Role of pH in Metal Adsorption from Aqueous Solutions Containing

(EDTA), citric acid, tartaric acid, and sodium gluconate were used. It was shown that the ... Copper(II)-chelated species were often present in the ma...
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Ind. Eng. Chem. Res. 1999, 38, 270-275

Role of pH in Metal Adsorption from Aqueous Solutions Containing Chelating Agents on Chitosan Feng-Chin Wu,† Ru-Ling Tseng,† and Ruey-Shin Juang*,‡ Department of Chemical Engineering and Environmental Engineering, Lien Ho Junior College of Technology, Maio-Li, Taiwan 360, Republic of China, and Department of Chemical Engineering, Yuan Ze University, Chung-Li, Taiwan 320, Republic of China

The role of pH in adsorption of Cu(II) from aqueous solutions containing chelating agents on chitosan was emphasized. Four chelating agents including ethylenediaminetetraacetic acid (EDTA), citric acid, tartaric acid, and sodium gluconate were used. It was shown that the adsorption ability of Cu(II) on chitosan from its chelated solutions varied significantly with pH variations. The competition between coordination of Cu(II) with unprotonated chitosan and electrostatic interaction of the Cu(II) chelates with protonated chitosan took place because of the change in solution pH during adsorption. The maximum adsorption capacity was obtained within each optimal pH range determined from titration curves of the chelated solutions. Coordination of Cu(II) with the unprotonated chitosan was found to dominate at pH below such an optimal pH value. Introduction Copper(II)-chelated species were often present in the manufacturing process or waste streams of printed circuit boards (electroless Cu(II) plating) and the electroplating industry.1 This is because the presence of chelating agents not only prevents Cu(II) from hydroxide precipitation under alkaline conditions but also keeps the solution pH nearly constant. The chelated species existing in a rather wide pH range can hardly be removed or destroyed by the conventional chemical precipitation/coagulation methods such as hydroxide, sulfide, etc.1-3 Many approaches were made for this purpose such as activated carbon adsorption,4 ion exchange,5,6 substitution by Ca(II) or Fe(II) sulfate,7 and electrochemical methods.8,9 However, no satisfactory physicochemical and/or biological methods have been developed to cheaply remove or recover metals from dilute chelated solutions.9 Chitosan is a partially acetylated glucosamine biopolymer and mainly results from deacetylation of chitin.10 It can be produced cheaply because chitin is the second most abundant biopolymer in nature next to cellulose. It is also a known adsorbent for transition metals since the amine groups (-NH2) on chitosan can serve as coordination sites for many metals such as Cu(II), Ni(II), Zn(II), Cd(II), Hg(II), Cr(III), U(VI), and V(IV).11-16 Many studies indicated that proper adjustment of the solution pH leads to an increase in the metal adsorption capacity on chitosan.17-22 For example, Inoue et al. found a steep rise of the adsorption distribution ratio of Cu(II) on chitosan, defined as the equilibrium adsorption divided by the equilibrium concentration of the solution, at pH 4-5.18 On the other hand, the adsorption ability of metals can be reduced in the presence of anionic ligands such as carbonates, bicarbonates, EDTA, phosphate, etc.20,23,24 However, little attention was paid * To whom all correspondence should be addressed. Email: [email protected]. † Lien Ho Junior College of Technology. ‡ Yuan Ze University.

to examination of the effect of pH on the metal adsorption ability from solutions in the presence of chelating agents. The aim of this paper was to explore the role of pH on the liquid-phase adsorption of Cu(II) in the presence of chelating agents using chitosan and to provide fundamental data of Cu(II) adsorption from such chelated solutions. Experimental Section Chitosan and Reagents. Chitosan produced from lobster shell wastes was offered as flakes from YingHuah Co., Kaohsiung, Taiwan, without further purification. Prior to use as the adsorbent, the raw chitosan flakes were ground, and a particle size between 420 and 590 µm was used in this work. The BET surface area of chitosan was measured to be about 13 m2/g from N2 adsorption isotherms using Micromeritics model ASAP 2000. The degree of deacetylation for chitosan was obtained as about 87 mol % following the procedure of Guibal et al.11 The molecular weight of chitosan was found to be 4.1 × 105 by the Mark-Houwink equation from viscosity measurements. Four chelating agents, ethylenediaminetetraacetic acid (EDTA), citric acid, tartaric acid, and sodium gluconate, were supplied by Merck Co. as analytical reagent grade. The overall formation constants of these ligands with Cu(II) are compiled in Table 1, which is useful in data explanation. The aqueous solutions were prepared by dissolving equimolar chelating agent and CuSO4 in deionized water (Millipore Milli-Q) without pH adjustment. In several cases, the solution pH was adjusted between 2 and 10 by the addition of a small amount of HCl or NaOH. It is known that non-crosslinked chitosan is soluble in more concentrated carboxylic acid solutions. For example, the concentration limit of the most effective acid for this purpose, formic acid, is about 0.2 wt %.10,16 In this work, chitosan was insoluble in citric and tartaric acid solutions because the maximum concentration of citric and tartaric acids studied was 5 mol/m3, which only corresponds to 0.096 and 0.075 wt %, respectively.

10.1021/ie980242w CCC: $18.00 © 1999 American Chemical Society Published on Web 12/04/1998

Ind. Eng. Chem. Res., Vol. 38, No. 1, 1999 271 Table 1. Overall Formation Constant (log Kf)a for Complexation of Cation and Anionic Ligand (L) at 298 K and Zero Ionic Strength34 species H+

Cu2+

L- (OH-)

L2- (SO42-)

L4- (EDTA)

HL (14.0)

HL (1.99)

CuL (6.3) CuL2 (11.8) CuL4 (16.4) Cu2L2 (17.7)

CuL (2.4)

HL (11.12) H2L (17.80) H3L (21.04) H4L (23.76) H5L (24.76) CuL (20.5) CuHL (23.9) CuOHL (22.6)

L2- (tartarate)

L- (gluconate)

HL (6.40) H2L (11.16) H3L (14.29)

HL (4.37) H2L (7.44)

HL (3.48)b

CuL (7.2) CuHL (10.7) CuH2L (13.8) CuOHL (16.4) Cu2L2 (16.3)

CuL (3.97)

CuL (1.82)c

L3- (citrate)

a K is defined as, for example, K ) [M H L ]/[M]x[H]y[L]z (in molar units) for the reaction xM + yH + zL S M H L . b Obtained at an f f x y z x y z ionic strength of 1.0 mol/kg. c Taken from the data of NiL at an ionic strength of 0.1 mol/kg.

taken at preset time intervals, and the concentration of Cu(II) was analyzed as described previously. Results and Discussion

Figure 1. Equilibrium adsorption of Cu(II) on chitosan from solutions containing various chelating agents.

Equilibrium Experiments. In adsorption experiments, a fixed amount of dry chitosan (0.1 g) and 0.1 dm3 of an aqueous phase were placed in a 0.25-dm3 glass-stoppered flask and shaken at 130 rpm for 5 days using a thermostated shaker bath (Firstek model B603, Taiwan). Preliminary runs showed that the adsorption studied was complete after 4 days. After filtration, the aqueous pH was measured with a pH meter (Horiba model F-23) and the concentration of Cu(II) was analyzed using an atomic absorption spectrophotometer (GBC model 932). Each experiment was duplicated under identical conditions. The amount of Cu(II) adsorbed qe (mol/kg) can be obtained as follows:

qe ) (C0 - Ce)V/W

(1)

where C0 and Ce are the initial and equilibrium solution concentrations (mol/m3), V is the volume of solution (m3), and W is the amount of dry chitosan used (kg). Batch Kinetic Experiments. The contact-time experiments were made in a Pyrex glass vessel of 100mm i.d. and 130-mm height, fitted with four glass baffles, of 10-mm width. In each run, an aqueous phase (0.8 dm3) was first poured and stirred using a ColeParmer Servodyne agitator with a six-bladed, flatbladed impeller (12 mm high, 40 mm wide). A stirring speed of 500 rpm was adopted because above that the stirring has little effect on the adsorption process. A given amount of dry chitosan (0.48 g) was then added into the vessel, and the timing was started. The whole vessel was immersed in a thermostat controlled at 298 K (Haake model K-F3, Germany). Samples (5 cm3) were

Adsorption of Cu(II) from the Chelated Solutions. Figure 1 illustrates the equilibrium adsorption of Cu(II) in the absence of chelating agents on chitosan. In this case, the solution pH was not adjusted. A maximum adsorption capacity of around 2.5 mol/kg is obtained, which is comparable to those obtained as 2.6 mol/kg for Cu(II) adsorption from sulfate solutions25 and 2.3 mol/kg from 1000 mol/m3 NH4NO3 solutions.18 The present result indicates that the chitosan flakes used are very suitable for this purpose. It is known that the uptake of transition metals is mainly effected via coordination with the amine groups (NH2) on chitosan.12-16 Two OH groups and one NH2 group are grabbed by Cu(II), and the fourth site is probably occupied by a water molecule or the OH group on the third carbon atom. Hence, we have

Cu2+ + RNH2 S Cu(RNH2)2+

(2)

It is also known that the amino groups of chitosan may react with H+ according to

RNH2 + H+ S RNH3+, Kp

(3)

Because the protonation constant (log Kp) equals 6.3,26 on the average, more than 20% and 90% of the amino groups of chitosan are protonated even at pH 6.9 and 5.0, respectively. The two reactions of eqs 2 and 3 do compete during adsorption, which is supported from experiments that the solution pH more or less increases after adsorption. Figure 1 also shows the adsorption of Cu(II) in the presence of equimolar chelating agents. Here, the initial Cu(II) concentration C0 is within 0.3-4.5 mol/m3 and the solution pH was still not adjusted. Unlike EDTA, chitosan reveals a higher affinity for Cu(II) in the presence of chelating agents at rather low Cu(II) concentrations (Ce < 0.3 mol/m3) (Figure 2). The different behavior of EDTA is probably due to the low pKa2 of EDTA (3.24) compared to those of the other three chelating agents, as discussed in the following section. The divalent EDTA anions H2L2- presented at pH beyond pKa2 would form the neutral Cu(II) chelates or react with chitosan via electrostatic attraction, thus decreasing the adsorption capacity of Cu(II) because near such pH values the amino groups in chitosan are mostly protonated.

272 Ind. Eng. Chem. Res., Vol. 38, No. 1, 1999

Figure 2. Comparison of equilibrium adsorption of Cu(II) on chitosan at relatively low Cu(II) concentrations. The meaning of each symbol is the same as that given in Figure 1.

Figure 3. Correlation between solution pH and Cu(II) adsorption capacity in the presence of EDTA.

Table 2. Adsorption Capacity and Solution pH before and after Adsorption for Solutions Containing Equimolar (2 mol/m3) Chelating Agent and Cu(II) chelating agent

pH0

pHeq

qe (mol/kg)

EDTA tartaric acid citric acid sodium gluconate

2.38 2.76 2.72 4.38

2.77 4.00 3.37 5.29

0.45 0.55 0.90 1.74

On the other hand, when concentration of the chelates increases, the adsorption ability of Cu(II) drops especially with EDTA, tartarate, and citrate. Also, the pH increases after adsorption (Table 2) as found in the case of Cu(II) adsorption without chelating agents, but the pH increment depends on the type of chelating agents. It is seen from experiments that the higher the initial concentration of the chelates, the higher the equilibrium pH (not shown). Therefore, the decrease in the adsorption ability by increasing the Cu(II) concentration is believed to be a result of the competition between coordination of Cu(II) with unprotonated chitosan and electrostatic interaction of the Cu(II) chelates with protonated chitosan. This action is strongly pH-dependent. Preliminary tests show that the equilibrium adsorption of Cu-EDTA chelates from an aqueous phase on chitosan is greatly affected by the presence of inorganic salts such as Na2SO4, especially at pH < 6 (not shown). It is thus expected that the chelated anions, e.g., CuEDTA, react with the amine groups of chitosan via electrostatic attraction according to27,28

CuL2- + 2RNH3+ S CuL(RNH3)2

(4)

This is also the case of copper citrate chelated anions. For copper tartarate solutions, the drop in qe at high Ce (Figure 1) may be due to the effect of anionic species HL- rather than neutral species CuL (Table 1). In fact, Jha et al. found that the adsorption capacity of Cd(II) on chitosan from dilute solution (0.05 mol/m3) at pH 6.5 sharply drops to zero by adding a large excess of EDTA (20 mol/m3).20 The trace Cd(II) in the solutions completely exists as the chelated anions CdL2-, and they cannot compete with the EDTA anions (high ionic charge and high concentration) although about 40% of the amine group in chitosan is protonated in this case.

Figure 4. Correlation between solution pH and Cu(II) adsorption capacity in the presence of citrate.

For the gluconate system, the effect of aqueous complexation with Cu(II) is less obvious because of a comparatively small overall formation constant (Table 1). Also, the equilibrium pH and adsorption ability remain rather high with gluconate, as shown in Table 2. Optimal pH Range and Titration Curve. An attempt was then made to find the optimal pH range from solution chemistry that gives a maximum adsorption capacity. A total of 100 mL of an equimolar solution of Cu(II) and chelating agent (20 mol/m3) in the absence of chitosan was titrated with 100 mol/m3 NaOH. Figures 3-6 show these results. The capacities measured at different conditions are also indicated for comparison. Variations of adsorption capacity with equilibrium pH are unexpectedly sharp, which are different from those observed earlier for adsorption of metals alone on chitosan and its chemically modified derivatives.17-22 This is because of the occurrence of aqueous complexation of Cu(II) with chelating agent in this work. The pH trends of metal adsorption, either physically or chemically, are traditionally explained by the changes in distribution of the species with the pH.17,19,21,22 To our best knowledge, the changes in solution environments including the presence of strong chelating agents are scarcely studied. Here, it is evident that coordination of Cu(II) with unprotonated chitosan, interaction of

Ind. Eng. Chem. Res., Vol. 38, No. 1, 1999 273

Figure 5. Correlation between solution pH and Cu(II) adsorption capacity in the presence of tartarate.

Figure 7. Equilibrium isotherm of Cu(II) adsorption from solutions containing different chelating agents under optimal pH conditions. Table 3. Langmuir Constants for the Adsorption of Cu(II) on Chitosan under Each Optimal pH Range

Figure 6. Correlation between solution pH and Cu(II) adsorption capacity in the presence of gluconate.

Cu(II) chelates with protonated chitosan, or both plays an important role. As shown in Figures 3-6, the pH values where the adsorption capacity reaches a maximum are near the ranges of 3.1-4.2, 5.0-5.8, 5.0-6.0, and 5.2-5.8 with EDTA, citrate, tartarate, and gluconate, respectively. The dotted curves in these figures are the titration results. It is interestingly found that the optimal pH presents a slightly acidic front of the equivalent points on titration curves. This means that the aqueous-phase complexation of Cu(II) and chelating agent starts to proceed after optimal pH. That is, the coordination of Cu(II) with unprotonated chitosan dominates under more basic solutions. With regard to the optimal pH, two points are mentioned. One, the optimal pH is just larger than the pKa2 of corresponding acids of the chelating agent except for gluconate. The pKa2 values for EDTA, citric, and tartaric acids are 3.24, 4.76, and 4.37, respectively, which are calculated from the overall formation constants (Table 1). For example, the pKa2 for EDTA is the difference between the log Kf for H3L and the of log Kf for H2L. It can be easily understood that Cu(II) readily starts to form chelates with negatively divalent ligands. Two, in equimolar dilute solutions of Cu(II) and EDTA (10 mol/m3), it was found from mass-balance calculations that the divalent species CuL2- dominates at pH

system

KL (m3/mol)

qmon (mol/kg)

R

pH range

Cu(II) Cu(II) + tartarate Cu(II) + citrate Cu(II) + gluconate Cu(II) + EDTA

3.83 2.19 × 101 3.70 × 101 3.80 1.43 × 101

2.75 4.29 2.60 2.48 0.59

0.998 0.999 0.999 0.997 0.999

4.7-5.4 5.0-6.0 5.0-5.8 5.2-5.8 3.1-4.2

4-12.6 This again verifies our argument because the optimal pH for adsorption of Cu(II) with EDTA locates within 3.1-4.2. Adsorption Isotherm. Figure 7 shows the equilibrium isotherms of the chelated species in each optimal pH range. The adsorption capacity decreases in the order tartarate > citrate > gluconate > EDTA. Evidently, the relation between the amount of adsorption, qe (mol/kg), and the liquid-phase concentration at equilibrium, Ce (mol/m3), is described by the Langmuir equation

qe ) KLqmonCe/(1 + KLCe)

(5)

where KL is the Langmuir constant and qmon is the amount of adsorption corresponding to complete monolayer coverage. If the adsorption obeys the Langmuir equation, KL and qmon can be evaluated by the plot of Ce/qe vs Ce. The values of KL and qmon are listed in Table 3. The fit of the Langmuir equation is quite well because the correlation coefficients all exceed 0.997. The maximum capacities (i.e., qmon) in the coexistence of citrate and gluconate are 2.60 and 2.48 mol/kg, respectively. Both are comparable to that obtained for Cu(II) adsorption without chelating agent (2.75 mol/kg). Strong chelation between EDTA and Cu(II) largely decreases qmon to 0.59 mol/kg, probably because of the less adsorption effectiveness of the chelates with greater size and smaller charge than hydrated ions.29 On the other hand, the maximum capacity of Cu(II) with tartarate in a proper pH range is highly more than that of Cu(II) alone (4.29 mol/kg). No satisfactory explanation can be given at this stage. Chemical modification of chitosan with tartarate under such gentle conditions may be a possible reason. In fact, the authors have recently examined the exchange equilibria of metal-EDTA chelates with a

274 Ind. Eng. Chem. Res., Vol. 38, No. 1, 1999

tartarate, the third stage appears after 30 min of contact. It takes 60 and 120 min with citrate and gluconate, respectively, to reach this stage. Conclusions

Figure 8. Adsorption rates of Cu(II) on chitosan from solutions at pH 3 containing various chelating agents.

weak-base ion exchanger, Amberlite IRA-68.6 The equilibrium exchange of the Cu-EDTA chelates slightly increases only with increasing the solution pH in the range of 3.0-4.0. For equimolar dilute solutions, the exchange capacities of the Cu-EDTA-chelated anions are obtained as 0.40 and 0.43 mol/kg at pH 3.1 and 3.5, respectively. This is less than the present results under comparable conditions, indicating the superior capability and potential of chitosan for this purpose. Adsorption Rate: Intraparticle Diffusion Model. Intraparticle diffusion was extensively studied, and the model chosen here refers to the theory developed by Weber and Morris.30 The fractional approach to equilibrium varies according to a function of (Dt/r2)1/2, where D is the diffusivity in the solid and r is the particle radius. The initial rate of intraparticle diffusion is thus calculated by linearization of the curve qt ) f(t1/2). Figure 8 shows the results. It is seen that the rate of adsorption of the chelates decreases in the order tartarate > citrate > gluconate, the same as that obtained for adsorption capacity. Previous studies showed that such plots may present a multilinearity.31 This indicates that two or more phenomena occur successively. The first, sharper portion is the external surface adsorption or instantaneous adsorption stage. The second, gradually linear portion is the gradual adsorption stage, where the intraparticle diffusion is the controlling factor of the adsorption process.32 The third, linear portion is the final equilibrium stage where the intraparticle diffusion starts to slow because of extremely low solute concentration in the solution.33 This behavior was also due to diffusion in pores with different sizes.19,21 As shown in Figure 8, the instantaneous adsorption in the first stage is not so apparent. The second stage due to intraparticle diffusion being rate-limiting is immediately attained at the beginning of the adsorption process. When the bulk and surface concentrations of the chelates start to decrease, the third section of the figures is due to a decrease in the diffusion rate. A good correlation of rate data with the linearized form of the model indicates that diffusion mechanisms are controlling the rates. The slope of the second, linear stage is the rate parameter of intraparticle diffusion, which is 0.023, 0.014, and 0.006 mol/kg‚s1/2 for solutions containing tartarate, citrate, and gluconate, respectively. Correlation coefficients all exceed 0.996. In the case of

The capacities and rates of Cu(II) adsorption using chitosan from aqueous solutions containing EDTA, citrate, tartarate, and gluconate have been measured. In the absence of chelating agents, chitosan shows an excellent ability for Cu(II) adsorption with a capacity of 2.5 mol/kg. The capacity significantly drops when the solutions contain chelating agents except for gluconate, which has a comparatively small formation constant. The variations of solution pH bring about competition between coordination of Cu(II) with unprotonated chitosan and electrostatic interaction of Cu(II) chelates with protonated chitosan. The optimal pH ranges for Cu(II) adsorption are 3.1-4.2, 5.0-5.8, 5.0-6.0, and 5.2-5.8 with EDTA, citrate, tartarate, and gluconate, respectively. Such an optimal pH range locates at slightly acidic ahead of titration equivalent points for each chelated solution and is just larger than the pKa2 of corresponding acids of the chelating agents except for gluconate. All isotherms fit the Langmuir equation at each optimal pH, and the maximum capacities are 0.59, 2.60, 4.29, and 2.48 mol/kg, respectively. Finally, a plot of qt vs t1/2 presents a linear relation, indicating the ratecontrolling nature of intraparticle diffusion. The rate parameter is the largest with tartarate (0.023 mol/kg‚ s1/2) and the smallest with gluconate (0.006 mol/kg‚s1/2). This work provides useful data for adsorption removal of Cu(II) from chelating-agent-bearing streams using chitosan. Nomenclature Ce ) liquid-phase Cu(II) concentration at equilibrium (mol/ m3) Ct ) Cu(II) concentration in the aqueous phase at time t (mol/m3) C0 ) initial Cu(II) concentration in the aqueous phase (mol/ m3) D ) diffusivity in the solid (m2/s) Kf ) overall formation constant KL ) parameter of the Langmuir equation (m3/mol) qe ) amount of Cu(II) adsorbed at equilibrium (mol/kg) qmon ) amount of Cu(II) adsorbed corresponding to monolayer coverage (mol/kg) qt ) amount of Cu(II) adsorbed at time t (mol/kg) r ) radius of the particle (m) R ) correlation coefficient t ) time (s) V ) volume of the solution (m3) W ) amount of dry chitosan used (kg) Subscripts eq ) equilibrium 0 ) initial

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Received for review April 20, 1998 Revised manuscript received October 15, 1998 Accepted October 18, 1998 IE980242W