Biodegradation of Nickel−Citrate and Modulation of Nickel Toxicity by

Environ. Sci. Technol. 1996, 30, 562-568

Biodegradation of Nickel-Citrate and Modulation of Nickel Toxicity by Iron AROKIASAMY J. FRANCIS,* GEETA A. JOSHI-TOPE, AND CLEVELAND J. DODGE Department of Applied Science, Brookhaven National Laboratory, Upton, New York 11973

Biodegradation of 1:1 nickel:citric acid by Pseudomonas fluorescens proceeded after a lag (∼17 h) at the rate of 11 ( 1 µmol h-1, with only partial mineralization of the complex. The incomplete degradation of the complex was not attributed to changes in its structure, but was due to the toxicity of the Ni released. Addition of 1:1 Ni:citric acid inhibited glucose metabolism by the bacterium. The toxicity of the released Ni was evident only when it attained a threshold concentration of >0.3 mM in the culture medium. Speciation calculations showed that Ni released after metabolism of the complex was present as Ni2+ ion and nickel carbonate. Addition of iron as a ferric hydroxide or 1:1 Fe:citric acid to 1:1 Ni:citric acid resulted in the complete metabolism of the Ni-citrate complex, with concurrent removal of the released Ni from solution by coprecipitation with iron.

Introduction Contamination of soil and water from the disposal of domestic, industrial, and radioactive wastes containing Ni is a major environmental concern (1-3). The primary sources of Ni contamination are from mining and smelting operations, electroplating, production of alloys and alkaline batteries, and combustion of fossil fuels (4, 5). Citric acid, a naturally occurring complexing agent, is used in electroplating baths for various Ni alloys (6), and the wastes generated from this process contain the soluble Ni-citrate complex. Mobilization of radionuclides and toxic metals from waste disposal sites due to the codisposal of naturally occurring and synthetic chelating agents has been reported (7, 8). Biodegradation of metal-organic complexes should result in the precipitation of the metal ions as insoluble hydroxides, oxides, or other salts and retard their migration in subsurface environments. Citric acid forms different types of complexes with metal ions, and the biodegradation of metal-citrate complexes depends on the type of complex formed (9). Previously we reported that Pseudomonas fluorescens completely mineralized Ca-, Fe(III)-, and Zn-citrate complexes, whereas * Corresponding author telephone: (516) 282-4534; fax: (516) 2822060.

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Ni-citrate was degraded only partially. Ni-citrate showed a lag (∼17 h) before the onset of biodegradation; this was attributed to an inducible transport factor involved in the uptake of the complex (10). The incomplete degradation of Ni-citrate may be due to either the toxicity of the nickel released to the bacterium after biodegradation or a change in the nature of the complex to a recalcitrant form during biodegradation. Although Ni is an essential trace element and a component of the enzymes urease, hydrogenase, methyl-S-coenzyme M reductase, and carbon monoxide dehydrogenase, at higher concentrations it is toxic to microorganisms (11-14). In this paper, we report that incomplete degradation of Ni-citrate was due to the toxicity of the Ni released from the biodegradation of the complex when its concentration in the medium reached the toxic threshold. Ni toxicity was ameliorated by coprecipitation of Ni by iron, which allowed the complete degradation of the complex.

Materials and Methods Culture Conditions. Pseudomonas fluorescens (ATCC No. 55241), a bacterium capable of metabolizing citric acid, was grown in a modified mineral salts medium to determine the biodegradation of the Ni-citrate complex. The medium was developed in accordance with equilibrium calculations, using a MINEQL computer program, so that the cation complex to be studied was the predominant species (15). The medium consisted of the following ingredients (per liter): NH4Cl, 35.80 mg; CaCl2‚2H2O, 2.75 mg; MgCl2‚6H2O, 6.25 mg; PIPES buffer (disodium salt, Sigma Chemical Co., St. Louis, MO), 1.47 mg; glycerol phosphate, 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; citric acid (anhydrous, Sigma Cell Culture Reagent, St. Louis, MO), 100 mg. The ionic strength of the medium was adjusted to 0.1 M with KCl, and the pH was adjusted to 6.1 with 1 M KOH. The culture was grown in the dark on a rotary shaker at 26 ( 1 °C. Preparation of Fe(III)- and Ni-Citrate Complexes. A stock solution of citric acid (13.0 mM) was prepared in deionized water, and an aliquot was standardized by titration with 0.1 M NaOH (Acculute). Stock solutions containing 13.0 mM Ni or Fe(III) were freshly prepared by dissolving Ni(NO3)2‚6H2O (Alfa Products, Danvers, MA) or Fe(NO3)3‚9H2O (Mallinckrodt, Paris, KY) in deionized water. A 1:1 metal:citric acid complex was prepared by slowly mixing equimolar amounts of citric acid (pH 6.1) and the respective metal solution in a sterile acid-washed beaker. The pH was adjusted to 6.1 with 1.0 M NaOH, and the ionic strength was adjusted to 0.1 M with KCl. The solution was stirred continuously and diluted with deionized water to obtain a stock solution containing 5.2 mM metal-citrate complex. Metal-citrate complexes were equilibrated for 24 h, and the pH was readjusted to 6.1. The complexes were filtered through a 0.22-µm-pore-size filter and aseptically added to the sterile culture medium to give a final concentration of 0.52 mM. Characterization of the Ni-Citrate Complex. Potentiometric Titration. The type of complex formed between Ni and citric acid was determined by potentiometric titration of the hydrogen ion content of the ligand in the

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presence and absence of Ni. The ionic strength was adjusted to 0.1 M by the addition of KCl. The change in pH of 0.52 mM citric acid and 0.52 mM 1:1 Ni:citric acid solutions upon incremental additions of 0.01 M NaOH was measured at 26 ( 1 °C by using a Futura II pH electrode (Beckman Instruments, Inc., Fullerton, CA). Spectrophotometric Analysis. Spectrophotometric analysis of the Ni-citrate complex was performed to determine changes in the nature of the complex as a function of pH. Equimolar and 1:2 Ni:citric acid complexes were prepared by using nickel perchlorate (Johnson Matthey, MA) and the pH was adjusted to 5.0, 6.0, 7.0, 8.0, 9.0, and 10.0. The ionic strength was maintained at 0.1 M with KCl, and the complexes were allowed to equilibrate in the dark for 24 h before analysis. Absorption spectra of the Ni-citrate complexes were obtained by scanning from 800 to 300 nm using an HP8452A diode array spectrophotometer in a 1-cm quartz cuvette. Speciation Calculations. Speciation of Ni in the mineral salts medium before and during the biodegradation of 1:1 and 1:2 Ni:citric acid was determined by using the MINTEQA2 computer modeling program (16). The following thermodynamic constants were used: log K ) 6.62 for [NiCit]-; log K ) 4.09 for [NiHCit]; log K ) 2.13 for [NiH2Cit]+; and log K ) 2.1 for the [NiCit2]2- (16, 17). Calculations included equilibration of the system with atmospheric carbon dioxide (pCO2 ) 3.5) at an ionic strength of 0.1 M. Biodegradation of 1:1 Ni:Citric Acid. The degradation of 1:1 Ni:citric acid was investigated under aerobic conditions. One milliliter of a late log phase bacterial culture (18 h old) grown in the defined mineral salts medium containing citric acid was inoculated into 200 mL of medium containing 0.52 mM citric acid or 0.52 mM Ni-citrate in 500-ml acidwashed Erlenmeyer flasks. Inoculated and uninoculated (control) samples, in triplicate, were incubated in the dark at 26 ( 1 °C. Aliquots were withdrawn periodically, filtered through a 0.22-µm-pore-size filter, and analyzed for citric acid, pH, and Ni. In addition, the number of bacteria in the samples at the end of the experiment was determined by epifluorescence microscopy (18). Biodegradation of 1:2 Ni:Citric Acid. The effect of adding 1-fold excess citric acid on the biodegradation of 1:1 Ni:citric acid was investigated. The treatments consisted of (i) 1.04 mM citric acid and (ii) 0.52 mM 1:1 Ni:citric acid plus 0.52 mM citric acid. Inoculated and uninoculated (control) samples, in triplicate, were incubated as described earlier, and aliquots withdrawn periodically were analyzed similarly. Nickel Toxicity. To determine the toxicity of Ni to the bacterium, 2.0 mL of a 24-h-old culture grown in the mineral salts medium containing 0.52 mM citric acid was used to inoculate 100 mL of the defined mineral salts medium containing 0.52 mM 1:1 Ni:citric acid plus 0, 0.1, 0.2, 0.3, 0.4, 0.5, or 1.0 mM Ni2+ added as NiCl2. Triplicate samples of each treatment were incubated in the dark on a rotary shaker at 26 ( 1 °C. Aliquots were withdrawn periodically, filtered, and analyzed for pH and citric acid. In addition, the effect of 1:1 Ni:citric acid on glucose metabolism by the bacterium was determined. The bacterium was grown in the defined mineral salts medium containing 0.52 mM glucose for 24 h, and 2 mL of this culture (early log phase culture) was used to inoculate 100 mL of the defined mineral salts medium containing 0.52 mM glucose in 250-mL Erlenmeyer flasks. After 23 h, 0.52 mM citric acid or 1:1 Ni:citric acid was added to the culture.

Inoculated and uninoculated (control) samples, in triplicate, were incubated in the dark on a rotary shaker at 26 ( 1 °C. Aliquots were withdrawn periodically, filtered, and analyzed for pH, glucose, and citric acid. Effect of Adding Ferric Iron on the Biodegradation of 1:1 Ni:Citric Acid. The degradation of 1:1 Ni:citric acid was investigated in the presence of ferric iron, added as ferric hydroxide or as a 1:1 Fe:citric acid complex. The treatments consisted of (i) 0.52 mM citric acid, (ii) 0.52 mM Fe(III)-citrate, (iii) 0.52 mM Ni-citrate, (iv) 0.52 mM Ni-citrate + 0.52 mM ferric hydroxide, and (v) 0.52 mM Ni-citrate + 0.52 mM Fe(III)-citrate. Ferric hydroxide was prepared by dissolving FeCl3‚6H2O (Mallinckrodt, Paris, KY) in deionized water, neutralizing it with sodium hydroxide, and washing the resulting precipitate several times with deionized water (2). The precipitate was refrigerated, aged for 6 months, and then resuspended in deionized water. The amount of iron added to the culture medium was determined by atomic absorption spectroscopy, after dissolving a known volume of the suspension in dilute hydrochloric acid. The filter-sterilized metal-citrate complexes were added aseptically to 200 mL of mineral salts medium in 500-mL acid-washed Erlenmeyer flasks. Inoculated and uninoculated (control) samples, in triplicate, were incubated in the dark on a rotary shaker at 26 ( 1 °C. Aliquots were removed periodically, filtered, and analyzed for pH and citric acid. Separate 1-mL aliquots were removed and centrifuged at 13000g for 5 min at room temperature, and the supernatant and the pellet were each acidified and analyzed for iron and Ni. Chemical Analysis. Citric acid was analyzed on an Aminex HPX-87H column (BioRad) using high-pressure liquid chromatography (Shimadzu) with a UV-vis detector at 210 nm. Glucose was analyzed with a RID-6A refractive index detector (Shimadzu). Iron was analyzed by atomic absorption spectroscopy or by the spectrophotometric method (19), and Ni was analyzed by atomic absorption spectroscopy or inductively coupled plasma mass spectrometry (ICP-MS).

Results Characterization of the Ni-Citrate Complex. Potentiometric Titration. The titration curve for uncomplexed citric acid exhibited a sharp inflection point at 3 mM OH-, indicating that the three acid hydrogens of citric acid were completely neutralized (Figure 1). As a result of similarities of their pK’s, the hydrogens are released in overlapping, indistinguishable steps. The alcoholic hydrogen was not titrated (pK > 11). The titration curve for 1:1 Ni:citric acid was displaced downward compared to that of citric acid due to replacement of the hydrogen ions of citric acid by Ni. An inflection point at 3 mM OH- at pH 6.5, similar to citric acid, was observed, indicating that it is a tribasic acid. Above pH 8, there was a broad inflection point due to the removal of a fourth proton from the acid, which suggests the participation of the hydroxyl group in the Ni-citrate complex. Spectrophotometric Analysis. Figure 2a shows the absorption spectra for the Ni2+ ion at pH 6.0 and for equimolar Ni:citric acid at pH 6.0, 9.0, and 10.0. Two peaks were observed: one at 394 nm with molar absoptivity () equal to 4.0 and another broad asymmetric peak at 720 nm; these reflect spin-allowed transitions that are characteristic of octahedral Ni2+ complexes (20). In the presence of 1:1

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FIGURE 1. Potentiometric titration of citric acid and 1:1 Ni:citric acid.

FIGURE 2. Spectrophotometric analysis of aqueous Ni2+ ion with (a) 1:1 Ni:citric acid and (b) 1:2 Ni:citric acid.

Ni:citric acid, there was a shift in absorbance of the first peak to 390 nm and a marked increase in the molar absorptivity to 10.7. A downfield shift in the asymmetrical

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absorption band to 650 nm was also observed. These changes in the absorption spectra indicate the interaction of the Ni ion with citric acid in forming a Ni-citrate complex. The absorption spectra of 1:1 Ni:citric acid at pH 5.0, 7.0, and 8.0 were identical to that at pH 6.0, indicating that there is no change in the nature of the complex in this pH range. However, at pH 9.0 and 10.0, the absorbance maximum shifted to 400 nm, with an increase in the molar absorptivity to 15.2 and 16.7, respectively. Also, a change in peak height and a shift in the asymmetric peak from 720 to 680 nm were observed, indicating changes in complex structure. The absorption spectra of 1:2 Ni:citric acid (Figure 2b) at pH 6.0 showed a marked shift from 394 to 386 nm compared to the Ni2+ ion, with an increase in molar absorptivity to 9.0. The asymmetric peak maximum shifted to 638 nm. As with the equimolar complex, spectra at pH 5.0, 7.0, and 8.0 were identical, but the molar absorptivity was lower. At pH 9.0 and 10.0, however, the molar absorptivity increased to 10.4 and 14.3, respectively, and was accompanied by a slight shift in the absorption maximum to 400 nm. The asymmetric peak underwent a shift to 656 nm at pH 9.0 and to 678 nm at pH 10.0. Biodegradation of 1:1 Ni:Citric Acid. Uncomplexed citric acid was degraded completely in 16 h at the rate of 70 ( 2 µmol h-1, whereas 1:1 Ni:citric acid was only partially degraded (70%) at a lower rate of 11 ( 1 µmol h-1 after an initial lag of 17 h (Figure 3a). The cell numbers increased from 3.2 x 102 to 6.4 × 105 and 4.3 × 105 cells ml-1 in the citric acid and 1:1 Ni:citric acid samples, respectively. During the biodegradation of citric acid and 1:1 Ni:citric acid, the pH of the culture medium increased from 6.1 to 7.8 and 7.5, respectively (Figure 3b). The concentration of Ni in the culture medium remained the same (0.52 ( 0.01 mM) throughout the experiment. The speciation of Ni during biodegradation of the 1:1 Ni:citric acid was determined by the MINTEQA2 program as a function of the increase in pH and the concomitant decrease in citric acid concentration (Figure 3c). Initially, 91% of the Ni (0.45 mM) was complexed with citric acid, but as it was degraded, the pH of the medium increased, and Ni was present predominantly as uncomplexed Ni2+ ion (0.34 mM) and, to a lesser extent, as the carbonate complex (0.06 mM). The total amount of Ni released after biodegradation of the 1:1 Ni:citric acid was 0.40 mM. Biodegradation of 1:2 Ni:Citric Acid. Compared to the equimolar complex, 80% of the 1:2 Ni:citric acid (0.52 mM Ni:1.04 mM citric acid) was degraded at the rate of 26 ( 2 µmol h-1 and there was a much shorter lag period (∼2 h) (Figure 4a). Analysis of the filtered culture supernatant at the end of the experiment showed that all of the Ni added to the medium (0.52 ( 0.01 mM) remained in solution. The number of bacteria increased from 3.2 × 102 to 1.4 × 106 and 9.0 × 105 cells mL-1 in the citric acid and 1:2 Ni:citric acid samples, respectively. The pH of the medium containing citric acid increased from 6.1 to 8.1, and that of the medium containing 1:2 Ni:citric acid increased to 7.8 due to biodegradation (Figure 4b). Speciation calculations at the beginning of biodegradation of 1:2 Ni:citric acid showed that 98% of the Ni (0.51 ( 0.02 mM) was complexed with citric acid. As citric acid was degraded and the pH of the medium increased, the concentration of Ni2+ and carbonate species increased. Biodegradation proceeded until 80% of the citric acid (0.83 ( 0.02 mM) was metabolized; this corresponded to 0.4 mM

FIGURE 3. Biodegradation of 1:1 Ni:citric acid: (a) citric acid degraded; (b) changes in pH; (c) Ni speciation using MINTEQA2 (error bars represent (1 standard error of the mean).

Ni (as 0.14 mM Ni2+ and 0.26 mM NiCO3) released in the medium (Figure 4c). The predominant Ni species was NiCO3 due to the higher final pH attained from the metabolism of excess citric acid. Nickel Toxicity. The effect of the addition of Ni2+on the biodegradation of 0.52 mM 1:1 Ni:citric acid in the defined mineral salts medium is shown in Figure 5a. Only 70% of 1:1 Ni:citric acid was biodegraded when no excess Ni was present, similar to Figure 3a. In the presence of 0.1 and 0.2 mM Ni2+, only 46% and 29% of the citric acid was degraded, respectively. Biodegradation of 1:1 Ni:citric acid was not observed in media containing 0.3, 0.4, 0.5, and 1.0 mM Ni2+. Speciation calculations of Ni after bacterial growth in medium containing 1:1 Ni:citric acid plus 0.1 mM Ni2+ showed 0.34 mM as Ni2+ and 0.28 mM as Ni-citrate, and in medium containing 1:1 Ni:citric acid plus 0.2 mM Ni2+, 0.33 mM was present as Ni2+ and 0.39 mM as Ni-citrate. These results show that the biodegradation of 1:1 Ni:citric acid was completely inhibited when the concentration of Ni2+ in solution exceeded 0.3 mM. Figure 5b shows the metabolism of glucose in the presence of 0.52 mM citric acid or 1:1 Ni:citric acid in the defined mineral salts medium. With citric acid present, glucose (0.52 mM) was metabolized at a rate of 57 ( 1 µmol h-1 in 36 h. The pH increased from 6.1 to 8.1. However,

FIGURE 4. Biodegradation of 1:2 Ni:citric acid: (a) citric acid degraded; (b) changes in pH; and (c) Ni speciation using MINTEQA2 (error bars represent (1 standard error of the mean).

with 1:1 Ni:citric acid, only 66% of the glucose (0.34 ( 0.02 mM) was metabolized in 40 h at the rate of 11 ( 1 µmol h-1. The 1:1 Ni:citric acid was only partially degraded (0.36 ( 0.03 mM) at the rate of 24 ( 1 µmol h-1 (data not shown). Speciation calculations showed that the concentration of Ni released in the medium was 0.34 mM, after which there was no further biodegradation of glucose or citric acid. Effect of Adding Ferric Iron on the Biodegradation of 1:1 Ni:Citric Acid. Biodegradation of 1:1 Ni:citric acid commenced after a 17-h lag at the rate of 9.6 ( 1.0 µmol h-1, with only partial mineralization (70%). The final concentration of Ni in solution was 0.52 ( 0.01 mM, which included both Ni complexed with citric acid (undegraded) and Ni released from metabolism of the complex (Figure 6a). The bacterium completely metabolized the 1:1 Fe:citric acid complex at the rate of 20 ( 2 µmol h-1 (Figure 6b), and all of the iron was precipitated from solution (Table 1). The pH of the medium increased from 6.1 to 7.2. The 1:1 Ni:citric acid was completely degraded in the presence of ferric hydroxide (Figure 6c). Citric acid was degraded slowly in the first 22 h (3.8 µmol h-1), in contrast to the lag seen with 1:1 Ni:citric acid, and then faster at a rate of 14.4 µmol h-1. The pH of the culture medium

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and 1:2 Ni:citric acid complexes was 7.5 and 7.8, respectively. Therefore, the incomplete degradation of the Ni-citrate complex cannot be attributed to changes in its structure caused by an increase in pH. The failure of the bacterium to completely degrade Ni-citrate was due to the toxicity of the Ni released in the culture medium. Several physicochemical factors, such as pH and the presence of other inorganic cations and anions, clay minerals, and chelating agents, affect the toxicity of Ni to microbes. The toxic thresholds of Ni depend on the types of organisms and growth media used (13, 24). Further, the chemical speciation of Ni affects bioavailability and toxicity. Our results show that Ni is not toxic when complexed with citric acid, but when it is released from the complex (in the ionic or carbonate forms) and reaches the toxic threshold (>0.3 mM), it inhibits the metabolism of citric acid and glucose. Addition of 1:1 Fe:citric acid completely removed Ni from solution, whereas addition of aged ferric hydroxide removed only one-third of the added Ni, but brought its final concentration below the toxic threshold. The removal of Ni from solution by ferric hydroxide and the iron released from the biodegradation of 1:1 Fe:citric acid is due to a coprecipitation process. In a previous study, we showed the removal of Ni from solution by ferric iron by coprecipitation (25). Although the nature of the coprecipitate was not determined, we propose the following mechanisms. The biodegradation of Ni- and Fe(III)-citrate can be represented as follows:

[NiC6H5O7]- + 9/2O2 f Ni2+ + 3CO2 + H2O + 3HCO3- (1) FIGURE 5. Effect of nickel on citrate and glucose metabolism: (a) Ni2+ ion plus 0.52 mM 1:1 Ni:citric acid; (b) 1:1 Ni:citric acid plus glucose (0.52 mM citric acid or 1:1 Ni:citric acid added at 23 h to a culture growing in mineral salts medium containing 0.52 mM glucose).

increased from 6.1 to 7.2. Finally, only 0.37 ( 0.01 mM Ni (70%) was detected in solution, while 0.15 ( 0.01 mM (30%) was associated with the iron precipitate (Table 1). In the uninoculated control sample, 25% (0.13 mM) of the iron was detected in solution after 80 h. Addition of 0.52 mM 1:1 Fe:citric acid to the medium containing 0.52 mM 1:1 Ni:citric acid facilitated the complete degradation of both complexes. Citric acid was degraded at the rate of 3.7 µmol h-1 up to 22 h, after which time the rate increased to 32.5 µmol h-1. The pH of the medium increased from 6.1 to 7.2. The concentration of Ni in the medium decreased as citric acid was degraded, and all of the Ni (0.50 ( 0.01 mM) precipitated from solution, along with iron (Figure 6d and Table 1).

Discussion Equimolar amounts of Ni and citric acid are present predominantly as a mononuclear bidentate [NiCit]- complex in the pH range 5-8 (17, 21). Above pH 8, the complex exists in a tridentate form involving the hydroxyl group of the citric acid, and above pH 9, it exists in a polymeric form [Ni4(OH)Cit3]5- (22, 23). A change in the nature of the 1:1 Ni:citric acid complex with pH was observed spectrophotometrically, and the formation of the tridentate complex above pH 8 was observed by potentiometric titration. The maximum pH recorded during the biodegradation of 1:1

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[Fe(OH)2C6H5O7]2- + 9/2O2 f Fe(OH)3 + 4CO2 + H2O + 2HCO3- (2) A combined equation for Ni-citrate plus Fe(III)-citrate can be written as

[NiC6H5O7]- + [Fe(OH)2C6H5O7]2- + 9O2 f tFesOsNi+ + 8CO2 + 3H2O + 4HCO3- (3) where tFe-O-Ni+ represents Ni coprecipitated with iron hydroxide. The biodegradation of Fe(III)-citrate begins before the biodegradation of Ni-citrate (10, 26). Thus, Ni is expected to be removed by coprecipitation involving the freshly formed ferric hydroxide. The rate of degradation of 1:1 Ni:citric acid plus 1:1 Fe:citric acid (32.5 µmol h-1) was similar to the sum of the rates for both complexes alone. Also, the amount of citric acid metabolized up to 22 h, until which time no degradation of 1:1 Ni:citric acid is expected to occur, correlates well with the amount of citric acid metabolized in the same time with 1:1 Fe:citric acid alone. The incomplete precipitation of Ni in the presence of ferric hydroxide may be due to changes in the numbers and types of sorption sites on the iron surface (27). Some Ni may be associated with the iron in the colloidal phase. In the presence of ferric hydroxide precipitate, the predominant sorption reactions (eq 4) include adsorption, absorption, and surface precipitation (28):

tFeOH0 + Ni2+ f tFesOsNi+ + H+ where tFeOH0 is the hydroxyl-bound iron surface.

(4)

FIGURE 6. Effect of iron on biodegradation of Ni-citrate: (a) 0.52 mM 1:1 Ni:citric acid; (b) 0.52 mM 1:1 Fe:citric acid; (c) 0.52 mM 1:1 Ni:citric acid plus 0.52 mM ferric hydroxide; and (d) 0.52 mM 1:1 Ni:citric acid plus 0.52 mM 1:1 Fe:citric acid (error bars represent (1 standard error of the mean). TABLE 1

Fate of the Metals after Biodegradation of the Metal-Citrate Complexes metal (mM) precipitatea

solution

a

complex

Ni

Ni-citrate Fe(III)-citrate Ni-citrate + Fe(III)-citrate Ni-citrate + ferric hydroxide

0.524 ( NAc 0.024 ( 0.001 0.372 ( 0.036 0.21b

Values normalized after dissolution of the precipitate.

b

Fe

Ni

0.11 ( 0.003 0.011 ( 0.001 0.053 ( 0.003

7) has been demonstrated (29). These results show that iron plays a major role in sequestering Ni by coprecipitation processes and, thus, modulates Ni bioavailability

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and toxicity during the bioremediation of such metalbearing waste streams.

Acknowledgments This research was performed under the auspices of the Environmental Sciences Division’s Subsurface Science Program, Office of Health and Environmental Research, Office of Energy Research, U.S. Department of Energy, under Contract No. DE-AC02-76CH00016. We thank Dr. F. J. Wobber, Program Manager, for continued support.

Literature Cited (1) Srivastav, R. K.; Gupta, S. K.; Nigam, K. D. P.; Vasudevan, P. Water Res. 1994, 28, 1631. (2) Benjamin, M. M.; Hayes, K. F.; Leckie, J. O. J. Water Pollut. Control Fed. 1982, 53, 1472. (3) DOE Basic Research Needs for Management and Disposal of DOE Wastes, DOE/ER-0492T, Office of Energy Research; U.S. Department of Energy: Washington, DC, 1991. (4) Richter, R. O.; Theis, T. L. In Nickel in the environment; Nriagu, J., Ed.; John Wiley and Sons, Inc.: New York, 1980; pp 189-202. (5) Schmidt, J. A.; Andren, A. W. In Nickel in the environment; Nriagu, J., Ed.; John Wiley and Sons, Inc.: New York, 1980; pp 203-218. (6) Brenner, A. In Electrodeposition of alloys: Principles and practice; Brenner, A., Ed.; Academic Press: New York, 1963; Vol. II, pp 239-314. (7) Killey, R. W. D.; McHugh, J. O.; Champ, D. R.; Cooper, E. L.; Young, J. L. Environ. Sci. Technol. 1984, 18, 148. (8) Means, J. L.; Crerar, D. A.; Duguid, J. O. Science 1978, 200, 1477. (9) Francis A. J.; Dodge, C. J.; Gillow, J. B. Nature 1992, 356, 140. (10) Joshi-Tope, G.; Francis. A. J. J. Bacteriol. 1995, 177, 1989. (11) Wackett, L. P.; Orme-Johnson, W. H.; Walsh, C. T. In Metal Ions and Bacteria; Beveridge, T. J., Doyle, R. J., Eds.; John Wiley and Sons, Inc.: New York, 1989; pp 166-206. (12) Avakyan, Z. A. In Microbiology; Smirnova, L. S., Ed.; PrenticeHall: Boston, 1974; Vol. 2, pp 3-28.

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Received for review May 9, 1995. Revised manuscript received August 24, 1995. Accepted September 8, 1995.X ES950312F X

Abstract published in Advance ACS Abstracts, December 1, 1995.