Phosphonate- and Carboxylate-Based Chelating ... - ACS Publications

This study explores environmentally relevant synthetic phosphonate-based chelators that solubilize (hydr)oxide-bound MnIII via ligand-promoted and/or ...
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Environ. Sci. Technol. 2008, 42, 4397–4403

Phosphonate- and Carboxylate-Based Chelating Agents that Solubilize (Hydr)oxide-Bound MnIII YUN WANG* AND ALAN T. STONE Department of Geography and Environmental Engineering, The Johns Hopkins University, 3400 North Charles Street, Baltimore, Maryland 21218

Received December 31, 2007. Revised manuscript received April 2, 2008. Accepted April 6, 2008.

Recent field studies suggest that dissolved MnIII should be ubiquitous at oxic/anoxic interfaces in all natural waters and may play important roles in biogeochemical redox processes. Here, we uncovered environmentally relevant synthetic phosphonate-based chelators that solubilize (hydr)oxide-bound MnIII via ligand-promoted dissolution at circum-neutral pHs and that their ability to release aqueous MnIII can be predicted based on the chemical structure. For two (hydr)oxides (manganite and birnessite) reacting with excess concentrations of pyrophosphoric acid (PP), methylenediphosphonic acid (MDP), and phosphonoacetic acid (PAA), ligand-promoted dissolution is predominant: from pH 6-8, initial dissolution rates and plateau concentrations for dissolved MnIII decrease in the order PP > MDP > PAA, and at pH 5, MDP reacts equally well (with birnessite) or more efficiently (with manganite) than PP, and PAA remains the least reactive chelator. For manganite reacting with an excess concentration of aminophosphonate/carboxylate-based chelators, the aminophosphonatecontaining iminodimethylenephosphonic acid and glyphosate yield appreciable amounts of dissolved MnIII, but the aminocarboxylate-based methyliminodiacetic acid yields solely dissolved MnII via MnIII reduction.

Introduction Dissolved MnIII is a powerful oxidant, which has been implicated in the microbial degradation of lignin (1) and xenobiotics such as alkyphenols (2) and in the abiotic degradation of a variety of contaminants including ethylenediaminetetraacetic acid (EDTA) (3) and nitrilotris(methylene)phosphonic acid (NTMP) (4). It is thought to adversely affect iron bioavailability under nutrient-limited environments via binding siderophores, such as pyoverdine (5) and desferrioxamine B (6), equal to or stronger than its FeIII counterpart. However, the role of dissolved MnIII in natural waters largely has been overlooked because it is unstable with respect to disproportionation to yield MnIV solids and dissolved MnII and persists only in the presence of chelators (e.g., pyrophosphate) in laboratory studies (3). Field studies recently have shown that micromolar concentrations of dissolved MnIII are present near oxic/anoxic interfaces in the Black Sea and Chesapeake Bay (7). These authors further suggested that dissolved MnIII should be ubiquitous at oxic/ * Corresponding author phone: (617)324-2772; e-mail: yunwang@ mit.edu. 10.1021/es7032668 CCC: $40.75

Published on Web 05/14/2008

 2008 American Chemical Society

anoxic interfaces in diverse aquatic settings and hence may have additional profound environmental and ecological consequences through their effects on contaminant and nutrient cycling (7). In natural environments, dissolved MnIII may arise from biotic or abiotic oxidation of MnII (6, 8) or from chelator driven MnIII detachment via nonreductive ligand-promoted dissolution of mineral surfaces (9–11). Synthetic chelators are introduced into the environment through their use as scale/corrosion inhibitors (e.g., pyrophosphate, 2-phosphonobutane-1,2,4-tricarboxylic acid (PBTC), methylenediphosphonic acid (MDP), and iminodi(methylphosphonic acid) (IDMP)) (12, 13), pharmaceutical reagents (e.g., MDP and phosphonoacetic acid (PAA)) (14, 15), and herbicidal agents (e.g., glyphosate) (16) (chemical structures shown in Figure 1). These chelators, mostly phosphonate- and carboxylatebased, are known to be effective at coordinating metal ions in solution, facilitating adsorption, and assisting dissolution (10, 13). Upon contact with MnIII-containing minerals, they also may promote dissolution with the concomitant release of dissolved MnIII. Nevertheless, because dissolved and solid phase MnIII are effective oxidants, there is a competitive reaction: the oxidation of coordinated chelators, yielding the reduction of MnIII to MnII. Reduction can out-compete ligandpromoted dissolution. EDTA (9, 17) and NTMP (18), for example, react with MnIIIOOH at circum-neutral pH values primarily via reductive dissolution. Depending on the chelator structure, mineralogy, and other factors such as pH, there are situations where appreciable amounts of dissolved MnIII can be formed via ligand-promoted dissolution of MnIII-containing (hydr)oxides. An excess concentration of the redox-inert chelator pyrophosphate completely solubilizes MnIIIOOH (feitknechtite) and releases primarily dissolved MnIII within the range of 6.5 e pH e 8 (9). The same excess concentration of citrate, on the other hand, only yields predominantly dissolved MnIII when the pH is near 8. Dissolved MnII via reductive dissolution becomes dominant at a lower pH near 6.5 (9). A similar pH dependence was observed for MnIIIOOH (manganite) dissolution by phosphonoformate (10) and desferrioxamine-B (19), where ligand-promoted dissolution is important at a pH near 7 and reductive dissolution grows in importance at lower pH values. When MnIIIOOH is replaced by MnO2 (a mixed MnIII,IV phase), phosphonoformate only leads to reductive dissolution (10). Our objectives here were to uncover additional synthetic chelators (Figure 1) that react with MnIII-containing (hydr)oxides via ligand-promoted dissolution and to learn more about reaction kinetics and governing mechanisms. The ability to monitor dissolved MnIII concentrations using capillary electrophoresis (20) facilitated these efforts.

Materials and Methods All solutions were prepared from reagent grade chemicals without further purification and distilled, deionized water (DDW) with a resistivity of 18 MΩ cm (Millipore Corp.). Filter holders (Whatman Scientific) were soaked in 1 N ascorbic acid (Aldrich) and rinsed with distilled water and DDW. All bottles and glassware in contact with MnO2 (birnessite) and MnOOH (manganite) were cleaned using a 1 N ascorbic acid soak, a distilled water rinse, and a 4 N nitric acid (J. T. Baker) soak overnight, followed by thorough rinsing with distilled water and DDW. Chemicals. PAA, methyliminodiacetic acid (MIDA), and MnIII(acetate)3 · 2H2O were purchased from Aldrich. IDMP, VOL. 42, NO. 12, 2008 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 1. Eleven ligands employed in this study. disodium dihydrogen pyrophosphate (PP), methylphosphonic acid (MP), and N-(phosphonomethyl)glycine (glyphosate) were purchased from Fluka. Iminodiacetic acid (IDA) was purchased from Sigma. 3-Phosphonopropionic acid (PPA) and 1,2-ethylenediphosphonic acid (EDP) were purchased from Lancaster Synthesis. MDP was purchased from Alfa Aesar. PBTC was purchased from Halag Chemie AG. HNO3, HCl, NaOH, and NaCl were purchased from J. T. Baker. Our MnIII-pyrophosphate stock solution, serving as a capillary electrophoresis (CE) standard for dissolved MnIII, was synthesized by dissolving MnIII(acetate)3 · 2H2O into a 20fold concentration excess of pyrophosphate solution at pH 7 (details described in ref 20). MnO2 herein refers to particles synthesized according to the method of Luo et al. (21). Details of the synthesis and characterization are provided elsewhere (20). X-ray diffraction lines are consistent with a distorted birnessite layered structure. Transmission electron microscopy (TEM) revealed that many of the MnO2 sheets were curled, attributable to edge sections being thicker than middle sections (Figure S1a, Supporting Information). A BET surface area of 174 m2/g was determined using a freeze-dried sample. An average Mn oxidation state of +3.78 was determined using iodometric titration (22). The (hydr)oxide consists of a mixture of 22% MnIII and 78% MnIV, assuming that the MnII content is negligibly low. Recent XANES work on similarly prepared birnessite (23) confirmed this assumption. MnOOH herein refers to particles synthesized according to the method of Giovanoli and Leuenberger (24); details of the synthesis and characterization are provided elsewhere (10). X-ray diffraction lines are consistent with the mineral manganite. TEM revealed needle-shaped crystals of uniform size (Figure S1b, Supporting Information). An iodometric titration of MnOOH was not performed. An earlier preparation synthesized using the same procedure exhibited an average Mn oxidation state of +3.0 (17). A BET surface area of 27.3 m2/g was determined using a freeze-dried sample, within the range of reported surface areas (10-50 m2/g) for synthetic manganite samples (25). Experimental Setup. All dissolution experiments were conducted in 100 mL polypropylene bottles in a constant temperature circulating bath. Ligand (plus additional constituents, as appropriate) were sparged with Ar (BOC gases) for 1 h prior to MnO2 or MnOOH addition. Sparging was continued during the experiments. The pH was monitored 4398

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during each reaction using a Fisher Accumet AR15 m with an Orion combination semimicro glass electrode and NISTtraceable standard buffer solutions. In some experiments, one or more pKa values of the chelator were close enough to the desired pH, and the chelator concentration was high enough to maintain a constant pH throughout the reaction time course. We referred to this as self-buffering. In other experiments, additions of HCl or NaOH were necessary during the reaction time course. A computer-controlled system consisting of two burettes, a pH meter, and controlling software, termed a pH-stat, was used for this purpose. The chelator concentrations were high enough to serve an additional purpose: maintenance of constant ionic strength conditions (20-200 mM). No additional electrolyte was added. Reaction suspension aliquots of 4 mL were collected at periodic intervals. Reactions were quenched by immediately filtering through 0.1 µM pore diameter track-etched polycarbonate filter membranes (Whatman). Total dissolved manganese (Mn(aq)) in the filtered solutions was analyzed using flame atomic absorption spectrophotometry (AAS: Aanalyst 100, PerkinElmer). Dissolved MnIII(MnIII(aq)) in the filtered solutions was analyzed using capillary electrophoresis (CE, see CE: MnIII(aq) Analysis). To avoid any MnIII(aq) degradation due to the delay between sampling and CE analysis, all analytes were injected into the CE system immediately after sampling. CE: MnIII(aq) Analysis. A CE unit from Beckman Coulter (P/ACE MDQ) with a diode-array UV-vis detector was employed. The detector bandwidth was set at 6 nm for 190 nm wavelength detection and 10 nm for all other wavelengths. Bare fused silica capillaries (Polymicro Technologies) with 75 µM i.d. × 60 cm total length were used for all separations. The effective length (from inlet to detector) was 52 cm. The capillary and sample board temperature was thermostatted to 10 °C during operation. Between separations, the capillary was sequentially rinsed by flushing with DDW for 0.5 min, 0.1 M NaOH for 1 min, DDW again for 1 min, and the CE for 2 min. Sample injection employed 0.5 psi of positive pressure for 15 s. Anion mode with a constant applied voltage (-22 kV) was employed for all separations. We elected to use a CE electrolyte consisting of 20 mM pyrophosphate, 0.4 mM TTAB, and 2 mM orthophosphate (pH 9.5) for dissolved MnIII (MnIII(aq)) analysis (20). For experiments involving MP, EDP, PPA, PP, MDP, PAA, and PBTC (Figure 1), pyrophosphate was a strong enough chelator

FIGURE 2. Illustrative dissolution of 200 µM (a) MnO2 and (b) MnOOH by 5 mM PP, MDP, and PAA at pH 7.0. Open symbols denote Mn(aq), and solid symbols denote MnIII(aq). and present in a high enough concentration to capture all the MnIII in the samples during electromigration. A sharp, symmetrical peak, easily discernible at a detection wavelength of 235 nm, provided the means of measuring MnIII(aq). MnIII concentrations in filtered reaction solutions were calculated using calibration curves obtained from MnIII-pyrophosphate standard solutions. The detection limit for MnIII(aq) was approximately 1.5 µM. For experiments involving IDMP, glyphosate, IDA, and MIDA (Figure 1), pyrophosphate was not a strong enough chelator to capture all of the MnIII during electromigration. A peak corresponding to dissolved MnIII could still be discerned at 235 nm, however. MnIII(aq) production by these ligands can be qualitatively compared using the CE signals (ratios of CE peak area over migration time) that appeared in the electropherograms given that MnIII complexed by these structural similar ligands should exhibit comparable molar extinction coefficients at 235 nm.

Results and Discussion Throughout the pH range considered (5.0 < pH < 8.0), ligandfree MnO2 suspensions did not yield dissolved Mn (Mn(aq)) above the AAS detection limit of 1 µM. The same holds for MnOOH experiments at pH g 6.0, consistent with previous observations from independent studies (26). At pH 5.0, however, 200 µM MnOOH supensions yielded 14 µM Mn(aq), or 7% of the total loading after 250 h of reaction. Slow disproportionation yielding MnII is believed to be responsible since no peak corresponding to MnIII was observed using CE. MP, EDP, and PPA failed to dissolve either MnO2 or MnOOH under the experimental conditions considered (5.0 < pH < 8.0, 5.0 mM total added ligand, 25 h reaction time with MnO2 and 250 h reaction time with MnOOH). PP, MDP, and PAA: MnO2. We investigated MnO2 dissolution by PP, the diphosphonate MDP, and the phosphonate-carboxylate PAA as a function of pH, ligand concentration, and MnO2 loading. In all dissolution experiments performed, the total dissolved Mn (Mn(aq)) measured using AAS and dissolved MnIII (MnIII(aq)) measured using CE were consistently equal to one another (Figure 2a). Since O2 was excluded from our experiments, any MnII generated by reductive dissolution should have been discernible. The lack of any dissolved MnII indicates that PP, MDP, and PAA are resistant to oxidation and do not yield reductive dissolution. Unreacted MnO2 has an average Mn oxidation state of +3.78, corresponding to 22% MnIII and 78% MnIV. When 200 µM MnO2 was reacted for 25 h with 5 mM PP at pH 7.0, 40 µM dissolved MnIII was released, close to the expected value of 44 µM. Remaining Mn particulates collected by filtration yielded an average Mn oxidation state of +3.96. The transfer of particulate MnIII to dissolved MnIII is therefore stoichio-

metric. Together, these results reveal that the release of MnIII(aq) by all three ligands should be via ligand-promoted dissolution. A typical time course exhibited an initial fast release of MnIII(aq), a subsequent diminished rate, and a final leveling out of MnIII(aq) at ∼20 h of reaction (Figure 2a and Figures S2 and S3, Supporting Information). To explore the initial fast release, we applied least-squares fits to data collected during the first 0.4 h of reaction, yielding initial dissolution rates derived from the slopes of concentration versus time plots (dMn(aq)/dt) with r2 values g0.9. To explore the final plateau values, the fraction MnIII(aq)/TOTMnIII was calculated from data collected after 20 h of reaction. To assess the effect of ligand concentration, we investigated the dissolution of 200 µM MnO2 by PP, MDP, and PAA in the concentration range of 1-100 mM at pH 7 (Figure 3). As indicated by Figure S4 (Supporting Information), very little Mn was dissolved by the 1 mM ligand, and the time course behavior was complex. With PP, raising its concentration to 5 mM led to a significant jump of the rate to 26 µM/h. The rate continued to increase as the PP concentration was further increased but became less sensitive to the concentration change: from 5 to 20 mM, the rate increased by ∼40%; from 20 to 100 mM, the rate only increased by ∼10%. Similar to the trend observed with the initial rate, raising the PP concentration from 1 to 5 mM led to a MnIII(aq)/ TOTMnIII jump from ∼10 to 90% at the end of each reaction, corresponding to MnIII(aq) increasing from nearly 4 to 40 µM. MnIII(aq)/TOTMnIII remained at 90% when the PP concentration was further increased. As compared to PP, it took higher concentrations of MDP to reach the maximal MnIII dissolution efficiency with respect to both the initial rate and the eventual MnIII(aq)/TOTMnIII: 5 mM MDP was ∼40% less efficient than PP; the difference became smaller as the ligand concentration was increased, and at 20 mM, MDP was equally good as PP; the MnIII dissolution efficiency remained nearly the same when the MDP concentration was continuously raised to 100 mM. As compared to PP and MDP, PAA was the poorest in dissolving MnIII. At 5 mM, the initial rate was below 1 µM/h, and the eventual MnIII(aq)/TOTMnIII was less than 10%. In the concentration range of 10-100 mM, the rate increased nearly linearly as the PAA concentration was increased and reached 35 µM/h at 100 mM PAAsclose to those of PP and MDP at 20 mM; MnIII(aq)/ TOTMnIII also increased as the PAA concentration was increased and reached 86% at 100 mM PAAsclose to those of PP at 5 mM and MDP at 20 mM. To assess the effect of pH, we investigated the dissolution of 200 µM MnO2 by 5 mM PP, MDP, and PAA in the pH range of 5-8 (Figure 4). PP was efficient in dissolving MnIII, and the dissolution was not sensitive to pH changes: the initial rates were slightly higher at pH 5-6 (32-35 µM/h) than those at VOL. 42, NO. 12, 2008 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 3. Ligand concentration effect on (a) initial rate and (b) fraction of MnIII dissolved at 20 h, for reaction of PP, MDP, and PAA with 200 µM MnO2 suspensions at pH 7.0.

FIGURE 4. pH effect on (a) initial rate and (b) fraction of MnIII dissolved at 20 h, for reaction of 5.0 mM PP, MDP, and PAA with 200 µM MnO2 suspensions.

pH 7-8 (26-28 µM/h); the eventual MnIII(aq)/TOTMnIII was ∼90-100%. With MDP, at pH 5, MnIII dissolution nearly matched that of PP but was less effective at higher pH values; the initial rate at pH 6 was higher than those at pH 7 and 8, but MnIII(aq)/TOTMnIII remained at ∼65% at all three pH values. PAA performed the poorest at all pH values, and raising the pH caused a continuous decrease in the MnIII dissolution efficiency: MnIII(aq)/TOTMnIII decreased from 50% at pH 5 to below the detection limit (5%) at pH 8, and the decrease in initial rates shared a similar trend. To assess the effect of MnO2 loading, in one experiment, the MnO2 loading was increased from 200 to 500 µM while keeping the ligand concentration at 5.0 mM and pH at 6.0 (Figure S5, Supporting Information). For each ligand, the initial rate and MnIII(aq) at the end of the reaction increased by ∼2.1-2.4-fold, commensurate with the 2.5-fold increase in the MnO2 loading (and hence the MnIII content in the mineral phase). The dissolution efficiency again decreased in the order PP > MDP > PAA. PP, MDP, and PAA: MnOOH. Dissolution of 200 µM MnOOH by 5 mM PP, MDP, and PAA was investigated in the pH range of 5-8. At pH g 6, Mn(aq) and MnIII(aq) were consistently equal to one another (Figure 2b). At pH 5, ligandfree MnOOH released MnII(aq) over time; therefore, Mn(aq) was constantly higher than MnIII(aq) (Figure S6, Supporting Information). For the same reasons as MnO2, all three ligands led to ligand-promoted dissolution when dissolving surfacebound MnIII from the MnOOH phase. Like MnO2, MnOOH dissolution also exhibited an initial fast release of MnIII(aq) followed by a slow release later on, but because the reaction

was slower, monitoring for nearly 200 h was necessary (Figure 2b). Consequently, initial rates were determined from the first 6.5 h of reaction to hold the same r2 criteria as MnO2, and MnIII(aq)/TOTMnIII values were determined after 200 h of reaction to illustrate final plateaus. As shown in Figure 5, with PAA, the pH dependence of MnOOH was similar to that of MnO2, where the dissolution ability decreased as the pH was increased, leaving the rate only measurable at pH 5. However, with PP and MDP, such a similarity was not observed. For each ligand, the pH effects on the initial rate and the eventual MnIII(aq)/TOTMnIII were different. For example, although PP yielded an initial rate 40% higher at pH 6 than at pH 7, the MnIII(aq)/TOTMnIII values were about the same (nearly 90%). The dissolution efficiency decreased in the order PP > MDP > PAA at pH g 6 and MDP > PP > PAA at pH 5. The initial rates with MnOOH ranged from 5 to 15 times lower than with MnO2, depending on the ligand structure and pH. PP, MDP, and PAA: Thermodynamics, Kinetics, and Mechanisms. Our results show that the redox-inert ligands PP, MDP, and PAA solubilize MnIII-bearing solid phases via ligand-promoted dissolution. The dissolution profile follows nonlinear kinetics: the increase of MnIII(aq) diminishes over time and eventually levels out to a plateau concentration. Such kinetic behavior may arise from (i) surface heterogeneity and a diminished number of surface sites as dissolution takes place (27) and (ii) for a strongly sorbing ligand, the metal-ligand complexes that are released to solution may quickly readsorb or precipitate back onto the surface, forming a species different from the precursor complex for dissolution

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FIGURE 5. pH effect on (a) initial rate and (b) fraction of MnIII dissolved at 200 h for reaction of 5.0 mM PP, MDP, and PAA with 200 µM MnOOH suspensions. (28). This backward readsorption/precipitation counteracts the dissolution reaction and grows in importance as the dissolution reaction proceeds. Another important feature is that an excess of ligand is necessary to ensure that the dissolved ligand concentration is high enough to support MnIII staying in solution. Concentrations of 1 mM PP and MDP and 5 mM PAA are necessary to yield only barely measurable MnIII(aq). These ligand concentrations (TOTL) are much greater than the estimated concentrations that are required to cover all the surface sites (ST ) ∼30 µM, see Supporting Information). Consistent with what we report here, a similar threshold effect was observed for dissolving Fe(OH)3 (ferrihydrite) by triphosphate, where measurable FeIII(aq) can only be observed when TOTL > ST (28). Together, these observations imply common characteristics for strongly sorbing ligands. Dissolution is ligand structure dependent: (i) MP and acetate (ligands with a single Lewis base group) cannot, but PP, MDP, and PAA (ligands with two Lewis base groups) can dissolve MnIII; (ii) PP, MDP, and PAA, with the potential to form six-membered chelate rings with MnIII, are much more capable of dissolving MnIII than their seven-membered ring counterparts EDP and PPA; and (iii) PAA dissolves MnIII less efficiently than PP and MDP, correlating with the fact that as compared to the carboxylate group, the phosphonate (or phosphate) group has a higher charge and higher basicity (29). All these suggest that ligands with higher equilibrium constants (log K) for complexing MnIII in solution lead to

more efficient dissolution with respect to initial rate and plateau MnIII(aq), as proposed for the ligand-promoted dissolution of other metal (hydr)oxides (30, 31). A correlation between log K and initial rates would suggest that the activated complex on the surface resembles the corresponding solution complex in significant ways (30, 32). Unlike PP and MDP, plateau values of MnIII(aq) fall short of TOTMnIII and increase in direct proportion to an increase in PAA concentrations. These observations suggest that thermodynamic constraints on dissolution are more pronounced for PAA than for PP and MDP. In principle, plateau values can be used to estimate complex formation constants and solubility product constants (Supporting Information). In practice, the lack of information regarding the protonation level and stoichiometry of major dissolved MnIII species precludes such estimates. Dissolution is dependent on pH and MnIII mineralogy. Protons compete with surface MnIII for available ligands, and the ligands compete with hydroxide ions for surface MnIII. The situation becomes more complex when we bring another oxide into consideration. Although both MnO2 and MnOOH exhibit nonlinear dissolution profiles, their initial rates and MnIII(aq)/TOTMnIII are quite different, and their dissolution has a different dependence on ligand structure and pH. The following different mineral properties may contribute to the differences in MnIII dissolution: MnO2 (birnessite) is a layered phase (33) with strong cation-exchange properties and a pHzpc of 2.3 (34), yet MnOOH (manganite) is a nonion exchanging phase with a pHzpc of 8.1 (25); the surface area loading in experiments involving MnO2 is 6.5 times higher than in experiments involving MnOOH; and MnO2 consists of 22% MnIII, yet MnOOH consists of 100% MnIII. In addition, here we show that it is easier to dissolve MnIII from MnO2 than from MnOOH, which leads us to speculate that MnIII may be thermodynamically more stable within the MnOOH structure than within the MnO2 structure. Given all these complications, it is perhaps not surprising that we cannot predict dissolution behavior based on a single property. IDMP, Glyphosate, IDA, and MIDA: MnOOH. Aminocarboxylates and aminophosphonates represent important synthetic chelators due to their applications in the agricultural, industrial, and pharmaceutical fields. Here, we selected four such chelators to quantify their abilities to dissolve MnIII by performing experiments employing 5.0 mM ligands and 200 µM MnOOH at pH 6.0. The time course plot (Figure 6a, top panel) shows that Mn(aq) production decreased in the order IDMP > MIDA > glyphosate . IDA. The low Mn(aq) released (∼1 µM) in the presence of IDA versus in the absence of IDA was the same, indicating that IDA cannot dissolve MnOOH. CE electropherograms provide evidence for MnIII(aq) production (Figure 6b). Samples were selected for CE analysis that contained 16-19 µM Mn(aq). Both IDMP and glyphosate yielded MnIII peaks, with the peak for IDMP significantly higher. MIDA did not yield a MnIII peak. For the reasons described in the Materials and Methods, a high signal corresponded to a high MnIII concentration. MnIII signals as a function of time (Figure 6a, bottom panel) showed that with IDMP, MnIII(aq) increased in the first 5 h of reaction and leveled out afterward; with glyphosate, MnIII(aq) leveled out at one-tenth the value for IDMP; and with MIDA, MnIII(aq) was below the detectable level. Note that Mn(aq) is from the sum of reductive and ligandpromoted dissolution, while MnIII(aq) is from ligandpromoted dissolution only. Together, our results indicate that IDMP yielded the most efficient ligand-promoted dissolution among the four ligands, and since IDMP is resistant to oxidative degradation by MnOOH in the circumneutral pH range (18), IDMP should dissolve MnOOH primarily via ligand-promoted dissolution. Glyphosate yielded both reductive and ligand-promoted dissolution, MIDA VOL. 42, NO. 12, 2008 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 6. Reaction of 200 µM MnOOH with IDMP, MIDA, glyphosate, and IDA at pH 6.0. (a) Total dissolved manganese (Mn(aq)) (top panel) and total dissolved MnIII (MnIII(aq)) (bottom panel) as a function of time. (b) For the three aliquots denoted using solid symbols, CE electropherograms also were collected (right). The migration time for MnIII(aq) was within the range of 2.5-2.6 min. yielded primarily reductive dissolution, and IDA was not able to dissolve MnOOH. As discussed in the previous section, ligands involving phosphonate groups should bring about ligand-promoted dissolution more effectively than their analogues involving carboxylate groups. From IDMP to glyphosate to MIDA (and IDA), a phosphonate group is continuously replaced by a carboxylate group. Therefore, the ability of these ligands to dissolve MnOOH via ligand-promoted dissolution decreases. MIDA is an aminocarboxylate with a tertiary amine group, and IDA is its analogue with a secondary amine group. MIDA is efficient at reductively dissolving MnOOH, but IDA is not. This is consistent with earlier studies showing that, in comparison to their analogues involving secondary amine groups, aminocarboxylate or aminophosphonate ligands involving tertiary amine groups are more subject to oxidation by MnIII,IV with concomitant release of dissolved MnII (4, 18, 31). MnII(aq) may arise from (i) the reduction of surfacebound MnIII and (ii) from ligand-promoted dissolution followed by MnIII reduction within the dissolved complexes. MIDA and IDA have the same Lewis base groups and hence should yield similar rates of ligand-promoted dissolution. Since IDA cannot dissolve MnOOH, it is unlikely that MIDA dissolves MnOOH via pathway ii. For glyphosate, both pathways are possible. This is different from an earlier study using a MnO2 phase where no MnIII(aq) has been observed (35), again indicating that MnIII(aq) production is dependent on MnIII,IV mineralogy. Environmental Relevance. In this study, we uncovered several synthetic chelators that solubilize (hydr)oxide-bound MnIII via ligand-promoted dissolution at circum-neutral pHs. Among these phosphonate-based chelators, MDP, PAA, PBTC, and IDMP are scale/corrosion control chemicals and/ or pharmaceuticals and are released into environments (12–15); glyphosate, one of the most widely used herbicides in the world, is applied directly to agricultural lands (16). MnIII-containing minerals are present not only in diverse aquatic settings (36) but also in water-based engineered 4402

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systems including water supply plants and power plant cooling circuits (37). Gradients in organic matter content, dissolved O2 concentrations, and pH can yield Mn(hydr)oxide enrichments in soil microenvironments (38) and biocorrosion films (39). Adding synthetic chelators with the right functional groups is all that is needed to generate soluble MnIII. With the recent finding that dissolved MnIII is present in natural waters (7), it is tempting to speculate that MnIII solubilization by synthetic phosphonates in this study and many others may prove to be an important source of this powerful oxidant in marine and freshwater sediments, soil interstitial waters, and engineered aqueous systems. Dissolved MnIII hence is expected to play important roles in biogeochemical redox processes in all these environments.

Acknowledgments We thank Anne-Claire Gaillot for help with MnO2 (birnessite) and MnOOH (manganite) characterization, Thanh (Helen) Nguyen for help with BET surface area measurements, Rodrigue Spinette for help with pH-stat experiments, James Morgan for stimulating discussions, and three anonymous reviewers for their insightful comments. This work was supported by Grant R-82935601 from the U.S. Environmental Protection Agency’s Science to Achieve Results (STAR) program. This article has not been subject to EPA review and therefore does not necessarily reflect the views of the agency, and no official endorsement should be inferred.

Supporting Information Available Detailed descriptions of stoichiometric estimation of surface complexation; thermodynamic information obtained from dissolution experiments; transmission electron micrographs of MnO2 and MnOOH; descriptions and figure of dissolution of 200 µM MnO2 by 1.0 mM PP, MDP, and PAA; figure of MnO2 loading experiment; figure of dissolution of 200 µM MnOOH by 5.0 mM PP, MDP, and PAA at pH 5 and 6; and descriptions and figure of dissolution of 200 µM MnO2 by 5.0

mM PBTC. This material is available free of charge via the Internet at http://pubs.acs.org.

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