Environmental Geochemistry of Sulfide Oxidation - American

Chapter 26

Kinetics of Oxidation of Hydrogen Sulfide in Natural Waters Downloaded by UNIV OF CALIFORNIA SAN DIEGO on November 15, 2014 | http://pubs.acs.org Publication Date: December 20, 1993 | doi: 10.1021/bk-1994-0550.ch026

Jia-Zhong Zhang and Frank J. Millero Rosenstiel School of Marine and Atmospheric Science, University of Miami, 4600 Rickenbacker Causeway, Miami, FL 33149-1098

Recently we have studied the oxidation of H S with O in natural waters as a function of pH (4 to 10), temperature (278.15 to 338.15 K) and salinity (0 to 36). The major products formed from the oxidation of H S were SO , S O and S O . A kinetic model was developed to predict the distribution of the reactants and products over a wide range of conditions. Dissolved and particulate metals have a significant effect on the rates of oxidation and the product formation. Field measurements made in the Black Sea, Framvaren Fjord, Chesapeake Bay and Cariaco Trench are in reasonable agreement with the values predicted from laboratory studies at the same concentration of F e . 2

2

2-

2-

3

2

2

2-

3

4

2+

The formation of hydrogen sulfide occurs in a variety of natural waters. The production in the pore water of sediments and stagnant basins (seas, lakes,riversand fjords) is due to biological processes while the production in hydrothermal systems is an abiotic process. In anoxic environments organic matter can be oxidized by bacterial anaerobic respiration using various oxidants as an electron acceptor. The preference of oxidant is related to the greatest free energy yield per mole of organic carbon oxidized (1). Molecular oxygen is the thermodynamically most favorable electron acceptor which, if available, will be used preferentially in any ecosystem. If the supply of organic matter exceeds that of oxygen, other electron acceptors (in the order of M n 0 , N0 ", NO3-, F e 0 and S 0 ) are used when oxygen has been depleted. When one oxidant is depleted, the oxidant with the next highest energy yield is consumed until every oxidant is removed, or until all the metabolizable organic carbon has been depleted. Dissimilatory sulfate reduction 2

2

2

2

3

4

2

SO4 - + 2 C H 0 HS- + 2 C 0 + H 0 + OH" (1) is most commonly observed in marine environments where water circulation, consequently oxygen availability, is limited, but where sulfate is easily available because of its relatively high concentration in seawater ( « 0.029 M). The production of hydrogen sulfide also occurs in hydrothermal systems. In hot vent waters hydrogen sulfide may be leached from crustal basalts or produced 2

2

2

0097-6156/94/0550-0393$06.00/0 © 1994 American Chemical Society In Environmental Geochemistry of Sulfide Oxidation; Alpers, C., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 1993.

394

ENVIRONMENTAL GEOCHEMISTRY OF SULFIDE OXIDATION 2 +

by reduction of sulfate from seawater coupled with oxidation of F e from basalt to F e . Part of hydrogen sulfide so produced reacts with metal ions depositing metal sulfide minerals, mainly as pyrite, the remainder stays in the vent solution. 3S0 - + 6H+ + 17Fe Si0 -» H S + FeS + 2 H 0 + H F e 0 + 17Si0 (2) Once the hydrogen sulfide has been formed in natural waters, the oxidation of hydrogen sulfide is an important pathway for its removal. When water containing H S mixes with oxygenated water at the oxic and anoxic interface, the hydrogen sulfide can be oxidized by a number of oxidants. This oxidation is frequently coupled to changes in the redox state of metals (2) and nonmetals (3). Dissolved oxygen, however, is the most important and abundant oxidant. This oxidation involves a complex mechanism that results in the formation of several sulfur species (i.e., S0 ", S 0 ", S and S ) as well as S 0 \ Although the formation of the resultant products has been studied by a number of workers (4-7). only a few of these product studies have been made over a wide range of experimental conditions and reaction media. In recent years we have studied the oxidation of FkS (8. 9) and H S 0 (10) with 0 in water and seawater in the laboratory (9.10) and in the field (11-14). We have attempted to characterize how the rates and distributions of products are affected by trace metals (15). A kinetic model has been developed (16) to predict the rates of oxidation and formation of products. The results of these studies are briefly reviewed in this paper. 3+

2

4

2

4

2

2

2

3

4

2

Downloaded by UNIV OF CALIFORNIA SAN DIEGO on November 15, 2014 | http://pubs.acs.org Publication Date: December 20, 1993 | doi: 10.1021/bk-1994-0550.ch026

2

2

2

3

2

2

3

n

2

4

2

3

2

Overall Rate of Oxidation of H S 2

The overall rate equation for the oxidation of hydrogen sulfide can be represented by

brackets represent molar concentrations (9). When the concentration of oxygen is in excess the rate equation can be reduced to - d[H S]/dt = k' [H S] (4) where k' is related to k (. ) by ^ = νπ)[0 Ρ (5) The order with respect to sulfide was determined byfittingthe data to various rate equations with different values of a (9. 16V Plots of In [H S] versus time were found (9,16) to give straight lines and indicated that a is equal to 1 or the reaction is first order with respect to the concentration of H S in agreement with earlier studies (4-7). The order of the reaction with respect to the oxygen concentration was determined from the values of k' at different concentrations of oxygen in seawater (16). These results gave a first-order dependency with respect to the concentration of oxygen. Our measured order of the oxidation with respect to oxygen is in close agreement with the value (0.8) obtained at mM levels of initial sulfide concentrations (7), but higher than the value of 0.56 found by Chen et al. (5). In summary all our studies mdicated that the rate equation for the oxidation of H S in water and seawater is given by a

2

2

S

D

2

2

2

2

been measured as a function of pH (1 -12), temperature (278.15 - 338.15 K), and ionic strength (I, 0 - 6M) (9). At pH 8.0, the rate constant (k (_ ), M mûr ) in 1

S

n

In Environmental Geochemistry of Sulfide Oxidation; Alpers, C., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 1993.

1

26.

Oxidation of Hydrogen Sulfide in Natural Waters 395

ZHANG & MILLERO

equation (6) is given by (Τ, K) 1 0

0 0

3

x

1

3

T

4 4 1 0 5

?

!og k _ii) = · - ( ·° ° ) / + ° () (σ = 0.18 in log k . ) . At 298.15 Κ the half time for the oxidation of H S with 0 was found to be 50 ± 16 h in water and 25 ± 9 h in seawater (9). These results are in good agreement with the results of Chen and Morris (5) and O'Brien and Birkner (7j. s(

s(

n)

2

2

The effect of pH in water at 328.15 Κ was found to be represented by (9) ks(-ii) = fera + kns K i / p i + M l + ^/[H*]) where k = 1.33 ± 0.28 M mûr for the oxidation of H S and k 0.12 M" min- for the oxidation of HS"


- 1

1

H 2 S

1

2

H S

(8) = 5.73 ±

1

kjj g 2

HS + 0

-> products (9) k HS- + 0 -* products (10) The value of Κχ is the thermodynamic constant for the ionization of H S (8). The effect of temperature and ionic strength on the rate constants k and k have been given by (9) log k = 7.44 - (2.4 χ 10 >T (11) tog k s · ° - P " ( -° ° ) + ( > These equations are valid from pH = 4 to 8, Τ = 278.15 to 338.15 K, and I = 0 to 6M. 2

2

H S

2

2

H 2 S

H S

3

H 2 S

=

8

7 2

+

1 6

H

3

x

1

3

/ T

0

M

1 0 , 5

12

H

Effect of Metals on the Rate of Oxidation As will be discussed later, field measurements (11-14^ made on the oxidation of H S in natural waters yielded half times that were much faster than determined in the laboratory on Gulf Stream seawater (9). To determine if this increase was due to trace metals, we have measured the rates of oxidation of H S in seawater with added transition metals (15). These studies have shown that at total dissolved con centrations below 300 nM, the rates are only affected by F e , C u and P b . At higher metal concentrations, the rates of oxidation of H S increase for all the met als except Z n . The order of the increase in the rates at higher concentrations for these metals is F e > Pb + > C u > F e > C d > NP+ > C o > M n 05). Only F e and M n have levels in anoxic basins high enough to affect the oxidation of H S. The relative effect of metals on the oxidation of H S with oxygen at 298.15 Κ and pH 8.1 can be estimated (15) from (Figure 1) log(k (.iiyks(-ii)°) = a + b log[M] (13) where ks(-u) * ks(-n) are the rate constants, respectively, with and without added metal and a = 6.55, b = 0.820 for Fe(II) from 10" to 10" - M a = 5.18, b = 0.717 forFe(III)from 10 · to 10 · Μ a = 1.68, b = 0.284 forMn(II)from 10 · to 10 · Μ. The larger effect of F e is probably related to the formation of dissolved F e from the rapid oxidation of F e with oxygen (17) Fe + 0 ^ Fe + 0 (14) The oxidation of F e provides a higher initial concentration of F e than can be 2

2

2+

2 +

2+

2

2 +

2+

2 +

2

2 +

3+

2 +

2 +

2 +

2 +

2

2

S

a n c

0

8

5 3

7

2

3

3

5

9

3

3

2 +

3 +

2+

2 +

3+

2

2

2+

3 +


396

ENVIRONMENTAL GEOCHEMISTRY OF SULFIDE OXIDATION 3 +

added from a stock solution of F e (which may be locally supersaturated). The peroxide generated by the oxidation of F e 0 - + H+ H0 (15) H0 + H0 H 0 + 0 (16) can also increase the rates since it has a higher rate of oxidation with sulfide than oxygen (18). The F e formed from the reaction of dissolved or particulate F e can regenerate F e to complete the catalytic cycle. The effect of F e on the rates at low concentrations may be related to the reduction of F e to F e with HS" and resultant oxidation of F e and generation of 0 . We presently are investigating this reduction process in the absence of 0 . The effect of the metals on the rates of oxidation of H S below the observable precipitation of metal sulfides (which may be a slow process) can be attributed to the formation of ion pairs (15) M + HS- -> MHS+ (17) 2+

2

2

2

2

2

2

2

2 +

3 +

2 +

3 +

3 +

2 +

2 +

_

2


2

2

2+

The overall rate constant is given by ks -ii)[HS-] = k [HS-] + k [MHS+] (18) where k and k ^ are the rate constants for the oxidation of HS and MHS+. From the mass balance of [ H S ] and [ M ] and the stability constant for the formation of M H S (

T

HS

MHS

-

H S

H S

2 +

X

T

+

2

iS Hs = [MHS+]/[M +][HS-] we have

(19)

M

k (-ii)[HS-] = k [HS-] + k ^ s [M2+h/(l + V / W H S ] ) S

T

(20)

HS

sc

o s e

t 0

If the value of / J is large enough (19), then (1 + 1//?MHS[ 1) * ' 1· Thus, a plot of k (. ) versus the total metal concentration can be used to estimate MHS- e s e plots give k = 14.7 ± 1.9 M " min" , k Q , = 53.5 ± 1.6 M " mini and k ^ s = 201.7 ± 16.5 M mini (15). h metals C o , P b , and N i the rates did not increase until visual precipitation of metal sulfides formed in the solutions. The decrease (15) in the rate of oxidation with added Z n above 2μΜ may be related to the formation of zinc sulfide ion pairs, Z n H S (19). Unlike the ion pairs of the other metals, ZnHS may be more stable than HS". If we assume that the Z n H S species is non-reactive, we obtain ks(-nyks -n)° = «HS = [HS-] /[HS] = (1 + / ? [Zn +] )-i (21) where a is the fraction of free HS", / ? is the stability constant for the forma tion of Z n H S and [ Z n ] is the free zinc concentration. The experimental results gave l o g / ? = 8.2 ± 0.5, which is higher than the value given by Dyrssen (12) of log /JznHS 6.5. These difference could be related to the formation of a kinetically stable product that is not in equilibrium. The presence of Fe and Mn in natural waters not only increases the rate of oxidation of sulfide, but also can have an effect on the oxidation of intermediates such as sulfite. This catalysis can change the distribution of the products formed during the oxidation. The effects of metals on the formation of products during the oxidation of H S are discussed in the next section. hs

M H S

S

k

n

Th

1

1

1

M n H S

HS

1

2+

F

o

r t

2+

e

2 +

2 +

+

+

+

2

(

F

T

H S

Z n H S

F

Z n H S

+

2+

F

ZnHS

=

2

Productsfromdie Oxidation of H S 2

The final product from the oxidation of sulfide is sulfate, the sulfur compound having the highest oxidation state and the most stable compound in oxic waters.


26.


ZHANG & MHUERO

Various intermediates, such as sulfite and thiosulfate, also can be formed during the course of the reaction. The products formed from the oxidation of H S in seawater have been studied (16) as a function of pH, temperature, salinity, and reactant concentration. To examine the mass balance of sulfur compounds during the oxidation the experiments were done (16) in pure water where SO^ * formed from the oxidation could be measured by an ion chromatographic technique. The major products formed were found to be S 0 , S 0 and S^0 (Figure 2V Elemental sulfur or polvsulfides were not found by spectroscopic techniques (5). The total equivalent sulfur of the products and reactants was constant indicating that S0 ", S0 " and S 0 " are the main products. The distribution of products from the oxidation of H S in seawater (Figure 3) is similar to the results in water. Sulfite has been proposed to be the initial product from the oxidation of sul fide with oxygen in alkaline solutions (4). Unfortunately sulfite was not measured in the earlier studies due to the limitation of methods used. Our measurements in water and seawater at pH = 8.2 support the contention that sulfite is the initial oxidation product (16). The concentration of sulfite increases rapidly at the beginning of the reaction and decreases slowly when the production is slower than the rate of oxidative removal (see Figure 2). The major product is S 0 in all of the seawater runs. The concentration of sulfate is higher at higher concentrations of oxygen as one might expect. Because we determined S 0 in seawater by dif ference, some of the assigned values could be elemental sulfur or polysulfides pre sent at levels below the detection limit of the spectroscopic method used (1 μΜ). The concentration of thiosulfate increases slowly throughout the reaction after an initial lag period. This observation suggests that thiosulfate is not the ini tial product of the oxidation and supports the previously made assertion that S 0 serves as the initial oxidation product (16). In earlier studies, the formation of sul fite and thiosulfate was treated as a parallel reaction based on 8-hour experiments (7). Under those experimental conditions (equal moles of initial sulfide and oxygen at 298.15 K) the sulfide oxidation was very slow and only 30% of the initial sulfide was oxidized. A longer period of reaction is necessary to show the true pat tern of the intermediate products. This requirement is, of course, the reason that we made most of our measurements at 318.15 K. Thiosulfate is a stable product in the absence of bacteria and little oxidation occurs over 80 hours. This observation is in agreement with the finding of earlier studies (4. 7). The effect of p H on the distribution of products has also been examined and the results nave been attributed to changes in the rates of the individual reaction steps (16). The effect of metals (Fe +, F e , M n , C u , P b ) and solids (FeOOH and M n 0 ) on the distribution of products has also been studied (16). The inter mediates formed (Figure 4) during the oxidation of Cariaco Trencn waters (350 nM F e ) clearly show that dissolved and particulate metals not only increase the rate of oxidation of H S, but also change the distribution of the products. 2

2

2_

4

2

4

2

2_

2_

3

3

2

3

2

?


2

2_

4

2_

4

2_

3

2

3+

2 +

2+

2+

2

2+

2

Oxidation of the Intermediate S(IV) Because the rate of oxidation of sulfite to sulfate affects the distribution of products formed during the oxidation of sulfide, we have studied its oxidation (10). The overall rate equation for the oxidation of sulfite in seawater can be written as

(io) - d[S(IV)]/dt = kgpv) [S(IV)] [O ]05 (22) where k ^ ™ is the overall rate constant; the brackets represent concentrations. All the kinetic measurements in seawater and seasalts (10) showed that the oxida tion reaction was second order with respect to sulfite. This relationship is in agreement with previous measurements in NaCl solutions (20) and in seawater (21). The order of the oxidation with respect to oxygen was found to be 0.5(± 2

2



398

ENVIRONMENTAL GEOCHEMISTRY OF SULFIDE OXIDATION

—I

I

I

1

-8

ι

-7

I

ι

I

-6

ι

-5

I

-4

ι

1

-3

log [M]

Figure 1. The effect of Fe and Mn on the rate of oxidation of H S in seawater at 298.15 Κ and pH 8.1 (15). 2

30 η

1

1

1

1

1

1

1

1

1

1

1

Γ

TIME (HOUR)

Figure 2. The sulfur balance during the oxidation of sulfide in water at 318.15 Κ and pH = 8.2. The smooth curves are calculated from the model (16).



ZHANG & MILLERO


TIME (HOUR)

Figure 3. The distribution of products from the oxidation of sulfide in seawater at S = 35.0, pH = 8.2 and Τ = 298.15 Κ. The smooth curves are calculated from the model (16).

TIME (HOUR)

Figure 4. The distribution of products from the oxidation of sulfide in the Cariaco Trench (mixture of surface and deep water) at 298.15 K. The smooth curves are calculated from the model (16).


400


0.03) (10). The values of log k ( ) as a function of temperature and ionic strength has been fitted to (10) log k = 19.54 - 5069.47/T + 1 4 . 7 4 - 2.93 I - 2877.01°-VT (23) where I is the molal ionic strength, Τ is the temperature (K), k is in Μ · min . The standard error is 0.05 in log k mn. This equation should be valid for most estuarine and sea waters. The effect of ionic strength on the energy of activation is in agreement with the earlier measurements made in NaCl solution (20). The major ionic components of seawater also have an effect on the rates of oxidation of sulfite (10). The rates measured in 0.57M NaCl solution were found to be higher than the rates in seawater. Measurements made in the major sea salts solution indicate that C a , M g and S 0 " added to NaCl solution cause the decrease. Measurements made in artificial seawater (Na , M g , C a , CI and S0 ") were found to be in good agreement with the measurements in real seawater (10). The effect of pH on the rate of oxidation was found to be significant (10). The rate increased from pH 4 to a maximum at pH 6.5 and decreased at higher pH. The effect of pH on the rates was attributed to the rate-determining step involving the combination of HS0 " and S0 ". This yields S

IV

s ( I V )

4

5

-1


s

2 +

2 +

2

4

+

2 +

2+

-

2

4

2

3

3

2

k iv) = k" a(HS0 -) HS0 (25) S0 " + 0.5 0 S0 " (26) S 0 - + HS- + 0.5 0 S 0 " + OH(27) S 0 - + 0.5 0 S0 - + S (28) With these overall reactions in mind, one can attribute the formation of S 0 " to the oxidation of HS" and the formation of S 0 to the oxidation of S 0 . The formation of S0 ~ from the oxidation of S 0 can be neglected (4) for solutions devoid of bacteria (13). The formation of S 0 can be attributed to the overall reaction of S 0 and HS" with 0 . This leads to the following overall rates of oxidation 2

2

2

2

3

2

2

3

4

2

2

3

2

2

3

2

4

2

2

3

2

2

2

2

3

2

3

2

4

2

3

2_

2_

4

2

3

2_

4

2

3

2_

2

3

2_

3

H S 4- 0 2

H S0 2

2

2

Products(S0 )

2

(29)

3

2

3

+ 0 ^ Products(S0 ) 2

(30)

4

2

H S + H S 0 + 0 ^ Products(S 0 ) The overall rate equations for H S, S0 ", S 0 ", S 0 are given by d[H S]/dt = - k![H S][0 ] - k [H S][S0 ][0 ] d[S0 ]/dt = k![H S][0 ] - k [S0 -] [0 ]°-5 - k [H S][S0 -][0 ] 2

2

3

2

2

3

2

2

2

3

2

3

4

2

2

2

2

3

2

3

2

3

2

2

2

2

3

2

2

2

2

(31)

2_

3

2

3

2


(32) (33)

26.


ZHANG & MILLERO

(34) d[S20 2-]/dt = k [H S][S0 2-][0 ] (35) d[S0 2-]/dt = k [ S 0 2 f [ 0 F where [i] is the total concentration of i. It should be pointed out that these are overall rate equations and they do not represent the mechanism of the reactions. The order of the rates of oxidation of H S (equation 29) and H S 0 (equation 30) are assumed to be equal to the order in equations 6 and 22, respectively. The order of the rate of formation of S 0 has been taken from the work of Avrahami and Golding (4) These rate equations have been integrated simultaneously to evaluate the values of k k and k using the experimental time dependence concentrations of all the reactants and products (16). The values of k k and k determined from these studies as a function of temperature, salinity and pH in water and seawater are given elsewhere (16). The experimentally measured concentrations of H S, S 0 , S 0 - and SO^ were found to be in good agreement with the model predictions up to reaction times of 80 hours (Figure 2 and 3). The values of ko in seawater needed to fit the data were slightly smaller than the values determined in our previous study (10), especially at higher temperatures. This difference is prob ably due to the inhibition of the oxidation of sulfite in the presence of sulfide (5). This finding is supported by the previous observations (6) that sulfite in the presence of H S is more stable in seawater than predicted by its rate of oxidation. Because the oxygen is in excess, the competition for oxidant is unlikely to cause this difference. The complexation of HS with trace metals which can catalyze the oxidation of sulfite in seawater may be a more likely cause. Trace metals have higher tendency to complex with HS" than S0 " (16). We found (16) that the formation of thiosulfate could be best represented with a rate equation which is zero-order with respect to oxygen, a finding in agree ment with earlier work (7). The effect of changes in the ratio of H S to 0 on the product distribution over the range of our measurements can be attributed to the order in the rate equations which is independent of the rate constants. The values of k k and k from the experiments in pure water as a function of pH (4 to 10) have beenfittedto smooth equations of pH (T = 318.15 K) (16) In k = - 4.71 + 0.914 pH - 0.0289 p H (36) In k = 3.87 + 1.51 pH - 0.103 p H (37) In k = - 9.09 + 3.01 pH - 0.177 p H (38) Assuming that the p H dependences of rate constants are independent of ionic strength and temperature (9.10). these equations can be used to estimate the rate constants for other natural waters. Further measurements are needed to examine the pH dependence of rate constants over a wide range of temperature and ionic strength to test the validity of this assumption. The values of k k and k as a function of salinity (S) and temperature (Τ, K) have been fitted to the equations (pH = 8.2) In k = 26.90 + 0.0322 S - 8123.21/T (39) In k = 14.91 + 0.0524 S - 1764.68A (40) In k = 28.92 + 0.0369 S - 8032.68/T (41) These equations should be valid for most estuarine and sea waters. This kinetic model can be used to predict the product distribution for the oxidation of sulfide in natural waters with low concentrations of trace metals. The agreement between the model and the observed distribution of reaction products does not provide conclusive proof that the reaction pathways of the over all model actually describe the series of elementary reactions that occur in an abiotic environment. The detailed mechanisms might involve many elementary 3

3

4

2

2

3

3

2

2

2

2

3

2 _

2


1?

2

3

3

h

2

3

2

2_

2

3

2

2-

3

2

-

2

3

2

h

2

3

2

x

2

2

2

3

h

2

3

x

1

2

3


2

402


reaction steps. At present there are two detailed mechanisms for the oxidation of H S with 0 in aqueous solutions (5,22). the polar mechanism and the free radical chain mechanism. The polar mechanism for the formation of H S 0 " is given below (22): HS" + 0 -> H S 0 (42) HS0 H+ 4- S0 " (43) S 0 - 4- 0 S0 - + 0 (44) S0 - + 0 -> S 0 + 0 (45) S 0 4- H 0 -> HS0 - 4- H+ (46) The overall reaction is given by HS- + 3 0 4- H 0 -> HSO3- + 2 H 0 (47) The slow step in the oxidation of HS" to HSO3- is given by equation 42. The superoxide ion formed can react with itself to form hydrogen peroxide which can also react with HS" (11^23). H 0 4- H 0 H 0 4- 0 (16) The intermediate H S 0 or S 0 can react with HS" and form thiosulfate HS0 - 4- HSS0- + H 0 (48) S 0 " + 0 -> S 0 (49) The initial reaction can also result in the formation of elemental sulfur HS- + 0 -> S 4- H 0 (50) which can react with sulfite to give thiosulfate S 4- S 0 S0 (51) or hydrogen sulfide to give polysulfides nS 4- HSHS (52) The formation of sulfate comes from the oxidation of sulfite with oxygen 2S0 - + 0 -» 2S0 (53) or hydrogen peroxide H S 0 - + H 0 -* HSO4- + H 0 (54) which involves the formation of the intermediate 0 S O O H (23). In the free radical mechanism, oxidation is initiated by an outer-sphere electron transfer from HS" to oxygen to form the HS- radical. HS" 4- 0 -> HS- + 0 (55) In the presence of trace metals in seawater, electron transfer from HS to the transition metal ions to form HS- radical is more favorable. HS" + M + -> HS- 4- M( -!) (56) The HS- radical is further oxidized to sulfite by a free radical chain sequence involving oxygen HS- 4- 0 -» HS0 (57) HS0 - -* H+ 4- S 0 (58) S0 - + 0 -> S 0 4- 0 (59) S 0 + H 0 ^ HS0 " 4- H+ (60) H S 0 - ^ S 0 - 4- H+ (61) 2

2

3

2

2

2

2

2

2


2

2

2

2

2

2

2

2

2

2

3

2

2

2

2

2

2

2

2

2_

2

2

2

2

2

2

2

2

2

2

2

3

2

2

2

2

3

2

3

n

2

2

3

2

3

4

2

2

2

2

2

2

-

n

n

2

+

2

2

2

2

2

2

2

2

2

3

2

3

3


26.


ZHANG & MILLERO

The sulfite formed can react with HS- to produce thiosulfate (62) HS + S 0 - -* HS2O3 H S 0 - + O -* H S 0 - + 0 (63) H S 0 -

Environmental Geochemistry of Sulfide Oxidation - American

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