Interactions II. Manganite Dissolution Promoted by Desferrioxamine B

we measured the steady-state dissolution rates of manganite (γ-MnOOH) in the presence of desferrioxamine. B (DFOB), a common trihydroxamate sideropho...
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Environ. Sci. Technol. 2005, 39, 6045-6051

Siderophore-Manganese(III) Interactions II. Manganite Dissolution Promoted by Desferrioxamine B OWEN W. DUCKWORTH* AND GARRISON SPOSITO Division of Ecosystem Sciences, University of California, Berkeley, California 94720-3114

Recent laboratory and field studies suggest that Mn(III) forms persistent aqueous complexes with high-affinity ligands. Aqueous Mn(III) species thus may play a significant but largely unexplored role in biogeochemical processes. One formation mechanism for these species is the dissolution of Mn(III)-bearing minerals. To investigate this mechanism, we measured the steady-state dissolution rates of manganite (γ-MnOOH) in the presence of desferrioxamine B (DFOB), a common trihydroxamate siderophore. We find that DFOB dissolves manganite by both reductive and nonreductive reaction pathways. For pH > 6.5, a nonreductive ligand-promoted reaction is the dominant dissolution pathway, with a steady-state dissolution rate proportional to the surface concentration of DFOB. In the absence of reductants, the aqueous Mn(III)HDFOB+ complex resulting from dissolution is stable for at least several weeks at circumneutral to alkaline pH and at 25 °C. For pH < 6.5, Mn2+ is the dominant aqueous species resulting from manganite dissolution, implicating a reductive dissolution pathway. These results have important implications for the biogeochemical cycling of both manganese and siderophoressas well as Fe(III)sin natural waters and soils.

Introduction Manganese is the third most abundant transition metal in the Earth’s crust and is important to both biological and environmental processes (1). In Nature, Mn exists in three oxidation states, Mn(II), Mn(III), and Mn(IV), which span a wide range of solubility, Brønsted acidity, and redox reactivity (2). Aqueous Mn(II) can be oxidized to Mn(III) or Mn(IV), resulting in the formation of more than 30 different (hydr)oxide mineral phases (3). The dissolution and precipitation of these minerals impact the chemistry of natural waters by consuming or releasing electrons or protons (4, 5), sequestering metal ions by adsorption or coprecipitation (6, 7), and oxidizing organic matter (8-10). Manganese is also intimately involved in biological processes, acting as either a cofactor or the active metal site in enzyme systems (11) and as a terminal electron acceptor for microbes under suboxic conditions (2, 12). Although Mn(IV) oxides are typically formed by biologically mediated Mn(II) oxidation (13), Mn(III) hydroxides are * Corresponding author phone: 510-643-9951; e-mail: owend@ nature.berkeley.edu.

10.1021/es050276c CCC: $30.25 Published on Web 07/12/2005

 2005 American Chemical Society

typically the initial product of the abiotic oxygenation of aqueous Mn(II) (2, 14-16). Under most conditions, manganite (γ-MnOOH) is the most stable of these Mn(III) minerals (3). In suboxic solutions at circumneutral to alkaline pH, manganite is stable with respect to Mn(IV) oxides; at acidic pH, manganite disproportionates to form Mn(IV) oxides and Mn2+ (17). However, this latter process is kinetically slow, and manganite may persist in the environment even under conditions where it is thermodynamically unstable with respect to disproportionation (15). Consequently, manganite is present in diverse aquatic environments and may play an important role in the chemistry of the natural waters (18, 19). Because Mn(III) is stable in solution only at extremely low pH (20), or when complexed by high-affinity ligands (21-25), most studies of the dissolution of Mn(III) (hydr)oxides have focused on reductive pathways (8, 9, 26-28). Recent research has shown that ligands having a high affinity for Mn(III) may promote either nonreductive ligandpromoted dissolution (29) or reductive dissolution (29, 30). Siderophores, biologically produced organic compounds that strongly complex Fe(III) (31-33) and other hard metal ions (34, 35), have also been shown to form stable aqueous complexes with Mn(III) (36-39). In a previous report (40), we characterized the stable complex, Mn(III)HDFOB+, formed by Mn(III) and desferrioxamine B (DFOB), a common trihydroxamate siderophore, as a product of air-oxidation of aqueous Mn(II) in the presence of DFOB. This complex can also be formed as a product of the DFOB-promoted dissolution of Mn(III) oxides (36, 39). The nonreductive dissolution of Mn(III)-bearing minerals releases reactive Mn(III) complexes which may influence the biogeochemical cycling of manganese and other redox-active elements (12, 38, 41-43). Although the siderophorepromoted dissolution of Fe(III)-bearing mineral phases has been extensively studied (44-52), Lloyd (37) presents the only study of siderophore-promoted dissolution of manganese(III,IV) oxides known to the authors. The manganite experiments, which are relevant to the present work, were performed between pH 7 and 9.5. Only ligand-promoted dissolution was observed, and the reaction rate was insensitive to pH. This topic is of special interest because Mn(III)siderophore complexes have stability constants that are nearly equal to or even greater than those for Fe(III)-siderophore complexes (38, 40), suggesting that Mn(III) may compete with Fe(III) for binding by siderophores. In this paper, we quantify DFOB-promoted dissolution rates of manganite under chemical conditions that are reasonable for environmental and biological systems (53).

Materials and Methods Materials. The DFOB utilized in this study is the mesylate salt [(C25H46N5O8NH3+(CH3SO3)-] produced under the trade name Desferal. The sample was a gift from the Salutar Corporation. All solutions were made with deionized water with a resistivity of 18.3 MΩ-cm. Unless otherwise specified, all other chemicals were ACS reagent grade. Manganite was synthesized by following a standard literature procedure (30, 54). Solutions of 0.2 M NH4OH (GFS), 0.06 M MnSO4 (Aldrich), and 30% H2O2 (Fisher) were purged for 1 h with N2. The MnSO4 solution (1 L) was continuously stirred in a 2-L Erlenmeyer flask, and 20.4 mL of purged H2O2 solution was added. To this solution 300 mL of purged NH3OH was slowly added with vigorous stirring, resulting in the formation of a dark-brown precipitate. The suspension was VOL. 39, NO. 16, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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rapidly warmed to the boiling point, and then heated under reflux with constant stirring at 98 °C for 6 h. The suspension was then filtered while hot and washed with an additional 500 mL of hot water during filtration. The resulting solid was then cold-washed by suspension and sonication in deionized water for 15 min, followed by centrifugation for 15 min at 24,300 RCF. The cold-wash process was repeated five times. The washed solid was then lyophilized until dry, disaggregated with a mortar and pestle, and stored in a freezer until use. Sample particles were characterized as to morphology, phase, and specific surface area. The particles were imaged using a FEI Tecnai 12 120-KV transmission electron microscope operated at 100 keV. Micrographs showed needle-like particles with dimensions of 300-500 nm × 10-30 nm; no other morphologies were evident (Supporting Information). X-ray diffraction patterns collected with an X’Pert Pro X-ray diffractometer equipped with a Co X-ray source (PANalytical, Almelo, The Netherlands) confirmed that the particles are manganite, as compared to the diffraction patterns of known standards, with no other phases detected (data not shown). The specific surface area of the sample was 39.4 ( 0.1 m2 g-1, as determined by a 9-point Brunauer-Emmett-Teller (BET) nitrogen adsorption isotherm (Micrometrics ASAP 2010). The average manganese oxidation number of the sample was determined by the oxalate titration method (55). Briefly, a small mass (∼50-100 mg) of MnOOH was added to 15 mL of a standardized 0.1 N Na2C2O4 solution (Aldrich) and 5 mL of 5.4 M H2SO4 (Fisher). The suspension was heated and stirred until the solid was completely dissolved. The solution was then cooled and diluted to 50 mL with deionized water. A 20-mL aliquot of solution was then heated and titrated with a standardized 0.1 N KMnO4 solution (Fisher) until a faint pink color resulting from excess MnO4- was detected by eye in solution. The remainder of the diluted sample solution was retained for total aqueous manganese quantitation with a Thermo-Jarrell-Ash Iris inductively coupled plasma atomic emission spectrometer (ICP-AES) using emission lines at 257.6, 259.3, and 260.5 nm. From a triplicate analysis, the average manganese oxidation number was found to be 3.04 ( 0.05, the value expected for manganite. Complexes of DFOB with Fe(III) and Mn(III) were synthesized to determine their absorption spectra and BeerLambert Law extinction coefficients. As reported previously (40), a green 1:1 Mn(III)HDFOB+ complex was stable for several weeks under ambient laboratory conditions for 7 < pH < 11. The [Mn(III)DFOB+] was quantified by UV-visible absorption spectroscopy at 310 nm using an extinction coefficient of 310 ) 2060 ( 20 M-1 cm-1 (36, 40). An orange 1:1 Fe(III)HDFOB+ complex was reported by Monzyk and Crumbliss (56). To synthesize this complex, 0.05 M FeCl3 (Fisher, laboratory grade) was added in 100-µL aliquots to 100 mL of 1 mM DFOB, and the absorbance at 425 nm was measured after each addition. The extinction coefficient we measured (425 ) 2600 ( 100 M-1 cm-1) agrees well with published values (425 ) 2460-2600 M-1 cm-1) for trihydroxamate complexes of Fe(III) (56, 57). Dissolution Experiments. Continuous-flow stirred tank reactors (CFSTR, Cole-Parmer, Vernon Hills, IL) were utilized to study the steady-state dissolution of manganite in the presence of DFOB (45). The volume of each of the reactors was 78.0 mL. A 0.22-µm mixed cellulose ester filter (Millipore, Bedford, MA) was placed “upstream” of the reactor outlet to contain the solid. To control temperature at 25 °C, the reactors were immersed in a recirculating water bath. In some experiments, the reactors were covered with aluminum foil to exclude stray light. However, ambient laboratory light had no detectable effect on the dissolution rate. To begin an experiment, manganite (0.70 g) was loaded into the reactor and sonicated for 10 min in 0.1 M NaCl 6046

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(Fisher) without DFOB. The influent solution was pumped at a rate of 5.0-7.5 mL min-1 through the reactors using laboratory grade Tygon tubing and Masterflex long-life Tygon precision peristaltic pump tubing (Cole-Parmer), together with a Masterflex 7524-10 microprocessor drive pump. Influent solutions contained 0-1000 µM DFOB, 0.1 M NaCl, and 10 mM buffer with pH set with NaOH (Fisher) or HCl (Fisher). Depending on the pH desired, acetate (Ac) (Fisher), MES (Sigma, 99.5%), MOPS (GFS), or HEPES (Sigma, 99.5%) were chosen to serve as buffers because similar concentrations of these buffers have been utilized previously in iron (hydr)oxide dissolution experiments without reported interferences (45, 58). Experiments conducted without buffer yielded dissolution rates consistent with results for the buffered systems (data not shown). We define nonreductive DFOB-promoted dissolution as the process of dissolving a manganese mineral without a change in Mn oxidation state following the reaction

MnOOH(s) + H4DFOB+ f Mn(III)HDFOB+ + 2H2O (1) To determine the steady-state, nonreductive dissolution rate, effluent samples were collected every 30 min. The flow rate q (L h-1) was determined for each sample by measuring the effluent mass collected for an allotted time (ca. 1-2 min). The [Mn(III)HDFOB+] was then immediately quantitated photometrically at 310 nm. The steady-state (ss), nonreductive ligand RL (mol kg-1 h-1) was calculated as follows:

RL )

[Mn(III)HDFOB+]ss × q m

(2)

where m (kg) is the mass of manganite in the reactor and [Mn(III)HDFOB+] is the concentration (M) in the effluent after the reactor has achieved a steady-state output concentration. On the basis of aqueous manganese flux, the mass of manganite in the reactor changes by less that 6% in all experiments. We define reductive dissolution in the presence of DFOB as the process of dissolving a manganese mineral with a reduction in Mn oxidation number following the overall (unbalanced) reaction scheme

MnOOH(S) + H4DFOB+ f Mn2+ + Oxidized DFOB (3) To determine the reductive dissolution rate RR, the total aqueous manganese concentration [MnT] was determined by ICP-AES. Because free aqueous Mn(III) does not typically occur over the pH range of study (2, 5), we assume that all dissolved manganese, except that complexed by DFOB, is Mn(II). The steady-state reductive dissolution rate (mol kg-1 h-1) is thus

RR )

[Mn(II)]SS × q m

(4)

where [Mn(II)]ss ) [MnT]ss - [Mn(III)HDFOB+]ss. Adsorption Isotherms. The adsorption of DFOB to manganite was measured at pH ) 8.0 (45, 46). Manganite (0.20 g) and 10 mL of a solution containing 1 mM HEPES buffer and 0.1 M NaCl were added to a 30-mL amber highdensity polyethylene bottle. The suspension was sonicated in a water bath at 25 °C for 10 min. The sample was spiked with 1 mL of DFOB solution (0-30 mM) and agitated with a magnetic stirrer for 10 min. The equilibration time was chosen so as to avoid the influence of initial dissolution transients, as determined from batch adsorption kinetics experiments (data not shown). The suspension was then

TABLE 1. Experimental Conditions for and Results of Adsorption Experimentsa [DFOB]total (µM)

pH

[Mn(III)HDFOB+] (µM)

[DFOB]free (µM)

n (mmol kg-1)

460 684 883 1313 1810 2802

7.7 7.9 8.1 7.8 7.7 7.7

290 ( 30 454 ( 4 537 ( 2 780 ( 20 820 ( 40 1090 ( 40

90 ( 20 120 ( 20 195 ( 6 330 ( 20 670 ( 50 1300 ( 100

4.3 ( 0.9 5.9 ( 0.9 8.3 ( 0.4 11 ( 2 17.7 ( 0.2 22 ( 6

a Shared conditions: Error is estimated from the standard deviation of replicate samples. Conditions: 18.2 g L-1 manganite, 0.1 M NaCl, 1 mM HEPES buffer, equilibration time ) 10 min.

a regression analysis (Figure 1B) by using a linear form of the Langmuir equation (45)

n ) K(b - n) [DFOB]

FIGURE 1. Adsorption of DFOB to manganite. (A) Surface excess (mmol kg-1) versus free DFOB (µM). Conditions: 18.2 g L-1 manganite, 1 mM HEPES buffer, 0.1 M NaCl. (B) Surface excess divided by free DFOB (m3 kg-1) versus surface excess (mmol kg-1). The line is a parametrized Langmuir model. The correlation coefficient for the least-squares best-fit line is R2 ) 0.98. Error estimates on the parameters represent 95% confidence intervals. Conditions: 18.2 g L-1 manganite, 1 mM HEPES buffer, 0.1 M NaCl. extracted by syringe and filtered through a 0.22-µm polyethersulfone syringe filter (Millipore), with the first 3 mL of filtrate discarded. The drift in pH was less than 0.3 log units during the course of any experiment. Each experiment was performed in duplicate, with two controls that contained identical solutions without manganite. The filtered solution was diluted (if necessary), and [Mn(III)HDOB+] in the filtrate was quantified at 310 nm. To analyze for [DFOB]total, a 5-mL aliquot of the filtrate was spiked with 1 mL of 5 mM FeCl3 solution. The mixing of acidic FeCl3 (pH ) 2.6) with the filtrate yielded a solution with a final pH ≈ 3. The addition of Fe3+ and the accompanying acidification causes the complete displacement of Mn3+ from the complex and the quantitative formation of Fe(III)HDFOB+. Although we do not know the exact mechanism for this displacement, analytical experiments (data not shown) conducted with solutions of known composition with known but varying [Mn(III)HDFOB+] and [DFOB]total yielded 97-103% recovery of the [DFOB]total when spiked with excess FeCl3. Furthermore, adsorption control experiments conducted without manganite yielded a recovery of 97-103% of the DFOB added (data not shown). The free DFOB concentration is defined as [DFOB]total from control experiments minus the experimental [Mn(III)HDFOB+]. Adsorbed DFOB (denoted by >MnDFOB) is defined as [DFOB]total from control experiments minus [DFOB]total from adsorption experiments.

Results and Discussion Adsorption of DFOB. Figure 1A shows an adsorption isotherm for DFOB on manganite at pH ) 8.0. The data (Table 1) show an L-type isotherm (59), which was parametrized by

(5)

where n (mmol kg-1) is the surface excess, K (mM-1) is the Langmuir binding constant, and b is the maximum surface excess [nmax (mmol kg-1)]. It is important to note that there is a significant concentration of Mn(III)HDFOB+ resulting from manganite dissolution; in our calculations, [DFOB] is the free concentration in solution and does not include aqueous Mn(III)HFDOB+. We find that nmax ) 32 ( 5 mmol kg-1 (0.8 ( 0.1 µmol m-2) and K ) 1.7 ( 0.3 mM-1 at I ) 0.1 M. On the basis of a surface site density of 7.9 µmol m-2, as measured by fluoride titration for a similarly produced manganite (60), we can estimate that the maximum surface coverage found corresponds to adsorption of DFOB on approximately 10% of all surface sites. Proton-Promoted Dissolution and Disproportionation of Manganite. Proton-promoted dissolution rates were nonnegligible for pH e 7.5 (Table 2). Because noncomplexed aqueous Mn(III) does not typically occur in this pH range (2, 5), we assume that all aqueous manganese is Mn(II), implying that dissolution occurs by a process that releases reduced manganese. [We do not consider buffer-promoted reductive dissolution (MOPS, MES, or Ac). Although this process cannot be completely ruled out, the buffers employed do not typically affect dissolution processes (45).] A possible mechanism for reductive dissolution is disproportionation of Mn(III) (29)

2γ - Mn(III)OOH + 2H+ T Mn2+ + MnO2 + 4H2O (6) This process has been implicated in the dissolution of manganite at pH < 6 (17, 19). We observe the onset of dissolution at pH ) 7.5, slightly more alkaline conditions than reported for similar experiments in 0.1 M NaNO3 (17) or 0.01 M NaCl (19). The slightly higher concentrations of [Mn(II)T] observed in our experiments in 0.1 M NaCl may be due to increased formation of Mn(II)-chloride complexes (14), increasing the net manganite solubility. Nonreductive DFOB-Promoted Dissolution of Manganite. Figure 2A shows data on the DFOB-promoted steadystate dissolution rate of manganite as a function of pH for various fixed [DFOB]. When [DFOB] ) 100 µM, the dissolution rate is approximately independent of pH for 6 < pH < 9, but it decreases slightly with pH for pH < 6. This decrease in the rate may be due to the oxidation of DFOB, thus lowering the effective aqueous concentration of the reactant (see the following section). For a given pH value, Figure 2A indicates that the steady-state dissolution rate increases with increasing [DFOB]. Our measured dissolution rates agree well with those reported by Lloyd (37) in the presence of 1 mM DFOB. VOL. 39, NO. 16, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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TABLE 2. Experimental Conditions for and Results of Manganite Dissolution Experimentsa [DFOB] (µM)

pH, buffer

q (mL min-1)

MnT (µM)

Mn(III)HDFOB+ (µM)

Mn(II) (µM)

[Mn(II)]eff/[DFOB]inf

RR (mol kg-1 h-1)

RL (mol kg-1 h-1)

0 0 0 0 100 100 100 100 100 100 100 100 100 100 100 100 100 250 250 250 250 500 500 500 500 1000 1000 (37) 1000 (37) 1000 (37) 1000 (37) 1000 (37)

5.7, Ac 6.5, MES 7.5, MOPS 8.4, HEPES 5.1, Ac 5.5, Ac 5.6, Ac 5.8, Ac 6.2, MES 6.4, MES 6.4, MES 6.6, none 7.2, MOPS 7.3, MOPS 8.0, HEPES 8.4, HEPES 8.5, HEPES 6.1, MES 6.2, MES 7.8, HEPES 7.9, HEPES 6.1, MES 6.2, MES 8.0, HEPES 8.1, HEPES 6.0, MES 7.0, TRIS 7.7, TRIS 8.3, TRIS 8.9, TRIS 9.2, TRIS

2.5 2.7 2.6 2.7 6.6 6.7 5.9 5.4 6.8 5.6 5.7 5.0 5.9 6.0 7.2 6.6 5.5 5.3 5.4 5.3 5.6 5.3 5.2 5.1 5.8 6.7 batch batch batch batch batch

50 43 3 0 432 334 381 374 116 144 146 83 100 82 45 61 78 223 277 164 160 447 425 210 194 492 N/A N/A N/A N/A N/A

0 0 0 0 31 36 33 44 62 61 67 70 74 65 54 55 66 136 153 131 147 255 257 212 205 324 N/A N/A N/A N/A N/A

50 43 3 0 401 298 348 330 54 83 79 13 26 17 0 6 12 86 124 33 13 192 178 0 0 168 N/A N/A N/A N/A N/A

No DFOB No DFOB No DFOB No DFOB 4.01 2.98 3.48 3.30 0.54 0.83 0.79 0.13 0.26 0.17 0.00 0.06 0.12 0.34 0.50 0.13 0.05 0.38 0.36 0.00 0.00 0.17 N/A N/A N/A N/A N/A

0.011 ( 0.002 0.010 ( 0.002 0.0006 ( 0.0004 0 0.22 ( 0.02 0.17 ( 0.04 0.18 ( 0.01 0.168 ( 0.007 0.029 ( 0.001 0.040 ( 0.001 0.040 ( 0.001 0.005 ( 0.001 0.012 ( 0.002 0.009 ( 0.002 0 0.003 ( 0.002 0.006 ( 0.002 0.04 ( 0.01 0.06 ( 0.02 0.008 ( 0.004 0.010 ( 0.09 0.09 ( 0.02 0.07 ( 0.02 0 0 0.10 ( 0.05 not reported not reported not reported not reported not reported

N/A N/A N/A N/A 0.017 ( 0.001 0.020 ( 0.001 0.017 ( 0.001 0.020 ( 0.001 0.036 ( 0.001 0.030 ( 0.001 0.033 ( 0.001 0.030( 0.001 0.037 ( 0.001 0.034 ( 0.001 0.034 ( 0.001 0.031 ( 0.001 0.030 ( 0.001 0.062 ( 0.004 0.070 ( 0.002 0.06 ( 0.01 0.07 ( 0.01 0.116 ( 0.005 0.113 ( 0.006 0.09 ( 0.01 0.101 ( 0.005 0.18 ( 0.02 0.135 0.139 0.135 0.150 0.126

a Error is estimated from the standard deviation of the rate calculated from the reactor output at steady-state. The net flow rate through the reactor is q (mL min-1). Conditions: 9.0 g L-1 manganite, 0.1 M NaCl, 10 mM buffer. N/A ) not applicable for batch experiments.

Figure 2B shows a log-log plot of the nonreductive steadystate dissolution rate versus the surface concentration of DFOB (mol m-2) on manganite measured at pH ) 8. For surface-controlled reactions, ligand-promoted dissolution rates often are related to the surface concentration of ligand [>M-L] by a power-law expression (61)

RL ) kL[>M - L]R

(7)

where kL is the ligand-promoted rate coefficient and R is the order of the reaction with respect to >M-L. In our analysis, we assume (1) that all DFOB adsorbed forms a single type of dissolution-reactive surface complex structure and (2) that there is no readsorption of Mn(III)HDFOB+ after dissolution. Under these assumptions, the surface excess of DFOB equals the concentration of the dissolution-reactive DFOB surface complex. A best-fit line regression for the data shown in Figure 2B yields

RL ) 108.6(0.3[>MnDFOB]0.9(0.3

(8)

where RL is expressed in units of mol m-2 s-1. Equation 8 indicates that, within experimental precision, the dissolution rate is directly proportional to the surface concentration of adsorbed DFOB; therefore, the dissolution rate varies as a function of DFOB concentration in a manner analogous to the DFOB-promoted dissolution of goethite (45), as well as the ligand-promoted dissolution of other metal (hydr)oxides (45, 62-64). The promotion of manganite dissolution by DFOB thus can be explained by the adsorption of ligands polarizing Mn-O bonds in the mineral structure and increasing the rate at which these bonds break (58, 65). Reductive DFOB-Promoted Dissolution of Manganite. Figure 3 shows the rate of reductive steady-state dissolution, 6048

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which is near the limit of detection for pH > 6.5 and increases monotonically with decreasing pH for pH < 6.5. This trend is consistent with those observed for the pH-dependence of reductive MnOOH dissolution by low-molecular mass organic acids (8, 26, 28). Reductive dissolution is the dominant dissolution reaction for pH < 6.5. The shift in mechanism to reductive dissolution at acidic pH can be rationalized by changes in the thermodynamic driving force. The Mn(III)HDFOB+ complex is unstable at pH < 7, decomposing by internal electron transfer to yield Mn2+ and oxidized DFOB (40). It is thus reasonable that, at pH < 7, Mn(III) centers in manganite may also participate in redox reactions with adsorbed DFOB, and reductive dissolution may be driven by a fundamental change in the thermodynamic driving force as pH drops into the acidic range. An additional complication is that the measured reductive dissolution rate is likely a composite of several interdependent elementary reactions. At pH > 5, [Mn(II)] in the effluent is 3- to 4-fold greater than the [DFOB] in the influent. This disparity suggests that the reaction stoichiometry between surface sites and DFOB is not 1:1. It is thus reasonable that, in addition to the reaction described in eq 3, fragments of oxidized DFOB react with the surface, yielding additional aqueous Mn2+. This latter reaction may be expressed by the overall (unbalanced) reaction scheme

Oxidized DFOB(z - 1) + MnOOH(S) f Mn2+ + Oxidized DFOB(z) (9) where z represents the zth time that DFOB fragments react with the manganite surface. In our experiments, this can be approximated to be z e 3 based on the ratio of [Mn(II)]eff to [DFOB]inf, as shown in Table 2. Because DFOB contains many

FIGURE 2. Nonreductive DFOB-promoted manganite dissolution rate. (A) Nonreductive dissolution rate as a function of pH (mol kg-1 h-1). Conditions: 9 g L-1 manganite, 0.1 M NaCl, 10 mM buffer. (B) Log10 of the nonreductive DFOB-promoted manganite dissolution rate (mol m-2 s-1) versus Log10 of adsorbed DFOB concentration (mol m-2) at pH ) 8.0 ( 0.3. The correlation coefficient for the best-fit regression line is R2 ) 0.90. The open points are from Lloyd (37). Error estimates on these coefficients represent 95% confidence intervals of the best-fit line.

FIGURE 3. Reductive DFOB-promoted manganite dissolution rate as a function of pH (mol kg-1 h-1). Conditions: 9 g L-1 manganite, 0.1 M NaCl, 10 mM buffer. possible sites for an oxidation reaction to occur (66, 67), oxidized fragments of DFOB may contain moieties that are susceptible to further oxidation, allowing them to react multiple times (z) with the MnOOH surface before a stable aqueous species is produced. It would be useful to determine the products of DFOB oxidation; however, because of the possibility of multiple reaction pathways, the overall oxidation of DFOB by manganite may result in many different products. Consequently, we have made no attempt to determine the identity of the DFOB oxidation products. Environmental Implications of Siderophore-Promoted Manganese (Hydr)oxide Dissolution. Figure 4A shows

FIGURE 4. Log10 of dissolution rates of Fe(III) and Mn(III) (hydr)oxides (mol m-2 s-1). (A) Nonreductive DFOB-promoted Mn(III) (closed markers) and Fe(III) (open markers) (hydr)oxides dissolution rates. γ-MnOOH dissolution rates: (b) 1000 µM DFOB, this study and ref 37; 500 (2), 250 (9), and 100 ([) µM DFOB, this study. Fe(III) (hydr)oxides: 1000 (0) and 100 (+) µM DFOB and r-Fe2O3, (37); 80 (×) µM DFOB and r-FeOOH (45); 100 (O) µM DFOB and r-FeOOH, (37); and 240 (2) µM DFOB and r-FeOOH, (46). (B) Reductive γ-MnOOH dissolution rates by siderophores (closed markers) and other organic compounds (open markers): 1000 (b), 500 (2), 250 (9), and 100 ([) µM DFOB, this study; 1000 (0) and 10 (]) µM ascorbate, (26); 1000 (4) µM oxalate, (28); 1000 (O) µM aminocarboxylate ligands (30). The dissolution rates of McArdell et al. (30) are estimated using a SA ) 30 m2 g-1. Error bars that intersect the x-axis indicate the dissolution rate is below the limit of detection. siderophore-promoted dissolution rates of Mn(III) (closed markers) and Fe(III) (open markers) (hydr)oxides. The siderophore-promoted MnOOH dissolution rates are >100 times greater than those typically measured for iron minerals. This observation, coupled with the high affinity of manganite for DFOB, suggests that DFOB-promoted iron and manganese mineral dissolution could be competing processes and that the presence of manganese (hydr)oxides may disrupt bacterial iron acquisition by siderophores (68). The large MnOOH dissolution rate also provides a source of reactivesand thus potentially environmentally importantsaqueous Mn(III) complexes (40, 69, 70). Figure 4B shows a compilation of reductive dissolution rates of MnOOH as promoted by organic ligands. The reductive dissolution rates for DFOB are lower than those for most of the other ligands; therefore, DFOB is a less-reactive reductant than common low-molecular mass organic acids and synthetic chelating agents. Reductive dissolution by DFOB may not be highly important for the solubilization of Mn. However, the reductive reaction pathway also results in the oxidative degradation of the siderophore. Thus, Mn(III) (hydr)oxides may be an important abiotic sink for microbial siderophores in natural environments (66, 71). VOL. 39, NO. 16, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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Acknowledgments We are grateful to Kenneth Raymond for providing our sample of Desferal. Michael Finnegan graciously assisted in the BET surface area determination. We thank Edward Bellfield and Andrew Yang for logistical support, and Jasquelin Pen ˜ a and Brandy Toner for valuable discussion. This work was funded by the National Science Foundation, Collaborative Research Activities in Environmental Molecular Science (CRAEMS) program (CHE-0089208).

Supporting Information Available Example of our dissolution data as a function of time, and a TEM micrograph and X-ray diffractogram of manganite particles. This material is available free of charge via the Internet at http://pubs.acs.org.

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Received for review February 10, 2005. Revised manuscript received June 7, 2005. Accepted June 7, 2005. ES050276C

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