Oxidation of Cysteine and Glutathione by Soluble Polymeric MnO2

Jun 26, 2003 - Oxidation of Methionine by Colloidal MnO 2 in Aqueous and Micellar Media: A Kinetic Study. Mohammad Altaf , Deogratius Jaganyi. Journal...
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Environ. Sci. Technol. 2003, 37, 3332-3338

Oxidation of Cysteine and Glutathione by Soluble Polymeric MnO2 J U L I AÄ N H E R S Z A G E A N D M A R IÄ A D O S S A N T O S A F O N S O * 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 GEORGE W. LUTHER, III College of Marine Studies, University of Delaware, Lewes, Delaware 19956

The kinetics of reduction of soluble polymeric MnO2 by cysteine and glutathione has been studied in the pH range of 4.0-9.0. The concentration of thiols was varied between 1 and 2 mM, while the MnO2 concentration was varied between 2 and 12 µM. In this pH range, the reaction products were identified as Mn(II) and the corresponding disulfides (cystine and glutathione disulfide). Cysteic or cysteinesulfonic acid was formed only when pH < 2. Experimental data indicate that the rate law over the pH range of 4-9 is first-order in both MnO2 and thiol concentration. Eyring plots for both thiols reacting with MnO2 indicate that the reaction is associative (∆Sq ∼ -160 J mol-1 K-1) and proceeds via an inner-sphere redox process. The reaction proceeds via the formation of two different inner-sphere complexes tMnIVSR- and tMnIVSR and their further reaction to products. Both surface species are linked to each other via acid-base equilibria, and the rate constant decreases as pH increases. The presence of two ligand surface species is determined using surface complexation modeling. A reaction mechanism in agreement with the experimental results is proposed.

Introduction Thiols are widespread in the environment. Many have been identified and quantified in a variety of samples from suboxic and anoxic waters and sediments, using electrochemical techniques (1-4) and/or chromatographic techniques (3, 5-9). Particularly in sediments, thiols are found in the nanoto micromolar concentration range (2, 3, 5). Thiols are suspected to affect the bioavailability of essential trace metals (e.g., Cu and Zn) in sulfidic environments where they are strongly bound as solid sulfide minerals (3). Among the nonvolatile thiols reported in the literature, cysteine and glutathione are two of the most frequently found. Glutathione is a tripeptide formed by the amino acids glycine, cysteine, and glutamic acid and is thought to be the most abundant nonprotein thiol in animals, plants, and several bacteria (10, 11). Glutathione is a redox cofactor for proteins involving disulfide bonds (12, 13), works as a cysteine repository, and is possibly involved in the S0 metabolism (12). The main * Corresponding author fax: +5411-4576-3341; e-mail: dosantos@ qi.fcen.uba.ar. 3332

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sources of cysteine, an essential amino acid for protein synthesis, are microbial protein degradation (14) and assimilatory sulfate reduction (15). In aquatic environments, hydrogen sulfide can be readily oxidized by manganese and/or iron oxides, either biotically or abiotically, leading to the formation of several oxidized sulfur species such as sulfite, thiosulfate and polysulfides (16-21). Although the reductive dissolution of different iron oxides with several thiols has been studied in water (22-25), this is not the case for the manganese oxides. To the best of our knowledge, the only report for this kind of reaction in the aqueous phase is included in the study of the reactivity of manganese(III/IV) oxides toward many different organic molecules done by Stone and Morgan (26). The authors used thiosalicylic acid as an example of a thiol and only reported an apparent second-order rate constant. There are a few other reports in the literature for the reductive dissolution of manganese oxides with thiols in organic solvents (27-29) and with metallic oxides other than manganese and iron oxides, such as PbO2, CrO3, and Co2O3 (30). The study of the reductive dissolution of manganese oxides with thiols in water could contribute to a deeper and more complete knowledge of both the Mn and S biogeochemical cycles. Luther and Church (31) proposed that thiols could be oxidized to disulfides and even to sulfonic acids by reaction with manganese and or iron oxides. Sulfonates have been confirmed in sediments in an XANES study (32). According to the results reported for the reaction of thiols with iron oxides in water, the only oxidation products are the corresponding disulfides (24, 25). The same products were reported in the oxidation with manganese oxides in organic solvents (27-29). In this work the study of the reduction of soluble polymeric manganese(IV) dioxide by cysteine and glutathione has been undertaken to find a rate law, identify the reaction products, and elucidate the elementary steps involved in this process over the pH range of 4.0-9.0.

Experimental Section The soluble manganese dioxide used was a stable colloidal suspension, which we term polymeric. It was prepared following the technique described by Perez-Benito et al. (33) by mixing the appropriate amounts of KMnO4 (Aldrich) and Na2S2O3 (Aldrich) stock solutions, according to the following stoichiometry:

3S2O32- + 8MnO4- + 2H+ f 8MnO2 + 6SO42- + H2O Although this soluble phase is unlikely to exist in marine environments for extended periods, it is likely present in freshwater systems (34, 35). The polymeric manganese dioxide was electrochemically inactive (35) but showed a large absorption band covering the whole visible region of the spectrum with absorbance uniformly increasing with decreasing wavelength that gave a broad maximum at 300400 nm (Figure 1). Dilutions of this polymeric oxide followed the Lambert-Beer law in the concentration range used (400nm ) 1.77 × 104 M-1 cm-1). The UV-Vis spectra were performed using a Hewlett-Packard 8453 diode array spectrophotometer. The stoichiometry of the oxide was determined to be MnO2 (average of three replicates) by iodometric techniques and flame atomic absorption spectroscopy using a Varian AAS5 instrument for total Mn measurement. The point of zero charge (pzc) of the oxide was determined to be 1.93 by potentiometric titrations using a 716 Metrohm Titrino equipped with a Metrohm pH glass combination electrode. 10.1021/es0340634 CCC: $25.00

 2003 American Chemical Society Published on Web 06/26/2003

FIGURE 1. UV-Vis spectra of soluble polymeric MnO2. This polymeric phase is easy to syntethize, is stable for several months, is electrochemically inactive, and has been previously characterized (33). Some other properties of the polymeric oxide were previously reported (33, 36, 37): Negative particles roughly spherical, 90% of the particles had a diameter in the range of 89-193 nm, most of them had a diameter of 120 nm, and ζ-potential was -47( 3 mV (33). Thiol stock solutions were prepared fresh daily from the corresponding solids, cysteine (CSH) (Aldrich) and glutahione (GSH) (Merck), using degassed distilled deionized water (DDDW) and were standardized using Ellmann’s reagent (38). Oxidation-reduction experiments were performed under a nitrogen atmosphere. The temperature was held constant to ( 0.1 °C with a constant temperature circulation bath. The pH was held constant between 4.0 and 9.0 using the appropriate buffer solutions (acetate, borate, and Tris). In all the experiments, the ionic strength was 0.02 M using NaClO4. In one set of experiments designed to evaluate the influence of thiol concentration on the rate law, thiol concentration was varied between 1 and 2 mM, while MnO2 concentration was held constant. In another set of experiments, the MnO2 concentration was varied between 2 and 12 µM, while thiol concentration was held constant. Under the experimental conditions used, the reaction was completed in less than 1 min. The course of the reaction was followed using a stopped-flow system coupled to a HP 8451 UV-Vis diode array spectrophotometer set to follow the decay in absorbance at 400 nm due to MnO2 consumption. The absorbance versus time data were fitted using the data analysis package included in the spectrophotometer software. The resulting rate constants for different experimental conditions are the average of at least three different runs. Square-wave voltammetry (SWV) was used in separate experiments to evaluate the formation of Mn(II), Mn(III) (35, 39), and cystine (cysteine disulfide) (1, 35) using an Analytical Instrument Systems, Inc. DLK 100 potentiostat with a EG&G PAR 303 static dropping mercury electrode stand (DME). The electrode stand was modified to use a saturated calomel electrode (SCE) rather than the Ag/AgCl reference supplied. The SCE electrode was connected to the solution through a NaCl salt bridge. Instrumental parameters for the SWV mode were typically 200 mV s-1 scan rate over the potential range of -0.1 to 1.5 V with a 24-mV pulse height, which gave a detection limit lower than 1 µM. Because pyrophosphate forms stable complexes with Mn(III) at near neutral or slightly alkaline pH (39), separate electrochemical and spectrophotometric experiments were done in the presence of pyrophosphate (pH 8) in order to detect the formation of Mn(III). Glutathione disulfide (GSSG) was identified and quantified according to the HPLC technique described by Winterbourn and Brennan (40) using a Shimadzu LC-6A instrument with

FIGURE 2. Initial rate (R0) of the reduction of MnO2 by thiols for different [MnO2] concentrations: cysteine (2) and glutathione (9), pH 4.1; T ) 25 °C; I ) 0.02 M, [thiol] ) 1 mM. a Shimadzu SPD-6AV UV-Vis detector. The samples were injected into a Rheodyne injector 77251 with a 50-µL loop and a J&W Scientific ACCU ODS 5-µm analytical chromatographic column. The retention time for GHS and GSSG under these operational conditions was 4.72 and 21.49 min, respectively. The analytical precision was usually within 1%. In experiments performed below pH 2, cysteine was totally oxidized to an electrochemically inactive compound identified as cysteic (cysteinesulfonic) acid. The amount of cysteic acid formed was evaluated by ion chromatography using a DIONEX DX-100 instrument with a conductivity detector, a sample injection valve, and a 25-µL sample loop. Two plastic anion columns were coupled in series to serve both as precolumn (DIONEX AG-9) and 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 HCO3-/CO3-2 4 mM was chosen as eluent with a flow rate of 1 mL/min. The retention time for cysteic acid under these operational conditions was 3.92 min. Standard cystine and cysteic acid solutions were prepared by dissolving the corresponding amount of solid (used as received from Sigma) in DDW. Glutathione disulfide standards were prepared according to Luther et al. (1) by oxidizing a standardized solution of glutathione with I2. Mn(II) standard solutions were prepared by diluting a 1000 ppm Mn(II) AAS standard (Fluka) in acidified distilled deionized water. Errors for the k values were obtained from the fitting tool included in the HP8453 (diode array spectrophotometer) software package. Errors for the fitting parameters were obtained using the nonlinear curve fit tool in Microcal Origin software using the Levenberg-Marquardt fitting routine.

Results and Discussion Experiments performed in the absence of thiols showed no change in the MnO2 UV-Vis spectra during the time scale of the kinetic experiments, indicating that there is no reaction of the MnO2 with the buffers used to keep pH constant. Similar results were obtained when cysteine was replaced by glycine, indicating that only the sulfhydryl group of cysteine is oxidized. At a given pH value and thiol concentration, the initial rate is proportional to the amount of MnO2 added (Figure 2). At constant amount of manganese oxide added and pH, a linear correlation between the initial rate and the thiol concentration is observed (Figure 3). The experimental rate law is expressed as follows:

Ro ) k[thiol][MnO2]T VOL. 37, NO. 15, 2003 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 3. Initial rate constant (k′) of reduction of MnO2 by thiols for different [thiol] concentrations: cysteine (2) and glutathione (9), pH 4.4; T ) 25 °C; I ) 0.02 M, [MnO2] ) 20 µM.

FIGURE 5. Variation of the rate constant with temperature. (A) cysteine, pH 7, I ) 0.02 M; (B) glutathione, pH 5.4, I ) 0.02 M.

TABLE 1. Ea, ∆Hq, and ∆Sq for Cysteine and Glutathionea

cysteine glutathione a

FIGURE 4. Rate constant for the reduction of MnO2 by thiols at different pH values for cysteine (2) and glutathione (9); T ) 25 °C; I ) 0.02 M; [thiol] ) 1 mM; [MnO2] ) 20 µM. The fact that the reaction rate for glutathione is lower than that for cysteine can be attributed to steric hindrance, since glutathione is a tripeptide formed by the amino acids glutamate, cysteine, and glycine. On the time scale used in the experiments, only the SH group is oxidized by MnO2. Figure 4 gives the dependence of the rate constant with pH at 25 °C and demonstrates that the rate increases with decreasing pH. These data indicate that the protonated forms of the thiols react more quickly with the anionic polymeric MnO2, which has a pzc of 1.93. Between pH 4.0 and pH 9.0, the oxidation products were identified as the corresponding disulfides: cystine (CSSC) and glutathione disulfide (GSSG). Similar products were obtained in the oxidation of different thiols with MnO2 in organic solvents (27-29) and in the oxidation of thiols with iron oxides in water (22-25). At pH values below 2, cysteine was completely oxidized to an electrochemically inactive product that was identified as cysteic (cysteinesulfonic) acid using ionic chromatography. The method used to quantify GSSG allows for the simultaneous determination of glutathionesulfonic acid (GSO3H) as well (40). No signal characteristic of GSO3H was observed in experiments done at pH 4 and above. Considering the results for cysteine, a lower pH is also apparently needed to oxidize GSH to GSO3H. The dependence of k with temperature was in agreement with the Arrhenius equation (k ) Ae-Ea/RT) where A is the preexponential factor, R is the gas constant, T is the absolute temperature, and Ea is the overall activation energy (Figure 3334

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pH

Ea (kJ mol-1)

∆Hq (kJ mol-1)

∆ Sq (J mol-1 K-1)

7.0 5.4

17.4 ( 0.1 19.0 ( 0.1

15.0 ( 1.0 16.6 ( 1.0

-156 ( 15 -160 ( 15

In both cases I ) 0.02 M.

5). The overall activation energy value obtained for cysteine was (17.4 ( 0.1) kJ mol-1 at pH 7 and I ) 0.02 M. For gluthatione, the overall activation energy value obtained was (19.0 ( 0.1) kJ mol-1 at pH 5.4 and I ) 0.02 M. The overall activation energy for the rate constant k will be a combination of the activation energies for elementary reactions. Preexponential factors for bimolecular reactions commonly exhibit A values between about 107 and 1012 M-1 s-1. The dissolution rate of polymeric MnO2 has a preexponential factor of 1.1 × 105 and 7.5 × 104 M-1 s-1 for cysteine and glutathione, respectively. The preexponential factor values lower than normal are showing the influence of steric effects due to the molecule size and the presence of substituents. The different values of A between both thiols would depend on the steric effects. Consequently, this is due to the tripeptide structure of glutathione where the large size of the glutathione molecule may contribute to some steric effects that increase the activation energy while decreasing the preexponential factor and the dissolution rate. To determine whether the reaction mechanism is associative or dissociative, the entropy of activation is needed (42), which can be evaluated from an Eyring plot [ln(k/T) versus 1/T ] based on

ln(k/T) ) (∆Hq/RT ) + ln(k/h) + (∆Sq/R)

(II)

where ∆Hq is the enthalpy of activation, ∆Sq is the entropy of activation, k is the Boltzmann constant (1.381 × 10-23 J K-1), and h is the Planck’s constant (6.626 × 10-34 J s). The slope (∆Hq/R) provides the enthalpy of activation, and the intercept {ln(k/h) + (∆Sq/R)} yields the entropy of activation. A large negative ∆Sq indicates an associative reaction whereas a large positive ∆Sq indicates a dissociative reaction (41). All kinetic activation parameters are given in Table 1, which shows that both thiols have a large negative ∆Sq for MnO2 reduction and that the reactions are associative and proceed via inner-sphere redox reactions. The values of activation energy, enthalpy, and entropy of activation for a given reaction are usually indicative of some

effects such as van der Waals, dipole-dipole, H-bonding, ionic interactions between charged species, and also specific bonding of reactive moieties. The processes that can be explained via van der Waals, dipole-dipole or other weak intermolecular forces are occurring by an outer-sphere mechanisms and are characterized by a very low activation energy, while specific bonding of reactive moieties occur by inner-sphere mechanisms such as the one already mentioned here and require more higher activation energies. 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-determing step, there should be no difference for the dissolution rate between both 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 transport of reduced species away from the surface (Mn(II) release into solution is not the rate-determing step), and the reactions at the surface (which are responsible for the activation energy) must be rate-controlling (42). Therefore, for both ligands, the dissolution would be surfacecontrolled. It has been found that comparatively simple rate laws are obtained if the observed rates are plotted against the concentrations of the adsorbed species and surface complexes (43, 44). Then, a possible reaction mechanism, considering Mn(II) and the disulfides as the main products, should go through the following steps: (i) Formation of a precursor complex on the soluble polymeric MnO2 (similar to an oxide surface, tMnIVSR): k1

tMnIVOH + RSH {\ } tMnIVSR + H2O k -1

(1)

Mn(IV) has a d3 (t2g3) electronic configuration in octahedral symmetry. Ions with this electronic configuration are inert and have a slow rate of ligand exchange (45). Therefore, this precursor complex is an inner-sphere complex (45), which is consistent with the large negative ∆Sq (Table 1). The formation of such complexes as the first step has been proposed in many reductive dissolution reaction schemes, including the reaction of solid manganese(IV) oxides with sulfide (21, 46) and with organic reductants (26, 47-49). (ii) Electron transfer from the thiol to the Mn (in two steps) followed by release to the solution of sulfhydryl radicals (RS•): k2

tMnIVSR 98 tMnIIISR k3

tMnIIISR + H2O 98 tMnIIIOH + RS• + H+ k4

tMnIIIOH + RSH 98 tMnIIISR + H2O k5

tMnIIISR 98 tMnIISR k6

tMnIISR + H2O 98 tMnIIOH2+ + RS•

(2) (3) (4) (5) (6)

We postulate two separate electron transfers because the simultaneous transfer of two electrons is very unlikely to happen. It should be noted that no Mn(III) was detected either by SWV and/or UV-Vis spectroscopy even in experiments done in the presence of pyrophosphate that stabilizes Mn(III) (39). This appears due to (I) the fast reaction rate with soluble polymeric MnO2, (II) the large excess of thiol to

MnO2, and (III) and the competition between thiols (with carboxylate and amino functional groups) and pyrophosphate for any Mn(III) formed. Nevertheless, there is clear evidence in the literature that Mn(III) is an intermediate in the reductive dissolution of solid-phase MnO2. Nico and Zasoski (50, 51) observed an inhibition of the reductive dissolution of birnessite with Cr(III), sulfide, and hydroquinone in experiments made in the presence of pyrophosphate. On the basis of their observations, they proposed a model where it is assumed that any Mn(III) formed on the oxide surface is complexed by pyrophosphate and this complexation inhibits the reaction. Also, Mn(III) has been detected by XPS as an intermediate in the reductive dissolution of birnessite with several reductants such as arsenite (52), Cr(III) (53), oxalate (54), selenite (55), and humic acids (56). Their results show that as the reaction proceeds Mn(III) is formed on the oxide surface. In the cases where the reductant does not stabilize Mn(III) through complexation [e.g., As(III), Cr(III), Se(IV)], Mn(III) behaves as an intermediate because the profile of Mn(III) concentration versus time has a maximum, i.e., it is formed and consumed in another reaction step (52, 53, 55). The formation of RS• radicals was proposed by Wallace (28) in his study of thiol oxidation with solid-phase MnO2 in organic solvents because, in the presence of olefins, sulfur addition products (addition to the double bonds) were formed. These results are indicative of RS• radical formation. These radicals were also detected by ESR in solutions resulting from the oxidation of CSH and GSH with peroxynitrite or OH radicals (57). Because of the relative stability of sulfur radicals (58), RS• radicals were proposed as intermediates in reaction mechanisms to explain the oxidation of mercaptocarboxylic acids with magnetite (24), the oxidation of thioglycolic acid and glutathione with Mn(III) complexes, and the oxidation of thiols with complexes of iron and another metals Ce4+, Co3+, V5+, and Cu2+ (30). (iii) RS• radicals dimerize to form the corresponding disulfide: k7

2RS• 98 RSSR

(7)

Sulfhydryl radicals might dimerize while the sulfur atoms are still attached to the polymeric MnO2. In that case, the distance between both radicals, which would be equivalent to the Mn-Mn distance on an oxide surface, should be similar to the bond length of the corresponding disulfide. The length for a S-S bond in most molecules having two or more sulfur atoms linked to each other is aproximately 2 Å (59). However, the distances between vicinal Mn atoms in most manganese(IV) oxides is slightly larger than 4 Å (60); therefore, dimerization of radicals while attached to polymeric MnO2 or a MnO2 surface is very unlikely to occur. Another possible dimerization mechanism (24) that cannot be ruled out as possible is that thiols coordinate to the surface through other complexing groups (e.g., carboxylate in cysteine). This second group could work as an anchor allowing the S-S bond to be formed in solution while the thiol is still attached to a MnO2 surface or polymeric MnO2 through the second coordinating group. This would also permit their further oxidation to sulfonates at low pH by allowing contact with soluble polymeric MnO2. The increase in the Ea for the oxidation of gluthathione over that for cysteine suggests that carboxylate ion complexation with MnO2 is most likely. (iv) Desorption of Mn(II) and generation of a new surface site: k8

tMnIIOH2+ 98 tMnIVOH + Mn2+(aq) + H+ VOL. 37, NO. 15, 2003 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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TABLE 2. Fitting Parameters Used in Figure 4 Corresponding to Eq 14 (See Text) for Cysteine and Glutathione parameters

k2(MnSRH)

thiol cysteine glutathione

2060 ( 20 222 ( 17

K °app

k2(MnSR-) 200 ( 2 12 ( 1

10-4

4.4 × ( 0.5 × 10-4 1.2 × 10-4 ( 0.2 × 10-4

Assuming that all the intermediates are in steady state, the initial reduction rate becomes first-order with respect to both reactants (the surface species {tMnIVOH} and the free thiol concentrations [RSH]) (eq 9). Then the stoichiometric rate law has the same shape as the experimental rate law (eq I):

R)

k1k2 {tMnIVOH}[RSH] k-1 + k2

mize the value of χ2, which was taken as a convergence criterium. Different initial scenarios with different values for the fitting parameters were used to verify the uniqueness of convergence. The protonation-deprotonation process for the reaction with cysteine would take place on the nitrogen atom of the amino group as previously proposed by Amirbahman et al. (25) in their study of the reductive dissolution of hydrous ferric oxide with cysteine. The authors suggested the formation of surface complexes where the metal is complexed either by the thiol through the sulfur, or the nitrogen, or the oxygen of the carboxylate groups. These surface complexes are the ones involved in the acid-base equilibrium already mentioned. Assuming that the residual charge on the metal center is +1, some structures for complexes of overall charge 0 and -1 can be postulated:

(9)

In the pH range studied, the rate constant for reduction of polymeric MnO2 was observed to decrease with increasing pH (Figure 4). These results can be explained considering that the species, formed in the first reaction step, are involved in acid-base equilibria characterized by the corresponding K°app value:

tMnIVSRH / tMnIVSR- + H+

(10)

and

K°app )

{tMnIVSR-}[H+] {tMn SRH} IV

(11)

Both species can undergo the reaction steps already mentioned (eqs 1-8), and their different reactivities cause the reaction rate to change with pH (49). Amirbahman et al. (25) successfully used a similar model in their study of the reductive dissolution of hydrous ferric oxide with cysteine. Therefore, the rate of step 2 would be

R ) k2(MnSRH){tMnIVSRH} + k2(MnSR-){tMnIVSR-} (12) A species distribution varying with pH is given by the values of the distribution coefficients (Ri) and total concentration (CT), which are a function of K°app and pH:

{tMnIVSRH} ) CTRMnSRH {tMnIVSR-} ) CTRMnSRCT ) {tMnIVSRH} + {tMnIVSR-}

(13)

Therefore, the rate for step 2 can be written as a linear combination of both distribution factors:

R ) CT[k2(MnSRH)RMnSRH + k2(MnSR-)RMnRS-]

(14)

The above equation accounts for the pH dependence of the overall rate constant k. Figure 4 shows the variation in k for cysteine and glutathione with pH. The symbols are experimental points, and the lines were plotted as the best fit assuming that CT is equal to the total available MnO2 surface sites and using three adjustable parameters in eq 14. These parameters were K°app, k2(MnSRH) and k2(MnSR-) accounting for the differences in reactivity of both surface complexes. Table 2 shows the results for these three parameters for both thiols. As shown in Figure 4, the proposed model is a good fit to the experimental data. The fit was done using the non-linear curve fit tool included in Origin 5.0 software (OriginLab software). Levenberg-Marquardt routine was used to mini3336

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In the neutral complex, the amino group is protonated; in the negatively charged complex, it is deprotonated. There are two possible structures for the latter complex if S binds to the Mn, one with the nitrogen coordinated to the Mn center and the other one with the oxygen coordinated to the metal center. Table 2 and Figure 4 demonstrate that the protonated species are the most reactive species for both thiols. Similar results were obtained for the reaction between solid MnO2 and sulfide (21). Also, the pK°app, corresponding to the protons on the amino groups, are dramatically shifted to more acidic values. Coordination of cysteine with Fe(III) on an hydrous ferric oxide surface causes the amino group proton to be 3 orders of magnitude more acidic than usual, given that Fe(III) acts as a Lewis acid (25). In our study, the shift in pK°app is even higher, which can be attributed to the fact that Mn(IV) is more acidic than Fe(III). In the reaction with glutathione, the group involved in that acid-base equilibria would be the amino group corresponding to the glutamic acid. In both cases, the pK°appvalues followed the same behavior than the free thiols pKa. The surface complexes involved in the acid-base equilibria are more labil when increase the strength of the free thiol. The tendency to form chemical linkages to solid surface atoms correlates with likelihood of forming comparable complexes in solution (61). Likewise, the exchange of protons by a dissociative or associative processes would follow a similar tendency. Unlike the reaction with H2S, the speciation of the oxidation products does not change significantly in the pH range typical of natural waters as the corresponding disulfides are the only products detected in this pH range. Sulfonic acids are formed at much more acidic conditions (pH < 2) such as those found in salt marshes and other oxidizing sediments (2); therefore, the reduction of MnO2 by thiols may not be a significant source of sulfonates in some natural systems. Given their widespread occurrence (32, 62), other mechanisms must also be operational for sulfonate formation such as sulfite and/or thiosulfate addition to R,β-unsaturated organic compounds (31, 32). For soluble polymeric MnO2, the rate constants at pH 10 follow the order ksulfide ) 103 M-1

s-1 (21), kcysteine = 200 M-1 s-1, and kglutathione = 20 M-1 s-1 (kcysteine and kglutathione were obtained by extrapolation of the plots in Figure 4). This rate constants are proportional to the thiol dissociation constants (RSH pKa values of 7.0, 8.18, and 8.66 with RdH, CS, and GS respectivily), the pkthiol versus pKa plot has a slope of 0.94 and R 2 ) 0.89. Then the reactivity decreases in the order HS- > CSH > GSH, reflecting the electronic and steric effects of the thiol substituents on the reaction considered. Although the oxidation of inorganic sulfide with manganese oxides has been studied extensively (16, 21, 46, 63), oxidation of thiols of environmental importance has not received much attention. This lack of information hampers our knowledge of the reaction pathways for several species in the sulfur biogeochemical cycle (31). Polymeric MnO2 may not represent the most typical MnO2 phases present in natural systems, but the results obtained with this polymeric phase are relevant and would contribute to understand the chemical behavior of the anoxic basins. Further work is needed with other solid manganese oxide phases typical of natural systems to further assess the contribution of the studied process to the biogeochemical cycles of Mn and S. As expected, several results have similarity to those obtained in a previous study on the reaction with H2S (21). Redox rate constants increase as pH decreases, and this fact can be explained by considering a model where reactant species, linked to each other via acid base equilibria, are formed. The proposed mechanism involves complexation to polymeric MnO2 (or reductant adsorbing on an oxide surface for solids) as a first step because the reaction is an associative process, followed by several electron transfers and the release of products to the solution. The use of soluble polymeric MnO2 provides the capability of writing a rate law in both reactants and calculating activation parameters (∆Sq, ∆Hq, ∆Gq) to establish whether the reaction is dissociative or associative, as was done with the reaction of MnO2 and nitrite species (64), another associative and inner-sphere redox reaction.

Acknowledgments The authors acknowledge Universidad de Buenos Aires, Secretaria de Ciencia y Te´cnica through Project UBACyT TW99, and the U.S. National Science Foundation (OCE0096365) for financial support of this work.

Literature Cited (1) Luther, G. W., III; Church, T. M.; Giblin, A. E.; Howarth, R. W. In Organic Marine Geochemistry; Sohn, M. L., Ed.; ACS Symposium Series 305; American Chemical Society: Washington, DC, 1986; pp 341-355. (2) Luther, G. W., III; Church, T. M.; Scudlark, J.; Cosman, M. Science 1986, 232, 746-749. (3) MacCrehan, W.; Shea, D. In Geochemical Transformations of Sedimentary Sulfur; Vairavamurthy, M. A., Schoonen, M. A. A., Eds.; ACS Symposium Series 189; American Chemical Society: Washington, DC, 1995; pp 295-310. (4) Al-Farawati, R.; van den Berg, C. M. G. Environ. Sci. Technol. 2001, 35, 1902. (5) Mopper, K.; Taylor, B. F. In Organic Marine Geochemistry; Sohn, M. L., Ed.; ACS Symposium Series 305; American Chemical Society: Washington, DC, 1986; pp 341-355. (6) Shea, D.; MacCrehan, W. A. Anal. Chem. 1988, 60, 1449. (7) Vairavamurthy, A.; Mopper, K. Anal. Chim. Acta 1990, 236, 363. (8) Ginzburg, B.; Dor, I.; Chalifa, I.; Hadas, O.; Lev, O. Environ. Sci. Technol. 1999, 33, 571. (9) Luther, G. W., III; Church, T. M.; Powell, D. Deep Sea Res. 1991, 38 (Suppl. 2A), S1121-S1137. (10) Kosower, E. M. In Glutathione: Methabolism and Functioning; Arias, I. M., Jakoby, W. B., Eds.; Kroc Foundation Series; Raven Press: New York, 1976; Vol. 6, p 1. (11) Giovanelli, J. Methods Enzymol. 1987, 143, 419.

(12) Maw, G. A. In Sulfur in Organic and Inorganic Chemistry; Senning, A., Ed.; Marcel Dekker: New York, 1972; pp 113-142. (13) Meister, A.; Anderson, M. Annu. Rev. Biochem. 1983, 52, 711. (14) Zinder, S. H.; Brock, T. D. Appl. Environ. Microbiol. 1978, 35, 344. (15) Trudinger, P. A.; Loughlin, R. E. In Comprehensive Biochemistry; Neuberger, A., et al., Eds.; Elsevier: New York, 1981; pp 165257. (16) Burdige, D. J.; Nealson, K. H. Geomicrobiol. J. 1986, 4, 361. (17) Yao, W.; Millero, F. J. Mar. Chem. 1996, 52, 1. (18) Peiffer, S.; dos Santos Afonso, M.; Wehrli, B.; Ga¨chter, R. Environ. Sci. Technol. 1992, 26, 2408. (19) dos Santos Afonso, M.; Stumm, W. Langmuir 1992, 8, 1671. (20) Herszage, J.; dos Santos Afonso, M. Colloids Surf. A 2000, 168, 61. (21) Herszage, J.; dos Santos Afonso, M. Submitted for publication in Langmuir. (22) Baumgartner, E.; Blesa, M. A.; Maroto, A. J. G. J. Chem. Soc., Dalton Trans. 1982, 1649. (23) Blesa, M. A.; Maroto, A. J. G.; Morando, P. J. J. Chem. Soc., Faraday Trans. 1 1986, 82, 2345. (24) Borghi, E. B.; Morando, P. J.; Blesa, M. A. Langmuir 1991, 7, 1652. (25) Amirbahman, A.; Sigg, L.; von Gunten, U. J. Colloid Interface Sci. 1997, 194, 194. (26) Stone, A. T.; Morgan, J. J. Environ. Sci. Technol. 1984, 18, 617. (27) Papadopoulos, K. P.; Jarrar, A.; e Issidores, C. H. J. Org. Chem. 1966, 31, 615. (28) Wallace, T. J. J. Org. Chem. 1966, 31, 1217. (29) Parida, K. M.; Samal, A.; Das, N. N. J. Colloid Interface Sci. 1998, 197, 236. (30) Capozzi, G.; Modena, G. In The Chemistry of the Thiol Group; Patai, S., Ed.; 1974; Vol. 2, Chapter 17, pp 785-839. (31) Luther, G. W., III; Church, T. M. In Sulphur Cycling on the Continents; Howarth, R. W., Stewart, J. B., Ivanov, M. V., Eds.; 1992 SCOPE; John Wiley & Sons Ltd.: New York, 1992; pp 125142. (32) Vairavamurthy, A.; Zhou, W.; Eglinton, T.; Manowitz, B. Geochim. Cosmochim. Acta 1994, 58, 4681. (33) Pe´rez-Benito, J. F.; Arias, C.; Amat, E. J. Colloid Intterface Sci. 1996, 177, 288. (34) Hem, J. D. In Particulates in Water: Characterization, Fate, Effects and Removal Interfacial and Interspecies Processes; Kavanaugh, M. C., Leckie, J. O., Eds.; ACS Symposium Series 189; American Chemical Society: Washington, DC, 1980; pp 45-72. (35) Luther, G. W., III; Ruppel, D. T.; Burkhard, C. In Mineral-Water Interfacial Reactions: Kinetics and Mechanisms; Sparks, D. L., Grundl, T., Eds.; ACS Symposium Series 715; American Chemical Society: Washington, DC, 1999; pp 265-280. (36) Pe´rez-Benito, J. F.; Brillas, E.; Pouplana, R. Inorg. Chem. 1989, 28, 392. (37) Pe´rez-Benito, J. F.; Arias, C J. Colloid Interface Sci. 1992, 152, 70. (38) Benedict; Stedman Analyst 1970, 95, 296. (39) Kostka, J. E.; Luther, G. W., III; Nealson, K. H. Geochim. Cosmochim. Acta 1995, 59, 885. (40) Winterbourn, C. C.; Brennan, S. O. Biochem. J. 1997, 326, 87. (41) Atwood, J. D. Inorganic and Organometallic Reaction Mechanisms; Brooks/Cole Publishing Co.: Monterey, CA, 1985. (42) Connors, K. A. Chemical Kinetics: The Study of Reaction Rates in Solution: VCH Publishers: New York, 1990. (43) Pulfer, K.; Schindler, P. W.; Westall, J. C.; Grauer, R. J. Colloid Interface Sci. 1984, 101, 554-564. (44) Furrer, G.; Stumm, W. Geochim. Cosmochim. Acta 1986, 48, 2405. (45) Luther, G. W. In Aquatic Chemical Kinetics; Stumm, W., Ed.; Wiley-Interscience: New York, 1990; pp 173-198. (46) Yao, W.; Millero, F. J. Geochim. Cosmochim. Acta 1993, 57, 3359. (47) Stone, A. T.; Morgan, J. J. Environ. Sci. Technol. 1984, 18, 450. (48) Stone, A. T.; Ulrich, H. J. J. Colloid Interface Sci. 1989, 132, 509. (49) Stone, A. T.; Godtfredsen, K. L.; Deng, B. In Chemistry of the Aquatic Systems: Local and Global Perspectives; Bidoglio, G., Stumm, W., Eds.; 1994; pp 337-374. (50) Nico, P. S.; Zasoski, R. J. Environ. Sci. Technol. 2000, 34, 3363. (51) Nico, P. S.; Zasoski, R. J. Environ. Sci. Technol. 2001, 35, 3338. (52) Nesbitt, H. W.; Canning, G. W.; Bancroft, G. M. Geochim. Cosmochim. Acta 1998, 62, 2097. (53) Banerjee, D.; Nesbitt, H. W. Geochim. Cosmochim. Acta 1999, 63, 3025. (54) Banerjee, D.; Nesbitt, H. W. Geochim. Cosmochim. Acta 1999, 63, 1671. (55) Banerjee, D.; Nesbitt, H. W. Am. Mineral. 2000, 85, 817. VOL. 37, NO. 15, 2003 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

3337

(56) Banerjee, D.; Nesbitt, H. W. Geochim. Cosmochim. Acta 2001, 65, 1703. (57) Kalyanaraman, B.; Karoui, H.; Singh, R. J.; Felix, C. C. Anal. Biochem. 1996, 241, 75. (58) Gangopadhyay, S.; Ali, M.; Dutta, A.; Banerjee, P. J. Chem. Soc., Dalton Trans. 1994, 841. (59) Cotton, F. A.; Wilkinson, G. In Advanced Inorganic Chemistry, 2nd ed.; John Wiley & Sons-Interscience: New York, 1967; pp 530-531. (60) Burns, R. G.; Burns, M. B. In Marine Minerals; Burns, R. G., Ed.; Reviews in Mineralogy 6; Mineralogical Society of America: Washington, DC, 1979; pp 1-46. (61) Stumm, W. In Chemistry of the Solid-Water Interface; John Wiley & Sons-Interscience: New York, 1992; pp 25-38.

3338

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 37, NO. 15, 2003

(62) Lie, T. J.; Leadbetter, J. R.; Leadbetter, E. R. Geomicrobiol. J. 1998, 15, 135. (63) 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; pp 260-279. (64) Luther, G. W., III; Popp, J. I. Aquat. Geochem. 2002, 8, 15-36.

Received for review January 23, 2003. Revised manuscript received April 30, 2003. Accepted May 5, 2003. ES0340634