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J. Phys. Chem. B 2000, 104, 5298-5301
Evidence for Fenton Photoassisted Processes Mediated by Encapsulated Fe ions at Biocompatible pH Values J. Fernandez, M. R. Dhananjeyan, and J. Kiwi* EPFL, Institute of Physical Chemistry II, 1015 Lausanne, Switzerland
Y. Senuma and J. Hilborn EPFL Materials Department, Polymer Laboratory, 1015 Lausanne, Switzerland ReceiVed: December 13, 1999; In Final Form: March 23, 2000
Iron alginate gel beads were prepared starting from sodium alginate solutions. Fe complexed with carboxylate is active during the Fenton-enhanced decoloration/degradation of Orange II via the encapsulated Fe catalyst. By energy-dispersive X-ray microanalysis (EDX) it was shown that a significant amount of Fe was on the catalyst surface and only a smaller fraction inside the cross-linked Fe alginate. The Fe alginate beads were ∼2 mm in diameter containing highly dispersed Fe species with sizes of about 0.5 nm. The Fe alginate mediated decoloration of Orange II takes place in less than an hour in the presence of H2O2 under visible light irradiation, at pH values between 5 and 8. Repetitive decoloration of Orange II solutions was observed by addition of the azo dye at the beginning of each cycle.
Introduction The Fenton1 and Fenton photoassisted degradation of aromatic compounds2 is an area of growing interest from the mechanistic3 and applied point of view. But drawbacks in the use of Fenton reactions are (a) they are limited to the acidic pH range, and (b) the Fe ions remain in aqueous solutions. To avoid the last limitation, Fe ions have been exchanged recently on Nafion membranes4 which are resistant to the attack of the hydroxyl radicals generated in solution. The preparation of Fe alginate in aqueous solutions capable of degrading organic compounds at biocompatible pH values is explored in this study via Fecross-linked alginate encapsulated materials in the presence of H2O2. Alginic acid is a naturally occurring polysaccharide that crosslinks in the presence of di- and trivalent cations.5 The poly(1,4-β-D-mannuronic) acid and poly(R-D-glucuronic acid) are polyelectrolytes with charged carboxylate groups. In recent years, there has been a growing number of applications of Na and Ca alginate in pharmaceutic and food industry and also in the uptake of metal ions in industrial wastewaters. To test the decomposition of organic compounds at biocompatible pH values in photoassisted Fenton reactions, nonbiodegradable azo dye Orange II3,4 was selected in the present study. Experimental Section The preparation of the Fe alginate was carried out starting from alginic acid sodium salt (Sigma). A 3% (w/v) solution of sodium alginate in distilled water was prepared 1 h before use. Droplets of the sodium alginate solution were generated in a syringe with a gauge 21 needle size. Size uniformity was attained by slowly displacing the syringe to attain an interval of 30 s between droplets. The sodium alginate drops were collected in a separate bath containing FeCl3‚6H2O (0.05 M) and Pluronic L 68 (0.25 g/L) both of Fluka AG, Buchs. The
Figure 1. Adsorption isotherm of Fe3+ ions on alginic acid. n2s refers to the number of Fe3+ ions adsorbed per gram of alginic acid.
batch was stirred at 60 rpm for 24 h and the alginate beads formed in solution were filtered off and stored in deionized water. Elementary analysis of the alginate beads revealed an Fe concentration of 8.5 × 10-2 M and a Na concentration of 8.2 × 10-2 M. This indicates that alginate chains are crosslinked by Fe3+ ions and that each Fe3+ replaces about two Na+. During the preparation of Fe alginate, the Fe3+ ions exchange significantly with the Na+ ions of the functional carboxylic groups of the alginate strands. Irradiation of the solutions used for the experiments reported in Figures 1 and 2 was carried out on Pyrex cylindrical flasks (60 mL volume) transmitting light at λ > 290 nm. The solutions were irradiated with a Hanau Suntest solar simulator with 80 mW/cm2. The total organic carbon (TOC) was monitored via a Shimadzu 500 unit. Electron microscopy was carried out by means of a Philips 20 MS instrument with a resolution limit of 3 Å. Photoelectron spectroscopy (XPS) was carried out using a Leybold-Heraeus instrument referenced to the Mg KR1,2 line at 1253.6 eV. The binding energies of the iron oxide surface species were referenced to the Au 4f 7/2 level of 83.8 eV. The
10.1021/jp9943777 CCC: $19.00 © 2000 American Chemical Society Published on Web 05/11/2000
Fenton Photoassisted Processes
J. Phys. Chem. B, Vol. 104, No. 22, 2000 5299
Figure 3. Variation of the pH within a 24 h period during the decoloration/degradation of Orange II (0.1 mM) under Suntest simulated solar light in the presence of H2O2 (4.85 mM) and two beads of Fe alginate (20 mg dry weight). The initial pH of the solutions were adjusted with HCl (0.01 M) or NaOH (0.01 M).
Figure 2. Orange II decoloration with Fe alginate under light and in the dark as a function of time and solution pH. The solutions used were not buffered. For other experimental conditions see Figure caption and text.
quantitative evaluation of the experimental data was carried out with a Shirley type background.7 This correction was necessary due to the electrostatic charging of the particles during the measurements. Energy-dispersive X-ray microanalysis (EDX) was used to determine the location of the Fe clusters in the alginate beads. It allowed the determination of the relative abundance of Fe in the inside and also in the outer shell of the alginate beads. Results and Discussion Figure 1 shows the adsorption isotherm when different concentrations of Fe3+ (from FeCl3) were stirred continuously with solutions containing 1.5 g/L of alginic acid for 30 min. The concentration of the Fe ion at the different concentration was determined spectrophotometrically with a Hewlett-Packard 8452A diode array. The Fe(III) in solution was determined with thiocyanate6 subtracting the blank in each case. In Figure 1, the n2s (number of dye molecules adsorbed per gram of alginic acid) is shown to increase up to Ceq of 7 × 10-3 M attaining equilibrium at this concentration. The cross-linking by iron during the beads formation attains a maximum as shown by the plateau in Figure 1. Figure 2 shows results observed during the decoloration of Orange II by Fenton and photo-Fenton reactions in the presence of H2O2 under visible light irradiation at initial pH values of 5.6 and 7.8 mediated by Fe alginate beads. Our laboratory has
recently reported Fe immobilized on Nafion membranes4 having suitable kinetics for the decoloration/degradation of this azo dye but this abatement was only observed at pH < 5. Complete decoloration was observed in less than 1 h under Suntest irradiation. In the dark at pH 5.6 a modest reduction of the dye was observed indicating that intermediates once formed in the dark preclude further decoloration. This occurs due to the unfavorable recycling of Fe3+/Fe2+ taking place in the absence of light irradiation as suggested recently by Bauer.8 The results obtained in Figure 2 with Fe alginate at pH 7.8 are close to the time of degradation observed for Orange II (0.1 mM) in homogeneous media at pH values 2 nm for Fe particles.9 Energy-dispersive X-ray microanalysis (EDX) was carried out to clarify the nature of the darker particles observed at the border as well as inside the Fe alginate. On the basis of the C/O/Fe ratios observed, Fe species were identified in both cases. The inside darker discrete spots would correspond to the crosslinking points of the Fe alginate network. Surface complex formation involving Fe carboxylate seems to be the mechanism for trivalent ion cross-linking of the alginate network. The detailed nature of the Fe alginate intervening in the Fenton process has yet to be examined. The results reported hereby have significant environmental implications since they lead the way to Fenton pretreatment of toxic organic pollutants without the need to adjust subsequently the pH for further less costly biological degradation. Acknowledgment. We thank P. Albers, Degussa-Hu¨ls AG, Hanau, for his help with the electron microscopy.
J. Phys. Chem. B, Vol. 104, No. 22, 2000 5301 References and Notes (1) Walling, Ch. Acc. Chem Res. 1975, 6, 125. (2) Halmann, M. Photodegradation of Water Pollutants; CRC Press: Boca Raton, FL, 1996. (3) (a) de Laat, J.; Gallard, H. EnViron. Sci. Technol. 1999, 33, 2726. (b) Pignatello, J.; Liu, Di; Huston, P. EnViron. Sci. Technol. 1999, 33, 1832. (c) Bossmann, H.; Oliveros, E.; Go¨b, S.; Siegwart, S.; Dahlen, E.; Payawan, L.; Straub, M.; Worner, M.; Braun, A. M. J. Phys. Chem. 1998, 102, 5542. (d) Nadtochenko, V.; Kiwi, J. Inorg. Chem. 1998, 37, 5223. (e) Nadtochenko, V.; Kiwi, J. EnViron. Sci. Technol. 1998, 32, 3273. (4) Fernandez, J.; Bandara, J.; Lopez, A.; Buffat, Ph.; Kiwi, J. Langmuir 1999, 15, 185. (5) (a) Martinsen, A.; Skjak-Braek, G.; Smidsrod, O. Biotechnol. Bioeng. 1989, 33, 70. (b) Ku¨htreiber, W.; Lanza, R.; Chick, W. Cell Encapsulation, Technology and Therapeutics; Birkhauser: Boston, MA, 1998. (6) Kolthoff, M.; Sandell, E.; Meehan, E.; Bru¨ckenstein, S. QuantitatiVe Chemical Analysis; Macmillan: London, 1969. (7) Shirley, A. Phys. ReV. 1979, B5, 4709. (8) Rupert, G.; Bauer, R.; Gisler, G. J. Photochem. Photobiol. A Chem. 1993, 73, 75. (9) Anderson, J. R. The Structure of Metallic Catalysts; Academic Press: New York, 1975.
