from 5.0 to 40.0 g/L (Figure 4). The graft yield increased with increasing monomer concentration. The enhancement in grafting with the increasing monomer concentration can be attributed to the following factors: First, the complexation of BS with monomer required for enhancing monomer activity would be favored at higher monomer concentration. A second cause might be the gel effect, resulting from an enhanced solubility of polyacrylamide in its own monomer.20 This would, consequently, increase the viscosity of the reaction medium and reduce the rate of termination through the coupling of the growing polymer chain. For the graft copolymerization of AAm onto BS, initiated by the Fe2+/H2O2 initiator system, the following mechanism is proposed

initiation

H2O2 + Fe2+ f Fe3+ + OH- + OH• (2) ∼BS- + Fe3+ f Fe2+ + ∼BS• •

•

∼BS + M f ∼ BS-Μ

(3) (4)

Adsorbent Characterization. FTIR spectra of BS and PGBS-COOH are shown in Figure 5. The spectra of BS exhibit a broad absorption peak at about 3400 cm-1 and a sharp peak at 2945 cm-1, indicative of the -OH group and C-H stretching from a -CH2 group, respectively. The spectra of PGBS-COOH exhibit a broad signal around 3300 cm-1 representing the overlap of O-H, C-H, N-H, and C-O stretching vibrations.21 The peaks at 1626 cm-1 (amide carbonyl group) and 1568 cm-1 for PGBS-COOH are due to the presence of an aliphatic amide group, and the weak band at 1180 cm-1 is due to the N-H stretching vibration. These observations clearly indicate the formation of a polymeric chain (backbone) in PGBS-COOH. The additional peaks at 1696 cm-1 (νCdO) and 1455 cm-1 (νC-O) indicate the presence of -COOH groups21 in PGBS-COOH. The thermal behavior of BS and PGBS-COOH was examined by primary TG and DT analyses (Figure 6). It was observed that the thermal stability of PGBSCOOH is higher than that of the original BS. The thermogravimetric curves obtained were characterized by determining the decomposition temperatures, i.e., the initial decomposition temperature, TD1, and the temperature at the maximum rate of weight loss, Tmax. The weight loss curves exhibit a number of kinetically distinct steps. The weight losses at these peak positions are noted. The final temperature for each step is taken as the valley point on the rate of weight loss versus temperature plot. In the case of original BS, TD1 is approximately 200 °C, and in the case of PGBS-COOH, TD1 is 280 °C. In general, the thermal stability of PGBS-COOH is comparable to that of typical phenolformaldehyde resins.22 The weight loss of BS was

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Figure 6. TG and DTA curves of BS and PGBS-COOH.

Figure 8. Scanning electron micrographs of BS and PGBSCOOH.

Figure 7. X-ray diffraction patterns of BS and PGBS-COOH.

relatively higher than that of PGBS-COOH. The TG curve exhibits an initial small weight loss of 10.0% for BS starting at 70 °C and ending at 120 °C, whereas, for PGBS-COOH, a weight loss of 7.0% due to the loss of physically adsorbed water starts at 90 °C and ends at 120 °C. BS exhibits an endotherm close to 100 °C. No other features are visible at higher temperatures. A temperature break observed for PGBS-COOH in the range of 120-180 °C indicates the loss of water, which is possibly associated (coordinated) with carboxylate groups on the adsorbent surface. The DTA curve shows a broad endothermic peak in the range of 90-180 °C, coinciding with the loss of both physically adsorbed and coordinated water in the adsorbent system, and a major endothermic peak at 420 °C. The temperature for 10% weight loss (T10) and Tmax are the two main criteria used to indicate the thermal stability of polymers.23 The higher the values of T10 and Tmax, the higher the thermal stability of the system. The values of T10 and Tmax for PGBS-COOH were found to be 190 and 450 °C, respectively, indicate the high thermal stability of the adsorbent. X-ray diffraction patterns of BS and PGBS-COOH are shown in Figure 7. It is obvious that no new phase formation due to the grafting of polyacrylamide onto BS occurs. However, the results suggest that BS and its

grafted products have corresponding crystalline structures. Native BS shows scattering at 2θ ) 13°, 27°, and 34°. XRD data for BS indicate that BS has a crystalline domain of cellulose substrate. In the crystalline regions, cellulose molecules (major component of BS) are arranged in ordered lattices in which hydroxyl groups are bonded by strong secondary forces. Diffraction maxima at 20° and 34° can be attributed to the crystalline region of cellulose. XRD studies have shown that, upon graft polymerization, a significant decrease in crystallinity occurs. The broad peak centered at 21° appears on PGBS-COOH because of various lattice planes and is a combination of peaks of various intensities (lowerordered fractions).24 Thus, some rearrangement in the morphology of cellulosic chains in PGBS-COOH occurs as a result of grafting. The decrease in crystalline domains in PGBS-COOH results in the loss of tensile strength of the grafted chain and consequently enhances the free mobility of grafted chain. This change is speculated to take place by small lateral shifts of the chain segments that are loosened from one another by grafting, giving rise to an increase in the lower-ordered fractions. The free mobility of grafted chains is also enhanced by the breakdown of the hydrogen-bonding network structure during the chemical reaction. SEM was used to examine the surface characteristics of BS and PGBS-COOH (Figure 8). The surface morphology of BS is different from that of PGBS-COOH. The SEM image of BS has unit cells that run longitudinally with parallel orientations. The intercellular gaps, in the form of longitudinal cavities, can be clearly marked as the unit cells are partially exposed. To hold the unit cells firmly in the stalk fibers, the intercellular gaps are filled by binder lignin and fatty substances.

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Figure 9. Effects of pH on the surface charge density of BS and PGBS-COOH at different ionic strengths.

The SEM image of PGBS-COOH clearly shows that polyacrylamide (PAAm) grafts have been deposited more on the surface of the unit cell than in the intercellular gaps. The PAAm units in PGBS-COOH are amorphous, reducing the degree of crystallinity and causing a reduction in tensile strength of the grafted chain. The overall fracture mode is ductile and clearly corresponds to the PAAm grafts phase.25 The deposits of PAAm grafts onto the surface of the unit cell, resulting in deep voids between the overgrowths of the unit cells. PGBSCOOH presents large pores, which are also attributed to rearrangements of the polymer chains during the chemical reaction. The zero point of charge pH, pHzpc, is defined as the pH of the suspension at which the surface charge density σ0 ) 0. σ0 was calculated from the potentiometric titration data using the equation

σ0 )

F{(CA - CB) + [OH-] - [H+]} A

(8)

where F is the Faraday constant; A is the surface area of the suspension (cm2/L); CA and CB are the concentrations of acid and base (equiv/L), respectively, after each addition during the titration; and [OH-] and [H+] are the numbers of equivalents of OH- and H+ ions, respectively, bound to the suspension surface (equiv/ cm2). The points of intersection of the σ0 versus pH curve (Figure 9) show that the pHzpc of BS and PGBS-COOH occurred at 8.0 and 5.5, respectively. The decrease in pHzpc after surface modification indicates that the surface becomes more negative after grafting and functionalization and, thus, facilitates the electrostatic interaction with metal cations. On the other hand, the low pHzpc value exhibited by PGBS-COOH indicates that it has more acidic functional groups than the original BS, or in other words, it is less basic, which is advantageous for the removal of metal cations. The characteristics of PGBS-COOH are as follows: surface area, 110.3 ( 3.8 m2/g; porosity, 0.51 ( 0.06 mL/g; pHzpc, 5.5 ( 0.3; range of particle size, 0.063-0.177 mm (average size, 0.096 mm); cation exchange capacity, 2.38 ( 0.12 mequiv/g; total acidity, 2.10 ( 0.36 mequiv/g; carboxylate content, 1.84 ( 0.41 mequiv/g; and apparent density, 0.72 ( 0.08 g/mL. Effect of pH on Hg(II) Adsorption. The pH of the aqueous solution is one of the important controlling

Figure 10. Effects of pH on the removal of Hg(II) from aqueous solution by BS, PGBS-COOH, and Ceralite IRC-50.

parameters in the adsorption process. This is partly because H+ ions themselves are strong competing adsorbates and partly because the solution pH influences the speciation of metal ions and the ionization of surface functional groups. The effect of pH on the adsorption of Hg(II) by PGBS-COOH, BS, and Ceralite IRC-50 is presented in Figure 10. It can be seen that the adsorption capacities increase with increasing pH, reaching a plateau value in the pH range of 6.0-9.0 for both PGBS-COOH and Ceralite IRC-50; however, BS is effective as a Hg(II) adsorbent only within the narrow range of pH 4.0-6.0. The maximum removals of 99.3 and 84.1% for PGBS-COOH, 75.1 and 69.3% for Ceralite IRC-50, and 49.4 and 39.0% for BS were observed at initial Hg(II) concentrations of 50 and 100 mg/L, respectively, at pH 6.0. In particular, an abrupt increase in the Hg(II) removal efficiency of Ceralite IRC50 occurred when the pH varied from 4.5 to 6.0. In this pH range, the Hg(II) removal increased from 31.5 to 75.1% and from 24.3 to 69.3% at initial concentrations of 50 and 100 mg/L, respectively. The adsorption of Hg(II) ions on original BS can be attributed to Coulombic interactions. Coulombic interactions can be reflected from the adsorption of cationic species rather than anionic species on an adsorbent. BS is basically a lignocellosic material. Lignocellulosic materials contain polar functional groups, which have exhibited relatively strong Coulombic adsorption to cations as well as intrinsic adsorption to anions.26 The data clearly show that PGBS-COOH is more effective than BS and Ceralite IRC-50 as a Hg(II) adsorbent. The decrease in adsorption with increasing H+ ion concentration (low pH) indicates that the adsorption process occurs via ion exchange. At lower pH values, the H+ ions compete with Hg2+ ions for the exchange sites in the system. It was shown that the final pH is always less than the initial pH. This indicates that, as the metal ions are bound to the adsorbent, H+ ions are released from the -COOH functional groups into the solution. This leads to the conclusion that PGBS-COOH probably acts as an acid-form ion exchanger. Perusal of the Hg(II) speciation diagram27 clearly indicates that, in the range of highest removal efficiency, the dominant species present is Hg(OH)+. The pHzpc of PGBS-COOH was found to be 5.5, and below this pH, the surface charge of the adsorbent is positive.

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Figure 11. Plot of amount of H+ released versus amount of Hg(II) adsorbed.

ferent initial concentrations plotted as a function of time. Hg(II) adsorption consisted of an initial fast process, which was completed after about 30 min of contact time, and a slow process that gradually approached an equilibrium condition. Depending on the initial concentration, about 70-85% removal of Hg(II) was achieved during the first 30 min of contact time, whereas only 15-20% of additional removal occurred in the following 4 h of contact time. The sharp increase in the removal rate during the initial stage might be because of the availability of sufficient exchange sites. On the basis of the results in Figure 12, the equilibrium period in forthcoming experiments was chosen to be 4 h. The adsorption density increased from 24.78 to 51.89 mg/g, but the removal efficiency decreased from 99.3 to 69.2% when the initial Hg(II) concentration was increased from 50 to 150 mg/L. This obviously shows that the removal of Hg(II) is concentration-dependent. The results suggest that PGBS-COOH can remove most of the Hg(II) from water if its concentration is below 50 mg/L. It can be concluded that higher removal at low concentration is important in terms of industrial applications. Adsorption Kinetics. Sorption kinetics ultimately controls the process efficiency. The kinetic constants of metal adsorption, which can be used to optimize the residence time required for completion of the sorption reaction, were calculated using the kinetic data. The mechanism of Hg(II) removal is thought, basically, to be complexation and ion exchange, and in the absence of stoichiometric data, it can be represented as k

S + M y\ z [SM] 1 k

Figure 12. Effects of contact time and initial concentration on the adsorption of Hg(II) on PGBS-COOH.

The sorption of Hg(II) onto the PGBS-COOH releases protons into the acidic medium by an ion-exchange mechanism, resulting in a decrease in the solution pH at equilibrium

PGBS-COOH + Hg(OH)+ h PGBS-COOHg(OH) + H+ (9) According to the above equation, plotting the number of H+ ions released against the Hg(II) uptake should give a straight line of slope 1.0. The plot of protons released from the adsorbent versus Hg(II) adsorbed was, indeed, found to be linear (Figure 11). The slope of the line was found to be 0.89, indicating that the ionexchange process is involved in Hg(II) uptake. The slight deviation from unity might be due to partial consumption of the H+ ions released from PGBS-COOH by OHcomplexed to Hg. At pH’s greater than 5.5, PGBSCOOH becomes negatively charged, and the Hg(II) species is still present mainly as cation [Hg(OH)+]. In this case, adsorption occurs by electrostatic attraction. At very high pH’s, a third process might be involved, namely, retention of the Hg(OH)2 species in the micropores of the adsorbent particles.28 Effect of Initial Hg(II) Concentration and Contact Time. Figure 12 shows the data from the kinetic studies of Hg(II) adsorption by PGBS-COOH for dif-

(10)

where S is an available adsorption site on the PGBSCOOH surface (PGBS-COO-dS), M is a dissolved metal ion, and [SM] is the adsorbed state. k and k1 are the adsorption and desorption rate constants, respectively. The rate equation for the adsorption of divalent metal ions onto PGBS-COOH can be developed by making certain assumptions. In deriving this kinetic expression, the following assumptions were made: (1) Sorption occurs only on localized sites and involves no interaction between sorbed ions. (2) The energy of adsorption is independent of surface coverage. (3) Maximum adsorption corresponds to a saturated monolayer of adsorbates on the adsorbent surface. (4) The process of metal uptake on PGBS-COOH is governed by a pseudo-second-order process. (5) The rate of desorption, k1, is negligible in comparison with rate of adsorption, k. The rate of a pseudo-second-order reaction might depend on the amount of metal ions on the surface of the adsorbent and the amount of metal ions sorbed at equilibrium. This means that the rate of reaction is directly proportional to the number of active sites on the surface of the adsorbent. Consider the pseudo-second-order reaction represented by the equation29

d(S)t ) k[(S)0 - (S)t]2 dt

(11)

where (S)t is the number of adsorption sites occupied on the adsorbent at time t and (S)0 is the total number of adsorption sites available on the adsorbent. In terms

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Figure 13. Pseudo-second-order kinetic plots for the adsorption of Hg(II) on PGBS-COOH at different (A) concentrations and (B) temperatures.

of adsorption quantity, the pseudo-second-order rate equation can be rewritten as

d(qt) ) k(qe - qt)2 dt

(12)

where k is the rate constant of sorption [g/(mg min)] and qe and qt are the amounts of metal ions adsorbed (mg/g) at equilibrium and at time t, respectively. Separating the variables in eq 12 and integrating for the boundary conditions t ) 0 to t ) t and qt ) 0 to qt ) qt gives29

t t 1 + ) qt kq 2 qe

(13)

e

The product kqe2 is actually the initial sorption rate represented as h ) kqe2. The straight-line plots of t/qt versus t with a correlation coefficient (r2) of more than 0.993 at different concentrations and temperatures indicate that the adsorption process follows a pseudosecond-order reaction (Figure 13). From the slopes and intercepts of the plots, the values of k, h, and qe were calculated, and they are presented in Table 1. The data clearly show that the initial sorption rate, h, increases with increasing initial metal concentration. For a decrease in the initial concentration of Hg(II) from 150 to 50 mg/L, the value of k was found to increase from 1.52 × 10-3 to 2.74 × 10-3 g/(mg min). Identical trends in the h and k values with the initial concentration have been reported also by earlier workers who studied the adsorption characteristics of Cu(II) on chitosan30 and

Hg(II) conc (mg/L)

k × 103 [g/(mg min)]

qe (mg/g)

h [mg/(g min)]

r2

50 75 100 150

2.74 2.48 1.86 1.52

24.84 28.95 43.21 53.64

1.69 2.08 3.47 4.38

0.997 0.999 0.999 0.998

temperature (°C)

k × 103 [g/(mg min)]

qe (mg/g)

h [mg/(g min)]

r2

30 40 50 60

1.86 2.40 4.16 7.64

43.21 44.33 45.29 50.33

3.47 4.72 8.53 19.34

0.999 0.999 0.993 0.997

Hg(II) on sulfurized activated carbons.31 The decrease in the k values with increasing concentration might be due to a progressive decrease in covalent interactions, relative to electrostatic interactions, of the sites with lower affinity for Hg(II) that occurs with increasing initial Hg(II) concentration. The values of qe were found to increase from 24.84 to 53.64 mg/g for an increase in the initial Hg(II) concentration from 50 to 150 mg/L. This occurs because more efficient utilization of the adsorptive capacity of the adsorbent is expected because of the greater driving force resulting from the higher concentration gradient. The results shown in Table 1 indicate that the values of qe, h, and k increased with increasing temperature. An increase in temperature favors the adsorption process by activating the adsorption sites. The increase in temperature also increases the chemical forces responsible for adsorption. The activation energy, Ea, was also calculated using the Arrhenius equation. A plot of the ln k versus 1/T was found to be linear (figure not shown). The Ea value calculated from the slope of the plot was found to be 41.05 kJ/mol. The positive value of Ea indicates that a higher solution temperature does favor the removal of Hg(II) by adsorption on PGBS-COOH. The relatively low value of Ea suggests that Hg(II) adsorption is not fully controlled by chemical forces, but that diffusion is also responsible for the adsorption process.32 Effect of Ionic Strength. The effect of ionic strength on the adsorption of Hg(II) on PGBS-COOH was studied at pH 6.0 with different ionic strengths varying from 0.001 to 1.0 M at a fixed initial metal concentration of 50 mg/L. It was found that the adsorption of Hg(II) from solution decreased with increasing NaCl concentration. The amounts of Hg(II) adsorbed per unit weight of adsorbent were found to be 24.95, 24.70, 23.30, 22.63, 22.60, and 17.18 mg/g at ionic strengths of 0.001, 0.01, 0.05, 0.10, 0.50, and 1.0 M, respectively. Adsorption is sensitive to the change in ionic strength if electrostatic attraction is a significant mechanism.33 Thus, it seems that electrostatic attraction plays an important role in the removal of Hg(II) by PGBS-COOH. Lee and Yang2 explain the reduction in metal adsorption as resulting from the presence of Na+ ions competing for metal binding, but the reduction can also be explained on the basis of the presence of different ionic species at different chloride concentrations. Perusal of the Hg(II) speciation diagram27 indicates that an increase in chloride concentration reduces the Hg2+and Hg(OH)+ species through the formation of chloro complexes. Mercury forms some stable chloro complexes, namely, HgCl2, HgCl+, HgOHCl, and (HgCl2)2, that do not appear to be adsorbed to the same extent as Hg2+and Hg(OH)+ ions.

Ind. Eng. Chem. Res., Vol. 41, No. 22, 2002 5349

Figure 14. Adsorption isotherms of Hg(II) for PGBS-COOH and Ceralite IRC-50. The experimental values are shown as points, and the curves were calculated according to the Langmuir model (solid and dashed lines).

Adsorption Isotherm. Adsorption isotherms were recorded to evaluate the potential of the adsorbent for commercial application. The shape of the isotherm, i.e., the plot of amount of metal adsorbed (mg/g) versus the equilibrium concentration (mg/L), gives an indication of whether the adsorption is favorable. Adsorption isotherms are regular, positive, and concave to the concentration axis at all temperatures (Figure 14). According to the slope of the initial portion of the curve, these isotherms were classified as L type in the Giles classification.34 This fact suggests that PGBS-COOH has a high affinity for Hg(II) ions, so that the ions are strongly adsorbed onto the adsorbent. Initially, the adsorption proceeds quite rapidly, after which a slow approach to equilibrium occurs at high concentrations. According to Giles et al.,34 saturation of the surface by Hg(II) ions seems to occur i.e., a complete monolayer covering of Hg(II) on PGBS-COOH is possible under the experimental conditions employed. The adsorption isotherm data at different temperatures were processed using the Langmuir isotherm model, which can be expressed as

Ce Ce 1 ) 0 + 0 qe Q b Q

(14)

where qe is, again, the amount adsorbed at equilibrium (mg/g) and Ce is the equilibrium concentration (mg/L). Q0 and b are the Langmuir constants related to the adsorption capacity and energy of adsorption or binding constant, respectively. The constants Q0 and b obtained from the isotherms, as well as the correlation coefficients (r2), are listed in Table 2. From the values obtained for these parameters, the theoretical Langmuir curves were calculated and are plotted in Figure 14. The good fit of the experimental data at all temperatures shows the applicability of the Langmuir model, thus suggesting the formation of a monolayer of the adsorbate on sorbent surface. It is clear that the Q0 and b

Figure 15. Plots of ln Ce for Hg(II) at constant amount adsorbed as a function of 1/T. Table 2. Langmuir Constants and Free Energy Values for the Adsorption of Hg(II) onto PGBS-COOH temperature (°C)

Q0 (mg/g)

b x 102 (L/mg)

r2

∆G° (kJ/mol)

30 40 50 60

137.89 161.74 191.03 210.50

0.92 1.20 1.70 2.94

0.998 0.991 0.992 0.997

-18.934 -20.255 -21.845 -24.037

values (Table 2) increase with increasing temperature, showing that the adsorption capacity and the intensity of adsorption are also increased with increasing temperature. The removal of metals by weak carboxylic cation exchangers such as Amberlite IRC-64, Amberlite IRC88, Amberlite CG-50, and Duolite ES-468 has been reported.35 All of these adsorbent materials are very expensive and are available for $190-390/kg of resin. Ceralite IRC-50 (similar to Amberlite IRC-50), produced by Central Drug House, Bombay, India, is the least expensive variety of commercial weak cation exchanger available in India; it is based on polystyrene-divinylbenzene copolymer having carboxylate functional groups. The cost of this material is $70/kg. An adsorption isotherm study was conducted to determine the adsorption capacity of Ceralite IRC-50. The experimental isotherm data are shown in Figure 14. The equilibrium isotherm data at pH 6.0 and 30 °C were also correlated using the Langmuir isotherm equation. The maximum adsorption capacity, Q0, and the binding constant, b, of Ceralite IRC-50 were calculated from the Langmuir isotherm (Figure 15) and were found to be 102.3 mg/g and 0.021 L/mg, respectively; these values are considerably lower than those for the newly developed adsorbent material. BS is a waste material obtained in large quantities at no cost in many tropical countries. No price estimate has yet been made on the process of modifying BS to include a carboxylate functionality. Even after considering the expenses for transportation, chemicals, and electrical energy, the relative cost of the final

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product (PGBS-COOH) will most likely be lower than those of commercially available polymeric cation exchangers. Thermodynamic Parameters. Thermodynamic parameters such as the standard Gibbs free energy change (∆G°), enthalpy change (∆H°), and entropy change (∆S°) of the adsorption of Hg(II) with PGBS-COOH were calculated using the equations

∆G° ) -RT ln b1 ln b1 )

∆S0 ∆H0 R RT

(15) (16)

where b1 is the Langmuir constant when the concentration terms are expressed in molarity (mol/L). The van’t Hoff plot ln b1 versus 1/T was found to be linear (figure not shown). The value of ∆H° as calculated from the slope of the plot by regression method was found to be 32.13 kJ/mol, which suggests the endothermic nature of adsorption. This is in accordance with the finding that, when a divalent ion binds onto an adsorbent, ∆H° will be positive36 when 2rH > rA, i.e., 2rH > rHg, where rH ) 0.154 nm is the ionic radius of H+ and rHg ) 0.112 nm is the ionic radius of Hg2+. The values of ∆G° (Table 2) are negative and small, which is an indication of the feasibility of the process. Further, the value of ∆G° becomes more negative with increasing temperature, indicating an increasing feasibility of adsorption at higher temperatures. The positive value of ∆S° [169.13 J/(K mol)] reflects an increase in randomness at the solid/solution interface during adsorption. The relatively small positive value in the system under investigation indicates that no significant structural change in the adsorbent material occurs as a result of adsorption. Isosteric Heat of Adsorption. The heat of adsorption determined at constant amount of sorbate adsorbed is known as the isosteric heat of adsorption (∆Hx). The magnitude of ∆Hx, and its variation with surface loading can provide useful information regarding the nature of the surface and the sorbed molecules. ∆Hx can be calculated by the well-known Clausius-Clapeyron equation

∆Hx d(ln Ce) )dT RT2

(17)

The equilibrium concentration, Ce, at constant amount of adsorbed Hg(II) is obtained from the adsorption isotherm data (Figure 14) at different temperatures. The plots of ln Ce versus 1/T (Figure 15) were found to be linear, and values of ∆Hx were calculated from the slopes of plots. The value of ∆Hx remains almost constant (49.89 kJ/mol) and is independent of surface loading, indicating the absence of lateral interactions between adsorbed Hg(II) ions.37 Testing with Industrial Wastewater. The utility of the adsorbent material was demonstrated by using it to treat wastewater from a chlor-alkali plant situated in Cochin (India). Industrial wastewater samples collected from the plant were characterized by standard methods.38 The compositions of the wastewater samples are reported in Table 3. Because the amount of Hg(II) in wastewater sample 1 was found to be very low (2.86 mg/L), it was spiked with a Hg(II) solution so that the final concentration of Hg(II) was 50 mg/L. Sample 2 was

Figure 16. Effect of adsorbent dose on the removal of Hg(II) from chlor-alkali industrial wastewater by PGBS-COOH. Table 3. Composition (mg/L) of Industrial Wastewater Hg Pb Cd Mg Ca Na PO4 NO3 NH4 Cl BOD COD SS

sample 1

sample 2

50.0 2.7 0.5 25.6 41.2 280.8 10.9 16.5 20.7 398.39 58.4 138.6 358.7

23.8 3.1 0.4 34.8 49.2 291.4 11.7 16.4 28.7 342.34 78.8 231.4 352.6

used as such without alteration of the composition. The typical concentration of Hg(II) in chlor-alkali industry wastewater is in the range of 7.0-35.0 mg/L (data obtained through personal communication with a Chief Chemist of the plant from which samples were collected). The effect of the adsorbent dose on Hg(II) removal by PGBS-COOH is represented in Figure 16. Almost complete (100%) removal of Hg(II) from 50 mL of wastewater was possible with 125 mg (2.5 g/L) and 80 mg (1.6 g/L) of adsorbent for samples 1 and 2, respectively, which is in good agreement with the results obtained from the batch experiments mentioned above. An examination of the curves plotted in Figure 16 also reveals that, at a fixed Hg(II) concentration, the percentage adsorption increases with increasing adsorbent dose; however, the unit adsorption or adsorption per unit mass of adsorbent decreases, with the effect being more pronounced the higher the adsorbent dose. The adsorption of Hg(II) per unit weight of adsorbent decreases from 57.50 mg/g (46.0%) at 0.4 g/L of adsorbent to 20.0 mg/g (100%) at 2.5 g/L for sample 1 and from 46.41 mg/g (78.0%) at 0.4 g/L of adsorbent to 14.88 mg/g (100%) at 1.6 g/L for sample 2. This is because, as the dose of adsorbent (g/L) is increased, there is smaller commensurate increase in adsorption resulting from lower utilization of the adsorptive capacity of the adsorbent. The results of the experiment can be used to develop a mathematical relationship between percentage removal (R) and adsorbent dose (ms) (using

Ind. Eng. Chem. Res., Vol. 41, No. 22, 2002 5351 Table 4. Desorption and Regeneration Data for the Hg(II)/PGBS-COOH System Using 0.2 M HCl cycle

adsorption (%)

desorption (%)

1 2 3 4

98.6 96.3 93.1 90.2

96.3 93.2 91.6 86.7

boundary conditions). This relationship, for which the correlation coefficient (r2) is >0.99, is

R)

R)

ms (5.80 × 10-3) + (7.56 × 10-3)ms for sample 1 (18) ms -3

(1.70 × 10 ) + (9.04 × 10-3)ms for sample 2 (19)

These types of equations can be used to predict the Hg(II) removal for any PGBS-COOH dose under the particular experimental conditions. To test the robustness of the relationships presented in eqs 18 and 19, values of R [percentage Hg(II) removal] for different values of ms were calculated from the above equations. A comparison of the observed and predicted values of R for different values of ms is shown in Figure 16. In each case, the experimental and predicted values are in excellent agreement. Desorption and Regeneration Studies. Desorption studies help to elucidate the mechanisms of adsorption and recovery of the adsorbate and recycling of the adsorbent. Regeneration of the adsorbent might make the treatment process economical. PGBS-COOH loaded with the maximum amount of Hg(II) ion from chloralkali industry wastewater (sample 1) was placed in the desorption medium, and the amount of Hg(II) desorbed in 4 h was measured. Solutions having different concentrations of NaNO3, Na2CO3, Na2SO4, NaCl, NaClHCl, and HCl were evaluated for the extraction of adsorbed Hg(II) from spent adsorbent. From these experiments, 0.2 M HCl was found to be a good reagent for the desorption of Hg(II) from the adsorbent. Desorption efficiency is defined as the percentage extraction of Hg(II) initially loaded onto the PGBS-COOH adsorbent . Table 4 clearly shows that PGBS-COOH can be used repeatedly without significantly losing its adsorption capacity. As shown in Table 4, for four adsorption/ desorption cycles, the Hg(II) adsorption capacity of the PGBS-COOH material decreased from 98.6% (cycle 1) to 90.2% (cycle 4). On the other hand, the recovery of Hg(II) ions in 0.2 M HCl decreased from 96.3% in the first cycle to 86.7% in the fourth cycle. The reduction in Hg(II) uptake in four cycles of adsorption/desorption indicates that the binding sites on the surface of the adsorbent were destroyed or morphologically altered upon repeated exposure to a strong acid environment. The initial mass of adsorbent used for the first cycle, 0.1 g, was reduced to 0.085 g over four cycles of adsorption/desorption, indicating an adsorbent loss of 15%. The loss of adsorbent might have contributed to the reduction in Hg(II) uptake, although it is not known whether HCl has the ability to rupture the structure of the adsorbent. Great care was taken to minimize the loss of adsorbent, so this issue requires further investigation.

Conclusions In this article, the adsorption behavior of polyacrylamide-grafted banana stalk bearing carboxylate functional groups (PGBS-COOH) with respect to Hg(II) ions was investigated by batch techniques. From the experimental results of our study the following conclusions were derived: We have, for the first time, synthesized a new adsorbent (PGBS-COOH) from banana stalk, a waste biomass. The adsorption Hg(II) on PGBS-COOH is dependent on the pH, initial concentration, and temperature of the solution. PGBS-COOH is effective for the quantitative removal of Hg(II) over the pH range 6.0-9.0. About 99.3% Hg(II) removal is possible from aqueous Hg(II) solutions containing 50 mg/L. A pseudosecond-order rate equation can be used to describe the kinetics of Hg(II) sorption on PGBS-COOH at different initial concentrations and temperatures. A temperature study revealed that the adsorption of Hg(II) on the new adsorbent is endothermic in nature. An increase in the NaCl concentration induced a decrease in Hg(II) removal. The adsorption isotherm data for this system can be interpreted using the Langmuir isotherm model. From the temperature dependence of the sorption data, the Arrhenius parameter and the thermodynamic parameters for the process of equilibrium sorption were estimated. The utility of the adsorbent was tested using chlor-alkali industry wastewater samples. The spent adsorbent can be regenerated and reused upon treatment with acid. Acknowledgment The authors are thankful to the Head, Department of Chemistry, University of Kerala, Kerala, India, and the Director, Sree Chithira Thirunal Institute for Medical Science & Technology, Trivandrum, Kerala, India, for providing laboratory and instrumental facilities. Literature Cited (1) Al-Asheh, S.; DuvnJak, Z. Adsorption of Copper by Canola Meal. J. Hazard. Mater. 1996, 48, 83-93. (2) Lee, S. H.; Yang, J. W. Removal of Copper in Aqueous Solution by Apple Wastes. Sep. Sci. Technol. 1997, 32, 1371-1387. (3) Gupta, V. K.; Mohan, D.; Sharma, S. Removal of Lead from Wastewater using Bagasse Fly AshsA Sugar Industry Waste Material. Sep. Sci. Technol. 1998, 33, 1331-1343. (4) Yu, B.; Zhang, Y.; Shukla, A.; Shukla, S. S.; Dorris, K. L. The Removal of Heavy Metals from Aqueous Solutions by Sawdust AdsorptionsRemoval of Lead and Comparison of Its Adsorption with Copper. J. Hazard. Mater. 2001, B84, 83-94. (5) Saito, K.; Iwata, H.; Furusaki, S. Adsorption Characteristics of an Immobilized Metal Affinity Membrane. Biotechnol. Prog. 1991, 7, 412-418 (6) Sing, D. K.; Tiwari, D. P.; Saksena, D. N. Removal of Lead from Aqueous Solution by Chemically Treated Used Tea Leaves. Indian J. Environ. Health 1993, 35, 169-177. (7) Baes, A. U.; Umali, S. J. P.; Mercado, R. L. Ion Exchange and Adsorption of Some Heavy Metals in a Modified Coconut Coir Cation Exchanger. Water Sci. Technol. 1996, 34, 193-200. (8) Baes, A. U.; Okuda, T.; Nishijima, W.; Shoto, E.; Okada, M. Adsorption and Ion Exchange of Some Ground Water Anion Contaminants in an Amine Modified Coconut Coir. Water Sci. Technol. 1997, 35, 89-95. (9) Simkovic, I.; Laszlo, J. Preparation of Ion Exchanger from Bagasse by Cross-linking with Epichlorohydrine-NH4OH or Epichlorohydrine-Imidazole. J. Appl. Polym. Sci. 1997, 69, 25612566. (10) Unnithan, M. R.; Anirudhan, T. S. The Kinetics and Thermodynamics of Sorption of Chromium(VI) onto the Iron(III) Complex of a Carboxylated Polyacrylamide-Grafted Sawdust. Indu. Eng. Chem. Res. 2001, 40, 2693-2699.

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Received for review April 1, 2002 Revised manuscript received August 15, 2002 Accepted August 17, 2002 IE020245F