Mechanism of Hydrogen Sulfide Oxidation by Manganese(IV) Oxide in

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Mechanism of Hydrogen Sulfide Oxidation by Manganese(IV) Oxide in Aqueous Solutions Julia´n Herszage and Marı´a dos Santos Afonso* INQUIMAE and Departamento de Quı´mica Inorga´ nica, Analı´tica y Quı´mica Fı´sica, Facultad de Ciencias Exactas y Naturales, Universidad de Buenos Aires, Ciudad Universitaria Pabello´ n II, C1428EHA Buenos Aires, Argentina Received January 6, 2003. In Final Form: August 16, 2003 The kinetics of reductive dissolution of Mn(IV) oxides by hydrogen sulfide has been investigated. The concentration of total H2S(aq) in the reaction system ranged between 1.0 and 4.0 mM and was adjusted keeping a constant partial pressure of the gaseous H2S phase above the solution. The reaction products were identified as Mn(II), sulfate, elemental sulfur, and small amounts of thiosulfate. The distribution of products changes with pH; sulfate is the main product at low pH values while elemental sulfur is the main product at near neutral pH values. The rate constant decreased with an increase in pH. Experimental data indicate that the rate law is first order on both the surface sites and the H2S concentration. The reaction proceeds via the formation of two different inner-sphere surface complexes tMnIVS- and tMnIVSH and their further oxidation to products. A mechanism in agreement with the experimental results is proposed.

Introduction Several different pathways are involved in the overall dissolution reactions of manganese oxides. It is known that dissolution processes are kinetically controlled by the reactivity of crucial coordinative arrangements around the metal centers in the lattice surface. At the anoxic zone of an aquatic system, manganese or iron oxides can oxidize hydrogen sulfide, either biotically or abiotically, leading to the formation of several oxidized sulfur species.1-5 Elemental sulfur6 and sulfate7,8 are the main oxidation products. Burdige and Nealson6 reported the formation of elemental sulfur in the course of Mn(IV) mineral dissolution by sulfide ions, whereas in the study of Yao and Millero7,8 the formation of sulfate, thiosulfate, and traces of sulfite was observed under similar experimental conditions. Significant amounts of thiosulfate have also been observed in the pore waters of sediments sampled from salt marshes.9-11 Luther et al.12 suggested that thiosulfate, originating from pyrite oxidation, can also form organic sulfur. Sulfite has been seldom found in nature. Boulegue et al.10 and Luther et al.11 have reported sulfite as a minor component of the natural products of sulfur oxidation, but its presence was not detected by Howarth et al.13 * To whom the correspondence should be addressed. Fax: (++5411) 4576 3341. E-mail: [email protected]. (1) Goldhaber, M. B.; Kaplan, I. R. In In the Sea; Goldberg, E. D., Ed.; Wiley: New York, 1975. (2) Pyzik, A. J.; Sommer, S. E. Geochim. Cosmochim. Acta 1981, 45, 687-698. (3) dos Santos Afonso, M.; Stumm, W. Langmuir 1992, 8, 16711675. (4) Pfeiffer, S.; dos Santos Afonso, M.; Wehrli, B.; Ga¨chter, R. Environ. Sci. Technol. 1992, 26, 2408-2413. (5) Yao, W.; Millero, F. J. Mar. Chem. 1996, 52, 1-16. (6) Burdige, D. J.; Nealson, K. H. Geomicrobiol. J. 1986, 4, 361-387. (7) Yao, W.; Millero, F. J. Geochim. Cosmochim. Acta 1993, 57, 33593365. (8) Yao, W.; Millero, F. J. In Geochemical Transformations of Sedimentary Sulfur; Vairavamurthy, M. A., Schoonen, M. A. A., Eds.; ACS Symposium Series 612; American Chemical Society: Washington, DC, 1995.

Although hydrogen sulfide oxidation has been the subject of several past studies, uncertainties still remain regarding rates, mechanisms of oxidation, and the formation of the oxidation products. In this work, the study of the reductive dissolution of Mn(IV) oxides by hydrogen sulfide has been undertaken to elucidate the elementary steps involved in this process in the pH range 3.0-10.0. A mechanism in agreement with the experimental results is proposed. Experimental Section All the solutions used were made using reagent-grade chemicals and distilled and deionized water (DDW) produced in a Milli-Q apparatus (conductivity less than 0.1 µS cm-1). All the glassware used was acid-washed for, at least, 24 h and rinsed with DDW several times before doing the experiments. The dissolution experiments were performed in a magnetically stirred 400-mL airtight cylindrical water-jacketed beaker. The temperature was held constant to (0.1 °C with a constanttemperature circulation bath. Three different solid phases of manganese dioxide were synthesized using the following methods: Murray14 for MnO2(A), Stone and Ulrich15 for MnO2(B), and Pe´rez Benito et al.16 for MnO2(C). The first two oxides were analyzed by X-ray diffraction (powder method) and transmission electron microscopy. The X-ray powder and oriented diffraction patterns were performed using a Siemens diffractometer D5000 equipped with a graphite monochromator and a Cu radiation tube. The operational conditions were 40 000 V and 30 mA. MnO2(A) and MnO2(B) were identified as vernadite (9) Howarth, R. W.; Teal, J. M. Limnol. Oceanogr. 1979, 24, 9991013. (10) Boulegue, J.; Lord, C. J.; Church, T. M. Geochim. Cosmochim. Acta 1982, 46, 453-464. (11) Luther, G. W., III; Varsolona, R.; Giblin, A. E. Limnol. Oceanogr. 1985, 30, 727-736. (12) Luther, G. W., III; Church, T. M.; Scudlark, J. R.; Cosman, M. Science 1986, 232, 746-749. (13) Howarth, R. W.; Giblin, A.; Gale, J.; Peterson, J.; Luther, G. W., III. Environmental Biogeochemistry Ecological Bulletin 1983, 35, 135152. (14) Murray, J. W. J. Colloid Interface Sci. 1974, 46, 357-371. (15) Stone, A. T.; Ulrich, H. J. J. Colloid Interface Sci. 1989, 132, 509-522. (16) Pe´rez Benito, J. F.; Arias, C.; Amat, E. J. Colloid Interface Sci. 1996, 177, 288-297.

10.1021/la034016p CCC: $25.00 © 2003 American Chemical Society Published on Web 10/11/2003

Mechanism of Hydrogen Sulfide Oxidation and akhtenskite, and the specific surface areas (BET) were 30.8 and 16.1 m2/g, respectively. MnO2(C) was a colloidal phase oxide and showed a large band covering the whole visible region of the spectrum with the absorbance uniformly decreasing with increasing wavelength, as well as a wide maximun at 300-400 nm. The UV-visible spectroscopy was performed using a Hewlett-Packard 8453 diode array spectrophotometer. The stoichiometry of the oxides was determined using iodometric techniques and flame atomic absorption spectroscopy (AAS) using a Varian AAS5 instrument for total Mn measurement. The concentration of the manganese oxide suspension stock was 0.12 M in MnO2. The point of zero charge (pzc) of the MnO2(C) was determined using potentiometric acid-base titrations (pzc ) 1.9). Therefore, at the experimental pH range the colloidal particles had a negative electrostatic charge probably due to the anion adsorption on the colloid surface. The dissolution experiments for MnO2(A) and MnO2(B) were initiated by adding a known volume of resuspended oxide suspension to 350 mL of solution in equilibrium with hydrogen sulfide. The stock MnO2 suspensions were resuspended before starting the experiments using a magnetic stirrer, and upon resuspension, there was no redissolution of MnO2. The manganese oxide concentrations were varied between 0.0089 and 0.140 g/L. The experiments with the colloidal phase oxide were made using a stopped flow system coupled to a Hewlett-Packard 8453 diode array spectrophotometer. The colloid UV-vis spectrum was previously recorded with the diode array spectrophotometer. In these experiments, the colloidal manganese oxide concentration was varied between 8.7 × 10-5 and 5.2 × 10-4 g/L. H2S was supplied to the reaction mixture by continuous bubbling of special H2S/N2 commercial mixtures provided by Alphagaz. The N2 was scrubbed through a V(II) solution (Jones’ Reagent) to ensure the anoxia of the system. The experiments were performed in the pH range of 3.0-8.0. During the course of each experiment, the pH was kept constant by adding small aliquots of 10-2 M HClO4 using a 682 Titroprocessor Metrohm with Dosimat pH-stat system. Light was excluded from all the experiments. To ensure that the solution was in equilibrium with hydrogen sulfide, the H2S/N2 mixture was bubbled longer than 1 h before adding the manganese oxide to the reaction vessel. Because HClO4 is a strong oxidant, some blank experiments were made mixing all the reactants without manganese oxide. The experiments done without MnO2 were longer than those done in the presence of MnO2. No H2S oxidation products were detected in experiments done in the absence of MnO2. With the experimental setup used, the partial pressure of H2S was maintained constant during the course of each experiment using mass flow controllers, and the H2S concentration was calculated taking into account the gas-liquid equilibrium [H2S(g) ) H2S(aq), KH ) 0.999]. Then, because the pH was held constant, [H2S], [HS-], [S2-], and [H2S]T were also constant because [H2S]T ) [H2S] + [HS-] + [S2-]. In different experiments, the equilibrium concentration of H2S(aq) varied between 1.0 × 10-3 and 4.0 × 10-3 M while the total pressure (N2 + H2S) was held constant. In all the experiments, the ionic strength was 10-2 M of NaClO4. Sodium sulfide stock solutions were used as the source of sulfide in the experiments done with the colloidal phase oxide. These solutions were prepared by dissolving crystals of Na2S‚9H2O (Aldrich) in degassed DDW. The stock solutions were standardized iodometrically and were freshly prepared every day. In this case, the experiments were performed in the pH range of 4.0-10.0. This set of experiments were not made at a constant partial pressure of H2S. Many experiments were carried out until the total dissolution of manganese oxide was achieved. No MnS precipitation was detected under these experimental conditions probably because the solubility product of amorphous MnS was not exceeded [MnS(s) + H+ a Mn2+ + HS-; log Ks ) 2.95). Small aliquots (5 mL) were removed periodically with a syringe from the reaction mixture and filtered through a 0.2-µm membrane filter.

Langmuir, Vol. 19, No. 23, 2003 9685 The progress of the dissolution was followed by measuring the concentration of dissolved manganese, SO42-, and S2O32- as a function of time. Mn3+ was excluded as a reduction product on the basis of the results of the electrochemical experiments, so the total dissolved manganese is Mn(II). The concentration of total dissolved manganese was determined by flame AAS using a Varian AAS5 instrument. The standard solutions were prepared by dilution of a 1000-ppm AAS standard provided by Merck. The concentration of the oxoanions of sulfur formed in the reaction was determined by ionic chromatography using a DIONEX DX-100 instrument with a conductivity detector,17-19 a sample injection valve, and a 25-µL sample loop. Two plastic anion columns were coupled in series to serve both as the precolumn (DIONEX AG-9) and as the analytical chromatographic column (DIONEX AS-9). The suppressor was regenerated with 50 mN H2SO4 with a flow rate of 12.5 mL/min. A mixture of 4 mM HCO3-/CO32- was chosen as the eluent with a flow rate of 1 mL/min. The retention time under these operational conditions and using an isocratic method were 1.9 and 4.3 min for SO42- and S2O32-, respectively. The standard solutions were prepared using analytical reagent-grade chemicals. The identification and quantification of S8 was made by HPLC20 using a Shimadzu LC-6A instrument with a Shimadzu SPD-6AV UV-vis detector. The samples were injected into a Rheodyne injector no. 77251 with a 50-µL loop and passed through a Supelco LC 18 analytical chromatographic column. A methanol solution containing 2% water was used as the eluent. A standard was prepared by dissolving powered elemental sulfur in toluene, and dilutions from of this stock solution were prepared in methanol. The analytical precision was usually within 1%. To investigate the possible formation of polysulfides, the UV spectra of filtered reaction mixtures were performed between 200 and 400 nm.21,22 Square wave voltammetry (SWV) was used in experiments made with the colloidal phase to evaluate the formation of polysulfides and sulfur.23 Both an EG&G Princeton Applied Research model 384B polarographic analyzer in conjunction with a model 303A static dropping mercury electrode and an Analytical Instrument Systems, Inc., model DLK 100 voltammetric analyzer were used for SWV measurements. The electrode stands were modified to use a saturated calomel electrode (SCE) rather than the Ag/AgCl reference supplied. Instrumental parameters for the SWV mode were typically a 200 mV s-1 scan rate over the potential range -0.1 to +1.3 V with a 25-mV pulse height; the detection limit was lower than 1 µM.

Results and Discussion Some experiments were done to evaluate the dissolution rate of MnO2 promoted by H+ at pH 3.0, without any H2S added. In these experiments, small aliquots were removed periodically and analyzed by AAS. The amount of dissolved manganese was not detectable by AAS after 20 h. Therefore, the contribution of this process to the overall amount of dissolved manganese for the dissolution process in the presence of the H2S was considered negligible. A typical experiment is shown in Figure 1 for MnO2(B) and MnO2(A). The amount of dissolved manganese oxide, in the presence of 1.69 mM H2S(aq), is significant, reaching eventually a plateau value after 1 h. Measurements done by monitoring the absorbance of the suspen(17) Sunden, T.; Lindgren, M.; Cedergren, A.; Siemer, D. D. Anal. Chem. 1983, 55, 2-4 (18) Lindgren, M.; Cedergren, A.; Linberg, J. Anal. Chim. Acta 1982, 141, 279-286. (19) Shpigun, O.; Zolotov, K. A. Ion Chromatography in Water Analysis; Wiley: New York, 1988. (20) Henneke, E.; Luther, G. W., III; De Lange, G. J.; Hoeps, J. Geochim. Cosmochim. Acta 1997, 61, 307-321. (21) Hoffmann, M. R. Environ. Sci. Technol. 1977, 11, 61-66. (22) Zhong, J.; Millero, F. J. In Environmental Geochemistry of Sulfide Oxidation; Alpers, C. N., Blowes, D. W., Eds.; ACS Symposium Series; American Chemical Society: Washington, DC, 1992. (23) Taillefert, M.; Bono, A.; Luther, G. W., III. Environ. Sci. Technol. 2000, 34, 2169-2177.

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Figure 1. Variation of the total manganese concentration in solution as a function of time. (A) MnO2(A): (b) acid dissolution at pH ) 3 and {MnO2} ) 0.9 × 10-3 M; (2) {MnO2} ) 0.9 × 10-3 M, [H2S] ) 1.69 × 10-3 M, and pH ) 4.0. (B) MnO2(B): {MnO2} ) 1.5 × 10-3 M, [H2S] ) 1.69 × 10-3 M, and pH ) 4.0. In all the cases, I ) 0.01 M NaClO4 and T ) 25 °C.

Figure 3. Dependence of the initial rate expressed as Mn2+ formation rate and MnO2 dissolution rate with the manganese oxide concentration at I ) 0.01 M NaClO4. (9) pH ) 3.5, [H2S]T ) 1.69 × 10-3 M, and T ) 25 °C. (2) pH ) 7.0, [H2S]T ) 3.4 × 10-3 M, and T ) 25 °C.

Figure 4. Dependence of k′ with the hydrogen sulfide concentration for MnO2(A) at pH ) 6.0, {MnO2} ) 1.0 × 10-3 M, I ) 0.01 M NaClO4, and T ) 25 °C. Figure 2. Typical experiment for colloidal manganese oxide, MnO2(C), at {MnO2} ) 12 × 10-6 M, pH ) 10.0, [Na2S] ) 1.00 × 10-3 M, I ) 0.02 M NaClO4, and T ) 25 °C. The dissolution rate is measured following the colloidal oxide concentration by analyzing the absorbance decay at the wavelength of 400 nm, where the colloidal phase oxide UV-vis spectrum has a maximun.

sions of MnO2(C) at 400 nm for a typical dissolution experiment in the presence of 1 mM of H2S are shown in Figure 2. The dissolution rate is measured following the colloidal oxide concentration by analyzing the absorbance decay at a wavelength of 400 nm, where the colloidal phase oxide UV-vis spectra has a maximun. At a given pH value and H2S partial pressure, the initial rate of dissolution expressed as dissolved Mn per time unit is linear with the amount of oxide added (Figure 3); therefore, the rate law should have the form

R0 ) k′{MnO2}T

(I)

where {MnO2}T denotes manganese oxide in mol m-2. On the other hand, when the pH and {MnO2}T are kept constant a linear relationship between k′ and [H2S]T is obtained (Figure 4). Thus, k′ ) k[H2S]T and

R0 ) k′{MnO2}T ) k[H2S]T{MnO2}T

(II)

The results of experiments done with the colloidal oxide using Na2S as the source of sulfide, at pH 10.0, to minimize the loss of H2S, are also in agreement with the rate law given by eq II (Figure 5). The oxidation products were identified as sulfate, sulfur, and small amounts of thiosulfate. No polysulfide or sulfite formation was observed. The percentages of sulfide oxidation products formed at the end of the reaction (i.e., complete dissolution of the oxide) are shown in Figure 6A,B. The data shown in Figure 6A correspond to the experiments done with suspensions of MnO2(A); nevertheless, similar results were obtained when the colloidal oxide was subject to reaction with Na2S solutions (Figure 6B). The fact that neither the product distribution nor the rate law changes when the colloidal phase is used suggests that the reaction mechanism is the same for both the solid and the colloidal phase oxides. As can be seen in Figure 6A,B the distribution of the oxidation products changes with pH, being sulfate the main product at low pH values and elemental sulfur the main product at near neutral pH. This fact can be explained considering that the thermodynamic driving force for the reaction increases with the decrease in pH. Burdige and Nealson6 studied the oxidation of H2S by MnO2 in aqueous solution at pH 8.0 in seawater. They reported elemental sulfur as the only product of H2S oxidation, although the amounts of sulfur formed accounted only for ∼50% of the Mn(II) formed.

Mechanism of Hydrogen Sulfide Oxidation

Figure 5. Dependence of the ratio between the initial rate of manganese solubilization (R0) and the hydrogen sulfide concentration with the colloidal manganese oxide concentration at pH ) 10.0, [Na2S] ) 1.0 × 10-3 M, I ) 0.02 M NaClO4, and T ) 25 °C. The straight line slope is k′.

Figure 6. Oxidation product percentages as a function of the pH for (A) MnO2(A), [H2S] ) 1.69 × 10-3 M, {MnO2} ) 1.0 × 10-3 M, T ) 25 °C, and I ) 0.01 M NaClO4 and (B) colloidal manganese oxide, [Na2S] ) 1.00 × 10-3 M, {MnO2} ) 1.0 × 10-3 M, T ) 25 °C, and I ) 0.01 M NaClO4. In both cases, the solid bars represent [SO42-], the dashed bars represent [S2O32-], and the crossed bars represent elemental sulfur.

In our study, elemental sulfur represents ∼80% of the oxidation products formed at pH 8.0 (Figure 6A,B). These differences may arise from the different experimental conditions or the different initial ratios of sulfide to oxide. Yao and Millero8 studied the changes in the distribution of reaction products with the sulfide-to-oxide ratio in seawater. Their results showed that, at pH 7.5 and for a {MnO2}/[H2S] ratio equal to 1, sulfur is the main product and that the amount of sulfate formed increases with the {MnO2}/[H2S] ratio. It should be noticed that the data in Figure 6A,B correspond to a {MnO2}/[H2S] ratio equal to 0.5. They also reported the formation of traces of sulfite; we did not find any sulfite, which is in agreement with the results of Petrie,24 who reported the fast oxidation of SO2 by MnO2.

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Figure 7. SWV at {MnO2}/[Na2S] of 1:60, pH ) 6.5, I ) 0.05 M NaClO4: (a, open circles) without N2 bubbling; (b, dashed line) after 4 min of N2 bubbling; (c, open circles with solid line) after 8 min of N2 bubbling; and (d, solid line) SWV at {MnO2}/ [Na2S] of 20:1, pH ) 6.5, and I ) 0.05 M NaClO4 without N2 bubbling.

In some SWV experiments performed at pH 6.5, using the colloidal oxide and an excess of Na2S, a signal at E ) -0.52 V (vs SCE) was observed (see Figure 7, line a). This signal did not disappear upon bubbling of the cell with N2 for several minutes (Figure 7, lines b and c) and decreased slowly with time. When a similar experiment is done in excess of MnO2, no signal is observed at all (Figure 7, line d). There are several sulfur species that show a signal around this potential (e.g.. sulfide, polysulfides, sulfur).25 This signal cannot be assigned to sulfide or polysulfides because at pH 6.5 sulfide is completely boiled out with N2 bubbling, and polysulfides show two signals under the experimental conditions used.25 The slow decrease with time indicates that this signal corresponds to some form of elemental sulfur that slowly polymerizes to S8 and precipitates. Other possible sulfur species such us sulfite, thiosulfate, and sulfate were not considered because they either do not show a signal around that potential value or are electrochemically inactive. On the other hand, the experiment made in excess of MnO2 demonstrates that the elemental sulfur species does not form or is rapidly oxidized to sulfate when the oxide is in excess. Taillefert et al.23 previously observed this behavior in colloidal suspensions of amorphous Fe(OH)3 when reacted with Na2S solutions. To determine whether the reaction mechanism is associative or dissociative, the entropy of activation is needed,26 which can be evaluated from an Eyring plot [ln(k/T) vs 1/T] based on eq III,

ln(k/T) ) (∆H‡/RT) + ln(k/h) + (∆S‡/R)

(III)

where ∆H‡ is the enthalpy of activation, ∆S‡ is the entropy of activation, k is the Boltzmann constant (1.381 × 10-23 J K-1), and h is Planck’s constant (6.626 × 10-34 J s); the slope (∆H‡/R) provides the enthalpy of activation; and the intercept ln(k/h) + (∆S‡/R) yields the entropy of activation. A large negative ∆S‡ indicates an associative reaction, whereas a large positive ∆S‡ indicates a dissociative (24) Petrie, L. M. Appl. Geochem. 1995, 10, 253-267. (25) Luther, G. W., III; Glazer, B. T.; Hohmann, L.; Popp, J. L.; Taillefert, M.; Rozan, T. M.; Brendel, P. J.; Theberger, S. M.; Nuzzio, D. B. J. Environ. Monit. 2001, 3, 61-66. (26) Atwood, J. D. Inorganic and Organometallic Reaction Mechanisms; Brooks/Cole Publishing Co.: Monterey, CA, 1985.

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Herszage and dos Santos Afonso Scheme 1. Dissolution Reaction Mechanism of Hydrogen Sulfide Oxidation onto Manganese Oxide

Figure 8. Arrhenius plot for MnO2(B) at pH ) 4.0 and I ) 0.01 M NaClO4. The units of the second-order contant, k, are M-1 s-1.

reaction.26 The ∆H‡ and ∆S‡ values for MnO2 reduction with H2S were -70.5 kJ mol-1 and -13.70 J K-1 mol-1, respectively. The ∆S‡ value suggests that the reaction is associative and proceeds via inner-sphere redox reactions. The ∆H‡ and ∆S‡ parameters are apparent because they depict the thermodynamic activation parameters of the overall reaction. Note that the rate constant employed in eq III does not correspond to a single-step process. The dependence of k with temperature was in agreement with the Arrhenius equation (Figure 8); the apparent activation energy value obtained for MnO2(A) was 73 ( 5 kJ mol-1, and this value is also indicative of a surfacecontrolled process.26,27 The apparent activation energy for the rate constant k will be a combination of the activation energies for the elementary reactions. Preexponential factors for bimolecular reactions commonly exhibit A values between about 107 and 1012 M-1 s-1. The dissolution rate of MnO2(B) at pH ) 4.0 and I ) 0.01 M NaClO4 has a preexponential factor of 4 × 109 M-1 s-1 that is a normal value for this kind of reaction. Yao and Millero7 found that the rate of sulfide oxidation by MnO2 is decreased approximately 50% in the presence of 10 µM phosphate and attributed this result to phosphate adsorption on manganese oxide blocking the surface sites that must be available for the reaction with sulfide. Later, these authors28 also showed that manganese oxides can act as important adsorbents of phosphate in natural waters, as well as in surface sediments, because the adsorption of phosphate on d-MnO2 and goethite in seawater are comparable. In our study, preliminary experiments made with MnO2 in the presence of H2S and phosphate showed that the rate decreases substantially with respect to the experiments made in the absence of phosphate. Therefore, the dissolution would be surfacecontrolled. Dissolution can proceed via several parallel pathways that involve labilization of bridging oxygens by ligands that are dynamically stable in the inner-coordination sphere of the detaching manganese. If any of the dissolution reactions were completely transport (diffusion)controlled, there should be no pH dependence of the oxidation rate. Despite this, if Mn(II) released into solution is the rate-determining step, there should be no difference (27) Lasaga, A. C. In Kinetics of Geochemical Processes, Reviews in Mineralogy; Lasaga, A. C., Kirkpatrick, R. J., Eds.; Mineralogical Society of America: Washington, DC, 1981. (28) Yao, W.; Millero, F. J. Environ. Sci. Technol. 1996, 30, 536-554.

in the dissolution rate between ligands. Then, the rate is controlled either by the binding of the ligand to the surface or the detachment of the activated surface complex from the surface. Thus, the dissolution rate is not controlled by the transport of reduced species away from the surface [Mn(II) release into solution is not the rate-determining step], and the reactions at the surface (which are responsible for the activation energy) must be ratecontrolling. These results would suggest that the dissolution reaction proceeds through an inner-sphere mechanism such us the mechanism outlined in Scheme 1. Thus, the dissolution reaction starts through a surface complex formation by the adsorption of HS- onto the manganese oxide surface (Scheme 1 eq 1). Later on, this surface complex is oxidized on the surface. Although the transfer of two electrons at a time was postulated by Luther,29,30 we assumed that all electron transfers are consecutive one-electron transfers. Therefore, the manganese oxide surface oxidizes the surface sulfide complex to adsorbed zero-valent sulfur (tMnIIS) in two steps (Scheme 1, eqs 2 and 3). During this process, a Mn(III)surface complex is formed as a reaction intermediate. There is clear evidence in the literature that Mn(III) is an intermediate in the reductive dissolution of MnO2. Mn(III) was detected by X-ray photoelectron spectroscopy as an intermediate in the reductive dissolution of birnessite with several reductants such as arsenite,31 Cr(III),32 oxalate,33 selenite,34 and humic acids.35 Nico and Zasoski36,37 observed an inhibition of the reductive dissolution of birnessite with Cr(III), sulfide, and hydroquinone in experiments made in the presence of pyrophosphate, and they justified their observations using a model where it is assumed that Mn(III) formed on the (29) Luther, G. W., III. Geochim. Cosmochim. Acta 1987, 51, 31933199. (30) Luther, G. W., III. In Aquatic Chemical Kinetics; Stumm, W., Ed.; Wiley: New York, 1990. (31) Nesbitt, H. W.; Canning, G. W.; Bancroft, G. M. Geochim. Cosmochim. Acta 1998, 62, 2097-2110. (32) Banerjee, D.; Nesbitt, H. W. Geochim. Cosmochim. Acta 1999, 63, 1671-1687. (33) Banerjee, D.; Nesbitt, H. W. Geochim. Cosmochim. Acta 1999, 63, 3025-3038. (34) Banerjee, D.; Nesbitt, H. W. Am. Mineral. 2000, 85, 817-825. (35) Banerjee, D.; Nesbitt, H. W. Geochim. Cosmochim. Acta 2001, 65, 1703-1714. (36) Nico, P. S.; Zasoski, R. J. Environ. Sci. Technol. 2000, 34, 33633367. (37) Nico, P. S.; Zasoski, R. J. Environ. Sci. Technol. 2001, 35, 33383343

Mechanism of Hydrogen Sulfide Oxidation

oxide surface is complexed by pyrophosphate and this complexation inhibits the reaction. Kostka et al.38 synthesized Mn(III)-pyrophosphate complexes and studied their chemical stability under anoxic conditions. They found that the chemical reduction of Mn(III)-pyrophosphate complexes by HS- occurred rapidly, being essentially complete in seconds. We could not detect Mn(III) by either SWV or UV-vis spectroscopy even in experiments done in the presence of pyrophosphate, suggesting that the Mn(III) concentration must be less than the detection limit of both analytical methods or that the chemical reduction of Mn(III)-pyrophosphate by HS- occurred very fast in our experimental conditions. The surface complex, tMnIIS, formed in Scheme 1, eq 3 could follow two different and competitive reaction pathways leading either to S8 or sulfate formation. S8 formation could be explained through a zero-valent sulfur intermediate as it was suggested by the electrochemical experiments. This zero-valent sulfur, S(0), is released from the surface into the solution (Scheme 1 Equation 4) followed by a new surface site generation as soon as a reduced metal ion is released (Scheme 1 Equation 5). The S(0) formed could enter into the solution and form S8 precipitate. Since pH and H2S partial pressure were kept constant during the whole reaction and H2S concentration was varied between 10-3 to 4 × 10-3 M, the HS- concentration was maintained at high values range then all S(0) formed would be fast converted into S8 even at low pH value (see pH 4 at Figure 6). S8 formation via polysulfides could not be discarded (Scheme 1, eqs 6 and 7) despite polysulfides not being detected (see Experimental Section) because they are unstable at the studied pH range. Hoffmann21 has explored the kinetics of the H2S and H2O2 reaction. In his careful study over a wide pH range, he found the formation of elemental sulfur as the major end product of the oxidation with smaller amounts of sulfate. He also noted the transient existence of polysulfide ions, especially at neutral and higher pH. He postulated the formation of a HSOH intermediate that can react with HS- to form polysulfides. Luther,30 using a frontier molecular orbital model, suggested the formation of zero-valent sulfur, S(0), as an intermediate for the same reaction system (H2S and H2O2), which would react with HS- to form polysulfides (see Scheme 1, eqs 6 and 7). In this alternative mechanism, there is no S-O bond formed because this bond is quite stable its formation should lead to thiosulfate, sulfite, or sulfate formation rather than polysulfides species. According to Luther,30 the MnO2 reduction by H2S should follow a similar reaction pathway to the one described for H2S oxidation with H2O2. In MnO2 reduction, there should be a transfer of two electrons from any of the sulfide p orbitals not bound to hydrogen to the Mn(IV) ion at the surface of the crystal, which would result in an electron being accepted by each of the degenerate eg orbitals on Mn(IV). This electron transfer (σ to σ) requires an inner-sphere mechanism. Polysulfide formation was observed during the reaction of H2S with hydrogen peroxide21 but not with MnO2, as was discussed before. A possible explanation for the absence of polysulfides in solution in our system should be due to the difference in reactivity between both reaction systems and the fact that the polysulfide formation is favored at high pH values (Table 1), being quite unstable at the studied pH range. (38) Kostka, J. E.; Luther, G. W., III; Nealson, K. L. Geochim. Cosmochim. Acta 1995, 59, 885-894.

Langmuir, Vol. 19, No. 23, 2003 9689 Table 1. Polysulfide Formation Constants at I ) 0 and T ) 25 °C reaction

log K

reference

+ HS- T S22- + H+ 2- + H+ 2/ S 8 8,rom + HS T S3

-14.88 -11.56 -13.19 -9.46 -9.74 -9.59 -9.50 -9.79

39 39 40 39 40 39 40 40

1/

8S8,rom

3/

8S8,rom

+ HS- T S42- + H+

4/

8S8,rom

+ HS- T S52- + H+

5/

8S8,rom

+ HS- T S62- + H+

Another possible pathway for S8 formation could be a diffusion-controlled process. In this case, the overall energy transfer is diffusion-controlled and eqs 6 and 7 of Scheme 1 should be changed to the following equation: S + 7S f S8 with a rate of R ) k6′. However, the slow step is not the second-order transfer, itself, but the diffusion of sulfur atoms toward each other. According to this, because k6′ is a diffusional rate constant, its value is of the order of 108-109 M-1 s-1.41 Unfortunately, our results do not allow us to rule out either of these two pathways. S(0) could be further oxidized to sulfate; again, sulfate formation could happen through two different indistinguishable pathways. The first one involves migration of sulfur species on the oxide surface to neighboring nonreduced sites with the oxidation of the sulfur species, which transfer two electrons to the metal center, on the new site until sulfate is formed and then released to solution (Scheme 1, eq 8). Each migration-oxidation step is followed by the generation of a new surface site (Scheme 1, eq 5). The second one involves the desorption of sulfur species and resorption from solution on a nonreduced site elsewhere on the oxide surface followed by a two-electron transfer from the newly sorbed species to the Mn(IV). This desorption-resorption-oxidation pathway goes on until sulfate is formed and finally released into solution; once again, for each oxidation step a new surface site is formed. Then, in both cases elemental sulfur is readsorbed before release in a fast process on a new surface site of the manganese oxide, and a total of four tMnIIOH2 surface sites are formed for each sulfate released into the solution. This process must be pH-dependent because the sulfate formation is enhanced at low pH (Figure 6). The only intermediate that we could detect was S(0), although some other intermediates for H2S oxidation with different oxidants have been proposed or found by other authors. The sulfur-surface complex formed in Scheme 1, eq 8 transfers two electrons to Mn(IV) and probably is released into the solution as HSOH. The HSOH formation was previously proposed by Hoffmann21 as an intermediate of the oxidation of H2S by H2O2. Capitani and Di Toro42 calculated for this reaction the free energy for the intermediate formation using theoretical approximations. These authors conclude that this mechanism should be possible because the calculated free energy is favorable. The S-O bond is quite stable, and its formation should lead to thiosulfate, sulfite, or sulfate formation rather than formation of polysulfides species (see below). (39) Giggenbach, W. Inorg. Chem. 1974, 13, 1724-1730. (40) Cloke, P. L. Geochim. Cosmochim. Acta 1963, 27, 1265-1298. (41) Katakis, D.; Gordon, G. In Mechanism of Inorganic Chemistry; Wiley: New York, 1987; p 72. (42) Capitani, J. F.; Di Toro, D. M. Preprints of Extended Abstracts, Computational Methods in Environmental Chemistry Simposia 219th ACS National Meeting; American Chemical Society: Washington, DC, 2000; Vol. 40, p 314.

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Herszage and dos Santos Afonso

Later on, HSOH is adsorbed on a new surface site and oxidized to another sulfur compound, likely S(II) species. Vairavamurthy and Zhou43 studied the oxidation of Na2S with oxygen catalyzed by Ni(II), and they determined the products and intermediate oxidation states for sulfur oxidation using X-ray adsorption near-edge spectroscopy and Fourier transform infrared spectroscopy. They found an intermediate with an oxidation state of +2 and suggested a symmetric structure for it (SO22-). Other structures, such as sulfoxilate (HOSO-) or sulfinate [HS(O)O-], were not considered because their IR spectra should be different from those of their experimental results. The SO22- intermediate is stable only at a high pH (pH ) 11.5-12), and at a pH slightly alkaline or acidic, it is very reactive. Perhaps the SO22- intermediate is also formed during the oxidation of HSOH by the surface through successive steps with two-electron transfer from sulfur to the superficial manganese. The oxidation process continues to sulfate through previous sulfite formation. Sulfite was not detected as a product because Mn(IV) oxides are very reactive in the presence of sulfite.24 However, our experimental results do not allow us to decide whether partially oxidized species migrate on the oxide surface while they are oxidized to sulfate or they are released into solution between each oxidation step. According to the mechanism, the rate of metal ion release is proportional to the reductant surface coverage, and considering the formation of surface complexes as a pre-equilibrium step and using the steady-state approximation for the intermediate species, the formation rate of Mn2+, in terms of the surface species tMnIVOH, can be derived as:)

R)

[

IV d[Mn(II)] k1k2{tMn OH}[HS ] 1+ ) dt k2 + k-1 +

3k8{tMn OH}[H ] IV

k6[HS-] + k8{tMnIVOH}[H+]

]

IV d[Mn(II)] 4k1k2{tMn OH}[HS ] ) dt k2 + k-1

IV d[Mn(II)] k1k2{tMn OH}[HS ] ) dt k2 + k-1

(9)

(10)

On the other hand, at neutral or an alkaline pH, the main oxidation product is elemental sulfur (see Figure 6) and also in this case the experimental initial rate is of first order on the manganese oxide and sulfide concentrations (Figures 3-5). The rate law obtained from the proposed mechanism is also similar to the experimental law (eq II) when k8{tMnIVOH}[H+]/k6[HS-] , 1: (43) Vairavamurthy, A.; Zhou, W. In Geochemical transformation of Sedimentary Sulfur; Vairavamurthy, M. A., Schoonen, M. A. A., Eds.; ACS Symposium Series; American Chemical Society: Washington, DC, 1995.

(11)

It can be seen from eqs 10 and 11 that there is a ratio of 4:1 between the rates in the two limit pH conditions. The dependence of the rate constant on the pH is shown in Figure 9. In addition, as was suggested previously,7-8,44 this dependence can be explained considering that the surface complexes proposed (Scheme 1, eq 1) can also suffer acidbase dissociation to form protonated or deprotonated surface complexes such as tMnIVS- and tMnIVSH, which are linked by the corresponding acid-base equilibrium constant K. Each of these surface complexes have a different reactivity given by the rate constants k2(S-) and k2(SH), but the formation rate of Mn(II) can also be derived for both surface species. The overall rate equation should be the sum of the each of partial contributions and can be written as

R ) k2(S-){tMnIVS-} + k2(SH){tMnIVSH} (12) The concentration of both the surface species can be related to the total concentration CT of the surface species using their distribution coefficients Ri: CT ) {tMnIVS-} + {tMnIVSH}; CT ) {tMnIVS-}/R(S-); and CT ) {tMnIVSH}/R(SH). From the above discussion and taking into account the ratio of 4:1 of the reaction rate at an acidic pH to that at a basic pH, it is clear that the rate should be affected by a term with a pH dependence such as

1+

where {tMnIVOH} is the concentration of the surface active sites. Note that the term k6[HS-] corresponds to the rate of the elemental sulfur formation, while k8{tMnIVOH}[H+] means the rate of the sulfate formation. At a low pH, the main oxidation product is sulfate (see Figure 6), then k8{tMnIVOH}[H+] . k6[HS-], and the experimental initial rate is first-order regarding to the manganese oxide and sulfide concentrations (Figure 3). The rate law obtained from the proposed mechanism is similar to the experimental law (eq II) when k8{tMnIVOH}[H+] . k6[HS-]:

R)

R)

3[H+]

(13)

b + [H+]

where b is an adjustable parameter. This term would reach a maximum value of 4 when sulfate is the main product and a minimum value of 1 when the main product is elemental sulfur. The overall rate will be the following expression:

[

R ) CT[k2(SH)RMnSH + k2(S-)RMnS-] 1 +

3[H+] b + [H+]

]

(14)

Note the similarity among this expression and the rate law calculated from the proposed mechanism (eq 9). Using the distribution coefficients at different pH values, the fitting parameters were calculated (Table 2). The fit was done using the nonlinear curve fit tool included in Origin 5.0 software (OriginLab software). The Levenberg-Marquardt routine was used to minimize the value of χ2, which was taken as the convergence criteria. Different initial scenarios with different values for the fitting parameters were used to verify the uniqueness of convergence. There were four fitting parameters, where k2(SH) and k2(S-) are the linear coefficients of the rate, K is the superficial complex acid-base equilibrium constant for tMnIVSH T tMnIVS- + H+, and b is an adjustment parameter. The calculated rate using the fitting parameters is in good agreement with the experimental results (Figure 9). (44) Amirbahman, A.; Sigg, L.; von Gunten, U. J. Colloid Interface Sci. 1997, 194, 194-206.

Mechanism of Hydrogen Sulfide Oxidation

Langmuir, Vol. 19, No. 23, 2003 9691

Figure 9. Dependence of the second-order rate constant on pH at T ) 25 °C, I ) 0.01 M NaClO4, [H2S] ) 2.0 × 10-3 M, and {MnO2} ) 1.00 × 10-3 M. The points are the experimental values for (2) MnO2(A) and (9) MnO2(B) and the solid lines were calculated using the fitting parameters from Table 2. Table 2. Fitting Parametersa

oxide

k2(HS) × 102 (s-1)

k2(S-) × 102 (s-1)

b (M)

K 10-6

MnO2(A) 28.0 ( 0.7 6.7 ( 1.7 (6.3 ( 0.3) × (4 ( 1) × 10-6 MnO2(B) 7.0 ( 0.4 5.8 ( 1.2 (5.3 ( 1.0) × 10-7 (6 ( 1) × 10-6 a

These parameters described the rate constant dependence on pH shown in Figure 9.

the acidity of the HS group. Similar results were observed in the reaction of iron oxides with several reductants (ref 44 and references therein). The values of fitting parameters also suggest that the neutral species is the more reactive species, in agreement with the experimental data because the dissolution reaction is faster at a low pH, where the ionic species is negligible. Similar results were found for hematite3 with k2(SH) ) 11.1 × 10-2 s-1 and k2(S-) ) 8.3 × 10-3 s-1. Comparing the kinetic constants is possible to observe that manganese oxide dissolution is faster than the reaction of Fe(III) with H2S, as Luther suggested before.30 Because different manganese oxide phases have different crystal structures, the geometry of their reactive sites will differ as well. The surface site-binding model proposed for the actual reduction process involves the binding of the reductant to a site on the manganese oxide surface prior to electron transfer. Therefore, differences in the surface properties of different phases (reactive site density, activation energy of complex formation at different types of sites) should have the greatest effect on the first of these two steps. Thus, the relative rates of these steps will determine whether differences in mineral phases will significantly affect the overall rate of reduction and dissolution by a given substrate. Finally, from the proposed mechanism, it is also possible to obtain a mathematical expression for the oxidation product formation rates (elemental sulfur and sulfate):

R)

k8{tMnIVOH}[H+]

The pK of tMnIVSH is close to the pK’s of similar systems such as tZnIISH+ 45 and tFeIIISH.3

tMnIVSH T tMnIVS- + H+ tMnIVSH T tMnIVS- + H+

pK ) 5.40 for MnO2(A) pK ) 5.22 for MnO2(B)

tFeIIISH T tFeIIIS- + H+

pK ) 5.52 for Hematite

tZnIISH+ T tZnIIS + H+

pK ) 6.91 for Synthetic ZnS

tZnIISH+ T tZnIIS + H+

pK ) 7.10 for Sphalerite

The pK values would depend not only on the difference between the crystal structures of the solid phases and their lattice stabilization energies but also on the acidity of the metal centers. The acidity of the metal centers was assessed using a qualitative model based on the chargeto-radius ratio of the metal (z+-/r, where z+- is the metal charge and r is the radius). The 0.54, 0.64, and 0.74 values were obtained for Mn(IV), Fe(III), and Zn(II), respectively. Then, it is reasonable to assume that the pK of tMnIVSH is somewhat lower than that of tFeIIISH and tZnIISH+. It can be noticed that the equilibrium constant for these processes are several orders of magnitude higher than the corresponding value of Ka2 for H2S, and this would be due to the fact that the sulfide ion is bound to the metal center, forming a surface complex, which would increase (45) Ro¨nngren, L.; Sjo¨berg, S.; Sun, Z.; Forsling, W.; Schindler, P. W. J. Colloid Interface Sci. 1991, 145, 396-404.

[

IV d[S0] k1k2{tMn OH}[HS ] ) 1dt k2 + k-1

k6[HS-] + k8{tMnIVOH}[H+] R)

]

d[SO42-] ) dt k1k2k8{tMnIVOH}2[HS-][H+] (k2 + k-1)(k6[HS-] + k8{tMnIVOH}[H+])

(15)

(16)

From the eqs 9, 15, and 16, it is possible to obtain the following product concentration ratio:

[Mn(II)] [SO42-]

)

[S0] [SO42-]

+4

(17)

Figure 10 shows that this relationship, obtained from the proposed mechanism, is valid for our experimental data, where this has a linear profile with a slope of 0.999, an origin ordinate of 4, and a correlation factor (r2) of 0.99. Some data from Yao and Millero7,8 are also included in Figure 10. Yao and Millero did not report the amount of Mn(II) formed during the dissolution process; Mn(II) concentrations were calculated assuming the stoichiometry [Mn(II)] ) 4[SO4-] + 4[S2O32-] + [S0]. It is possible to observe that these data are in good agreement with the expected values calculated from eq 17. The proposed mechanism, the rate equation, and their combination seem to be good tools to explain the experimental oxidation product distribution and the ratio of the reduction to the oxidation products for samples from natural environments, such as those mentioned in these papers. Possible oxidants for hydrogen sulfide in natural waters would be oxygen or metal oxides such us iron or manganese oxides. Then, Fe(III) and Mn(IV) oxyhydroxides may also contribute to the reoxidation of hydrogen sulfide and FeS

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Figure 10. Ratio between total dissolved manganese and SO42concentrations as a function of the ratio of the S(0) and SO42concentrations at pH ) 7.0, T ) 25 °C, and I ) 0.01 M NaClO4. The slope, origin ordinate, and correlation factor are 0.999, 4, and 0.99, respectively. (2) This work and (9) data from Yao and Millero.7,8

in the chemocline of modern anoxic basins46-49 and nearsurface sediments.50-52 In the sedimentary system, dissolved Mn(II) is much more mobile than Fe(II), leading to a flux of Mn(II) from reduced sediments into the bottom waters, whereas iron is easily precipitated in the anoxic part as iron sulfide.53 However, the pool of Mn(IV) in freshwater sediments is usually much smaller than that of iron(III) oxides. The much higher turnover rates of manganese54 may also lead to a more pronounced importance for the interaction with the biogeochemical sulfur cycle. The relative rates may additionally be influenced by the reactivity of the sedimentary oxides, which is related to their mineralogy (ref 55 and references therein). It is, (46) Jacobs, L.; Emerson, S.; Skei, J. Geochim. Cosmochim. Acta 1985, 49, 1433-1444. (47) Millero, F. J. Limnol. Oceanogr. 1991, 36, 1007-1014. (48) Millero, F. J. Estuarine, Coastal Shelf Sci. 1991, 33, 521-527. (49) Millero, F. J. Deep Sea Res. 1991, 38, S1139-S1150. (50) Canfield, D. E.; Thamdrup, B. Hannsen, J. W. Geochim. Cosmochim. Acta 1993, 57, 2563-2570. (51) Aller, R. C. J. Mar. Res. 1994, 52, 259-295. (52) Moeslund, L.; Thamdrup, B.; Jørgensen, B. B. Biogeochemistry 1994, 27, 129-152. (53) Calvert, S. E.; Peterson, T. F. Econ. Geol. 1996, 91, 36-47. (54) Thamdrup, B.; Fossing, H.; Jørgensen, B. B. Geochim. Cosmochim. Acta 1994, 58, 5115-5129. (55) Thamdrup, B. Adv. Microbial. Ecol. 2000, 16, 41-84.

Herszage and dos Santos Afonso

therefore, expected that in coastal and estuarine sediments with high biological activity and a very dynamic sulfur and manganese cycle56-58 the influence of manganese oxides may be increased. The hydrogen sulfide concentration range used in this study is realistic in terms of the concentration of the hydrogen sulfide commonly found in natural environments. In the environment, the hydrogen sulfide probably diffuses both up and down in the sediment. Urban59 pointed to the fact that all the measured sulfate reduction rates in freshwater indicated a much higher turnover of sulfate than would be predicted by the calculation of diffusion fluxes from the concentration gradients. More than 50% of the sulfate reduction occurs below a depth of 2 cm, where diffusion gradients are negligible. Urban also concluded that only sulfate regeneration resulting from the reoxidation of hydrogen sulfide can explain sulfate reduction rates as high as those found under marine conditions, despite the low sulfate concentration in freshwater systems. Manganese oxides were shown to oxidize sulfide very easily, and the fact that the oxidized and reduced forms of manganese generally occur at different phases not only suggests a probable mechanism for the redox reaction in nature but also indicates how this reaction may be coupled to electron flow between aerobic regimes and reduced sulfide. The anaerobic sulfide oxidation associated with bacteria in the presence of MnO2 was broadly studied,60 and to understand the mechanisms of bacterial reactions, it is necessary to know the mechanisms of the abiotic reactions. Then, the present experimental findings have strong implications for understanding sulfur and manganese cycles on the sedimentary environments, as well as figuring out the coupled mechanisms of bacterial reactions. Acknowledgment. The authors acknowledge Universidad de Buenos Aires, Secretarı´a de Ciencia y Te´cnica, for financial support of this work through Project Nos. UBACyT Ex-037 and TW99. The authors also thank to Prof. Dr. George Luther, III, for helping us with the SWV measurements. LA034016P (56) Canfield, D. E.; Thamdrup, B. FEMS Microbial. Ecol. 1996, 19, 95-103. (57) Huettel, M.; Ziebis, S.; Forster, S.; Luther, G. W., III. Geochim. Cosmochim. Acta 1998, 62, 613-631. (58) Kristensen, E.; Bodenbender, J.; Jensen, M. H.; Rennenberg, H.; Jensen, K. M. J. Sea Res. 2000, 43, 93-104. (59) Urban, N. R. In Environmental Chemistry of Lakes and Reservoirs; Baker, L. A., Ed.; Advances in Chemistry Series: American Chemical Society: Washington, DC, 1994. (60) Bo¨ttcher, M. E.; Thamdrup, B. Geochim. Cosmochim. Acta 2001, 65, 1573-1581.