9684
Langmuir 2003, 19, 9684-9692
Mechanism of Hydrogen Sulfide Oxidation by
Manganese(IV) Oxide in Aqueous Solutions
Julián Herszage and Marı́a dos Santos Afonso*
INQUIMAE and Departamento de Quı́mica Inorgánica, Analı́tica y Quı́mica Fı́sica,
Facultad de Ciencias Exactas y Naturales, Universidad de Buenos Aires,
Ciudad Universitaria Pabellón 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.; Gächter, 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 Pérez 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) Pérez 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.
9686
Langmuir, Vol. 19, No. 23, 2003
Herszage and dos Santos Afonso
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.
Langmuir, Vol. 19, No. 23, 2003 9687
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.
9688
Langmuir, Vol. 19, No. 23, 2003
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.
9690
Langmuir, Vol. 19, No. 23, 2003
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) Rönngren, L.; Sjöberg, 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
9692
Langmuir, Vol. 19, No. 23, 2003
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, Secretarı́a de Ciencia y Técnica,
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) Böttcher, M. E.; Thamdrup, B. Geochim. Cosmochim. Acta 2001,
65, 1573-1581.