Kinetic Controls on Thermochemical Sulfate Reduction as a Source of

Chapter 23

Kinetic Controls on Thermochemical Sulfate Reduction as a Source of Sedimentary H S 2

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Martin B. Goldhaber and Wilson L. Orr

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Downloaded by RUTGERS UNIV on March 12, 2016 | http://pubs.acs.org Publication Date: May 5, 1995 | doi: 10.1021/bk-1995-0612.ch023

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U.S. Geological Survey, Mail Stop 973, Denver Federal Center, Denver, CO 80225 Earth and Energy Science Advisors, P.O. Box 3729, Dallas, TX 75208

Laboratory experiments with aqueous ammonium sulfate in the presence of H S and toluene are reported which show measurable SO reduction in 4 to 30 days at 175-250° C. Reduction rates increase with both increasing temperature and H S pressure but reduction was not measurable on our experimental timescale without H S initially present. An activation energy of 96 (±16) kJ/mole was estimated from the data. These results (and other published studies) indicate that thermochemical sulfate reduction (TSR) is difficult to document below 200° C on a laboratory time scale unless ΣS (i.e. SO + H S ) is initially very high (>.5M ) and pH is low (250°C the available thermal energy is sufficient to allow rapid chemical and isotopic equilibrium between S0 ' and H S (3). The subject of this paper is H S generation in the intermediate 4

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0097-6156/95/0612-0412$12.00/0 © 1995 American Chemical Society Vairavamurthy et al.; Geochemical Transformations of Sedimentary Sulfur ACS Symposium Series; American Chemical Society: Washington, DC, 1995.

23. GOLDHABER & ORR

413

Thermochemical Sulfate Reduction

temperature range of 100-200° C; too warm for bacterial processes and too cool for rapid chemical equilibrium. This intermediate temperature range is of interest because it characterizes the formation of sour (H S-bearing) gas fields (4) and the genesis of many large sediment-hosted ore deposits (5). Because of their geologic similarities, we will consider the origin of H S in sour gas fields and the so-called Mississippi Valley Type (MVT) ore deposits; both are hosted in carbonate rocks and formed over the temperature range of interest. The process generating H S in this intermediate temperature range has been termed thermochemical sulfate reduction (TSR; 4). TSR is analogous to the bacterial reduction pathway in that it is a kinetically controlled reaction involving the coupled oxidation of organic matter and reduction of S0 ". Although considerable success has been attained in documenting the occurrence and consequences of TSR from field studies (e.g. 4 6-8\ understanding the critical geologic and geochemical controls has been hindered by the inability to experimentally reproduce TSR under laboratory conditions comparable to those deduced from field sites. Previous successful experimental studies required temperatures above 220° C (9), and many involved pH values and reactant concentrations outside the range for most natural systems. One study (9) was designed to mimic natural systems. Temperatures rangedfrom90-190 °C, and a variety of reductant concentrations and, several possible catalytic agents, were employed. Yet no evidence for TSR was found despite using a sensitive radiosulfur tracer technique to detect low extent of reaction. In this paper we present experimental results with conditions selected to allow measurable rates of TSR on a convenient time scale at 175-250° C. These data were presented previously in abstract form (10). We consider the significance of the data in comparison with other experimental studies in the literature and the extrapolation of these studies to conditions characteristic of naturally occurring TSR in geologic settings. Specifically, we attempt to reconcile apparent differences in laboratory and field observations. 2

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y

Experimental TSR was studied at 175 to 250° C in a 500 cc stainless steel autoclave containing (NH ) S0 , H S, H 0 , and Toluene with reaction times of 92 to 620 hours. The vessel initially contained 100 cc of sulfate solution whose concentration ranged from .05 to .48 molar. A large molar excess of toluene (20 cc) relative to sulfate was added. Nitrogen was used to flush airfromthe system, to vary total charge pressure, and to flush H S from the autoclave following the completion of each run. Initial H S was added from a cylinder to the nitrogen filled reactor to yield final pressures of 10 to 225 psi. Most runs were made with 200 psi H S. The total charge pressure was adjusted with N to 300 to 1000 psi at room temperature. The initial and final quantities of 4

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Vairavamurthy et al.; Geochemical Transformations of Sedimentary Sulfur ACS Symposium Series; American Chemical Society: Washington, DC, 1995.

414

GEOCHEMICAL TRANSFORMATIONS OF SEDIMENTARY SULFUR

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S0 " were measured gravimetrically with an accuracy of ±1%. Quantities of H S produced were not measured because amounts formed were small relative to starting amounts and losses due to corrosion of the stainless steel vessels prohibited accurate measurements. The reactor was controlled at the desired temperature (±3°C) for the desired time. Heating and cooling times were short (1-2) hours with respect to reaction times. 4

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Results Results of runs conducted at 250° C are contained in Table I. Also shown in this table are calculated first order rate constants and the half life of the reaction in hours. The treatment of the results as first order in sulfate is an assumption because only initial and final concentrations were measured. However, previous time series studies (e.g. 11-13; see also 3) have deduced

Table I Experiments with Variable H S Pressure at 250 °C 2

Time; Hours

Initial HS

Initial S 0 " mmoles

Final S 0 " mmoles

% so reduction

Rate Constant

HalfLife

128

0

48.2

48.8

~0

120

10

48.1

-

15

1.4xl0"

-

40.8

120

25

48.1

36.9

23

2.2x10"

3

315

120

50

48.2

22.1

54

6.5xl0"

3

106

120

100

48.1

24.7

49

5.6xl0"

3

124

141

100

46.9

19.4

59

3

6.3x10"

111

240

200

48.2

12.1

75

5.8xl0"

3

120

240

200

24.9

4.5

82

7.1xl0-

92

200

24.4

11.5

53

8.2xl0"

120

200

48.2

15.6

68

168

(-225)?

5.1

(0.5)?

89

2

2

4

2

4

2

4

3

500

3

98

3

85

9.4xl0"

3

74

--

--

2

first order kinetics with respect to S0 ". With this assumption, the first order rate constant k is given by: 4


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415

1 n *=(-)ln(3 t C, c

where C and C are the initial and final concentrations respectively. The half life of the reaction is: 0

t

,

0.693

tl =


2

k

Reduction rates were quite rapid at 250° C and dependant on the pressure of H S. Without initial H S, the reaction was not measurable ( H S + 1.33 C 0 + 2 OH" 2

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(3)

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This pathway was in part deduced from earlier work by Toland (14-15), who verified equations 1 and 2 experimentally. More recent work has reconfirmed these results (73). The S0 VH S reaction (equation 1) was previously postulated to be the slow or rate deteimining step in TSR (6) when adequate amounts of suitable organic matter were present. The reactive S° intermediate may actually be any one of a number of intermediate oxidation state sulfur species such as polysulfides/hydropolysulfides (S 7HS ), or thiosulfate (HS 0 VS 0 ) in addition to elemental sulfur. The reaction between organic matter and S°" does not directly produce C 0 ; it leads initially to organic acids (14), which then may undergo subsequent decarboxylation reactions (7), to yield the observed association in sour gas fields between H S and C 0 (8). 2

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X

X

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3

2

3

H

2

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2

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In a detailed mechanistic analysis of isotope exchange between S0 " and H S (which is related to the reduction pathway), Ohmoto and Lasaga (3) concluded that the overall rate of reduction was proportional to the product of Σ [S0 ]- £ [ H S ] and the exchange rate was also proportional to the sum of the concentrations of [S0 ] + [H S], (i.e., £[$])· T h exchange was also a complex function of pH, with substantially increased reaction rates at low pH. The role of total £ [ S ] in solution and pH controls on the overall reaction was also discussed by Trudinger et al. (9) who pointed out that if equation 1 is rate controlling, the stability of S° (or related polysulfide compounds) would be critical to achieving rapid reduction rates. They emphasized that the elemental sulfur stability field expands with increasing temperature, increasing £ [ S ] , and decreasing pH. 4

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e

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Evaluation of the Experimental Results Based on the foregoing, three of the major variables controlling the rate of reduction are related to the initial total sulfur concentration in solution (£[S]),


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417


temperature, and pH. In this section, we compare experimental data on TSR on the basis of these variables. We make no claim that £ [ S ] is itself a kinetic parameter, but rather, that it is a convenient variable closely related to ratecontrolling parameters allowing comparison of a diverse group of experiments. Figure 2 shows schematically the available experimental data including the results from this study on a plot of £ [ S ] versus the temperature of the runs. Experiments which successfiilly produced H S are shown in a cross hatched pattern and unsuccessful experiments are shown as a solid field. For plotting purposes, the added solid S° in (73) was recalculated to its equivalent solution concentration. The justification for this is given below. Thisfigureshows that the successful experiments represent either higher total initial aqueous sulfur or higher temperatures or both compared to those that failed to yield observable reduction. It is evident that the 175° C run of this study (Table II) which resulted in measurable S0 " reduction, occurred at substantially higher £ [ S ] than the unsuccessful experiments of (9) at similar temperatures. It is also noteworthy that the unsuccessful runs of Drean (16), although conducted at elevated temperature, employed low initial £ [ S ] conditions. An important distinction not evident from Figure 2 is whether H S, a required reactant in equation 1, was present at the beginning of the experiments in the various studies. Most successful experiments, this study, and (75), started with initial sulfide and showed the necessity of reduced sulfur species to obtain relatively rapid reaction. Likewise, Kiyosu and Krouse (73) initially added S° which produced H S quickly at temperature by reaction with acetic acid (equation 2); the rate of H S formation exceeded that of sulfate reduction and the conversion of S° to H S was nearly quantitative. Most unsuccessful experiments in Figure 2 (11-12, 16) contained only S0 " with various types of organic matter. The unsuccessful experiments by Trudinger et al. (9) had either no initial sulfide or very low concentrations as well as a low £ [ S ] . The limited success of Kaiser (17) with high £ [ S ] and no initial sulfide suggests slow H S production directly from dextrose and sulfate but, in general, the direct reduction of sulfate by organic matter is known to be orders of magnitude slower than observed TSR rates below 300 ° C (4,6, 14-17, 19, 211 Also shown on Figure 2 is a field for the conditions of genesis of Mississippi Valley Type (MVT) sedimentary ore deposits. This field (crosshatched rectangle) has been generalized from data contained in a number of fluid inclusion studies of these ores (18). The sulfide component of a subset of these ores has been inferred to have formed via TSR based on sulfur isotope systematics (5); although non-TSR sulfide sources are implicated in many cases this distinction is not reflected in the M V T field in Figure 2. The negative experimental results of Trudinger (9) directly overlap the field for M V T ores. A possible resolution of this seeming paradox is discussed below.


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418

GEOCHEMICAL TRANSFORMATIONS OF SEDIMENTARY SULFUR


•

-3' 0.9

1.1

ι •... ι 1.3 1.5

1.7

1.9

ι 2.1

• .1 2.3

Log (PH2S) Figure 1. Effect of initial H S pressure on first-order rate constants for S0 " removal during thermochemical sulfate reduction. All data at 250 C°. 2

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Figure 2. Summary of experimental data on thermochemical sulfate reduction expressed as fields on a plot of total S (H S +S0 ") in solution vs. temperature of the experiments. Successful experiments are shown in a cross-hatched pattern, unsuccessful ones in a solid darker grey. Likely fields for the formation of M V T type ore deposits and sour gas deposits are shown respectively by a checkerboard rectangle and a gradationally colored gray rectangle. The H S concentration for the experiments of this study was calculated using the data of (27). 2

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419

A corresponding field for the condition of origin of sour gas fields analogous to the M V T field is also shown as gradational shades of gray. Though the temperatures involved can be constrained by fluid inclusion and burial history data (4,19) and are in fact similar to those shown on Figure 2 for M V T ore formation, the corresponding £ [ S ] is poorly known, and that is the reason for the shading. It is likely that £ [ S ] may be somewhat higher than those shown for M V T ores because most sour gas fields are associated with sulfate-bearing evaporites (79), and would, after only a moderate extent of reaction, also contain significant H S concentrations. There may, thus, exist a difference in a major parameter, £ [ S ] controlling the conditions of H S generation in these two environments. A second way to look at the experimental TSR data is in terms of overall rates, in this case expressed as moles of sulfate reduced/liter/year (Figure 3). In order to facilitate this comparison, the data have been recalculated to a common set of conditions; 0.025M S0 ' and pH 4. These conditions were chosen to be representative of ore fluids implicated in the formation of M V T ore deposits (20), and the pH may also be appropriate for oil field brines such as those of sour gas fields (e.g. 7). The assumptions involved in the recalculation of the data deserve elaboration. The recalculation of rates to the stated S0 " concentration is straightforward, and is based on the assumption of first order kinetics (i.e. the rate is directly proportional to initial concentration of S0 ). The pH correction is much more problematical. We have assumed that the pH dependence is similar to that estimated for the isotope exchange reaction (3), which decreases by one order of magnitude per pH unit between pH 1 and 3 and is constant above 4. We have simply assumed a one order of magnitude decrease in rate per pH unit below 4, although this will introduce some error between pH 3 and 4. The assumed pH dependence was applied to initial run pH's as calculated by computer modeling of the run conditions from calculations in (3) and (7). In die absence of other information, the pH values for (73) were assumed to be 2 based on 25° C measurements. Although this pH correction adds considerable uncertainty to the plotted rates, it is our feeling that it, nonetheless, gives a more reasonable picture of in situ rates than the very low pH values employed in many of the experimental procedures. The effects of variable concentrations of H S and nature of organic matter were not explicitly accounted for, although the role of these variables is addressed qualitatively below. Four separate studies are represented in Figure 3; Kiyosu (77) and later Kiyosu and Krouse (72) completed two series of experiments at various temperatures using dextrose and acetic acid respectively as the organic reductants and no initial added H S . A later set of runs by Kiyosu and Krouse (73) with acetic acid employed similar conditions to (72) but included initial elemental sulfur in molar amounts similar to the initial S0 ". The fourth set of experimental data (this study), used toluene as the organic reductant. The 2


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>v

1.8

.

1993

3

2.0

^Jjt^

(deg

Λ #

3

10 /T

%

290

**··

Τ

···.

: ·

2.2

M

C)

*· '*···...

ι X

•

1

•

190

2.4

2.6

T r u d i n g e r et a l .

M

W C C C field Ψ

ipo

Figure 3. Plot of rate of sulfate reduction expressed as moles reduced per liter of water per year vs. 10 /T. The experiments have all been normalized to a sulfate concentration of 0.025M and pH 4 as discussed in the text. The rates for the Whitney Canyon Carter Creek field are taken from (7). The results from this study are shown as filled circles; the results Kiyosu and Krouse as X's.

.6

^Kiyo^u^^

1990

2£0 Kiyosu and Krouse

Kiyosu and Krouskv

A.

300


23. GOLDHABER & ORR


results from this study are presented as 5 points. The two points plotted at 250° C correspond to the extremes in initial H S pressures employed (Table I), with the highest initial H S pressure resulting in the higher rate. This is equivalent to the overall spread in rates as a function of H S pressure shown in Figure 1. Thus, although initial H S concentration is not explicitly taken into account in our overall rate, Figure 1 and Figure 3 together suggest the range of rates which can be expected when P ^ is varied with other parameters held constant. Additional points are plotted for our results over the temperature range 175-250° C and 200 psi H S , which corresponds to the higher range of H S pressures we studied. Also plotted as a generalized region are the results of Trudinger et al., (9). Even though these were negative runs, it is possible to calculate a detection limit for H S for their experimental design and therefore a maximum possible rate for their experiments. The arrow towards lower rates indicates that these are maximum values. Although not specifically corrected to pH 4 and 0.025M S0 ", many of the experiments did, in fact, fall in this general range. The final data plotted is from a sour gas field; the Whitney Canyon Carter Creek (WCCC) field in the overthrust belt of Wyoming, USA. This rate was derived from geologic constraints and the temperatures from fluid inclusion studies on phases tied paragenetically to TSR. Details are given in (7). The range of values plotted for WCCC in Figure 3, to our knowledge, are the only estimates of TSR rates in a sour gas field. They are minimum values because they do not reflect probable H S losses by leakage from the field or to form FeS . Several interesting points emerge from Figure 3. The rates calculated from Kiyosu'sfirsttwo data sets (11-12), although substantially different at elevated temperatures (>200°C), both extrapolate to near the maximum permitted values for the Trudinger et al.'s experiments (9) at lower temperatures in the 150°C range. The data for runs with the lowest P ^ from this study are comparable to results of Kiyosu and Krouse (12) at 250° C, and thus approach values derived from H S -free conditions. When extrapolated to 150° C, the rates for the two data sets lacking initial H S (11-12) are on the order of 10" moles/liter/year, and

Kinetic Controls on Thermochemical Sulfate Reduction as a Source of

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