Biosorption of Manganese(II) Ions from Aqueous Solution ... - CiteSeerX

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Biosorption of Manganese(II) Ions from Aqueous Solution by Glutaraldehyde Cross-linked Chitosan Beads: Equilibrium and Kinetic Studies Madala Suguna, Alla Subba Reddy, Nadavala Siva Kumar and Abburi Krishnaiah* Biopolymers and Thermophysical Laboratory, Department of Chemistry, Sri Venkateswara University, Tirupati–517502, Andhra Pradesh, India. (Received 20 November 2009; revised form accepted 4 July 2010)

ABSTRACT: The biosorption of manganese(II) ions from aqueous solution by glutaraldehyde cross-linked chitosan (GCC) was studied under equilibrium conditions. The biosorbent was characterised by FT-IR spectroscopy and scanning electron microscopy (SEM) methods. The effects of the variable experimental parameters such as pH, metal ion concentration, adsorbent amount, contact time and temperature on the adsorption process were investigated. The equilibrium adsorption data were interpreted via the Freundlich, Langmuir and Sips isotherm models. Based on the error function values, the kinetic data were fitted to a better extent by the pseudo-second-order kinetic model and the chemisorption model, relative to the pseudo-first-order, fractional order and Weber–Morris models. The monolayer adsorption capacity of GCC as obtained from the Langmuir isotherm at 25 oC was found to be 278 mg/g. Thermodynamic studies indicated that the adsorption process was spontaneous (∆G0 < 0) and exothermic (∆H0 < 0).

INTRODUCTION Wastewater and effluents contaminated with heavy metal ions cause serious environmental problems. Even at low concentrations, heavy metal ions are highly toxic and not biodegradable. They must be removed from polluted streams in accordance with stringent environmental legislation and enforcement. In addition, groundwater often contains iron and manganese ions which not only affect the flavour and colour of water but also accumulate to cause obstruction in water equipment and pipelines (Gotoh et al. 2004). Several methods exist for the removal of toxic metal ions from aqueous solution, such as ion-exchange, reverse osmosis, adsorption, complexation and precipitation (Crist et al. 1996; Huang and Blankenship 1984; Juang and Shao 2002). Adsorption is considered to be an effective and economical method for the removal of pollutants from wastewater (Ng et al. 2002). Manganese is an essential metal for the human system and many enzymes are activated by this element. Manganese has a variety of applications in ceramics, dry battery cells, electrical coils and many alloys (Sharma et al. 2007). In addition to the disposal of untreated discharge from the above applications into water bodies, another major source of manganese pollution is the burning of coal and oil (Kannann 1995; Forstner and Wittmann 1979). The intake of higher concentrations of manganese causes manganese psychosis, an irreversible neurological disorder. It is characterised by uncontrollable laughter, sexual excitement and impotence (Kannann 1995). *

Author to whom all correspondence should be addressed. E-mail: [email protected].

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M. Suguna et al./Adsorption Science & Technology Vol. 28 No. 3 2010

Heavy metal ion biosorption by biological materials such as bacteria and fungi presents few problems when operated in a continuous mode; however, solid/liquid separation is a major constraint in such methods. Even though immobilisation may solve this problem, chemical costs and mechanical strength should be taken into consideration. For these reasons, recent research has focused on the use of low-cost and waste materials as adsorbents (Vijayaraghavan et al. 2005). The adsorption capacity of several low-cost adsorbents has been investigated, mainly using biopolymers which are obtained from renewable sources and adsorb metallic ions selectively (Hsien and Rorrer 1995). Of such biopolymers, chitosan is a linear polysaccharide based on a glucosamine unit. It is obtained from the de-acetylation of chitin, which is the major component of crustacean shells. As a consequence, it is one of the most available biopolymers in Nature. The potential use of chitosan as an adsorbent has been demonstrated, particularly for the removal of transition and post-transition elements from aqueous waste (Zamin et al. 2004), with the use of chitosan as a chelating ion-exchanger for the removal of metal ions being well established in the literature. Although its amino and hydroxy groups can act as chelating sites (Merrifield et al. 2004; Qi and Xu 2004; Dambies et al. 2000), these active binding sites are not readily available for sorption when it exists as a gel or in its natural form. Several investigators have attempted to modify chitosan to facilitate mass transfer and to expose the active binding sites in order to enhance its adsorption capacity (Hasan et al. 2006; Boddu et al. 2003). Hasan et al. (2003) and Guibal et al. (1998) noted that the maximum uptake of chitosan flakes towards molybdate ions was approximately one-half that obtained with chitosan beads. Some researchers have recognised that this biosorbent requires further modification and development before it can be used commercially. Thus, Boddu et al. (2008) and Gupta et al. (2009) evaluated the sorption of As(III) and As(V) ions by a chitosan-coated biosorbent and an iron–chitosan composite, respectively. Similarly, Tao et al. (2009) studied the removal of Pb(II) ions by a chitosan/TiO2 composite. However, for such materials, it is also necessary to provide a physical support and to increase the accessibility of the metal ion-binding sites for process applications. To overcome some of the problems associated with chitosan, several workers have reported the use of cross-linking treatments using several chemical reagents such as glutaraldehyde (Wan Ngah et al. 2005; Jeon and Wolfgang 2003), carboxymethyl-chitosan (Sun et al. 2006), epichlorohydrin (Wan Ngah et al. 2005; Vieira and Beppu 2006), etc. Hence, an attempt has been made in the present investigation to prepare cross-linked chitosan for the removal of Mn(II) ions, thereby preventing the dissolution of chitosan in acidic solutions and to improve its metal ion adsorption properties, viz. to increase its adsorption capacity or to enhance its selectivity. The objectives of the present study were to prepare glutaraldehyde cross-linked chitosan to remove Mn(II) ions from aqueous solutions. The effects of contact time, solution pH, initial metal ion concentration, temperature and amount of biomass employed on the extent of adsorption were also studied. Fitting of the kinetic data was attempted employing the pseudo-first-order, pseudosecond-order, chemisorption and fractional order kinetic models and the rate constants calculated. The adsorption capacity of the sorbent was evaluated by studying the equilibrium adsorption isotherms of Mn(II) ions in batch mode. The equilibrium data were fitted by the Langmuir, Freundlich and Sips isotherm models. In addition, the biosorbent was characterised by scanning electron microscopy (SEM) and FT-IR analysis to examine metal ion accumulation due to the presence of different functional groups present in the biosorbent.

Biosorption of Mn(II) Ions by Glutaraldehyde Cross-linked Chitosan Beads

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MATERIALS AND METHODS Materials Chitosan with a molecular weight of 9.9 × 105 g/g mol was purchased from Sigma-Aldrich, St. Louis, MO, U.S.A. All other chemicals were of analytical grade. Aqueous solutions were prepared using doubly distilled water. Preparation of chitosan beads A solution of chitosan was prepared by dissolving 2.00 g of chitosan flakes in 60 m of 5% v/v acetic acid solution. The chitosan solution was sprayed into a precipitation bath containing 500 m of 0.5 M NaOH which neutralised the acetic acid within the chitosan gel and thereby allowed the gel to coagulate into uniform spherical gel beads. A magnetic stirrer was used to stir the aqueous NaOH solution. The wet chitosan gel beads were extensively rinsed with distilled water to remove any adhering NaOH, filtered and finally air-dried to remove the water from the pore structure. Preparation of glutaraldehyde cross-linked chitosan beads A known amount of newly prepared wet chitosan beads was suspended in a 0.025 M glutaraldehyde solution to obtain a 1:1 chitosan/glutaraldehyde ratio. The chitosan beads were allowed to stand in the resulting solution for 24 h at room temperature after which the cross-linked chitosan (GCC) beads were intensively washed with distilled water, filtered and air-dried. The newly formed beads were ground and sieved to a constant size before use. Preparation of manganese(II) ion solution Test Mn(II) ion solutions were prepared using A.R. MnSO4• H2O (Merck Ltd., Darmstadt, Germany). Stock adsorbate solutions of 1000 mg/ Mn(II) ion concentration were prepared by dissolving 3.077 g of MnSO4•H2O in doubly distilled water. A range of Mn(II) ion solutions with concentrations within the range 100–400 mg/ was prepared by appropriate dilution of the stock solution. Doubly distilled water was used for preparing the stock solutions employed throughout subsequent experimental analyses. Batch studies Batch adsorption studies were carried out by adding 100 mg of GCC beads to 100 m of an Mn(II) ion solution contained in a series of 250 m Erlenmeyer flasks. The initial pH value in each flask was adjusted by the addition of either 1 M HCl or 1 M NaOH. Equilibration was achieved by shaking the flasks for 4 h at an agitation speed of 200 rpm on a Lab line rotary shaker. Equilibrium isotherm measurements were carried out by maintaining both the amount of GCC employed and the solution volume (100 m) constant and varying the initial concentration of Mn(II) ions employed. For kinetic studies, samples were withdrawn at periodic time intervals and filtered via Whatman filter papers. The concentration of metal ions in the filtrate was analysed spectrophotometrically employing an Elico-SL 177 UV–vis spectrophotometer at 545 nm. The

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effects of contact time, metal ion concentration, adsorbent dosage, pH and temperature were studied. The amount adsorbed per unit mass of adsorbent (Qe) was obtained from the equation:  C − Ce  Qe =  i  V  m

(1)

where Qe (mg/g) is the adsorption capacity at equilibrium, Ci and Ce are the initial and equilibrium concentrations of metal ion (mg/), respectively, V is volume of the solution () and m is the mass of adsorbent (g) employed. The effect of pH on the adsorption process was studied by carrying out the experiment at different pH values, keeping the concentration, volume of adsorbate solution and amount of adsorbent constant. The effect of adsorbent dosage on the adsorption of Mn(II) ions was studied by agitating 100 m of a 100 mg/ Mn(II) ion solution with different amounts of adsorbent. Statistical evaluation of the kinetic parameters The Marquardt percentage standard deviation (MPSD) error function (Marquardt 1963) was employed in this study to determine a suitable kinetic model for representing the experimental data: Ferror (%) = 100 ×

 q i model − q i exp.  ∑i   q i exp.  p

2

 1   p − 1

•

(2)

where qi model is the value of q predicted by the fitted model, qi exp. is the experimental value of q and p is the number of experiments performed. The MPSD error function has been used previously by a number of researchers in this field (Vaghetti et al. 2009). Adsorption kinetics Pseudo-first-order kinetics It should be emphasised that metal ion–sorbent interactions can be different for different biosorbents. This means that an adsorption process on a particular biosorbent may be kineticallycontrolled, diffusion-controlled or even a combination of the two (Cheung et al. 2003). With this point in mind, attempts were first made to fit the kinetic adsorption data for Mn(II) ions onto GCC with the pseudo-first-order rate expression as given by:  k  log(Q e − Q t ) = log Q e −  1  t  2.303 

(3)

where Qe and Qt are the amounts of metal ions adsorbed per unit mass of adsorbent at equilibrium and at time t (min), respectively, and k1 is the pseudo-first-order rate constant for the adsorption process. Linear plots of log(Qe – Qt) versus t at different initial metal ion concentrations indicate the applicability of this equation (Shekinaah et al. 2002). Fractional order kinetics Although the pseudo-first- and pseudo-second-order kinetic models have been used for fitting purposes in most adsorption kinetic studies for the determination of the kinetic parameters,


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information is still lacking in the literature concerning possible changes in the adsorption rates as a function of initial metal ion concentration and adsorption time, as well as regarding the determination of fractional kinetic orders. For this reason, the alternative fractional order kinetic equation proposed by Avrami (Lopes et al. 2003; Cestari et al. 2004) has been used to analyze the data obtained in the present study. This equation may be expressed as: α = 1 − exp [ −( k AV t )]

n

(4)

where α is the fraction adsorbed (qt /qe) at time t, kAV is the Avrami kinetic constant (min–1) and n is a fractional reaction order. Pseudo-second-order kinetics The experimental data were also analysed employing the pseudo-second-order kinetic model which may be expressed in the following form: t Qt

=

1 k 2 Q e2

+

1 t Qe

(5)

where k2 [g/(mg min)] is the rate constant in the pseudo-second-order equation, Qt (mg/g) is the amount adsorbed at time t (min) and Qe (mg/g) is the amount adsorbed at equilibrium. Chemisorption The Elovich equation has also found general application in chemisorption kinetics (Parez-Marin et al. 2007). This equation has been applied satisfactorily to some chemisorption processes and has been found useful in covering a wide range of slow adsorption rates. The equation, often valid for systems in which the adsorbing surface is heterogeneous, can be written as:  1  1 Q t =   ln(αβ) +   ln( t)  β  β

(6)

where α is the adsorption rate [mg/(g min)] and β (g/mg) is related to the extent of surface coverage and the activation energy involved in the chemisorption process. Weber–Morris method The intra-particle diffusion model proposed by Weber and Morris can be expressed by the equation: Qt = kidt1/2 + C

(7)

where Qt (mg/) is the amount adsorbed at time t (min), kid [mg/(g min)] is the rate constant for intra-particle diffusion and C is the value of the intercept of the related linear plot.

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Adsorption isotherms Langmuir isotherms The Langmuir adsorption model has often been used to describe equilibrium adsorption isotherms. This model is valid for monolayer sorption onto a surface with a finite number of identical sites and may be written as: Qe =

Q 0 bCe 1 + bCe

(8)

where Q0 is the maximum amount of metal ions adsorbed per unit weight of GCC to form a complete monolayer on the surface at equilibrium concentration, while b is related to the affinity of the binding site. The Langmuir parameters were calculated in the same way as employed in the literature by other workers (Nuhoglu and Malkoc 2009; Khambhaty et al. 2009; Vaghetti et al. 2009). Freundlich isotherms The Freundlich isotherm model is expressed as an exponential equation and assumes that the concentration of the adsorbate on the adsorbent surface increases as the initial adsorbate concentration increases. In theory, this expression suggests the possibility of an infinite amount of adsorption. The corresponding relationship may be expressed as: Q e = K f C1/e n

(9)

where Kf and n are the Freundlich constants which indicate the adsorption capacity and adsorption intensity, respectively. Sips isotherm This model is a combination of the Langmuir and Freundlich isotherm models. The Sips model (1948) takes the form: qe =

Q max K s C1e/n 1 + K s C1e/n

(10)

where Ks is the Sips constant related to the affinity constant (mg/)–1/n, Qmax is the Sips maximum adsorption capacity (mg/g) and ns is the Sips exponent (dimensionless). At low sorbate concentrations, this equation effectively reduces to the Freundlich isotherm, while at high sorbate concentrations it predicts a monolayer adsorption capacity characteristic of the Langmuir isotherm. Thermodynamic parameters On the basis of fundamental thermodynamics concepts, it may be assumed that energy cannot be gained or lost in an isolated system and that the entropy change is the sole driving force. In environmental engineering practice, both energy and entropy factors must be considered in order


219

to determine which process will occur spontaneously. In the present study, an increase in temperature led to a decrease in the Mn(II) ion biosorption rate. Thermodynamic parameters such as the standard enthalpy change (∆H0), the standard entropy change (∆S0) and the standard Gibbs’ free energy change (∆G0) for the sorption process may be calculated using the following wellknown relationships (Sharma et al. 2007; Catena and Bright 1989): ∆G0 = –RT ln KL

(11)

∆G0 = ∆H0 – T∆S0

(12)

∆S0 R

(13)

ln K L =

−

∆H0 RT

where R is the gas constant and KL is the Langmuir constant. RESULTS AND DISCUSSION Characterisation of the biosorbent The FT-IR spectra of GCC in the virgin form and after loading with Mn(II) ions are shown in Figure 1. The FT-IR spectrum of GCC shows the presence of predominant peaks at 3422.0 cm–1 (–OH and –NH2 stretching vibrations), 2926.5 cm–1 (–CH stretching vibration in –CH and –CH2), 1641.7 cm–1 (–NH bending vibration in NH2) and 1071.4 cm–1 (–CO stretching vibration in –COH). The intensity of the peaks in the transmittance spectra was relatively greater for the metal ion-loaded GCC beads relative to the unloaded GCC beads. This higher intensity may be attributed to the presence of a smaller number of functional groups in the loaded GCC beads. This observation accords with the view that functional groups such as –NH2, –OH and –CO– are involved in binding the metal ions to the GCC beads (Boddu et al. 2008). The scanning electron micrograph of glutaraldehyde cross-linked chitosan beads (GCC) depicted in Figure 2 shows no particular crystalline structure, with the beads appearing to be approximately spherical in shape. In addition, the texture of the beads appears to consist of holes and small openings on the surface. These aspects result in an increase in contact area, thereby facilitating pore diffusion during the adsorption process. Effect of pH In order to optimise the pH for maximum removal efficiency, experiments were conducted at room temperature employing 100 m of an Mn(II) ion solution of 100 ppm initial concentration containing 0.1 g of glutaraldehyde cross-linked chitosan over the pH range 3–8. This pH range was chosen so as to avoid chitosan dissolution, which occurs within the pH range 3–8 with consequent precipitation of Mn(II) ions. From Figure 3, it will be seen that the uptake capacity towards Mn(II) ions increased with increasing solution pH, with the adsorption capacity being a maximum at pH 6. The pH at the point of zero charge (pHpzc) of aminated chitosan beads is 5.53 (Jeon and Holl 2003). The pH value primarily affects the degree of ionisation of the metal ion and

220


b

98 97

a

96 95 94 93

% Transmittance

92 91 90 89 88 87 86 85 84 83 82 81 80 4000

3500

3000

2500

2000

1500

1000

500

Wavenumber (cm–1 ) Figure 1. FT-IR spectra of (a) glutaraldehyde cross-linked chitosan beads (GCC) and (b) GCC loaded with Mn(II) ions.


221

Figure 2. Scanning electron micrograph of glutaraldehyde cross-linked chitosan beads.

90 Adsorption capacity (mg/g)

80 70 60 50 40 30 20 10 0 0

2

4

6

8

10

pH Figure 3. Effect of pH on the adsorption capacity of GCC towards Mn(II) ions from aqueous solution.

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the surface properties of the chitosan. This may be due to the formation of soluble hydroxy complexes (Nacer Kocer et al. 2008). In acidic solutions, strong competition exists between the metal ions and protons for the available sorption sites on the adsorbent and this leads to a decrease in the uptake capacity of Mn(II) ions. Ionisable functional groups, such as –NH2 or –OH, on the surface of cross-linked chitosan may gain or lose protons, resulting in a surface charge that varies with pH. At low pH, the surface sites are protonated and the surface becomes positively charged; in contrast, at high pH, the ionisable groups lose their protons and the surface becomes negatively changed. This means that, at pH values lower than pHpzc, the surface of the chitosan membrane exhibits a positive charge. Since there are fewer accessible binding groups remaining, the uptake capacity of GCC decreases. The maximum removal of Mn(II) ions by adsorption at pH 6.0 can be understood on the basis of Mn(II) ions being present in the solution as Mn(H2O)62+. At pH > 8, Mn(OH)2 is formed as observed experimentally. This is further supported by earlier reports (Tiwary et al. 1993). Effect of adsorbent dosage Increasing the amount of adsorbent employed in the experiments led to an increase in the extent of Mn(II) ion adsorption. For a given initial Mn(II) ion concentration, the removal efficiency increased up to an optimum adsorbent dosage, beyond which any further increase in removal efficiency became negligible. This is to be expected because, for a fixed initial solute concentration, increasing the adsorbent dosage led to an increase in the surface area and hence a greater number of adsorption sites. The results depicted in Figure 4 indicate that GCC removed 97% of the Mn(II) ions initially present in solution.

% Removal of Mn(II) ions

120 100 80 60 40 20 0 0

0.1

0.2 0.3 0.4 Adsorbent dosage (g/)

0.5

0.6

Figure 4. Effect of adsorbent dosage on the percentage removal of Mn(II) ions from aqueous solution by GCC.

Adsorption kinetics data Application of the pseudo-first-order, pseudo-second-order and fractional order kinetic models The applicability of pseudo-first-order kinetic treatment of the data for the adsorption process was examined by plotting log(Qe – Qt) versus time (min) to evaluate the rate constant k1, while the


223

possibility of fractional order was examined by plotting ln[–ln(1 – Qe)/Qt] versus time (min) to evaluate kAV. The latter plots are depicted in Figure 5. It will be seen that the fits were approximately linear. The corresponding correlation coefficients, rate constants k1 and kAV, and MPSD values were calculated and are summarised in Table 1. The high error function values obtained for the application of the pseudo-first-order and fractional order models demonstrate that these models were not applicable in the present study. 1.4 1.2

ln[ ln(1 Qt)/Qe]

1.0 0.8 0.6 0.4 0.2 0 3.0

3.5

4.0

4.5

5.0

ln t Figure 5. Fractional order kinetic plots for the adsorption of Mn(II) ions onto GCC. Data points correspond to the following initial Mn(II) concentrations: , 100 mg/; , 200 mg/; , 300 mg/; , 400 mg/.

The pseudo-second-order rate constant, k2, calculated from the slopes of the plots of t/Qt versus t, are shown in Figure 6, with the corresponding values of the correlation coefficient and equilibrium adsorption capacity, Qe, also being listed in Table 1. From the data in Table 1, it may be concluded that the experimental data were well fitted by the pseudo-second-order kinetic model. This is supported by the high R2 and low error function values, with the values of q as calculated by this model being in good agreement with the experimental data. The error function evaluates the differences associated with each individual points fitted by the model in relation to each experimental point measured. This observation supports the contention that the adsorption of Mn(II) ions onto GCC followed pseudo-second-order kinetics. Application of the intra-particle diffusion model Values of the intra-particle rate constant kid [mg/(g min)] and intercept C (mg/g) are given in Table 1. The intercept provides information about the boundary layer thickness, i.e. the larger the intercept the greater the boundary layer effect. Using this model, linear plots of Qe versus t1/2 passing through the origin indicate that intra-particle diffusion alone determines the overall rate of adsorption. However, for the data obtained in the present study, the linear plots did not pass through the origin thereby indicating that intra-particle diffusion was not the sole rate-determining factor and that other factors appeared to be operating simultaneously. This indicates that the

224


TABLE 1. Parameter Values of Kinetic Models Applied to the Adsorption of Mn(II) Ions onto GCC

Kinetic model

Parameter

Initial concentration of Mn(II) ions (mg/) 100

200

300

400

Pseudo-first-order

k1 (min–1) R2 MPSD (%)

0.012 0.984 86

0.0081 0.994 94

0.0087 0.998 97.3

0.0092 0.993 106.4

Pseudo-second-order

k2 [g/(mg min)] R2 Qe MPSD (%)

0.001 0.998 75.2 4.1

0.0004 0.998 169.5 4.5

0.0004 0.998 270.2 4.3

0.0004 0.999 344.8 3.2

Chemisorption

α [mg/(g min)] β (g/mg) R2 MPSD (%)

595.8 0.136 0.962 1.9

2591.5 0.06 0.948 2.5

6.12 × 104 0.05 0.956 2.1

6.4 × 105 0.05 0.967 1

Fractional order

kAV (min–1) Qe N R2 MPSD (%)

0.001 13.5 0.33 0.9855 100

0.001 26.7 0.4 0.9766 102.3

0.001 32.2 0.3 0.9913 99.2

0.001 36.7 0.3 0.9477 103.6

Weber–Morris

kid [mg/(g min)] R2 C

1.59 0.953 48.8

3.48 0.923 113

4.29 0.916 202.8

4.69 0.957 277.2

3.5 3.0

t/Qt

2.5 2.0 1.5 1.0 0.5 0 0

50

100

150

200

250

Time (min) Figure 6. Pseudo-second-order kinetic plots for the adsorption of Mn(II) ions onto GCC. Data points correspond to the following initial Mn(II) concentrations: , 100 mg/; , 200 mg/; , 300 mg/; , 400 mg/.


225

mechanism of Mn(II) ion adsorption by GCC is complex and both surface adsorption as well as intra-particle diffusion contribute to the rate-determining step. Application of the Elovitch chemisorption model To examine the application of this model, plots of ln t versus Qt were constructed for the data arising from each of the initial Mn(II) ion concentrations employed. These are presented in Figure 7, while the corresponding values of α, β and the error function are listed in Table 1. It will be seen that the plots were linear with good correlation coefficients. The equilibrium concentrations calculated from this model were closely related to the values obtained experimentally. This suggests that, although the sorption system studied exhibited pseudo-secondorder kinetics, the rate-determining step may involve chemisorption which occurs through the sharing or exchange of electrons between the adsorbent and the adsorbate.

Adsorption capacity (mg/g)

400 350 300 250 200 150 100 50 0 3.0

3.5

4.0

4.5

5.0

5.5

ln t Figure 7. Elovich chemisorption model plots for the adsorption of Mn(II) ions onto GCC. Data points correspond to the following initial Mn(II) concentrations: , 100 mg/; , 200 mg/; , 300 mg/; , 400 mg/.

Equilibrium isotherm data The Langmuir, Freundlich and Sips isotherm models were used to fit the equilibrium adsorption data obtained at 25 oC, 30 oC and 35 oC, respectively. Of these models, the Langmuir isotherm model showed a good fit of the experimental sorption data with R2 values close to unity, allowing the values of Q0 and b to be determined from the linear plots of 1/Ce versus 1/Qe. The linearised Langmuir adsorption isotherms for Mn(II) ions onto GCC at 25 oC, 30 oC and 35 oC are depicted in Figure 8. It will be seen from the figure that the value of Q0 decreased with increasing temperature, with the corresponding values of 278, 99 and 39 mg/g being obtained at 25 oC, 30 oC and 35 oC, respectively. This result confirms that the adsorption of Mn(II) ions onto GCC beads was exothermic in nature.

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0.40 0.35 0.30

1/Qe

0.25 0.20 0.15 0.10 0.05 0 0

0.05

0.10

0.15

0.20

0.25

0.30

0.35

1/Ce Figure 8. Langmuir isotherms for the adsorption of Mn(II) ions onto GCC at () 25 oC, () 30 oC and () 35 oC, respectively.

Thermodynamics The values of ∆G0 for the sorption of Mn(II) ions onto GCC at different temperatures are listed in Table 2, from which it will be seen that the magnitude of ∆G0 decreased with increasing temperature. The negative values of ∆G0 indicate the feasibility and spontaneity of the adsorption process. The smaller values of ∆G0 at higher temperatures confirm the exothermic nature of the removal process. The values of ∆H0 and ∆S0 were determined from the slope and intercept of the plot of ln KL versus 1/T, with ∆H0 for the sorption of Mn(II) onto GCC being found to be –107.15 kJ/mol while ∆S0 was –0.3193 kJ/(mol K). The negative value of the standard enthalpy change also provides evidence that the adsorption process was exothermic in nature, while the negative value of the standard entropy change indicated that adsorption was favoured in the system (Sharma et al. 2007). TABLE 2. Thermodynamic Parameters for the Adsorption of Mn(II) Ions onto GCC

Temp. (oC)

∆G0 (kJ/mol)

∆H0 (kJ/mol)

∆S0 [J/(mol K)]

25 30 35

–12.189 –9.8 –9.7

–107.2

–0.3193

CONCLUSIONS The results obtained in the present study show that glutaraldehyde cross-linked chitosan may be considered as a potential biosorbent material for the removal of Mn(II) ions from aqueous media. The maximum uptake of Mn(II) ions occurred at pH 6, with increasing amounts of biosorbent


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leading to an increase in the percentage removal of ions. The biosorbent was characterised by FT-IR spectroscopy and scanning electron microscopy (SEM) methods, the equilibrium adsorption data were interpreted via the Langmuir, Freundlich and Sips models while the kinetic adsorption data were studied using the pseudo-first-order, pseudo-second-order and Elovitch isotherm models. Of these models, that of Langmuir provided the best fit of the equilibrium adsorption data while the kinetic data were best fitted by the pseudo-second-order kinetic and Elovitch chemisorption models. The maximum monolayer adsorption capacity of GCC was 278 mg/g at 25 oC. Thermodynamic investigations indicated that the adsorption process was spontaneous and exothermic.

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