On-line version ISSN 1816-7950
Water SA vol.36 n.3 Pretoria Apr. 2010
Sunil Kumar BajpaiI,*; Arti JainII
IPolymer Research Laboratory, Department of Chemistry, Govt. Model Science College (Auton.), Jabalpur (M.P.) 482001, India
IIDepartment of Chemistry, Shri Ram Institute of Science and Technology, Jabalpur (M.P.) 482001, India
In this work spent tea leaves were used as a non-conventional, cost-effective sorbent for removal of Cu(II) from aqueous solutions in batch systems. The sorbent was characterised with respect to surface area, pore volume, density, etc. The equilibrium sorption data were applied to various sorption isotherm models, and the order of fitness was: Langmuir > Temkin > Freundlich. The maximum sorption capacity Qo was found to be almost 90.9 and 68.4, as evaluated using Langmuir isotherms at 27ºC and 37ºC respectively. The observed decrease in sorption capacity with temperature indicated the exothermic nature of the uptake process. The kinetic uptake data were best interpreted by a pseudo second-order kinetic model with values of rate constants of adsorption of 1.47 x 10-2 and 3.01 x 10-2 g/mg·min, respectively, for the initial sorbate concentrations of 10 and 20 mg·ℓ-1 at 27ºC. The sorption mean free energy was determined from the Dubinin Radushkevich (DR) isotherm model and was found to be 9.91 kJ·mol-1, indicating ion exchange/chemisorption nature of uptake process. The Cu(II) uptake was found to increase with the pH of the sorbate solution and maximum sorption was observed in the pH range of 1.0 to 4.8. Finally, thermodynamic parameters like ΔGº, ΔHº, ΔSº were also evaluated.
Keywords: Spent tea leaves, copper(II), adsorption, Langmuir model
The presence of heavy metal ions from the transition metal series, viz, Ni, Fe, Cu, etc., in aquatic systems poses a major threat to the environment, due to their acute toxicity to many life forms. Unlike organic pollutants, the majority of which are susceptible to biological degradation, metal ions do not degrade into harmless end-products (Gupta et al., 2001). These metals enter into urban wastewater which, on being used in irrigation, results in accumulation of these metals in the surface soils. After saturation, surface soils release these contaminants into groundwater and soil solution for plant uptake. Wastewater irrigation further enhances the heavy metal content of soils (Singh et al., 2004).
Copper is one of the major contaminants released from metal-finishing, electroplating and electrical industries. In humans, copper toxicity causes itching and dermatisation and keratinisation of the hands and soles of feet (Huang et al., 2007). Severe gastro-intestinal irritation and possible changes in the liver and kidneys occur due to intake of large doses of copper (Ajmal et al., 1998). Inhalation of copper spray increases the risk of lung cancer among exposed workers (Aydin et al., 2008). According to the World Health Organisation (WHO) and United State Environmental Protection Agency (USEPA), the maximum permissible limit of copper in drinking water is 1.3 mg·ℓ-1 (Hach, 2003). Hence the removal of copper from wastewater before its discharge into aquatic systems is extremely important and deserves immediate attention.
Out of various methods employed for the removal of Cu(II) from aqueous solution, adsorption appears to be the most frequently adopted technique (Bajpai and Rohit., 2007). In addition, continuous attempts have been made to develop low cost non-conventional adsorbents for removal of toxic metal ions including Cu(II). These adsorbents include newspaper pulps (Chakravarty et al., 2008), pomegranate peel (ElAshtoukhy et al., 2008) waste tea fungal biomass (Razmousky et al., 2008), potato peels (Aman et al., 2008), soybean hulls (Yadav et al., 2008), pine cone powder (Ofomaja et al., 2009), areca a food waste (Zhang et al., 2008), orange peel (Ningchuan et al., 2008), pecan nutshell (Vaghetti et al., 2009), modified bagasse (Jiang et al., 2009), maple wood sawdust (Rahman and Islam, 2009), etc.
In a continuous attempt to search for potential cost-effective sorbents for removal of toxins from wastewater, we hereby propose spent tea leaves (STL) as a novel sorbent for effective copper removal. Tea is basically the dried and processed leaves of only one species of plant, Camellia sinesis (Mokgalaka et al., 2004). It is consumed by the largest number of people in the world and considered the second-most popular beverage in the world. It is estimated that somewhere between 18 and 20 billion cups of tea are drunk daily on our planet (Razmovski et al., 2008). Canned or bottled tea drinks as well as instant tea drinks are produced on industrial scale by hot water extraction of tea leaves, and producers face a great problem in disposing of the spent tea leaves after extraction. Therefore utilisation of STL is also desirable. The insoluble part of the spent tea leaves consists of mainly cellulose (37%), hemicellulose and lignin (14%), and polyphenols (25%). According to chemical analyses reported by Harler (1963), these components constitute almost 80% of the insoluble portion of spent tea leaves. The presence of hydroxyl and phenolic groups led us to think that these polar groups could work as active sorption sites for the uptake of transition metal ions like Cu (II).
Chemicals copper sulphate, sodium hydroxide and potassium ferrocyanide were obtained from Hi Media Laboratories, Mumbai, India and used as received.
Preparation of spent tea leaves (STL) sorbent
In order to prepare sorbent spent tea leaves (STL), dust tea (Flora Tea Company, Jabalpur, India) was purchased from a local shop and was boiled repeatedly with water until the filtrate was almost clear. The residue was oven-dried at 60ºC for 48 h. The dried sample was ground and sieved to obtain a particle size of 250 to 211 µm and stored in plastic bottles for further use. No other physical or chemical treatment was performed on the spent tea leaves (STL) thus obtained.
Characterisation of STL sorbent
The specific surface area of the sorbent was determined using methylene blue (MB) sorption method, as described elsewhere (Unnithan et al., 2002). In brief, methylene blue aqueous solutions, in the concentration range of 10 mg·ℓ-1 to 50 mg·ℓ-1, were allowed to agitate with 0.1 g of sorbent till the attainment of equilibrium. The concentration of MB solutions was determined at 664 nm. The adsorption isotherm data thus obtained (Fig. 1) was used to calculate specific surface area as per the equation below:
where Mf is the amount of MB (in μmol) adsorbed per 100 g of the sorbent when the surface is completely covered with a monolayer of MB, N is the Avogadro number, and Am, is the cross-sectional area per molecule on the surface (130 A2).
The value of Mfwas obtained by extrapolating the isotherm to the qe axis, and thus taken as point of monolayer coverage. In addition, other physical parameters such as moisture content, specific gravity, pore volume, true density and bulk density were also determined, as described elsewhere (Bajpai and Tankhiwale, 2008). The physical characteristics of STLs are given in Table 1.
Cu(II) uptake studies
The stock solution of Cu(II) was prepared by dissolving a pre-calculated quantity of CuSO4 in double-distilled water to obtain a final concentration of Cu(II) ions of 100 mg·ℓ-1. This stock solution was further diluted to obtain standard solutions with concentrations in the range of 10 to 60 mg·ℓ-1. To carry out adsorption experiments, 50 ml of Cu(II) solution of desired concentration was placed in a 125 ml Erlenmeyer flask containing 0.01 g of STL sorbent, and was agitated in a thermostat-controlled water bath (Rivotek, India) at a constant speed of 50 r/min for 2 h. At the end of the experiment the solution was filtered and supernatant was analysed spectrophotometrically (APHA,1989) to determine the concentration of Cu(II). The per cent sorption and amount of sorbate sorbed in mg per gram of sorbent i.e. qe (mg·g-1) were evaluated using the following equations:
where Co and Ce are the concentration (mg·ℓ-1) of Cu(II) solution at the beginning and at the end of the adsorption experiment, respectively, V is volume of sorbate solution (ℓ) and Wo is the amount (g) of sorbent
Results and discussion
Cu(II) uptake by spent tea leaves (STL)
The chemical composition of the insoluble portion, obtained after hot water extraction of green tea leaves, is cellulose (37%), lignin and structure proteins (14.7%) hot water insoluble proteins (12%) and polyphenols (25%) as major constituents. Being polar in nature, these may have specific binding sites available for adsorption of other molecular or ionic species. For example, polyphenols may bind with Cu(II) ions as shown in Scheme 1. In addition, electron-rich oxygen atom of OH groups of cellulose can also act as binding sites for uptake of Cu(II) ions. Hence, there are good prospects for of adsorption of various types of inorganic/organic compounds onto waste tea residues.
Effect of adsorbent dose on Cu(II) uptake
In order to investigate the effect of sorbent dose on the amount of Cu(II) sorbed per gram of sorbent (i.e. qe,, mg·g-1), experiments were conducted at initial Cu(II) concentration of 40 mg·ℓ-1, by agitating sorbate solutions with varying amounts of sorbent. The results, as depicted in Fig. 2, clearly indicate that for a sorbent dose of 10 mg the Cu(II) uptake was around 40.0 mg·g-1, and continued to decrease with further increase in sorbent dose. This may simply be attributed to the fact that increase in sorbent dose occurs in a greater proportion relative to a corresponding enhancement in Cu(II) uptake. This finally results in a decrease in the amount of Cu(II) sorbed per gram of STL sorbent. Since various parameters involved in the interpretation of kinetic and equilibrium uptake data are evaluated using expressions involving the term q, it is important to assess the conditions for optimum uptake in terms of mg·g-1; therefore, all further experiments were carried out by using 0.010 g of sorbent particles.
Effect of sorbate solution pH on Cu(II) uptake
In the adsorption studies, final pH of the sorption system plays an important role in affecting the extent of sorption. The variation in pH of the sorbate solution not only affects the charge profile of the sorbent but also adds additional H+/OH- ions to the sorption system. These ions may sometimes compete with the sorbate in occupying active sites.
The effect of solution pH on Cu(II) uptake was studied by agitating 0.01 g of STL sorbent in 50 ml of 40 mg·ℓ-1 concentration of Cu(II) solutions with final pH range of 1 to 4.8 at 270C. The results, as shown in Fig. 3, indicate that Cu(II) uptake increases with pH of the solution and later attains near-saturation approaching a pH value of 4.8. The observed findings may be explained as follows: When the pH of the copper (II) solution is quite low (i.e. 1.0) there is an excess of H+ ions in the solution, which compete with Cu(II) ions for adsorption onto binding sites available in the STL sorbent particles. As a result, Cu(II) uptake is quite low. However, as the pH of the solution is increased, the number of H+ ions in the sorbate solution continues to decrease, thus providing greater opportunity for Cu(II) ions to occupy available active sites. As the pH of the sorbate solution reaches 4.8, maximum copper (II) uptake is obtained. Similar observations have been reported by Razmouski et al. (2008) for adsorption of Cu(II) and Cr(VI) onto tea fungal biomass. Here it is noteworthy that the effect of pH could not be studied beyond a pH of 4.8 as the Cu(II) ions began to precipitate as hydroxides in the form of visible flocs. Therefore, beyond pH 4.8, the removal of Cu(II) from the solution is mainly due to precipitation, not adsorption (Alimohamadi et al., 2005).
Equilibrium Cu(II) uptake studies
Studies of adsorption isotherms are essential for the description of how a sorbate distributes itself between the liquid phase and the solid phase, and are useful to optimise the experimental conditions for obtaining maximum sorbate uptake (Emmanuel et al., 2007). The Langmuir model (Langmuir, 1918) was applied to the equilibrium uptake data obtained for sorption of Cu(II) ions from their solutions, with initial concentrations in the range of 5 to 45 mg ℓ-1, onto STL sorbent particles at 27 ºC and 37ºC. The Langmuir isotherm model may be given as:
where Ce is the equilibrium concentration in mg·ℓ-1, Qe (mg·g-1) is the amount of Cu(II) sorbed at equilibrium, b (ℓ·mg-1) is the coefficient related to the affinity between sorbate and sorbent and Qo (mg·g-1), is the maximum sorbate uptake that could be obtained under the given conditions. The isotherms plotted are shown in Fig. 4. The relatively higher regression values of 0.9663 and 0.9477, obtained for Langmuir plots at 27ºC and 37ºC, respectively, indicated fair suitability of this model to the present sorbent-sorbate system. The values of Langmuir constant b and maximum sorption capacity Qo were obtained using slopes and intercepts of the linear plots and are given in the Table 2, along with their corresponding theoretical values obtained by regression analysis. It is quite clear from the data given in Table 2 that experimental values of b and Qo are in fair agreement with those obtained by regression analysis. Here, it is also worth mentioning that the reason for selecting 27ºC and 37ºC as the experimental temperature was that these are the room temperatures frequently found in different parts of the country India, for most of the year, and so experiments can conveniently be carried out at these temperatures.
It is also evident that the value of maximum sorption capacity Qodecreases with temperature thus suggesting that sorption is favoured at lower temperatures. This may simply be explained on the basis of the fact that at higher temperature kinetic energy of sorbate Cu(II) ions is so high that they do not bind firmly with the active sites available on the sorbent surface. Similar types of results have also been reported by other workers (e.g. He et al., 2008; Papageorgiou et al., 2008) It is noteworthy to mention here that we also applied the Freundlich and the Temkin isotherm model on the equilibrium uptake data but the regressions were quite poor (i.e.< 0.8000) and hence data have not been reported. We also compared the maximum sorption capacity Qo, obtained in the present study, with those obtained using other sorbents (see Table 3). A close look at the values displayed reveals that the present sorbent (STL) has a fair maximum Cu(II) uptake value as compared to the other sorbents used in the recent past.
According to Venkata et al. (2002), the essential charactersitics of Langmuir isotherm can also be expressed in terms of a dimensionless constant separation factor RL that is given as:
where Co is the highest initial concentration of sorbate (mg·ℓ-1), and b (ℓ·mg-1) is the Langmuir constant. The value of RL indicates the shape of the isotherm to be either unfavourable (RL >1), linear (RL =1), favourable (0 < RL < 1) or irreversible (RL = 0). RL values between 0 and 1 indicate favourable adsorption. The RL values were calculated using the equilibrium sorption data obtained at 27 ºC and 37ºC, and variation of RL with initial concentration Coare shown in Fig. 5. It can be seen that RL values are in the range of 0.631 to 0.159 and 0.271 to 0.039 for 27 ºC and 37ºC, respectively, thus indicating that although adsorption is favourable at both temperatures, it is more favourable at 27ºC. In addition, at high initial Cu(II) concentrations it is almost irreversible. Similar results have also been reported elsewhere (Razmovski et al., 2008).
Finally, in order to find out the nature of the uptake process, i.e. whether physical or chemical, the equilibrium data, obtained for 27ºC, were analysed using the Dubinin Radushkevich (D-R) isotherm model which is given as:
where Cad is amount of Cu(II) sorbed in mol·g-1, Cm is the maximum amount of Cu(II) that could be sorbed under optimised experimental conditions, B is a constant with a dimension of energy, and ε, Polyanyi potential, is given as:
where R is the gas constant in kJ·mol-1·K-1, T is absolute temperature and Ce is equilibrium concentration in mol·ℓ-1. The linearised form may be given as:
The equilibrium Cu(II) uptake data, obtained at 27ºC, was used to draw a plot between ln Cad and ε2 as shown in
Fig. 6. A higher regression of 0.9672 indicates suitability of D-R isotherm on equilibrium uptake data. The computed value of B, obtained from slope of the linear plot, was 5.1 x 10-3 kJ·mol-2. From the calculated value of B, the mean sorption energy was computed as E = 1/ √ -2B, which is the free energy of transfer of 1 mole of solute from infinity to the surface of the sorbent. The numerical value of B was found to be 9.910 kJ·mol-1, which lies within the prescribed range of 8 to 16 kJ·mol-1 for ion-exchange or chemisorption. Bajpai and Rohit (2009) report a similar finding in a study on uptake of Cr(VI) onto a sawdust ion-exchanger.
Kinetic studies of Cu(II) uptake
The effect of contact time and initial concentrations of Cu(II) solutions on dynamic uptake of Cu(II) was investigated with Cu(II) solutions of initial concentrations of 10 and 20 mg·ℓ-1at 27ºC The results, as depicted in Fig. 7, indicate that for a given initial concentration of sorbate solution, the amount of Cu(II) sorbed per unit mass of sorbent (i.e. qt in mg·g-1) increases with time and then attains saturation value. Moreover, for a given time, the amount of Cu(II) uptake increases with increase in initial concentration of solution. The amounts of Cu(II) sorbed at equilibrium were found to be 36.40 and 62.55 mg·g-1 for the Cu(II) solutions with initial concentrations of 10 and 20 mg·ℓ-1, respectively. Indeed, the initial concentration provides an important driving force to overcome all mass transfer resistances of Cu(II) between solution and solid phase. So, a higher initial concentration shall always enhance the sorption process. It is also indicated from Fig. 7 that maximum Cu(II) uptake takes place in the first 10 minutes and then the sorption process becomes slower and finally attains equilibrium.
The main issue when searching for an appropriate sorption mechanism is to select a mathematical model that not only fits the data with satisfactory accuracy but also complies with a reasonable sorption mechanism. In the present studies 2 kinetic models, namely, the pseudo second-order kinetic model and simple Elovich model, were applied to the kinetic Cu(II) uptake data to make quantitative interpretations
The pseudo second-order model can be presented in the following form (Ho and McKay, 1999):
where Ks is the rate constant of pseudo-second-order model (g·mg-1·min). On definite integration of Eq. (9) for boundary conditions qt = 0 when t = 0, the following form of equation can be obtained.
The initial sorption rate constant, h (mg·g-1 min), at t = 0 can be defined as (Ho and McKay, 1999):
The kinetic Cu(II) uptake data displayed in Fig. 7, were applied to Eq. (10) and plots were obtained between t/qt and t which were almost linear with fairly high regressions (see Fig. 7). The Elovich equation is used successfully to describe second-order kinetics, assuming that the actual solid surface is energetically heterogeneous (Rudzinski and Panezyk, 2002). The Elovich equation is given as
where ά and β are Elovich constants. Figure 8 shows the qt versus ln t plots obtained using kinetic Cu(II) uptake data for Cu(II) solutions with initial concentrations of 10 and 20 mg·ℓ-1 at 27ºC . The plots obtained were quite linear with relatively lower regressions relative to the pseudo second-order model. Finally various kinetic parameters employed in the above models were obtained by using slopes and intercepts of linear plots depicted in Figs. 8 and 9; these are given in Table 4. A close look at the data displayed reveals that experimental qe values (i.e. Cu(II) sorbed in mg·g-1) are in close agreement with the theoretical values, thus establishing the fair fitness of the pseudo second-order kinetic model to uptake data. The Elovich constants ά and β have also been depicted. In addition, regression values o>btained also indicate that the pseudo second-order model indicates better suitability.
The uptake mechanism of a sorbate onto the adsorbent involves film diffusion, pore diffusion and intra-particle transport (Nadeem et al., 2006), of which the slowest is rate limiting and controls the overall sorption process. Generally, the pore diffusion and the intra-particle diffusion are often rate-limiting in a batch reactor (Goswami and Ghosh, 2005). In order to confirm the occurrence of the intra-particle diffusion phenomenon, we followed the observations of Raghuvanshi et al. (2004) who suggested a linear plot between log (time) and log (per cent adsorption) in the case of intra-particle diffusion. The plots of log (time) versus log (per cent adsorption), as shown in Fig. 10, clearly confirm that intra-particle diffusion does occur in the present study. An empirically functional relationship, common to most adsorption processes, is that the sorbate uptake varies almost proportionally with t1/2, the Weber Morris plot, rather than with the contact time, t (Weber and Morris, 1963):
where kid is the intraparticle diffusion rate constant. According to Eq. (13), a plot of qt versus t1/2 should be a straight line with a slope kid and intercept C when the adsorption mechanism follows the intra-particle diffusion process. The value of the intercept gives an idea about the thickness of the boundary layer, i.e. the larger the intercept the greater will be the boundary layer effect (Unnithan et al., 2002). In Fig. 11, plots of mass of Cu(II) ions sorbed per unit mass of STL sorbent, qt versus t½, are presented. It is evident that the plots are not only linear but also pass through the origin, thus indicating that the Cu(II) uptake by STL sorbent is almost intra-particle diffusion controlled. The values of kid, obtained using slopes of the linear plots for 10 and 20 mg·ℓ-1Cu(II) solutions were found to be 841.2 x 10-2 and 1637.0 x 10-2 mg·g-1·min1/2, respectively. Higher values of kid for higher concentrations of Cu(II) solutions may simply be due to the fact that higher sorbate concentration adds to a greater driving force for diffusion of sorbate molecules into the pores.
Evaluation of thermodynamic parameters
The thermodynamic parameters that must be considered to characterise the uptake process are the enthalpy of adsorption (ΔHº), the standard Gibbs free energy (ΔGº), and the entropy (ΔSº) changes due to transfer of 1 mole of solute from solution to solid-liquid interface. The standard free energy change (ΔGº) was evaluated by the relation:
where R is the ideal gas constant (8.314 J mol-1·K-1) and T is the absolute temperature (K). The apparent equilibrium constant (Kc) of the adsorption process is defined as:
where Cad,eq and Ceq are the concentrations of Cu(II) at the adsorbent (mg·ℓ-1) and equilibrium concentration (mg·ℓ-1) respectively. The standard enthalpy ΔHº is related to the Langmuir constant b by the well-known Vant Hoff equation
where b1 and b2 are Langmuir constants at temperature T1 and T2, respectively. The above relation was used to calculate ΔHº. Finally the standard enthalpy change ΔSº was calculated using the well-known relationship:
The various thermodynamic parameters, calculated using the above-mentioned relationships, are given in the Table 5. The negative standard free energy change ΔGº is indicative of the spontaneous nature of the Cu(II) sorption process. In addition, negative value of ΔHºindicate the exothermic nature of the sorption process. The negative value of ΔSº indicates greater order during the sorption process. This may probably be explained on the basis of the fact that during the sorption process the co-ordinated water molecules which are displayed by copper ions gain less translational entropy than is lost by Cu(II) species, resulting in decreased randomness. Similar negative entropy change has also been reported elsewhere (Unnithan et al., 2002).
From the above study it may be concluded that spent tea leaves (STL) bear great potential to be used as a cost-effective adsorbent material for effective removal of Cu(II)ions from aqueous solutions. The sorbent does not require any chemical modification and hence may be considered as an environment-friendly sorbent.
The authors are thankful to Dr. O P Sharma for providing necessary laboratory facilities.
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Received 20 November 2009; accepted in revised form 23 February 2010.