Removal of some divalent cations from water by membrane-filtration assisted with alginate

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Removal of some divalent cations from water by membrane-filtration assisted with alginate Nicolas Fatin-Rouge, Alexandra Dupont, Alain Vidonne, Je´rome Dejeu, Patrick Fievet, Alain Foissy Laboratoire de Chimie des Mate´riaux et Interfaces, Universite´ de Besanc- on, 25030 Besanc- on, France

art i cle info

A B S T R A C T

Article history:

The removal of divalent metal ions from hard waters or galvanic wastewater by polymer-

Received 12 September 2005

assisted membrane filtration using alginate was investigated. The ability of this natural

Received in revised form

polymer to form aggregates and gels in presence of metal ions was studied, in order to carry

4 January 2006

out metal removal by ultra or micro-filtration. Alginate titrations have shown the presence

Accepted 18 January 2006

of amine groups in addition to carboxylates onto the polymer backbone. The binding properties of alginate with divalent cations have been studied, showing an increasing

Keywords: Ultrafiltration complexation Polymer-assisted filtration Wastewater treatment Water softening

affinity for Ca2+ over Mg2+ as polymer concentration increases, and the relative affinity Pb2+XCu2+4Zn2+4Ni2+. The softening of hard natural waters was achieved successfully and easily, but needs an optimal alginate concentration
4  102 M. The alginate powder can be directly added to hard waters. Except for Ni2+, metal-removal was efficient. Polymer regeneration has shown that Cu2+-complexes are labiles.

Metal removal

& 2006 Elsevier Ltd. All rights reserved.

Alginate

1.

Introduction

Nowadays, water treatment is one of the main important fields of studies, due to the increase of the world population and industrial activities. Efforts are devoted to optimise the techniques for the prevention and control of pollutions through purification and recycling of wastewaters. Metal contamination of waters can be dramatic: Pb2+ causes poisoning and accumulates in bones. Toxic concentrations of Cu2+ or Zn2+ are low and Ni2+ ions are very toxic and carcinogen. Therefore, the presence of these metal ions commonly used in electroplating technologies, produces wastewaters with a significant risk of contamination for the natural environment. On the other hand, other divalent cations like Ca2+or Mg2+, which are daily needed to prevent cardiovascular diseases, may cause damages to heat-equipment or washing machines, because the low solubility of their Corresponding author. Tel.: +33 3 81 66 20 91; fax: +33 3 81 66 20 33.

E-mail address: [email protected] (N. Fatin-Rouge). 0043-1354/$ - see front matter & 2006 Elsevier Ltd. All rights reserved. doi:10.1016/j.watres.2006.01.026

carbonates. Moreover, in the latter case, they lead to an overconsumption of washing powders and to pollutions from polycarboxylates, surfactants that are used to stabilise colloidal metal carbonates produced during the water heating. A variety of separation processes can be used for the removal of metal ions from water: precipitation by hydroxydes, sulfides, sulfates or carbonates, electrochemical reduction, adsorption onto silica gel or inorganic colloids, ionexchange and filtration. In the present cases, [M2+] are in the range 104–103 M in galvanic wastewater and the solubility product of their hydroxides often make the precipitation process not sufficiently efficient. Moreover, this does not allow easy recovery of metal, because it produces large amounts of sludges (Kruithof and Kopper, 1989) which need to be treated further. Adsorption onto mineral surfaces is often strongly pH-dependent and ion exchange is long to

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mannuronate moiety

[M2+] PAF r Rf S t TOC UF Vp Vrw

microfiltration molecular weight

n

Glossary

AA Cp Cr G Jp Jw M MF Mw

atomic absorption photometry M2+-concentration in the permeate (mol L1) M2+-concentration in the retentate (mol L1) guluronate moiety permeate flux of contaminated water treated by PAF (L m2 h1) permeate flux of pure water (L m2 h1)

operate and not adapted to the treatment of large volumes. Separation of small ions from solutions has been achieved by using electro-reverse osmosis or nanofiltration. But, these need high pressures (e.g. energy consumption) while having small fluxes. In order to increase the fluxes and decrease the transmembrane pressure, a modified use of ultra-filtration (UF) or micro-filtration (MF) has appeared some years ago, called polymer-assisted filtration (PAF) (Spirakov et al., 1996). Removal of ions is obtained by introducing a binding polymer in a large reservoir that contains the wastewater. The polymer is too large to cross the membrane and it reduces free metal ion concentrations. Upon filtration, metal ions are removed from wastewater, while the retentate is enriched in metal ions. In that process, it is necessary that (i) the polymer present a good chemical affinity for metal ions, (ii) complexes are formed rapidly and (iii) they are labile in order to regenerate the polymer easily. In typical de-pollution plants dealing with filtration, two independent units work alternatively in metal removal and in polymer regeneration modes. Different kinds of polymers have been studied for this purpose, depending on the hardness of metal ions to remove. For soft metal ions, polyimines were used (Barron-Zambrano et al., 2002; Molinari et al., 2004), while for hard metals, mainly polyacrylic acids and parent compounds were tested (Bodzek et al., 1999; Zhang and Xu, 2003; Rivas et al., 2004; Canizares et al., 2004). But, some drawbacks are their toxicity and their low solubility. Natural polymers such as Chitosan have displayed interesting properties (Rhazi et al., 2002), but their solubility is low and they are more adapted to soft metal ions because binding groups are amines. In this work, we have used alginates, a natural abundant and not toxic polysaccharide bearing carboxylate groups, extracted from Brown algae (Mc Hugh et al., 2001). It is a copolymer made of (1–4) linked a-L-guluronate (G) and b-D-mannuronate (M) randomly arranged. Main applications of this polymer are gums or thickeners in food industry, or biomedical applications of its calcium gels. It is able to form gels with alkalineearth ions heavier than Mg2+ and many other multivalent cations. Previous studies have shown the interest of alginates for the recovery of metal ions from liquid effluents (Apel and Torma, 1993; Aderhold et al., 1996; Romero-Gonzalez et al., 2001; Hyun and Myeong, 2004). This investigation addresses the removal of metal ions from galvanic wastewater and alkaline-earth cations from hard

divalent metal ion concentration (mol L1) polymer assisted filtration ion radius (pm) retention factor effective area of the membrane (m2) time (h) total organic carbon analysis Ultrafiltration volume of permeate (L) relative volume of solution extracted from M2+alginate hydrogels after sieving through a napkin shaking frequency of the retentate solution (s1)

waters by means of PAF. In a preliminary section, we analyse the acid–base properties of Na-alginate and we measure the complexation equilibria of Na-alginate with Ca2+ in order to set a reference using a common divalent cation. In the second part, we analyse the binding and the removal of the metal ions. Finally, we deal with the regeneration of the polymer.

2.

Experimental section

2.1.

Reagents and solutions

All the solutions used for the study of metal alginate interactions were prepared from fresh Ultrapure MilliQ water. Solutions used for filtration were prepared from de-ionised water or natural water. Sodium alginate was obtained from Fluka Biochemika. Moisture (11%) was measured by a halogen moisture analyser HR73 Mettler Toledo and the alginate concentration in solutions was corrected for the water content. The ratio M:G obtained from Fluka is 1:1.5 and the alginates range in molecular weight from 12 to 80 kDa. Major cations found in alginate were Na+, Ca2+, Mg2+, Sr2+, Ba2+ and Al3+ at concentrations 4.769374  104; 8.105  10277  105; 1.98  10271  104; 4.01  10479  106; 3.43  10576  106 and 3.260  10371  106 mol/Kg of dry alginate, respectively. Other reagents were of analytical grade. Free Ca2+ and Mg2+ were titrated using EDTA with calcein (Merck) at pH 12.3 and with T Black Eriochrom at pH 9.5 in ammoniacal buffer, respectively (Charlot, 1966). To maintain pH at 5.0, 7.0 and 9.0 for interaction measurements, HCl, 3-(N-Morpholino)propanesulfonic acid (MOPS) sodium salt (Yu et al., 1997) and H3BO3 were used, respectively. For H+-titrations, KOH was standardised with potassium hydrogen phthalate. In the following, the alginate concentration is given as in terms of the equivalent monomer concentration (guluronate or mannuronate sodium salt, Mw ¼ 198:1 g mol1 ).

2.2.

H+-titration of alginate

H+-Potentiometry and conductimetry were performed in 60 mL of diluted Na-alginate solutions (1.78  103 M) with HCl (solution 0.547 M) in a thermostated cell at 20 1C. The glass electrode was a Radiometer PHC 3011-9 connected to a Meterlab PHM210. The conductivity was measured with an

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Inolab 2P apparatus using a WTW TetraCon 325 cell. The solution was initially set in alkaline media with 5 mL of a 0.115 M KOH solution, 15 min prior to the start of the titration. In order to calculate conductivity differences a reference solution, without alginate, was prepared in similar conditions.

M2+ binding titrations

The calcium binding by alginate was followed by potentiometric titrations of free (unbound) calcium ions using a selective electrode (ELIT Nico 2000). The electrode was calibrated with increasing aliquots of Ca(NO3)2 at the same pH and ionic strength (I ¼ 1:0 KCl) than the alginate solution. Solutions were buffered at either pH 5.0, 7.0 or 9.0. The binding ratio Ca2+/alginate was calculated with respect to the equivalent concentration of monomer units. Turbidimetry measurements were carried out at 600 nm with a Metrohm 662 photometer in 100 mL buffered alginate solutions at 20 1C, in order to follow M2+-complexation and formation of aggregates (M¼Ca, Pb, Cu, Zn).

Filtration

Filtration experiments were carried out with an ORELIS module, using a 15 kDa molecular weight cut-off Carbosep inorganic membrane from Orelis Novasep. The membrane was 0.4 m long, with a 0.01 m external diameter and a 0.008 m2 effective active area. The trans-membrane pressure was varied between 0 and 4 bars, but most measurements were made at 2 bars. The temperature was in the range 20–24 1C. pH adjustments were made with diluted solutions of either HNO3 or NaOH. The permeate flux was calculated from Jp ¼

Vp ðLÞ , Sðm2 ÞtðhÞ

Cp , Cr

(2) 2+

where Cp and Cr are total M -concentrations in the permeate and in the retentate, respectively. For Pb2+, Cu2+, Zn2+ and Ni2+ ions, diluted solutions of these cations put together were added slowly under vigorous stirring to the alginate solution and pH was adjusted 30 min before filtrations started. Permeates and retentates were collected 30 min after the beginning of the filtration. Membrane performances were checked daily by measuring water fluxes.

2.5.

Results and discussion

3.1.

H+-titration of alginate

Potentiometric and conductimetric titration curves of Naalginate solutions were compared with a reference solution without polymer (Fig. S1, Supplementary Material). The curves look very close, but differences were significant and highly reproducible. Figure 1 presents the difference in conductivity between the alginate and the reference solutions vs. HCl addition, together with the pH of the alginate solution. The conductivity features a positive excess and a peak followed by a sudden decrease towards negative values. The quantitative analysis of the conductivity is detailed in supplementary material, but the trend may be simply explained as follows.

∆χ2,1/2

(1)

where Vp, t and S are the permeate volume, the time and the effective active area of the membrane layer, respectively. Retention factors (Rf) was defined as follows: Rf ¼ 1 

3.

Other measurements

Relative viscosity of sodium alginate and calcium alginate solutions were carried out with a falling sphere viscosimeter at 20.070.5 1C. Zinc, nickel, lead and cooper concentrations in retentates and permeates were measured with an atomic absorption (AA) spectrometer Varian Spectra 50B. Permeates and retentates were immediately acidified before analysis. Before filtrations, alginate solutions were circulated within

∆χ0

0.1 0.0

∆χ1,1/2 -0.1

12 11 10 pKA2 9 8 7 6 5 4 pK A1 3 2 pH

2.4.

the filtration module to remove the fraction smaller than the molecular weight cut-off. This fraction (1.5%) was measured by TOC with a Shimadzu TOC-5050. Alginate regeneration kinetics from Cu2+-complexes by H+ or EDTA were recorded with a Secoman S250 I spectrophotometer at 720 nm, on the d–d band of Cu-EDTA2. Cu2+-alginate complexes ([Cu2+]tot ¼ 5  103 M, [alginate]tot ¼ 2.0  102 M) were made in a large glass vessel under diluted conditions and fast stirring. 5–10 mL were taken off and immediately filtered off with Millipore filters (0.22 mm). For dissociation assisted by H+, identical amounts of solution containing EDTA in large excess and buffered at pH ¼ 7.0 with urotropine were added to the filtrate before spectrophotometric measurements.

∆χ (test-reference) / mS cm-1

2.3.

1305

4 0 (200 6) 130 3 – 130 9

-0.2 -0.3

∆χinf

-0.4 0.0

0.5

1.0

1.5 2.0 VH+ / ml

2.5

3.0

Fig. 1 – Plot of the excess of conductivity between the solution of Na-alginate and a reference-solution vs. added acid, (’). Plot of the pH for the solution of alginate vs. added acid, (K). Each pKA-value was obtained as the pH corresponding to the volume at the half-equilibrium obtained from the plot of the excess of conductivity. V0 ¼ 60 mL; [H+]tit ¼ 0.547 M; [alginate]0 ¼ 1.78  103 M; T ¼ 20.070.5 1C. The starting pH was reached by adding 5 mL of [NaOH] ¼ 0.115 M to 55 mL of solution before titrations. Standard deviations are about 0.01 pHunit for the pH-titration and 5 lS cm-1 for the excess of conductivity.

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In the first part of the titration, the conductivity of the alginate solution is higher due to an excess of Na+ and alginate ions originating from the polymer to titrate. The difference in conductivity remains constant until protonation of a Lewis base (NH2 groups) begins (pH
11). At this point, the strong increase of the difference in conductivities is due to an excess of OH ions in the alginate solution (see Eq. (3)). The decrease after the peak and the negative excess of conductivity result from the lower H+-concentration in the alginate solution due to the protonation of the carboxylate groups (Eq. (4)). B þ H2 O Ð BHþ þ OH , 

þ

A þ H Ð AH:

(3) (4)

The first and the second derivatives of the curve in Fig. 1 were calculated in order to find the pH at half-protonation of the NH2 and COO groups (i.e. apparent pKA values for RNH+3 / RNH2 and R0 CO2H/R0 CO 2 ); as shown in the Fig. 1 they read 9.0970.03 and 3.5770.02, respectively. These values agree well with those reported by Taillefert and Gaillard (1999) and Romero-Gonzalez et al. (2001) for alginate; it is also inside the range of pKA for carboxylic groups in mannuronic (3.38) and guluronic (3.65) acids (Haug, 1961). The presence of amine groups (4 6.5% w of nitrogen per alginate monomer, since acid hydrolysis of alginate carried out at 150 1C for a day was still not complete) was certified by nitrogen titration using the Kjeldhal technique (Shugar and Dean, 1989). Although it was not usually reported, there was indeed a significant amount of aminated groups in alginate salts obtained from Fluka and SKW Biosystems. Titration curves provided ways (see Supplementary Material) to determine the concentrations of the carboxylic and the amine groups which were (1.7870.05)  103 and (5.370.1)  104 M, respectively. These were compatible with the total concentration of the major counter ions (Na+, Ca2+, Mg2+, Sr2+, Ba2+ and Al3+, see Section 2.1) in the sample. Using the characterisation, the charge density of the alginate could be evaluated as a function of pH using ! 10pKA2 10pKA1 rn ¼ R 1  pK , (5)  pK pH A2 10 þ 10 10 A1 þ 10pH where R is the ratio of amine to alginate monomer, equal to 0.3. The isoelectric pH of alginate is 3.2 (Fig. S2, Supplementary Material).

3.2.

Interactions with M2+ ions (M¼Ca, Mg, Pb, Cu, Zn)

3.2.1.

Ca- and Mg-binding

The calcium binding with alginate was studied in the range pH 5–9, which is the usual range of natural waters. The free Ca2+ activity was determined with the selective electrode from calibration curves obtained in the same conditions (Fig. S3, Supplementary Material). Calcium ions present in Na-alginate were accounted for in the titration as added calcium salt. No calcium binding could be detected below the alginate concentration 103 M. For [alginate] ¼ 7.0  103 M, the fractions of bound monomer to Ca2+ as a function of Ca2+loading at different pHs are displayed in Fig. 2. These fractions converge towards 1 in excess of metal, showing

1


= Ca2+ / alginate bound

1306

0 -4.0

-3.5

-3.0 -2.5 log [Ca2+]added

-2.0

-1.5

Fig. 2 – Fraction of alginate bound to Ca2+ ions vs. log [Ca2+]added at pH ¼ 5.0 (K), 7.0 (Mops, .) and 9.0 (borate buffer, ’). V 0 ¼ 60 mL; [Ca2+]tit ¼ 0.33 M; [alginate]0 ¼ 7.0  103 M ; T ¼ 20:0  0:5 1C; I ¼ 1:0 M (KCl).

that a Ca2+ ion could be bound by a single monomer in excess of metal. Experimental results have shown a close behaviour of the alginate binding strength for pH ranging between 5 and 9. Small discrepancies can be explained by the inaccuracy of potentiometric measurements and the possibility that borate buffer interacts with Ca2+ ions. The binding of calcium with alginate caused precipitation of the complex. Turbidimetric Ca2+-titrations of alginate carried out at different polymer concentrations (Fig. S4, Supplementary Material) showed that they should be X 2.8  102 M to soften hard waters substantially. Moreover, formation of kinetics of Ca-alginate assemblies at different Ca2+:alginate ratios display (Fig. S5, Supplementary Material): (i) a fast binding of Ca2+ by alginate due to the fast waterexchange rate of the metal ion (Table S1, Supplementary Material); (ii) a release of Ca2+ from [Ca2+]free
5  105 to 103 M, which occurs on 15 min., from the re-organisation of the Ca-alginate complexes to form gels by cross-linking, and in some cases, followed by a small re-loading of Ca2+. Thus, the binding ability of alginate is maximal immediately after mixing and a fast removal of the Ca2+-polymer should to be made to minimise free Ca2+ in waters. In natural hard waters Ca2+ and Mg2+ co-exist roughly in the ratio 3:1. Calcium is more inconveniencing with regards to scale formation due to the low solubility of calcium carbonate (103 fold that of MgCO3), but Mg2+ was reported as an anti-gelling ion in alginate solutions (Larsen et al., 2003) which finally is helpful in filtration processes. Ca2+/Mg2+ competition for alginate was investigated (see Fig. S6, Supplementary Material). The unbound concentration of Ca2+ ions was slightly above that of Mg2+ ions at lower alginate concentration, but the reverse was seen above some 2  102 M alginate concentration, which was four-times the concentration of each M2+ ion. We attribute the different binding trends to the difference in hardness between the two ions (Sillen. and Martell, 1964): The harder Mg2+ binds more strongly carboxylate groups of alginate, but the larger Ca2+ more easily forms

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intermolecular complexes at higher polymer concentration (Fraser and Bickerstaff, 1997). Mg2+ inhibits the cross linking of alginates.

3.2.2.

Metal-induced aggregation of Na-alginate solutions

In view of the subsequent filtration and regeneration of the alginate solution, it is important to control the aggregationgelling phenomena upon addition of metal salts. Optimal filtration performances is usually obtained with small aggregates, since de-watering is easily achieved. For that purpose the precipitation of the polymer was investigated at pH 5 as a function of metal/alginate ratio for Pb2+, Cu2+ and Zn2+. For a 7  103 M Na-alginate solution, Fig. 3 shows that cations may precipitate the polymer. Therefore, the alginate concentration should be increased in order to make aggregates for [M2+] ¼ 104–103 M. The onset of aggregation ranges in the order Cu2+pPb2+oZn2+.

3.3.

Filtration

Measurements of permeate volume fluxes through the membrane indicated no effect of the sodium alginate concentration in the range (1.4–2.8)  102 M (not shown), despite an important change of the solution viscosity. Fluxes were 50% lower than for pure water. This can be explained by a partial filling of the membrane and the presence of a gel

100

T (%)

90

80

70 1E-3

layer against it. From the comparison of free cations concentrations in solutions and in permeate, it was observed that this gel layer improves the removal of M2+-ions, probably because the ligand concentration being higher in this layer, complexation equilibria were shifted toward complexes.

3.3.1.

0.1

1

10

nM2+ / nalign.

Fig. 3 – Transmittance of Na-alginate solutions vs. M2+loading (molar ratio between metal ions and alginate monomer units). M ¼ Zn(’, solid line), Cu (E, dashed line), Pb (m, dotted line). pH ¼ 5.0. [alginate]0 ¼ 7.0  103 M; T ¼ 20 1C.

Treatment of hard waters

Synthetic hard waters ([CaCl2] ¼ 1.5  102 M, [MgCl2] ¼ 5  103 M) were softened with three alginate concentrations in order to evaluate the best conditions for treatment. After mixing alkaline earth ions with alginate solutions, a fast sieve (
10 s) of alginate gels was made through a thin napkin before UF of the permeate. The goal of this procedure is to estimate Rf in MF and improve UF fluxes of permeate, while getting a good removal of Ca2+ and Mg2+ ions. Key parameters are reported in Table 1. This procedure in 2-steps allowed rapidity and efficiency. Rf after sieving showed that the softening can be efficiently carried out in MF. The larger was the alginate concentration, the more important was the permeate flux after sieving, but the relative volume of solution extracted (Vrw) decreased. The optimisation of the main treatment parameters i.e. {Rf(Ca2+), Vrw, Jp/Jw} was achieved for [alginate] ¼ 4.2  102 M. Using this alginate concentration, the treatment of an hard water was further tested in real conditions, by mixing a natural water having the above Ca2+ and Mg2+-concentrations with a concentrated Na-alginate solution or with Na-alginate powder under fast stirring. Previous UF results were confirmed, but adding alginate as a powder to the hard water to soften improved permeate fluxes by 25%, while keeping same Rf for Ca2+ and Mg2+. In the two cases, direct UF was made after homogenisation. In the former one sieving through the napkin would be needed, because of the co-existence of a gel phase with the solution. In the second one it was not needed, because many aggregates poorly cross-linked were formed.

3.3.2.

0.01

1307

4 0 (200 6) 130 3 – 130 9

Treatment of galvanic wastewater

Effects of metal ion concentrations on their removal efficiencies at two alginate concentrations are shown in Fig. 4. Pb2+, Cu2+, Zn2+ and Ni2+ were introduced in Na-alginate solutions and pH were adjusted at 5.0 and 7.0. Filtration experiments were carried out and the initial permeate concentration of each ion was measured as a function of its concentration in solution. Experiments showed an acceptable flux of the permeate in presence of metal salts (see Fig. S7, Supplementary Material), very close to that measured for pure

Table 1 – Relative volume of solution extracted (Vrw) after sieving through a thin napkin, relative permeate fluxes (Jp/Jw) in UF and retention factors (Rf) for Ca2+ and Mg2+ ions as a function of the initial Na-alginate concentration [alginate]/M 2

3.5  10 4.2  102 4.9  102

Vrw

Jp/Jw

Rf(Ca)napkin

Rf(Ca)UF

Rf(Mg)napkin

Rf(Mg)UF

0.90 0.65 0.42

0.22 0.39 0.46

0.84 0.66 0.80

0.87 0.99 0.99

0.47 — 0.46

— 0.96 0.95

[Ca2+]tot ¼ 1.5  102 M; [Mg2+]tot ¼ 5.0  103 M; DP ¼ 2 bars; u ¼ 40 Hz; pH ¼ 7.0; T ¼ 22 1C.

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0.10 1 Cp (mg / L)

Cp (mg / L)

0.08 0.06 0.04

0.1

0.01

0.02 1E-3

0.00 0

50

100

(a)

150 200 Cr (mg / L)

250

0

300

20

(b)

40 Cr (mg / L)

60

80

10 10 Cp (mg / L)

Cp (mg / L)

1 0.1

1

0.01 0.1

1E-3 0

(c)

20

40 Cr (mg / L)

60

80

0

20

(d)

40 Cr (mg / L)

60

Fig. 4 – M2+-concentrations in permeate as a function of their initial concentration in retentate at different pHs and Naalginate concentrations. Horizontal dotted lines represent the legal disposal limit for the ions; they are 0.05, 0.04, 0.3 and 0.5 mg/L for Pb2+, Cu2+, Zn2+ and Ni2+, respectively. (a) Pb2+, (b) Cu2+, (c) Zn2+, (d) Ni2+. [alginate]tot ¼ 1.4  102 M (’) and 2.8  102 M (m); pH ¼ 5.0 (solid lines), 7.0 (dashed lines); DP ¼ 2 bars; t ¼ 40 Hz; T ¼ 22 1C. Metal ions were added together into Na-alginate solutions.

Na-alginate solutions. The removal of metal ions was due to the combination of complexation in solution and to an additional retention crossing a binding gel layer formed at the membrane interface. For Pb2+ions, concentrations in the permeate were close to the legal disposal limit, whatever the concentration in retentate. Fluctuation of concentrations about the mean value were attributed to analytical uncertainties due to interferences with other ions in atomic absorption spectrometry. Cu2+ ions passed the legal threshold at pH 7 and lower alginate concentration. At pH 7 and for the larger sodium alginate concentration, Zn2+-concentrations in the permeate were below the legal limit until the retentate concentration is o55 mg/L. Finally, Ni2+ displayed the lower retention and this procedure is not suitable for its removal at concentrations X40 mg/L. Data in Fig. 4 show a retention sequence decreasing in the order Pb2+XCu2+4Zn2+4Ni2+, in agreement with other reports (Smidsrod and Draget, 1996; Chen et al., 2002), but slightly different from that established in turbidity measurements (Cu2+XPb2+4Zn2+ at pH 5). We attribute this to a size effect: the larger Pb2+ (r ¼ 120 pm) more easily forms intermolecular complexes at higher ligand concentration than Cu2+ (r ¼ 72 pm). Concerning Ni2+, the low retention (i.e. the weak binding strength with alginate) is correlated with its slow ligand exchange rate constant (see Table S1) which is 3–5 orders of magnitude lower than for Pb2+, Cu2+ and Zn2+. In addition, Ni2+ is a quite soft ion which prefers to bind with softer ligands than carboxylate.

3.4.

Alginate regeneration

The full cycle of treatment involves the metal removal by filtration and then the regeneration of the polymer. In order to evaluate the complete cycle, kinetic measurements of the Cu2+-alginate dissociation were carried out in excess of competitor, examining two routes: ligand competition with EDTA and cation competition using H+. In the two cases, metal ions cross easily the membrane while the polymer cannot. In addition, in the latter case the separation is easier because alginate flocculates. Cu2+ was selected because it displays one of the largest affinities for alginate and its dissociation kinetics can be easily monitored by UV-vis absorption spectrophotometry, by following the concentration of the Cu-EDTA2 complex on its d–d band. At pH 7, ligand exchange is completed in 14 min (Fig. S8, Supplementary Material), while cation exchange is complete within 2 min at pH0 ¼ 1.0 and the equilibrium is reached (75% Cu2+ released) in 20 min at pH0 ¼ 1.6.

4.

Conclusion

The use of alginate to soften hard waters and remove Pb2+, Cu2+, Zn2+and Ni2+ from galvanic wastewater using polymerassisted UF was successfully tested, except for Ni2+. Treatment in real conditions of natural hard waters, by adding directly Na-alginate powder to water and having an alginate

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4 0 (200 6) 130 3 – 130 9

concentration
4  102 M was very efficient in terms of permeate fluxes and retention. In that case, poorly crosslinked agregates were formed and alkaline earth concentrations were about 8  104 M after treatment. The softening can be efficiently carried out in MF too. Metal ions have shown the following affinity for alginate: Pb2+XCu2+4Zn2+4 Ni2+. Very good retentions were obtained for Pb2+, Cu2+ and Zn2+. Alginate regeneration can easily and rapidly be achieved in strong acid medium.

Appendix A.

Supplementary materials

Supplementary data associated with this article can be found in the online version at doi:10.1016/j.watres.2006.01.026 R E F E R E N C E S

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Further reading Emsley, J., 1991. The Elements. Oxford University Press, Oxford. Margerum, D.W., Cayley, G.R., Weatherburn, D.C., Pagenkopf, G.K., 1978. Kinetics and mechanisms of complex formation and ligand exchange. In: Martell, A.E. (Ed.), Coordination Chemistry. ACS Monograph 174, vol. 2. Americal Chemical Society, Washington, DC, pp. 1–220.