Equilibrium, kinetic and thermodynamic studies on perchlorate adsorption by cross-linked quaternary chitosan

Chemical Engineering Journal 192 (2012) 269–275

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Equilibrium, kinetic and thermodynamic studies on perchlorate adsorption by cross-linked quaternary chitosan Yanhua Xie a,b, Shiyu Li b, Guangli Liu b,⇑, Jun Wang b, Ke Wu b a b

State Key Laboratory of Geohazard Prevention and Geoenvironment Protection, Chengdu University of Technology, Chengdu 610059, China School of Environmental Science and Engineering, Sun Yat-sen University, Guangzhou 510275, China

a r t i c l e

i n f o

Article history: Received 23 December 2011 Received in revised form 3 April 2012 Accepted 4 April 2012 Available online 13 April 2012 Keywords: Cross-linked quaternary chitosan salt Adsorption Perchlorate

a b s t r a c t Cross-linked quaternary chitosan salt (QCS) was used for perchlorate removal from aqueous solution. The effect of pH and co-existing anions on the adsorption process, adsorption isotherms, kinetics and thermodynamics was investigated. The thermo stability and the solubility in water were improved after QCS was cross-linked by glutaraldehyde. The cross-linked QCS showed maximum monolayer adsorption capacity of 119.0 mg/g, which was 2.5 times higher than protonated cross-linked chitosan. The adsorption process was almost independent of pH in a range from 4.0 to 10.3, which was much wider than using protonated cross-linked chitosan with applicable pH range of 4.0–6.0. The effect from co-existing anions on perchlo2 rate adsorption followed the order: Cl > NO 3 > SO4 . The adsorption data fitted Langmuir and Tempkin isotherm models well. Adsorption kinetic analysis showed that chemical adsorption was a rate-limiting stage during the adsorption process. The adsorption reaction was exothermic and spontaneous according to the calculated result of adsorption thermodynamics. The spent adsorbents could be effectively regenerated by NaCl solution with the concentration higher than 0.3%. Fourier transformation infrared spectra and the change of Cl concentration in solution demonstrated that the ion-exchange interaction might occur between the chloride ions on the quaternary ammonium salt groups of the cross-linked QCS and the perchlorate ions in solution, resulting in the perchlorate adsorption onto the cross-linked QCS. Ó 2012 Elsevier B.V. All rights reserved.

1. Introduction Perchlorate has been widely used in the solid propellants of rockets/missiles, safety flares, air bags, fireworks, and batteries, etc. [1–4]. However, excess perchlorate exposure to human bodies can block the uptake of iodine in the thyroid gland, resulting in a decreased production of thyroid hormones and then the retardation of physical and mental growths [5]. Many researchers have shown that perchlorate has been widely and increasingly detected in ground and surface water, foodstuffs, and even in human breast milk [6–10]. Given the public health risk, the United States Environmental Protection Agency established an official reference dose of 0.0007 mg/kg/day for perchlorate, corresponding to a drinking water equivalent level of 24.5 lg/L in 2005 [11]. Perchlorate removal is a great challenge to the traditional drinking-water treatment process because it is non-volatile, highly soluble, and kinetically inert in water [12,13]. Various treatment technologies including adsorption, ion exchange, membrane filtration, biological treatment, and chemical/catalytic reduction have been developed for perchlorate removal. Among those technologies, adsorption is regarded as one of the most promising methods, ⇑ Corresponding author. Tel.: +86 20 39332690; fax: +86 20 84110267. E-mail address: [email protected] (G. Liu). 1385-8947/$ - see front matter Ó 2012 Elsevier B.V. All rights reserved. http://dx.doi.org/10.1016/j.cej.2012.04.014

due to its advantages of low investment, easy design and operation [14]. Many novel adsorption materials have been tested for perchlorate removal from the aqueous solution. For example, the activated carbon (AC) which was tailored by cationic surfactants showed 30 times higher adsorption capacity of perchlorate than virgin AC [15,16]. Kumar et al. [14] demonstrated that granular ferric hydroxide could adsorb the perchlorate ions with the capacity of 20.0 mg/g at pH 6.0–6.5 and 25 °C. Lien et al. [17] showed that 95% of perchlorate with the initial concentration of 10 mg/L could be removed within 24 h using acidified zero-valent aluminum and aluminum hydroxide with the dosage of 35 g/L (calculated by aluminum) at pH 4.5. Giant reed modified by quaternary amine (QA) functional groups showed high adsorption capacity of perchlorate but its optimal pH value was narrowed to the range of 3.5–7.0 [18]. The adsorption process for perchlorate removal is still limited by low adsorption capacities, and narrow pH range for stable performance. As one of abundant natural organics, chitosan has attracted our attention for its high hydrophilicity, non-toxicity, and biodegradability. Protonated cross-linked chitosan showed a maximum monolayer adsorption capacity of 45.455 mg/g for perchlorate at the pH 4.0 [19]. Coupled with ultrafiltration process, i.e., hybrid adsorption/ultrafiltration process, chitosan with dosage of 0.8 g/L could remove 92% of perchlorate with the initial concentration of

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10 mg/L in less than 10 min at pH 4.3 [20]. Although the chitosan showed a promising result as an adsorption material for perchlorate removal, a narrow range of optimal pH value (4.0–6.0) and low selectivity to the perchlorate ion would limit its application in the future. In order to further improve the adsorption capacity, quaternary chitosan salt (QCS) was used to test perchlorate adsorption as its quaternary ammonium salt group shows a strong basic exchanger. The QCS was cross-linked by glutaraldehyde to make it insoluble in aqueous solution and reusable. The adsorption process of perchlorate onto the cross-linked QCS from water was investigated in details. 2. Materials and methods 2.1. Preparation of cross-linked QCS QCS with the substitution degree of 83% was purchased from Aoxing Biotechnology Co. (Zhejiang, China). The cross-linked QCS was prepared through following procedures. QCS (10 g) was suspended in 200 mL anhydrous alcohol. Then 50 mL glutaraldehyde (25% w/v) was added to the suspension. The mixture was continuously stirred for 24 h at 25 ± 1 °C [21]. Being separated, the product was washed by 0.1 mol/L NaOH solution and deionized water, filtered and finally dried at 100 °C. The reagents used in this study, including NaClO4H2O, NaCl, NaNO3, Na2SO4, Na2CO3, HCl, NaOH, AgNO3, acetate acid, anhydrous alcohol, and glutaraldehyde were of analytical grade supplied by Guangzhou Chemical Reagent Factory, China. 2.2. Analysis methods 

2 were deterThe concentrations of ClO4 , Cl, NO 3 , and SO4 mined by an ion chromatography (IC) system (ICS 3000, Dionex), which was equipped with a set of 2  250 mm AS19 column and guard column, a 2-mm ASRS suppressor, an electrical conductivity detector, and an auto-sampler. Sodium hydroxide (50 mM) was  used as the eluent for ClO4 detection, while 15 mM sodium 2 hydroxide for the analysis of Cl, NO 3 , and SO4 . With the injection  loop of 250 lL, the detection limit was 5 lg/L for ClO4 and 2 lg/L 2 for Cl, NO , and SO . All samples were filtered with a 0.22 lm fil3 4 ter before measurement. The cross-linked QCS was characterized by Fourier transformation infrared spectrometer (FTIR, Equinox55 Bruker, Germany). KBr pellets were prepared through mixing the samples with KBr. Thermo gravimetric analysis (TGA) was performed using a thermo gravimetric analyzer (TG-209/Vector-22, Bruker, Germany). The heating rate was kept at 10 °C/min in N2 atmosphere.

In equilibrium experiments, 0.02 g adsorbent was placed in a set of 50 mL perchlorate solution with the initial concentration varying from 5 to 200 mg/L. The equilibrium isotherms were determined for 12 h at 20 ± 1, 35 ± 1 and 50 ± 1 °C, respectively. All the experiments were conducted in triplicates. The amount of perchlorate adsorbed at equilibrium (qe, mg/g) was calculated by the following equation:

qe ¼

ðC 0  C e ÞV W

ð1Þ

where C0 is the initial and Ce is the equilibrium concentrations of perchlorate in solution (mg/L), V is the solution volume (L), and W is the dry mass of adsorbent (mg). The adsorption kinetics of perchlorate was studied with different initial concentrations (5, 10, and 25 mg/L) at 25 ± 1 °C. 0.1 g cross-linked QCS was added in 250 mL perchlorate solution and continuously stirred for 24 h. About 5 mL samples were collected at the 5th, 10th, 15th, 30th, 60th, 120th, 240th, 480th, 720th and 1440th min, then analyzed by IC. 3. Results and discussion 3.1. Characterization of adsorbents The FTIR spectra of cross-linked QCS are shown in Fig. 1(a). After QCS was cross-linked by glutaraldehyde, an important peak appeared at 1652 cm1, which should be attributed to –C@N– (Schiff alkali) band [22,23]. The cross-linking reaction might take place between the amine functions on QCS and the aldehyde groups on glutaraldehyde, and Fig. 2 showed the possible cross-linking reaction scheme [21–24]. The cross-linking process made the adsorbent insoluble in water, which was favorable for its separation from solutions. Besides, the thermo stability was improved after QCS was cross-linked. The TGA curves of QCS and cross-linked QCS are shown in Fig. 3. Both QCS and cross-linked QCS began dehydration at about 100 °C. However, the second degradation stage reached a maximum at 270 °C for the cross-linked QCS, which was higher than that of QCS whose maximum temperature was 220 °C. The burning residue of cross-linked QCS was more than the QCS about 10%. The TGA results indicated that the cross-linked QCS was more stable than QCS itself. The difference of the TGA curve at 300–450 °C between QCS and cross-linked QCS might come from the cross-linking reaction of the glutaraldehyde with QCS.

a

618 1375

1652

2933 3407

1066

688

b 1492

About 0.02 g of cross-linked QCS was placed in a series of conical flasks, which contained 50 mL perchlorate solution. The perchlorate solution was prepared by dissolving NaClO4H2O with required concentrations. The initial pH of the solution, adjusted by adding HCl (1.0 and 0.1 M) and NaOH (0.5 M), ranged from 3.0 to 10.3, while the initial perchlorate concentration was kept at 10 mg/L. Solutions were mixed in a shaking incubator at 180 rpm and a constant temperature at 25 ± 1 °C. After 12 h of reaction, 10 mL sample was taken with a single injector, filtrated, and analyzed by IC system. The effect from co-existing anions (initial concentration ranging from 5 to 200 mg/L), including Cl, NO 3, and SO2 4 , on perchlorate adsorption was examined. The experiment was conducted with the same as the pH effect experiment procedure described previously with no pH adjusting.

1492

2.3. Adsorption experiments

1377 2934

4000

3500

624

1653

3406

1108

3000

2500 2000 Wav (cm-1)

1500

1000

500

Fig. 1. FTIR spectra of (a) cross-linked QCS and (b) cross-linked QCS loaded with perchlorate.

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(HOC–(CH2)3–COH)n Reactive

Fig. 2. Reaction scheme of cross-linked QCS.

30

100 25

a

60

qe (mg/g)

Weight (%)

80 b

40

20

15

20 10

0

150

300 450 600 Temperature (̊C)

750

900

Fig. 3. TGA curves of (a) QCS and (b) cross-linked QCS.

Sruface charge (mol/kg)

16 8 0 -8 -16 -24

2

4

6

pH

8

10

12

Fig. 4. Potentiometric titration curve of the cross-linked QCS.

3.2. Effect of pH The pHzpc (pH at the surface has a net zero charge) of the crosslinked QCS was 10.6, which was determined by the potentiometric titration method (see in Fig. 4) [25,26]. When the pH value of the

2

4

6

8

10

12

pH Fig. 5. Effect of pH on adsorption capacity of perchlorate onto cross-linked QCS.

solution was less than the pHzpc (10.6), the surface of cross-linked QCS was positively charged, which was beneficial for anion adsorption. So the adsorption ability of the cross-linked QCS for perchlorate might be independent of the solution pH when the pH was less than 10.6. The effect of pH on perchlorate adsorption by crosslinked QCS was conducted, and the result is showed in Fig. 5. The amount of perchlorate adsorbed at equilibrium varied slightly from 23.5 to 24.3 mg/g with a pH range from 4.0 to 10.3. The results indicated that the adsorption capability of the adsorbent was nearly independent of pH in a range from 4.0 to 10.3, which was consistent with the pHzpc discussed above. While the adsorption ability of the cross-linked QCS decreased slightly when the pH further reduced as excessive Cl (introduced by acid adding) would compete with perchlorate to occupy the adsorption sites on the adsorbent. But the cross-linked QCS still exhibited an adsorption capacity of 22.4 mg/g at about a pH value of 3, which indicated that the adsorbent could endure strong acidic condition and show relatively good adsorption ability for perchlorate. So the optimal pH range of the cross-linked QCS was 4.0–10.3, which was much wider than that of the protonated cross-linked chitosan (4.0–6.0). The amount of perchlorate adsorbed at equilibrium of the cross-linked QCS (24.3 mg/g) was twice as much as that of the protonated cross-linked chitosan (11.9 mg/g) under the same reaction conditions [19].

Y. Xie et al. / Chemical Engineering Journal 192 (2012) 269–275

25

Cl NO3

20

SO4

(a)

-

1.8 1.5

2-

1.2 Ce/qe g/L

qe (mg/g)

272

15 10

0.9 0.6 0.3

5

293 K 308 K 323 K

0.0

0

0

50

100 C0 (mg/L)

150

-0.3 -30

200

Fig. 6. Effect of co-existing anions on adsorption capacity of perchlorate onto the cross-linked QCS.

(b)

0

30

60 90 Ce (mg/L)

120

150

180

2.4 2.1

The effects of the co-existing anions, including Cl, NO 3 , and with the initial concentrations ranging from 0 to 200 mg/L were investigated. As shown in Fig. 6, the adsorption capacity of perchlorate decreased with increasing anion concentrations. With co-existing Cl, the perchlorate adsorption capacity on cross-linked QCS decreased from 24.3 mg/g (initial perchlorate solution of 10 mg/L, no Cl) to 15.0 mg/g (Cl concentration of 20 mg/L). The adsorption capacity of percholrate further decreased to 8.7 mg/g when Cl concentration increased to 200 mg/L. Similar adsorption 2  inhibition was observed in NO 3 and SO4 . With co-existing NO3 of 200 mg/L, the adsorption capacity decreased to 11.3 mg/L. But less adsorption inhibition was observed in SO2 4 , which remained 14.38 mg/g adsorption capacity with 200 mg/L co-existing SO2 4 . The inhibiting effect from the co-existing anions on the adsorption of perchlorate to the cross-linked QCS followed the order: 2 Cl > NO 3 > SO4 . This phenomenon was different from that of using the protonated cross-linked chitosan as adsorbent, where SO2 showed the greatest inhibition than the others. In addition, 4 the cross-linked QCS exhibited relatively higher adsorption capacity for perchlorate than the protonated cross-linked chitosan when the co-existing anions existed at the same reaction conditions. The reason might be that the quaternary ammonium salt group has more strong exchange or adsorption force for perchlorate.

lg qe

3.3. Effect of co-existing anions 1.8

SO2 4 ,

3.4. Adsorption equilibrium isotherm The adsorption of perchlorate on the cross-linked QCS was investigated at different temperatures (20, 35 and 50 °C), and the result was shown in Fig. 7. The equilibrium adsorption data were analyzed using the Langmuir, Freundlich, and Tempkin isotherm models, and their linear forms were given as follows [27–29]:

Ce 1 Ce ¼ þ qe Q 0 b Q 0 lg qe ¼ lg K f þ

ð2Þ 1 lg C e n

qe ¼ B ln A þ B ln C e

ð3Þ ð4Þ

where Ce is the equilibrium concentration of perchlorate (mg/L), qe is the amount of adsorbate adsorbed per unit mass of adsorbent at equilibrium (mg/g), Q0 (mg/g) and b (L/mg) are the Langmuir constants reflecting the adsorption capacity and the energy of adsorption, respectively, Kf and 1/n are the Freundlich constants, with Kf referring to the adsorption capability of the adsorbent and 1/n

1.5 293 K 308 K 323 K

1.2 -1.50

-0.75

0.00

0.75

1.50

2.25

3.00

lg Ce

(c)

150

qe (mg/g)

120 90 60 30

293 K 308 K 323 K

0 -4

-2

0

2 lnCe

4

6

Fig. 7. Isotherm plots for the adsorption of perchlorate onto cross-linked QCS (a) Langmuir isotherm, (b) Freundlich isotherm, and (c) Tempkin isotherm.

giving an indication of how favorable the adsorption process was, A and B are the Tempkin constants related to the maximum binding energy and the heat of adsorption, respectively. The fitting values of the three models along with regression coefficients (R2) are listed in Table 1. As shown in Table 1, the equilibrium data fitted the Langmuir and Tempkin isotherm models well with the R2 being higher than 0.99 at different temperatures. It indicated that perchlorate was adsorbed mainly in the form of monolayer coverage on the surface of the adsorbent. Tempkin isotherm assumed that the adsorption heat of the molecules in the layer decreased linearly with coverage, because of adsorbate–adsorbent interactions [27]. So the adsorption process occurred along with the change of the adsorption heat. The values of Q0 increased with decreased temperature, which indicated that low temperatures were favorable for the adsorption process. A decrease in the values of b and A versus temperature

273

Y. Xie et al. / Chemical Engineering Journal 192 (2012) 269–275 Table 1 Langmuir, Freundlic and Tempkin isotherm model constants and correlation coefficients for adsorption of perchlorate onto cross-linked QCS. T (K)

293 308 323

Langmuir isotherm

Freundlich isotherm 2

(11/n)

Q0 (mg/g)

b (L/mg)

R

KF (mg

119.048 116.279 113.636

0.341 0.332 0.206

0.998 0.999 0.997

34.546 32.900 27.593

L

Tempkin isotherm

1/n

/g)

2

1/n

R

0.286 0.294 0.316

0.945 0.944 0.950

A (L/g)

B

R2

27.487 21.438 10.849

14.558 14.716 15.302

0.993 0.991 0.994

Table 2 Kinetic model constants and correlation coefficients for adsorption of perchlorate onto cross-linked QCs. C0 (mg/L)

5 10 25

Pseudo-first-order

Pseudo-second-order

Intra-particle diffusion

k1 (h1)

R2

k2 (g/mg h)

R2

kp,1 (mg/g h1/2)

R2

kp,2 (mg/g h1/2)

R2

Kp,3 (mg/g h1/2)

R2

0.547 0.525 0.429

0.640 0.750 0.692

1.689 0.883 0.249

0.999 1.000 1.000

17.110 34.358 75.055

0.966 0.955 0.956

0.117 0.267 1.359

0.946 0.969 0.984

0.006 0.009 0.139

0.997 0.983 0.993

showed that the bond was weak at high temperatures, which further confirmed that low temperatures were favorable for the adsorption process [28]. The essential features of the Langmuir isotherm could be expressed by a dimensionless constant named equilibrium parameter RL [29,30], which was defined as:

RL ¼

1 ð1 þ bC 0 Þ

ð5Þ

where b (L/mg) is the Langmuir constant and C0 is the initial concentration of perchlorate (mg/L). RL values denoted whether the adsorption was irreversible (RL = 0), favorable (0 < RL < 1), linear (RL = 1), or unfavorable (RL > 1). All of the calculated RL values (data not shown) varied from 0.01 to 0.5 with the initial perchlorate concentration ranging from 0.5 to 60.0 mg/L at different temperatures, which indicated that the perchlorate adsorption on the adsorbent was favorable. The maximum monolayer adsorption capacity of the cross-linked QCS reached to 119.0 mg/g, which was 2.5 times higher than that of using protonated cross-linked chitosan (45.0 mg/g) as sorbent. 3.5. Adsorption kinetics Three kinetic models were used to investigate the mechanism of perchlorate adsorption on cross-linked QCS. Firstly, the pseudo-first-order and pseudo-second-order models were discussed in the following equations [31]:

lnðqe  qt Þ ¼ ln qe  k1 t

ð6Þ

t 1 t ¼ þ qt k2 q2e qe

ð7Þ

where qe and qt are the amounts of perchlorate adsorbed at equilibrium (mg/g) and at time t (h), respectively, k1 and k2 are the rate constants of the pseudo-first-order and pseudo-second-order models, respectively. Values of k1 and k2 could be calculated from the slopes of their linear equations. The calculated results are listed in Table 2. From Table 2, the experimental data fitted the pseudo-secondorder model better than the pseudo-first-order model. The pseudo-second-order model was based on two assumptions: (1) the adsorption followed the Langmuir isotherm, and (2) the rate-limiting process should be chemical adsorption [32,33]. Adsorption isotherm analysis (see the Section 3.4) demonstrated that the adsorption process did follow the Langmuir isotherm. The stimulated results gained from the pseudo-second-order model

indicated that the adsorption of perchlorate onto cross-linked QCS would take place mainly through chemical adsorption. Secondly, intra-particle diffusion model was used to identify the diffusion mechanism as follows [34,35]:

qt ¼ kp;i t1=2 þ C i

ð8Þ

where kp,i (i = 1–3) is the intra-particle diffusion rate constant (mg/ gh1/2), i.e., kp,1 is the constant in the first stage called the external surface adsorption, kp,2 is the constant in the second stage as the gradual adsorption, and kp,3 is the constant in the third stage as the final equilibrium [36,37]. The external surface adsorption stage could be finished within 30 min as shown in Fig. 8. The kp,1 increased significantly with the increase of the initial perchlorate concentration due to the enhancement of the concentration gradient force [31]. The linear lines of stages 2 and 3 did not pass through the origin, which indicated that intra-particle diffusion was not the rate limiting mechanism for this adsorption process [28,29]. In addition, the kp,1 was much greater than the kp,2 and kp,3 (see Table 2), which indicated that external surface adsorption was much faster than intra-particle diffusion. The results further proved that the intra-particle diffusion was not the rate-limiting step in the whole adsorption process. 3.6. Adsorption thermodynamics Thermodynamic characteristics of perchlorate adsorption on the cross-linked QCS were determined, including the standard free energy change(DG0), standard enthalpy change (DH0), and standard entropy change (DS0) [38,39]. The values of DG0 were estimated according to the Gibbs free energy equation as follows:

DG0 ¼ RT ln K d

ð9Þ

where R is the universal gas constant (8.314 J/mol K), T is the absolute solution temperature (K), and Kd is the variation of the thermodynamic equilibrium constant. The value of Kd obtained by plotting ln (qe/Ce) versus qe and extrapolating qe to zero. Its intercept with the vertical axis gave the values of Kd. Where, qe is the amount of  ClO4 adsorbed by per mass of cross-linked QCS (mmol/g) and Ce is  the ClO4 concentration in solution at equilibrium (mmol/L) [38,39]. The standard enthalpy change could be calculated according to the Van’t Hoff equation [28,40]:

ln K d ¼ 

DH 0 þ const RT

ð10Þ

where DH0 is assumed to be independent of temperature. The plot of ln Kd versus 1/T fitted a straight line and its slope gave the value of DH0. The standard entropy change could be calculated as:

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Y. Xie et al. / Chemical Engineering Journal 192 (2012) 269–275

(a)

6

12 10

4 5 mg/L 10 mg/L 25 mg/L

0 -2

2

t/qt(h g/mg)

4

6 t/h

8

10

12

5 mg/L 10 mg/L 25 mg/L

0.4 0.2 0.0 0

2

4

6 t/h

8

10

12

80 60

qe (mg/g)

300

600 900 t(min)

1200

1500

3.7. Adsorption mechanism

0.6

40 20 5 mg/L 10 mg/L 25 mg/L

0

0

1

2

1/2

1/2

3

4

5

t /h

Fig. 8. Kinetic plots for the adsorption of perchlorate onto cross-linked QCS (a) Pseudo-first-order model, (b) Pseudo-second-order model, and (c) Intra-particle diffusion model.

Table 3 Thermodynamic parameters for adsorption of perchlorate onto cross-linked QCS. T (K)

Kd

DG0 (kJ/mol)

293 308 323

5.884 5.611 4.936

4.317 4.417 4.288

0

0

Fig. 9. Changes of chloride and perchlorate concentrations in solution during the adsorption process.

0.8

0

4

-2 0

1.0

-20

6

0

-6

(c)

Cl

2

-4

(b)

-

8 Ct(mg/L)

ln(qe-qt)

2

-

ClO4

DG ¼ DH  T DS

0

DH0 (kJ/mol)

DS0 (J/mol K)

4.569

0.858 0.493 0.870

Compared with that of the virgin cross-linked QCS, the FTIR spectrum of the cross-linked QCS loaded with perchlorate showed no significant changes (see Fig. 1). However, a new acute peak appeared at 624 cm1, which should be attributed to the perchlorate anions adsorbed by cross-linked QCS [41]. Adsorption kinetic analysis indicated that the adsorption of perchlorate onto cross-linked QCS occurred mainly through chemical adsorption. It was assumed that the adsorption process might take place via the ion-exchange reaction between the Cl on the cross-linked QCS and the perchlorate in solution. Therefore, the variation of Cl concentration in the solution was measured for 12 h in batch tests during the adsorption process as shown in Fig. 9. The results demonstrated that the Cl concentration in the solution increased with decreasing of the perchlorate concentration. The molar quantity of the Cl released into the solution almost equaled to the perchlorate adsorbed by the adsorbents, which confirmed that the perchlorate adsorption occurred mainly through ion-exchange between the Cl on the cross-linked QCS and the perchlorate in solution. In addition, the regeneration of cross-linked QCS was conducted using NaCl (0.1%, 0.3%, and 0.5%), NaOH (0.01 and 1 M) and NaCl (0.1%)/NaOH (0.01 M) solutions as eluents. The results demonstrated (data not shown) that NaCl solutions exhibited relatively more effectiveness than others for displacing perchlorate anions from the adsorbents. And the regeneration ratio achieved nearly 100% when the concentration of NaCl solution was higher than 0.3%. Meanwhile, there was no significant deterioration of the adsorbent performance after the regeneration. The results further supported that the perchlorate adsorption might occur mainly through ion-exchange between the Cl on the cross-linked QCS and the perchlorate in solution. Therefore, NaCl solution with relatively high concentration could regenerate the exhausted adsorbent effectively. 4. Conclusions

ð11Þ

The thermodynamic parameters at different temperatures are listed in Table 3. The changes in enthalpy were negative, which indicated that the adsorption process was exothermic. Therefore, the maximum monolayer adsorption capacity of perchlorate on the adsorbent increased slightly with decreased temperature. The negative values of standard free energy indicated that the adsorption process could happen spontaneously.

The QCS was cross-linked by glutaraldehyde and tested for perchlorate removal from aqueous solutions. The cross-linked process improved its thermo stability, and made it insoluble in aqueous solutions and reusable. The optimal adsorption pH value was ranged from 4.0 to 10.3. The co-existing anions inhibited the adsorption processes, but the adsorption capacity still reached up to 14.38 mg/g with 200 mg/L of co-existing sulfate anion (initial perchlorate concentration of 10 mg/L), which was much higher than that of using chitosan. The adsorption data fitted the Langmuir and Tempkin isotherm models well, which indicated that

Y. Xie et al. / Chemical Engineering Journal 192 (2012) 269–275

perchlorate was adsorbed mainly in the form of monolayer coverage on the surface of the cross-linked QCS. Adsorption kinetics followed the pseudo-second-order model, which demonstrated that chemical adsorption was a rate-limiting stage. The calculated thermodynamic parameters showed that the adsorption process of perchlorate onto the cross-linked QCS was exothermic and could occur spontaneously. The spent adsorbents could be effectively regenerated by NaCl solution with the concentration higher than 0.3%. Perchlorate anions were adsorbed mainly through the ion-exchange reaction between the Cl ions on quaternary ammonium salt groups of the cross-linked QCS and the perchlorate ions in solution. The cross-linked QCS provides a promising candidate for perchlorate removal from the contaminant water. Acknowledgements This work was partly supported by grants from the Prominent National Projects of Science and Technology (No.2009ZX07528001), and the Integration Project of Production, teaching and research from Guangdong Province and Ministry of Education (No.2009B090300324, 2008B090500197). This work was also supported by grants from the State Key Laboratory of Geohazard Prevention and Geoenvironment Protection, and the Chinese National Natural Science Foundation (No. 51109019). References [1] R. Srinivasan, G.A. Sorial, Treatment of perchlorate in drinking water: a critical review, Sep. Sci. Technol. 69 (2009) 7–21. [2] P.B. Hatzinger, Perchlorate biodegradation for water treatment, Environ. Sci. Technol. 39 (2005) 239A–247A. [3] P.L. McCarty, T.E. Meyer, Numerical model for biological fluidized-bed reactor treatment of perchlorate-contaminated groundwater, Environ. Sci. Technol. 39 (2005) 850–858. [4] X.Y. Yu, C. Amrhein, M.A. Deshusses, M.R. Matsumoto, Perchlorate reduction by autotrophic bacteria attached to zerovalent iron in a flow-through reactor, Environ. Sci. Technol. 41 (2007) 990–997. [5] F.X. Li, L. Squartsoff, S.H. Lamm, Prevalence of thyroid diseases in Nevada counties with respect to perchlorate in drinking water, J. Occup. Environ. Med. 43 (2001) 630–634. [6] E.T. Urbansky, Perchlorate as an environnemental contaminant, Environ. Sci. Pollut. Res. 9 (2002) 187–192. [7] D.R. Parker, A.L. Seyfferth, B.K. Reese, Perchlorate in groundwater: a synoptic survey of ‘‘pristine’’ sites in the coterminous United States, Environ. Sci. Technol. 42 (2008) 1465–1471. [8] A.B. Kirk, E.E. Smith, K. Tian, T.A. Anderson, P.K. Dasgupta, Perchlorate in milk, Environ. Sci. Technol. 37 (2003) 4979–4981. [9] C.A. Sanchez, K.S. Crump, R.I. Krieger, N.R. Khandaker, P. Gibbs, Perchlorate in leafy vegetables of North America, Environ. Sci. Technol. 39 (2005) 9391–9397. [10] A.B. Kirk, P.K. Martinelango, K. Tian, A. Dutta, E.E. Smith, P.K. Dasgupta, Perchlorate and iodide in dairy and breast milk, Environ. Sci. Technol. 39 (2005) 2011–2017. [11] J.D. Roach, D. Tush, Equilibrium dialysis and ultrafiltration investigations of perchlorate removal from aqueous solution using poly(diallyldimethylammonium) chloride, Water Res. 42 (2008) 1204–1210. [12] I.H. Yoon, X.G. Meng, C. Wang, K.W. Kim, S. Bang, E. Choe, L. Lippincott, Perchlorate adsorption and desorption on activated carbon and anion exchange resin, J. Hazard. Mater. 164 (2009) 87–94. [13] B.E. Logan, Assessing the outlook for perchlorate remediation, Environ. Sci. Technol. 35 (2001) 482A–487A. [14] E. Kumar, A. Bhatnagar, J.A. Choi, U. Kumar, B. Min, Y. Kim, H. Song, K.J. Paeng, Y.M. Jung, R.A.I. Abou-Shanab, B.H. Jeon, Perchlorate removal from aqueous solutions by granular ferric hydroxide (GFH), Chem. Eng. J. 159 (2010) 84–90.

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