Biotransformation of Binary and Ternary Citric Acid Complexes of Iron

Although the biotransformations of several binary metal−citrate complexes .... of a 18-h-old culture of P. fluorescens grown to middle log-phase (op...
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Environ. Sci. Technol. 1997, 31, 3062-3067

Biotransformation of Binary and Ternary Citric Acid Complexes of Iron and Uranium CLEVELAND J. DODGE AND AROKIASAMY J. FRANCIS* Department of Applied Science, Brookhaven National Laboratory, Upton, New York 11973

Citric acid forms ternary mixed-metal complexes with various metal ions involving the hydroxyl and carboxyl groups of citric acid. The coordination of the metal to citric acid has been shown to affect the biodegradation of the metal-citrate complexes and metal mobility in the environment. We investigated the formation and biodegradation of the ternary mixed metal Fe-U-citric acid complex. The presence of 1:1:2 Fe-U-citric acid in solution was confirmed by potentiometric titration, UV-vis spectrophotometry, gel-filtration chromatography, and extended X-ray absorption fine structure (EXAFS) analysis. Comparison of the EXAFS spectra shows that the 1:1:2 Fe-U-citric acid complex has structural characteristics similar to the 1:1 U-citric acid complex. Biotransformation studies of Fe(III)-citrate, U(VI)-citrate, and Fe-U-citrate complexes by Pseudomonas fluorescens showed that the binary 1:1 Fecitric acid was readily biodegraded, whereas the 1:1 U-citric acid complex and the ternary 1:1:2 Fe-U-citric acid complexes were recalcitrant. Adding excess citric acid to the 1:1 U-citric acid complex resulted in the formation of a 2:3 U-citric acid complex after metabolism of the excess citric acid. When 1-fold excess citric acid was added to the 1:1:2 Fe-U-citric acid complex, the excess citric acid was completely degraded with no change in the stoichiometry of the complex. However, in the presence of 2-fold excess citric acid, a 1:1:1 Fe-U-citric acid complex remained in solution after the excess citric acid had been biodegraded.

Introduction Citric acid, a naturally occurring complexing agent, is used in the separation of actinides and in the extraction of toxic metals and radionuclides from wastes, sludges, sediments, and contaminated soils (1, 2). It forms soluble metal-citrate complexes with transition metals and actinides (3, 4). In the presence of more than one metal, a ternary complex is formed with citric acid due to the bonding of the metals with both the carboxyl and hydroxyl groups (5). The bioavailability of metals may be affected by the formation of mixed-metal complexes (6, 7). Several metals such as Cr with In(III) (5); Cd with Ni, Mn, or Zn (8); and U(VI) with Al, Cu, Fe(III), and In(III) (9, 10) were reported to form mixed-metal complexes with citric acid. Mixed-metal complexes with citric acid are of particular interest because of their unique chemical properties compared to binary forms. Mixed-metal-citrate complexes of Cd, Zn, or Ni with Cu are more stable than the respective * Corresponding author phone: (516) 344-4534; fax: (516) 3447303; e-mail: [email protected].

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binary forms due to entropic effects (11). The stability constant for the mixed-metal Fe(III)-Sn(II)-citrate complex is K > 1024 (12), and it shows spectral enhancement below 650 nm as compared to the binary Fe(III)-citrate complex (13). Formation of the ternary Cu-Ni-citrate2 complex was favored when citric acid was present as the tetra-ionized form (14). Although the biotransformations of several binary metal-citrate complexes have been investigated (15-19), little is known of the biodegradation of mixed-metal-citrate complexes. The persistence of mixed-metal-citrate complexes in wastes and contaminated environments may result in the mobilization of radionuclides and toxic metals. In this paper, we report the biotransformation of Fe-U-citric acid mixed-metal complex by Pseudomonas fluorescens.

Materials and Methods Culture Conditions. A Pseudomonas fluorescens (ATCC 55241) capable of degradation of several binary metal-citrate complexes was grown in a mineral salts medium containing the following ingredients (per liter): NH4Cl, 35.8 mg; CaCl2‚ 2H2O, 2.75 mg; MgCl2‚6H2O, 6.25 mg; PIPES buffer (disodium salt, Sigma Chemical Co., St. Louis MO), 1.47 mg; β-glycerophosphate, 1.74 mg; FeSO4‚7H2O, 1.49 mg; MnSO4‚H2O, 1.15 mg; CuCl2‚2H2O, 0.101 mg; Na2MoO4‚2H2O, 0.094 mg; ZnSO4‚H2O, 0.103 mg; CoCl2‚6H2O, 0.151 mg; and citric acid (anhydrous, Sigma Cell Culture Reagent, St. Louis, MO), 100 mg (15, 17). The ionic strength of the medium was adjusted to 0.1 M by adding 7.4 g KCl, and the pH was adjusted to 6.1 with KOH. The culture was incubated in the dark on a rotary shaker at 26 ( 1 °C. Preparation of Mixed-Metal-Citrate Complexes. Citric acid (anhydrous, Sigma) was prepared and standardized by potentiometric titration. Ferric nitrate [Fe(NO3)3‚9H2O], Mallinckrodt, St. Louis, MO, and uranyl nitrate [UO2(NO3)2‚ 6H2O], BDH Chemicals, Analar, Poole, England, were analyzed for Fe and U content by the o-phenanthroline (20) and bromoPADAP methods (21), respectively. Binary Complexes. Fe(III)-Citrate. A 1:1 Fe(III)-citric acid complex (6.5 mM) was prepared by mixing continuously equal molar amounts of ferric iron and citric acid solutions (16). The ionic strength of the complex was adjusted to 0.1 M with KCl, and the pH was adjusted to 6.1 with KOH. U(VI)-Citrate. Equimolar 1:1 U(VI)-citric acid complex (6.5 mM) was prepared in a similar fashion to the Fe(III)citrate complex. Ternary Complex. Fe(III)-U(VI)-Citrate. The mixedmetal-citrate complex was prepared by combining equimolar solutions of ferric nitrate and uranyl nitrate in a beaker. The mixed metal solution was then slowly added to citric acid to obtain a molar ratio of 1:1:2 Fe-U-citric acid at a final concentration of 6.5 mM. The ionic strength of the solution was adjusted to 0.1 M with KCl, and the pH was adjusted to 6.1. The complexes were exposed to minimal light during preparation and allowed to equilibrate for 24 h (22). The pH was readjusted to 6.1 before addition to the growth medium. Potentiometric Titration. The number of hydrogens released from citric acid as a result of complex formation with the metal was determined by potentiometric titration. A 0.52 mM solution of each of the following complexes was made up in the mineral salts medium without citric acid: (i) 1:1 Fe-citric acid, (ii) 1:1 U-citric acid, and (iii) 1:1:2 FeU-citric acid. The change in pH due to the incremental addition of 0.01 M NaOH was measured in duplicate samples at 26 ( 1 °C using a Futura II pH electrode (Beckman Instruments, Inc., Fullerton, CA).

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Gel Filtration Chromatography. The Fe-, U-, and FeU-citrate complexes were fractionated on Sephadex G-15 gel ( 11) (25). Titration of the 1:1 Fe-citric acid complex showed a gradually increasing inflection point with an increase in pH. The three hydrogens released from citric acid were titrated up to pH 3.2. Continued addition of the base up to 5 mM OH-/mM Fe corresponds to the titration of two hydrogen ions released by a two-step hydrolysis of ferric ion. The hydroxyl group does not appear to participate in metal bonding (16). Titration of the 1:1 U-citric acid complex showed two inflection points, the first corresponding to dissociation of the three carboxylic acid hydrogens of citric acid as a result of complex formation with uranium. The second inflection point, at 4.7 mM OH-/mM U at pH 7.5, is due to the formation of a polymeric species (4, 22). The participation of the hydroxyl group of citric acid in complex formation has been established (26). It was shown that the U-citrate complex is in equilibrium as the mononuclear [UO2cit]- (20%) and a [(UO2)2(cit)2]2- (80%) species (4). Titration of the mixed metal 1:1:2 Fe-U-citric acid complex showed a sharp inflection point at 4 mM OH-/mM metal, indicating complete neutralization of the three acid hydrogens of citric acid and deprotonation of the R-hydroxy hydrogen resulting from complexation with Fe (10). This inflection point is unique as compared to the inflection points for the Fe- and U-citrate complexes, indicating that the two complexes interact, thereby suggesting that a ternary complex is present. Continued titration of the mixed-metal complex above pH 6.5 gave similar results to those of the 1:1 U-citric acid complex due to polymerization reactions. Gel Filtration Chromatography. Figure 2 shows the elution profiles of citric acid, 1:1 Fe-citric acid, 1:1 U-citric acid, and 1:1:2 Fe-U-citric acid complexes on Sephadex G-15 gel. Citric acid and the 1:1 Fe-citric acid complex showed similar elution pattern. The 1:1 U-citric acid and the 1:1:2 Fe-U-citric acid complexes were eluted in separate fractions. All of the uncomplexed citric acid, the 1:1 Fe-citric acid

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FIGURE 3. Comparison of UV-vis absorption spectra for (A) uncomplexed citric acid, (B) 1:1 U-citric acid, (C) 1:1:2 Fe-Ucitric acid, and (D) 1:1 Fe-citric acid. Each metal was present at a concentration of 0.52 mM, I ) 0.1 M.

FIGURE 2. Elution profile of metal-citrate complexes (6.5 mM) by G-15 Sephadex gel. The complexes were eluted in the order: citric acid ) Fe-citrate > U-citrate > Fe-U-citrate.

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complex and the 1:1:2 Fe-U-citric acid complexes were completely recovered (100%), and about 85% of the 1:1 U-citric acid was accounted for. These results indicate the presence of a 1:1:2 Fe-U-citric acid complex. Uncomplexed ferric ion and uranyl ion were not eluted from the column. Spectrophotometric Measurements. Figure 3 shows the UV-vis absorption spectra of 0.52 mM citric acid, 1:1 Fecitric acid, 1:1 U-citric acid, and 1:1:2 Fe-U-citric acid complexes between 350 and 500 nm. The Fe-citrate complex exhibited a broad absorption band that sharply increased at lower wavelengths, with a small shoulder at 454 nm. The absorption spectrum of the U-citrate complex exhibited a peak maximum at 437 nm, a minimum at 392 nm, and fine structure on either side of the maximum at 426 and 448 nm. For the Fe-U-citrate complex, three shoulder peaks were observed at 427, 440, and 456 nm, and the absorbance spectrum was less than the sum of the absorbances of the binary complexes. These peaks had similar positions to the absorbances of the 1:1 U-citric acid complex, indicating that there was little change in uranium bonding to citric acid. The change in shape, the change in absorbance, and the shift in peak positions of the observed spectrum confirm the interaction of Fe and U, with the formation of a ternary complex with citric acid. Citric acid does not absorb in this region. Adding excess citric acid to the ternary complex shifted the spectral position to lower energy with loss of fine structure (data not shown). Extended X-ray Absorption Fine Structure (EXAFS). Figure 4A shows the Fourier transform at the Fe K edge for the aqueous complexes 1:1 Fe-citric acid and 1:1:2 Fe-Ucitric acid at pH 6.1 over the k range 3.5-13.3 Å-1. Comparison of these spectra show slight phase and amplitude changes up to 4.0 Å, indicating that both complexes have similar near-range structure. The presence of the first shell oxygen atoms at 1.6 Å as well as coordination shells resulting from scattering from the carbon and oxygen atoms of citric acid at 2-4 Å are clearly shown. Between 4.0 and 5.6 Å, a difference in peak intensity but not in shape or phase was observed. Beyond 5.6 Å, a significant shift in phase was noted, indicating a difference in atomic structure or atom type between the two complexes. Figure 4B compares the Fourier transforms for the aqueous complexes of 1:1 U-citric acid and 1:1:2 Fe-U-citric acid at pH 6.1 over the k range of 3.7-12.1 Å-1 at the U LIII edge. The first coordination shell at 1.4 Å corresponds to the two

FIGURE 5. Biodegradation of 0.52 mM of uncomplexed citric acid, 1:1 Fe-citric acid, 1:1 U-citric acid, and 1:1:2 Fe-U-citric acid complexes.

FIGURE 4. Comparison of Fourier transforms of backgroundsubtracted and k3-weighted EXAFS spectra of (A) 10 mM 1:1 Fecitric acid and 1:1:2 Fe-U-citric acid complexes; (B) 10 mM 1:1 U-citric acid and 1:1:2 Fe-U-citric acid complexes at pH 6.1. collinear axial oxygen ligands of the uranyl ion, and the second shell at 1.9 Å corresponds to scattering from the equatorial oxygen atoms surrounding the uranium. The similarity of the pseudo-radial distribution functions (PRDF) for the two complexes up to 5.8 Å as a result of scattering from carbon and oxygen atoms suggests a similar structure for both complexes. Beyond 5.8 Å, there is a phase shift to greater distance (approximately 0.1-0.2 Å) for the 1:1 U-citric acid complex as compared to the 1:1:2 Fe-U-citric acid complex. Similar results for the solid phase of the complexes confirmed this shift. The resolution of the PRDFs to near-baseline levels indicates real interaction between the atoms and is not due to noise in the data. Additional data analysis and further refinement are in progress to obtain structural information for each complex. Biodegradation of Binary and Ternary Complexes. Figure 5 shows the biodegradation of citric acid, 1:1 Fe-citric acid, 1:1 U-citric acid, and 1:1:2 Fe-U-citric acid complexes. Uncomplexed citric acid was completely metabolized by P. fluorescens in less than 20 h at a rate of 40 µmol h-1, while the pH of the medium increased from 6.1 ( 0.0 to 7.9 ( 0.1. The citric acid was converted to carbon dioxide and water, with no intermediate products detected. The 1:1 Fe-citric acid complex was metabolized in about 60 h at a rate of 8.0 µmol h-1, and the pH of the medium increased to 7.3 ( 0.1 from 6.1 ( 0.1. Ferric iron precipitated from solution as ferric hydroxide after the citric acid was degraded (27). The 1:1 U-citric acid complex was not degraded by the bacterium during 144 h of incubation, and the pH decreased to 5.9 ( 0.1 from 6.1 ( 0.0 as a result of equilibrium of the complex. However, with the 1:1:2 Fe-U-citric acid complex, there was a lag up to 43 h after which 0.07 ( 0.01 mM (7%) of the citric acid was degraded at a rate of 3.5 µmol h-1 up to 60 h; then

there was no further degradation. The pH of the medium increased slightly to 6.3 ( 0.0 from 6.2 ( 0.1, and uranium and iron remained in solution. Effect of Excess Citric Acid on Biodegradation of Binary and Ternary Complexes. Adding 1-fold excess citric acid to 1:1 Fe-citric acid resulted in metabolism of all the citric acid (1.04 mM) at a rate of 25 µmol h-1, and the pH increased to 7.8 ( 0.1 from 6.1 ( 0.0. Although the 1:1 U-citric acid complex was not biodegraded, in the presence of 1-, 2-, and 3-fold excess citric acid, the citric acid remaining in each after biodegradation was 0.75, 0.80, and 0.83 mM, respectively (Figure 6A). The final stoichiometry of U-citric acid in all three treatments was approximately 2:3. The U remained in solution in all treatments. The biodegradation rates for citric acid in the presence of 1-, 2-, and 3-fold excess citric acid was 3.5, 16, and 28 µmol h-1, respectively. Figure 6B shows the effect of adding excess citric acid on the biodegradation of the 1:1:2 Fe-U-citric acid complex. A small amount (0.10 ( 0.01 mM, 10%) of the citric acid present in the 1:1:2 Fe-U-citric acid complex was degraded, giving a final stoichiometric ratio of 1:1:1.8. However, adding 1-fold excess citric acid (0.52 mM) to the complex (1:1:3 Fe-Ucitric acid) caused the degradation of 0.63 ( 0.02 mM (40%) of the citric acid at a rate of 26 µmol h-1. The pH of the medium increased from 6.2 ( 0.0 to 7.9 ( 0.0. The final stoichiometry of Fe-U-citric acid was 1:1:1.8, similar to the equimolar complex. The absorption spectra of the 1:1:3 FeU-citric acid complex before and after biodegradation (at 144 h) were similar, indicating that no detectable change in the complex had occurred (data not shown). Adding 2-fold excess citric acid (1.04 mM) resulted in the degradation of 1.56 ( 0.03 mM (75%) of the citric acid at a rate of 53 µmol h-1. The pH of the medium increased to 8.1 ( 0.1 from 6.3 ( 0.0, and the final stoichiometry of Fe-U-citric acid was 1:1:1. There was a lag period of 26 h in all excess citric acid treatments with little degradation after the first 50 h, similar to that of the 1:1 U-citric acid complex. At the end of the experiment, all of the iron (0.50 ( 0.01 mM) and uranium (0.51 ( 0.04 mM) remained in solution.

Discussion Characterization of several ternary mixed-metal-citrate complexes has been reported by a number of investigators using indirect methods (5, 10, 12, 28, 29). For example, UVvis spectrophotometry and polarographic and amperometric titrations of U(VI)-Al(III)-citrate2 suggested a structure

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FIGURE 6. Citric acid remaining during biodegradation of (A) 1:1 U-citric acid; (B) ternary 1:1:2 Fe-U-citric acid in the presence of excess citric acid.

FIGURE 7. Proposed structure for the mixed-metal 1:1:2 Fe-U-citric acid complex. involving U symetrically bonded to two citrates, with Al bonding to the carboxylate ligands closest to U (28). Infrared measurements of a bis(trilaurylamine) melt of U(VI)-In(III)citrate2 indicated that carboxyl groups participate in bonding as well as the hydroxyl group of citric acid (29). Calorimetric and potentiometric investigations of a mixture of Fe(III), U(VI), and citric acid showed the presence of a 1:1:2 Fe-U-citric acid complex involving the hydroxyl group of citric acid in metal chelation (10). The hydroxyl group was thought to become deprotonated during complex formation. Spectrophotometric analysis of 1:1:2 Cr(III)-In(III)-citrate complex showed di-µ-OR bonding between the metals (5). From the potentiometric analysis of 1:1:1 Sn(II)-Fe(III)-citrate and Sn(II)-Cu(II)-citrate complexes, Smith (12) proposed a novel structure involving the chelation of Sn to a carboxyl and hydroxyl group of citric acid and Fe or Cu to the two remaining carboxylic acid groups of citrate. We used a combination of analytical techniques to confirm the presence of the ternary 1:1:2 Fe-U-citric acid complex. Spectrophotometric analysis established the existence of the ternary complex, potentiometric titration indicated that the hydroxyl group of citric acid is involved in Fe bonding, and gel-filtration chromatography revealed the elution profile and stoichiometry of the complex. Comparison of the EXAFS spectra of 1:1:2 Fe-U-citric acid complex with 1:1 Fe-citric acid at the Fe edge and 1:1 U-citric acid at the U edge show the U-Fe distance to be greater than 4 Å and most probably between 5.8 and 6.8 Å. The U is coordinated to the first citric acid through the carboxylate and hydroxyl group and to the second citric acid through two carboxylate groups; the Fe in the mixed metal complex is bound to citric acid through a carboxyl and hydroxyl group with chelation of Fe to U

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occurring through a µ-citrato bridge (Figure 7). Axial oxygens of U and water molecules bound to U and Fe are not shown (U is thought to have pentagonal coordination about the equatorial axis in solution, and Fe exhibits octahedral coordination). These results are similar to the dimeric structure proposed for the 1:1 U-citric acid complex (4). We previously showed that the lack of biodegradation of the 1:1 U-citric acid complex is due to the hydroxyl group participating in complex formation, rendering citric acid unavailable to the microorganism. It is due to lack of transport inside the cell and not to uranium toxicity or inhibition of cell function (15, 17). However, in the presence of 1-3-fold excess citric acid, biodegradation of the U-citric acid complex results in a final U-citric acid stoichiometry of 2:3. This is in agreement with the studies of Heitner and Bobtelsky (30), who reported the formation of a 2:3 U-citric acid complex between pH 7 and pH 9. The 1:1:2 Fe-U-citric acid complex showed ∼7% degradation, indicating the presence of free citric acid or Fecitrate complex that are readily biodegradable. Adding 1-fold excess citric acid to 1:1:2 Fe-U-citric acid complex resulted in the biodegradation of citric acid, giving a final stoichiometric ratio of 1:1:1.8 Fe-U-citric acid. The rate of citrate biodegradation (26 µmol h-1) was similar to that of the treatment containing 1-fold excess citric acid with 1:1 Fecitric acid (25 µmol h-1), indicating the presence of free citric acid and 1:1 Fe-citric acid (16). Adding 2-fold excess citric acid, however, increased the rate and extent of citric acid degradation, suggesting that uncomplexed citric acid was the major component undergoing biodegradation. After degradation of excess citric acid, Fe and U remained in solution, resulting in a 1:1:1 Fe-U-citric acid complex. Although the

structure of this complex was not determined, it may involve chelation similar to that proposed by Smith (12) for the 1:1:1 Sn(II)-Fe(III)-citrate and the Sn(II)-Cu(II)-citrate complexes. These results show that the potential exists for the formation of the Fe-U-citrate mixed-metal complex in the environment, especially when citric acid is used in the remediation of uranium and toxic metal-contaminated soils and wastes. Similar to the U-citrate complex, the Fe-Ucitrate complex is recalcitrant to biodegradation.

Acknowledgments We thank Lars Furenlid (NSLS) for helpful discussions on interpreting EXAFS data, Kirk Mantione for technical assistance, and F. J. Wobber, Program Manager, for continued support. This research was performed under the auspices of the Environmental Sciences Division’s Subsurface Science Program, Office of Health and Environmental Research, U.S. Department of Energy, under Contract DE-AC02-76CH00016.

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(10) Manzurola, E.; Apelblat, A.; Markovits, G.; Levy, O. J. Chem. Soc., Faraday Trans. 1 1989, 85, 373. (11) Daniele, P. G.; Ostacoli, G.; Zerbinati, O.; Sammartano, S.; De Robertis, A. Transition Met. Chem. 1988, 13, 87. (12) Smith, T. D. J. Chem. Soc. 1965, 2145. (13) Binder, B. Inorg. Chem. 1971, 10, 2146. (14) Amico, P.; Daniele, P. G.; Ostacoli, G.; Arena, G.; Rizzarelli, E.; Sammartano, S. Inorg. Chim. Acta 1980, 44, 2219. (15) Francis, A. J.; Dodge, C. J.; Gillow, J. B. Nature 1992, 356, 140. (16) Francis, A. J.; Dodge, C. J. Appl. Environ. Microbiol. 1993, 59, 109. (17) Joshi-Tope, G.; Francis, A. J. J. Bacteriol. 1995, 177, 1989. (18) Brynhildsen, L.; Rosswall, T. Appl. Environ. Microbiol. 1989, 55, 1375. (19) Madsen, E. L.; Alexander, M. Appl. Environ. Microbiol. 1985, 50, 342. (20) Standard methods for the examination of water and wastewater, 14th ed.; American Public Health Association (APHA): Washington, DC, 1975; p 208. (21) Johnson, D. A.; Florence, T. M. Anal. Chim. Acta 1971, 53, 73. (22) Dodge, C. J.; Francis, A. J. Environ. Sci. Technol. 1994, 28, 1300. (23) Dodge, C. J.; Francis, A. J.; Clayton, C. R. In Application of Synchrotron Radiation in Industrial, Chemical, and Material Science; Terminello, L. J., D’Amico, K. L., Shuh, D. K., Eds.; Plenum Publishing Co.: New York, 1996; pp 159-168. (24) Sayers, D. E.; Stern, E. A.; Lytle, F. Phys. Rev. Lett. 1971, 27, 1204. (25) Martell, A. E. Special Publication of the Chemical Society, No. 17; The Chemical Society: London, 1964; pp 477-481. (26) Nunes, M. T.; Gill, V. M. S. Inorg. Chim. Acta 1987, 129, 283. (27) Baes, C. F.; Mesmer, R. E. The hydrolysis of cations; John Wiley and Sons: New York, 1976; p 237. (28) Booman, G. L.; Holbrook, W. B. Anal Chem. 1959, 31, 10. (29) Markovits, G.; Klotz, P.; Newman, L. Inorg. Chem. 1972, 11, 2405. (30) Heitner, C.; Bobtelsky, M. Bull. Soc. Chim., Fr. 1954, 174, 356.

Received for review December 19, 1996. Revised manuscript received May 21, 1997. Accepted July 21, 1997.X ES961058+ X

Abstract published in Advance ACS Abstracts, September 15, 1997.

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