Biosorption of Manganese(II) Ions from Aqueous Solution by Glutaraldehyde Cross-linked Chitosan...

213

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

Biosorption of Manganese(II) Ions from Aqueous Solution by GlutaraldehydeCross-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 solutionby glutaraldehyde cross-linked chitosan (GCC) was studied under equilibriumconditions. The biosorbent was characterised by FT-IR spectroscopy andscanning electron microscopy (SEM) methods. The effects of the variableexperimental parameters such as pH, metal ion concentration, adsorbent amount,contact time and temperature on the adsorption process were investigated. Theequilibrium adsorption data were interpreted via the Freundlich, Langmuir andSips isotherm models. Based on the error function values, the kinetic data werefitted to a better extent by the pseudo-second-order kinetic model and thechemisorption model, relative to the pseudo-first-order, fractional order andWeber–Morris models. The monolayer adsorption capacity of GCC as obtainedfrom the Langmuir isotherm at 25 oC was found to be 278 mg/g. Thermodynamicstudies indicated that the adsorption process was spontaneous (∆G0 < 0) andexothermic (∆H0 < 0).

INTRODUCTION

Wastewater and effluents contaminated with heavy metal ions cause serious environmentalproblems. Even at low concentrations, heavy metal ions are highly toxic and not biodegradable.They must be removed from polluted streams in accordance with stringent environmentallegislation and enforcement. In addition, groundwater often contains iron and manganese ionswhich not only affect the flavour and colour of water but also accumulate to cause obstruction inwater equipment and pipelines (Gotoh et al. 2004). Several methods exist for the removal of toxicmetal ions from aqueous solution, such as ion-exchange, reverse osmosis, adsorption,complexation and precipitation (Crist et al. 1996; Huang and Blankenship 1984; Juang and Shao2002). Adsorption is considered to be an effective and economical method for the removal ofpollutants from wastewater (Ng et al. 2002).

Manganese is an essential metal for the human system and many enzymes are activated by thiselement. Manganese has a variety of applications in ceramics, dry battery cells, electrical coils andmany alloys (Sharma et al. 2007). In addition to the disposal of untreated discharge from theabove applications into water bodies, another major source of manganese pollution is the burningof coal and oil (Kannann 1995; Forstner and Wittmann 1979). The intake of higher concentrationsof manganese causes manganese psychosis, an irreversible neurological disorder. It ischaracterised by uncontrollable laughter, sexual excitement and impotence (Kannann 1995).

Heavy metal ion biosorption by biological materials such as bacteria and fungi presentsfew problems when operated in a continuous mode; however, solid/liquid separation is amajor 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 beeninvestigated, mainly using biopolymers which are obtained from renewable sources andadsorb metallic ions selectively (Hsien and Rorrer 1995). Of such biopolymers, chitosan is alinear 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 mostavailable biopolymers in Nature. The potential use of chitosan as an adsorbent has beendemonstrated, particularly for the removal of transition and post-transition elements from aqueouswaste (Zamin et al. 2004), with the use of chitosan as a chelating ion-exchanger for theremoval of metal ions being well established in the literature. Although its amino and hydroxygroups 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 itsnatural form.

Several investigators have attempted to modify chitosan to facilitate mass transfer and to exposethe 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 chitosanflakes towards molybdate ions was approximately one-half that obtained with chitosan beads.Some researchers have recognised that this biosorbent requires further modification anddevelopment 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 aniron–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 aphysical support and to increase the accessibility of the metal ion-binding sites for processapplications.

To overcome some of the problems associated with chitosan, several workers have reportedthe 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 hasbeen made in the present investigation to prepare cross-linked chitosan for the removal ofMn(II) ions, thereby preventing the dissolution of chitosan in acidic solutions and to improveits metal ion adsorption properties, viz. to increase its adsorption capacity or to enhance itsselectivity.

The objectives of the present study were to prepare glutaraldehyde cross-linked chitosan toremove Mn(II) ions from aqueous solutions. The effects of contact time, solution pH, initial metalion concentration, temperature and amount of biomass employed on the extent of adsorption werealso studied. Fitting of the kinetic data was attempted employing the pseudo-first-order, pseudo-second-order, chemisorption and fractional order kinetic models and the rate constants calculated.The adsorption capacity of the sorbent was evaluated by studying the equilibrium adsorptionisotherms 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 scanningelectron microscopy (SEM) and FT-IR analysis to examine metal ion accumulation due to thepresence of different functional groups present in the biosorbent.

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

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 wereprepared 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/vacetic acid solution. The chitosan solution was sprayed into a precipitation bath containing500 m� of 0.5 M NaOH which neutralised the acetic acid within the chitosan gel and therebyallowed the gel to coagulate into uniform spherical gel beads. A magnetic stirrer was used to stirthe aqueous NaOH solution. The wet chitosan gel beads were extensively rinsed with distilledwater to remove any adhering NaOH, filtered and finally air-dried to remove the water from thepore structure.

Preparation of glutaraldehyde cross-linked chitosan beads

A known amount of newly prepared wet chitosan beads was suspended in a 0.025 Mglutaraldehyde solution to obtain a 1:1 chitosan/glutaraldehyde ratio. The chitosan beads wereallowed to stand in the resulting solution for 24 h at room temperature after which the cross-linkedchitosan (GCC) beads were intensively washed with distilled water, filtered and air-dried. Thenewly 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 bydissolving 3.077 g of MnSO4•H2O in doubly distilled water. A range of Mn(II) ion solutions withconcentrations within the range 100–400 mg/� was prepared by appropriate dilution of the stocksolution. Doubly distilled water was used for preparing the stock solutions employed throughoutsubsequent 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 flaskwas adjusted by the addition of either 1 M HCl or 1 M NaOH. Equilibration was achieved byshaking the flasks for 4 h at an agitation speed of 200 rpm on a Lab line rotary shaker. Equilibriumisotherm measurements were carried out by maintaining both the amount of GCC employed andthe solution volume (100 m�) constant and varying the initial concentration of Mn(II) ionsemployed. For kinetic studies, samples were withdrawn at periodic time intervals and filtered viaWhatman filter papers. The concentration of metal ions in the filtrate was analysedspectrophotometrically employing an Elico-SL 177 UV–vis spectrophotometer at 545 nm. The

Biosorption of Mn(II) Ions by Glutaraldehyde Cross-linked Chitosan Beads 215

effects of contact time, metal ion concentration, adsorbent dosage, pH and temperature werestudied. The amount adsorbed per unit mass of adsorbent (Qe) was obtained from the equation:

(1)

where Qe (mg/g) is the adsorption capacity at equilibrium, Ci and Ce are the initial and equilibriumconcentrations of metal ion (mg/�), respectively, V is volume of the solution (�) and m is the massof adsorbent (g) employed. The effect of pH on the adsorption process was studied by carrying outthe experiment at different pH values, keeping the concentration, volume of adsorbate solution andamount of adsorbent constant. The effect of adsorbent dosage on the adsorption of Mn(II) ions wasstudied 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) wasemployed in this study to determine a suitable kinetic model for representing the experimental data:

(2)

where qi model is the value of q predicted by the fitted model, qi exp. is the experimental value of qand p is the number of experiments performed. The MPSD error function has been usedpreviously 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 differentbiosorbents. This means that an adsorption process on a particular biosorbent may be kinetically-controlled, diffusion-controlled or even a combination of the two (Cheung et al. 2003). With thispoint in mind, attempts were first made to fit the kinetic adsorption data for Mn(II) ions onto GCCwith the pseudo-first-order rate expression as given by:

(3)

where Qe and Qt are the amounts of metal ions adsorbed per unit mass of adsorbent at equilibriumand at time t (min), respectively, and k1 is the pseudo-first-order rate constant for the adsorptionprocess. Linear plots of log(Qe – Qt) versus t at different initial metal ion concentrations indicatethe 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 fittingpurposes in most adsorption kinetic studies for the determination of the kinetic parameters,

log( ) log.

Q Q Qk

te t e− = −

1

2 303

Fq q

q perror

i

ii

p

% model i exp.

exp.

( ) = ×−

−

∑1001

1

2

•

QC C

mVe

i e=−


information is still lacking in the literature concerning possible changes in the adsorption rates asa function of initial metal ion concentration and adsorption time, as well as regarding thedetermination of fractional kinetic orders. For this reason, the alternative fractional order kineticequation proposed by Avrami (Lopes et al. 2003; Cestari et al. 2004) has been used to analyze thedata obtained in the present study. This equation may be expressed as:

(4)

where α is the fraction adsorbed (qt /qe) at time t, kAV is the Avrami kinetic constant (min–1) and nis a fractional reaction order.

Pseudo-second-order kinetics

The experimental data were also analysed employing the pseudo-second-order kinetic modelwhich may be expressed in the following form:

(5)

where k2 [g/(mg min)] is the rate constant in the pseudo-second-order equation, Qt (mg/g) is theamount 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-Marinet al. 2007). This equation has been applied satisfactorily to some chemisorption processes andhas been found useful in covering a wide range of slow adsorption rates. The equation, often validfor systems in which the adsorbing surface is heterogeneous, can be written as:

(6)

where α is the adsorption rate [mg/(g min)] and β (g/mg) is related to the extent of surfacecoverage 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 theequation:

Qt = kidt1/2 + C (7)

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

Qt =

+

1 1

βln( ) ln( )αβ t

β

t

Q k Q Qt

t e e

= +1 1

22

α = − −[ ]1 exp ( )k tAV

n


Adsorption isotherms

Langmuir isotherms

The Langmuir adsorption model has often been used to describe equilibrium adsorptionisotherms. This model is valid for monolayer sorption onto a surface with a finite number ofidentical sites and may be written as:

(8)

where Q0 is the maximum amount of metal ions adsorbed per unit weight of GCC to form a completemonolayer on the surface at equilibrium concentration, while b is related to the affinity of the bindingsite. The Langmuir parameters were calculated in the same way as employed in the literature byother 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 theconcentration of the adsorbate on the adsorbent surface increases as the initial adsorbateconcentration increases. In theory, this expression suggests the possibility of an infinite amount ofadsorption. The corresponding relationship may be expressed as:

(9)

where Kf and n are the Freundlich constants which indicate the adsorption capacity and adsorptionintensity, respectively.

Sips isotherm

This model is a combination of the Langmuir and Freundlich isotherm models. The Sips model(1948) takes the form:

(10)

where Ks is the Sips constant related to the affinity constant (mg/�)–1/n, Qmax is the Sips maximumadsorption capacity (mg/g) and ns is the Sips exponent (dimensionless). At low sorbateconcentrations, this equation effectively reduces to the Freundlich isotherm, while at high sorbateconcentrations it predicts a monolayer adsorption capacity characteristic of the Langmuirisotherm.

Thermodynamic parameters

On the basis of fundamental thermodynamics concepts, it may be assumed that energy cannot begained or lost in an isolated system and that the entropy change is the sole driving force. Inenvironmental engineering practice, both energy and entropy factors must be considered in order

qQ K C

K Ces e

n

s en

=+max

/

/

1

11

Q K Cf en

e = 1/

QQ bC

bCe

ee =

+

0

1


to determine which process will occur spontaneously. In the present study, an increase intemperature led to a decrease in the Mn(II) ion biosorption rate. Thermodynamic parameters suchas 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 well-known relationships (Sharma et al. 2007; Catena and Bright 1989):

∆G0 = –RT ln KL (11)

∆G0 = ∆H0 – T∆S0 (12)

(13)

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 inFigure 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 vibrationin –COH). The intensity of the peaks in the transmittance spectra was relatively greater for themetal ion-loaded GCC beads relative to the unloaded GCC beads. This higher intensity may beattributed to the presence of a smaller number of functional groups in the loaded GCC beads. Thisobservation accords with the view that functional groups such as –NH2, –OH and –CO– areinvolved 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 beapproximately spherical in shape. In addition, the texture of the beads appears to consist of holesand small openings on the surface. These aspects result in an increase in contact area, therebyfacilitating pore diffusion during the adsorption process.

Effect of pH

In order to optimise the pH for maximum removal efficiency, experiments were conducted atroom temperature employing 100 m� of an Mn(II) ion solution of 100 ppm initial concentrationcontaining 0.1 g of glutaraldehyde cross-linked chitosan over the pH range 3–8. This pH rangewas chosen so as to avoid chitosan dissolution, which occurs within the pH range 3–8 withconsequent precipitation of Mn(II) ions. From Figure 3, it will be seen that the uptake capacitytowards Mn(II) ions increased with increasing solution pH, with the adsorption capacity being amaximum 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

ln KS

R

H

RTL = −∆ ∆0 0



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

4000 3500 3000 2500 2000

Wavenumber (cm–1 )

1500 1000 500

80

81

82

83

84

85

86

87

% T

rans

mitt

ance

88

89

90

91

92

93

94

95

96

97

98

b

a


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

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

0 2 4 6 8 10

pH

Ads

orpt

ion

capa

city

(m

g/g)

0

10

20

30

40

50

60

70

80

90

the surface properties of the chitosan. This may be due to the formation of soluble hydroxycomplexes (Nacer Kocer et al. 2008).

In acidic solutions, strong competition exists between the metal ions and protons for theavailable sorption sites on the adsorbent and this leads to a decrease in the uptake capacity ofMn(II) ions. Ionisable functional groups, such as –NH2 or –OH, on the surface of cross-linkedchitosan may gain or lose protons, resulting in a surface charge that varies with pH. At low pH, thesurface sites are protonated and the surface becomes positively charged; in contrast, at high pH, theionisable 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 ofMn(II) ions being present in the solution as Mn(H2O)6

2+. At pH > 8, Mn(OH)2 is formed as observedexperimentally. 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 extentof Mn(II) ion adsorption. For a given initial Mn(II) ion concentration, the removal efficiencyincreased up to an optimum adsorbent dosage, beyond which any further increase in removalefficiency became negligible. This is to be expected because, for a fixed initial soluteconcentration, increasing the adsorbent dosage led to an increase in the surface area and hence agreater number of adsorption sites. The results depicted in Figure 4 indicate that GCC removed97% of the Mn(II) ions initially present in solution.


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

0 0.1 0.2 0.3 0.4 0.5 0.6

% R

emov

al o

f Mn(

II) io

ns

0

20

40

60

80

100

120

Adsorbent dosage (g/�)

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 wasexamined by plotting log(Qe – Qt) versus time (min) to evaluate the rate constant k1, while the

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


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

3.0 3.5 4.0 4.5 5.0

ln t

ln[l

n(1

Qt)/

Qe]

0

0.2

0.4

0.6

0.8

1.0

1.2

1.4

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 equilibriumadsorption capacity, Qe, also being listed in Table 1. From the data in Table 1, it may be concludedthat the experimental data were well fitted by the pseudo-second-order kinetic model. This issupported by the high R2 and low error function values, with the values of q as calculated by thismodel being in good agreement with the experimental data. The error function evaluates thedifferences associated with each individual points fitted by the model in relation to eachexperimental point measured. This observation supports the contention that the adsorption ofMn(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 interceptthe greater the boundary layer effect. Using this model, linear plots of Qe versus t1/2 passingthrough the origin indicate that intra-particle diffusion alone determines the overall rate ofadsorption. However, for the data obtained in the present study, the linear plots did not passthrough the origin thereby indicating that intra-particle diffusion was not the sole rate-determiningfactor and that other factors appeared to be operating simultaneously. This indicates that the


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) 0.012 0.0081 0.0087 0.0092R2 0.984 0.994 0.998 0.993MPSD (%) 86 94 97.3 106.4

Pseudo-second-order k2 [g/(mg min)] 0.001 0.0004 0.0004 0.0004R2 0.998 0.998 0.998 0.999Qe 75.2 169.5 270.2 344.8MPSD (%) 4.1 4.5 4.3 3.2

Chemisorption α [mg/(g min)] 595.8 2591.5 6.12 × 104 6.4 × 105

β (g/mg) 0.136 0.06 0.05 0.05R2 0.962 0.948 0.956 0.967MPSD (%) 1.9 2.5 2.1 1

Fractional order kAV (min–1) 0.001 0.001 0.001 0.001Qe 13.5 26.7 32.2 36.7N 0.33 0.4 0.3 0.3R2 0.9855 0.9766 0.9913 0.9477MPSD (%) 100 102.3 99.2 103.6

Weber–Morris kid [mg/(g min)] 1.59 3.48 4.29 4.69R2 0.953 0.923 0.916 0.957C 48.8 113 202.8 277.2

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

0

0.5

1.0

1.5

2.0

2.5

3.0

3.5

0 50 100 150 200 250

Time (min)

t/Qt

Equilibrium isotherm data

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


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

0

50

100

150

200

250

300

350

400

3.0 3.5 4.0 4.5 5.0 5.5ln t

Ads

orpt

ion

capa

city

(m

g/g)

mechanism of Mn(II) ion adsorption by GCC is complex and both surface adsorption as well asintra-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 dataarising from each of the initial Mn(II) ion concentrations employed. These are presented inFigure 7, while the corresponding values of α, β and the error function are listed in Table 1. It willbe seen that the plots were linear with good correlation coefficients. The equilibriumconcentrations calculated from this model were closely related to the values obtainedexperimentally. This suggests that, although the sorption system studied exhibited pseudo-second-order kinetics, the rate-determining step may involve chemisorption which occurs through thesharing or exchange of electrons between the adsorbent and the adsorbate.

Thermodynamics

The values of ∆G0 for the sorption of Mn(II) ions onto GCC at different temperatures are listed inTable 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. Thesmaller 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 versus1/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 evidencethat the adsorption process was exothermic in nature, while the negative value of the standardentropy change indicated that adsorption was favoured in the system (Sharma et al. 2007).


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

0 0.05 0.10 0.15 0.20 0.25 0.30 0.35

1/Ce

1/Q

e

0

0.05

0.10

0.15

0.20

0.25

0.30

0.35

0.40

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 –12.18930 –9.8 –107.2 –0.319335 –9.7

CONCLUSIONS

The results obtained in the present study show that glutaraldehyde cross-linked chitosan may beconsidered 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

leading to an increase in the percentage removal of ions. The biosorbent was characterised by FT-IRspectroscopy and scanning electron microscopy (SEM) methods, the equilibrium adsorption datawere interpreted via the Langmuir, Freundlich and Sips models while the kinetic adsorption datawere studied using the pseudo-first-order, pseudo-second-order and Elovitch isotherm models. Ofthese models, that of Langmuir provided the best fit of the equilibrium adsorption data while thekinetic 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. Thermodynamicinvestigations indicated that the adsorption process was spontaneous and exothermic.

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