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Sulphur isotope equilibrium during sulphur hydrolysis at high temperatures
EARTH AND PLANETARY SCIENCE LETTERS 18 (1973) 443-450. NORTH-HOLLAND PUBLISHING COMPANY
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SULPHUR SULPHUR
ISOTOPE EQUILIBRIUM
HYDROLYSIS
DURING
AT HIGH TEMPERATURES*
Brian W. ROBINSON Institute of Nuclear Sciences, Department of Scientific and Industrial Research, Lower Hutt, New Zealand
Revised version received 6 February 1973 Studies of the sulphur hydrolysis reaction, 4S + 4H20 ~ 3H2S + HSO4 + H~, were conducted between 200 and 320°C in sealed silica glass tubes. The isotope exchange reaction: H2 180 + HS 160~- # H2 160 + HS 180 4 is so rapid at the low pH (1.5-3) as to be unquenchable. However, the sulphur isotope exchange reaction: H2 34S + H 32SO4-# H2 32S + H 3450 4 gave t~ values of 0.1, 0.3 and 1.7 days at 320, 260 and 200°C respectively and equilibrium H2S - HSO4 sulphur isotope fractionation values of 20.9, 22.4, 24.8, 26.7 and 29.3%o at 320, 290, 260, 230 and 200°C respectively. This latter data is represented by: 1000 In a(HSO4_H2S) = 5.07 (10 6 T -2) + 6.33, and has valuable applications in geothermal and ore deposit studies.
1. Introduction Dissolved sulphate is one of the most abundant constituents of hot spring and geothermal waters. It is generally believed to have been produced from the oxidation of hydrogen sulphide and sulphur dioxide by atmospheric oxygen in the near surface environment [ 1,2]. However, concentrations of sulphate in the order o f several grams per litre often contained in hot spring water are difficult to reconcile with surface oxidation alone. Thus, both the disproportionation reaction of sulphurous acid and the reaction of sulphur with water (sulphur hydrolysis) have been pro: posed as important mechanisms for the occurrence of sulphate in volcanic thermal water [3,4]. The sulphur hydrolysis mechanism has often been regarded as improbable due to an insufficient excess of elemental sulphur. However, in the Tatun area, N. Taiwan and at Rotokaua, New Zealand, high temperature waters come intc contact with elemental sulphur and produce acidic, sulphate solutions [4]. In both areas sulphur containing beds occur to depths of a few hundred meters and produce acidic, sulphate rich waters, whereas deeper drillholes encounter neutral pH, hot waters low in sulphate. * I.N.S. contribution no. 563.
Furthermore, the above mentioned sulphur hydrolysis and disproportionation of sulphurous acid reactions, which are respectively: 4S + 4H:O ~ 3H2S + H2504
(1)
3HzSOs ~ 2HzSO4 + S + H20
(2)
also prompted sulphur isotope studies and Oana and Ishikawa [3] investigated sulphur isotope fractionation between sulphur and sulphuric acid whilst Oana and Mizutani [5] published data on the fractionation between hydrogen sulphide and sulphuric acid formed by reaction (1). Neither of these studies yielded a correlation between temperature and the isotope fractionatiop factors, probably because of slow reaction rates. The present study was therefore initiated to measure the sulphur isotope fractionation factor between hydrogen sulphide and sulphuric acid for reaction (1) using much longer reaction times. Also, the isotope exchange reaction between the species was studied at different time intervals to determine the rate of the reaction. The results are of relevance in geothermal studies where the sulphur isotope ratios of hydrogen sulphide and sulphuric acid from a surface discharge can indicate the water temperature at depth.
444
B. I4. Robinson, Sulphur isotope equilibrium
In addition, sulphur isotope exchange reactions are now of relevance in the study of hydrothermal ore deposits. Ohmoto [6] has shown that changes in the oxidation state of an ore fluid, in particular a change from oxidised to reduced sulphur solute species, or vice versa, can have a profound effect on the 8 34S values of precipitating sulphides, providing that equilibrium conditions prevail. If the temperature and chemistry of the ore solution and the 8 34S values of precipitated sulphides can be determined, then knowledge of the relevant sulphur isotope fractionation factors can be applied to calculate the 8 34S value of the total sulphur in solution (8 34Sxs ). This value, and not the 8 34S value of the sulphides, reflects the origin of the sulphur.
For simple isotope exchange reactions involving gaseous molecules, the equilibrium constants can be calculated by the methods given in Urey [7], Bigeleisen and Mayer [8] and Bigeleisen [9]. Tudge and Thode [ 10] calculated partition function ratios for sulphur compounds to determine the equilibrium constants for exchange reactions at 0°C and 25°C. Sakai [2,11], Melhuish in [1] and Thode et al. [12] extended this study to high temperatures. Theory predicts that the concentration of the heavy isotope increases with increase in the oxidation state [7] and [10]. Also considering the exchange equilibrium constant, K, l n K varies in proportion to l I T at low temperature, to l I T 2 at high temperature and approaches zero at very high temperatures. Fig. 3 gives a theoretical curve for sulphur isotope fractionation between SO4 and H2S on a 1000 l n a versus l I T z plot [11]. The fractionation factor, a, between two species A and B i s R a / R B, where R is the ratio of the heavy to the light isotope. Also the relative enrichment of one isotope in A with respect to B is given by the enrichment factor CAB = a -- 1 or AAB = 1000 (c~--l). In 6 * terminology it can be shown that: * The isotopic compositions of a sample (i) and the differences between samples are reported in permil deviation from a standard (s): Ri - Rs
R ~ - X 1000.
1 +
8A/1000
1 +
8B/lO00
and
8 A --8 B l O 0 0 ( a - 1 ) = 1 + 8B/1000 (4)
Hence AAB ~ 8 A - 8 B .
(3)
(5)
Fractionation factors obtained from theoretical calculations tend to be expressed in the form of 1000 in a, which can be expanded to give the following approximation: 1000
In a
~ 6n - 8 B "
(6)
For the sulphur hydrolysis reaction above 200°C the sulphate produced is present essentially as HSO~ions [ 13]. Thus the reaction can be expressed as 4S0) + 4H200) -~ 3H:S(aq) + I f + H S 0 4
2. Theoretical background
8i(%e) =
=
(7)
and the sulphur isotope exchange reaction being studied is H2 34S + H 32SO4 ~ H2 32S + H 34SO4
(8)
where all species are aqueous (aq). The oxygen isotope exchange reaction studied is HE x80 + HS 1604 ~- HE 160 + HS 1804 .
(9)
For these reactions a = K after isotopic equilibrium is attained, where K is the equilibrium constant. The half time, t~, for the reaction is defined as the time required for the value of 1 - F to fall to one half of the initial value, where F = A 34St/A 34Soo, the fraction of the isotope exchanged at time t [ 14]. A positive linear relationship on a In ( l - F ) versus time plot characterises a first order reaction where: k = in 2/t+(sec) and k is the rate constant for the reaction. The logarithm of the rate constants plotted against the reciprocal of the absolute temperature (Arrhenius plot) normally gives a straight line. For this line: Slope = - E / 2 . 3 0 R, where E is the activation energy and R the gas constant.
3. Experimental
techniques
3.1. Hydrolysis reactions
Silica glass tubes (35 ml) with a break seal at one end and loading-sealing facilities at the other were
B. W. Robinson, Sulphur isotope equilibrium
445
used for the reactions. Known amounts of pure sulphur (between 0.3 and 0.8 g weighed to -+0.1 mg) and boiled, distilled water (between 18 and 24 g weighed to -+0.01 g) were placed in the tubes. Remaining air was flushed out with nitrogen and the tubes were sealed under 1 mm nitrogen total pressure with the water frozen. 100 ml stainless steel reaction vessels carried the tubes as well as water to balance the steam pressure. Horizontally mounted furnaces, whose temperatures could be controlled to within -+l°C held the reaction vessels which had two chromel-alumel thermocouples strapped to the outside for a continuous record of temperatures and a monitor on thermal gradients. Internal thermocouples showed no difference between inside and outside temperatures. For quenching, the reaction vessels were cooled in air, then hot and finally cold water for 1 rain intervals each. More rapid cooling fractured the tubes.
the oxygen isotopes are given relative to the starting water and the overall error for the method is taken as +0.3%o. The sulphur isotope analyses of the silver sulphur and the sulphur were performed using the mass spectrometric techniques described in [ 18] on SO2 prepared by burning the silver sulphide and the sulphur in a stream of pure tank oxygen at 1100° and 800°C respectively [ 19]. Mixing and memory effects of the mass spectrometer were found to be negligible. Since A 34 S(HSO4_H2S) values are used, no further corrections to the data need be considered. All the analyses are given with respect to the sulphur starting material and the reproducibility for the whole procedure is normally -+0.1%o and never greater than -+0.2%o. The silver sulphide and elemental sulphur burnings have been shown to give equivalent 6 34S values within the quoted errors [20].
3. 2. Sulphur species separation
4. Results
dreakseals on the tubes were carefully opened under excess acidified silver nitrate solution, where the H2S(gas+aq) was precipitated as silver sulphide. At the dilutions used (total volume about 300 ml) silver sulphate is soluble and the excess silver ions were precipitated as silver chloride and filtered off. The sulphate was precipitated as barium sulphate and, after filtration, the remaining solution was tested for the presence of other sulphur species by oxidation with bromine water and the further addition of barium chloride solution. No other sulphur species were detected, confirming previous work [4,5]. After washing inside the tube the main body was broken and the sulphur filtered off. Between 10.9 and 29.3 mg of H2S and 11.9 to 29.4 mg of H2SO4 were produced in the reactions and the concentration of H2SO4 and the HzS/H2SO4 molar ratios are given in table 1.
All the experiments were run from the same water batch which had a 6 180 value of - 5 . 7 + 0.2%0 (SMOW). The 8 180 values of the bisulphate are given in table 1 with respect to the starting water, and are thus equivalent to A 180(HSO2~_H20) values. These values can be seen to increase with the run time and, for the maximum time at each temperature, to increase with decreasing temperature. However, under the conditions of high temperature and low pH of this study, reaction (9) is so rapid as to be unquenchable [21 ]. Thus the results are not equilibrium values for the run temperature, but equilibrium values frozen in at temperatures between 210 and 160°C [22]. Since no additional useful data on this system could be obtained without changing the chemistry of the runs, only half of the sulphate samples were analysed for the oxygen isotopic composition. For the sulphur isotope exchange reaction (8) runs at 320°C were conducted for 0.5, 2.5 and 5 days, at 260°C for 1, 2 and 8 days and at 200°C for 2.5, 5, 10, 21, and 42 days. The results of these runs are shown in table 1, and as data points in fig. 1 where the straight line plots demonstrate a first order reaction. Also the half times of the reactions can be determined as 1.7 + 0.1, 0.3 --- 0.05, 0.1 +- 0.03 days at 200, 260 and 320°C respectively. All reactions were run in ex-
3.3. Isotope techniques Oxygen isotope analysis of the sulphate was performed on CO2 evolved by the graphite reduction method [ 15], using an internally heated reaction vessel [ 16]. The barium sulphide formed is converted to silver sulphide for the analysis of the sulphur isotopes of the sulphate [ 17]. All the sulphate results for
4 4
1 2 8 8
10 10
2.5 5 10 21 42
290 290
260 260 260 260
230 230
200 200 200 200 200
0.0084 0.0056 0.0054 0.0046 0.0061
0.0088 0.0078
0.0092 0.0117 0.0107 0.0098
0.0135 0.0122
0.0121 0.0128 0.0129
ftzSO4 formed (moles/kg)
2.92 2.89 3.05 2.63 2.82
2.95 2.96
2.74 2.91 3.00 2.95
2.84 2.97
2.87 2.91 2.83
Molar ratio H2S/HzSO 4
38.6 30.2 23.8 20.7 27.5
38.5 37.6
28.1 33.6 39.9 32.6
49.6 40.1
31.8 49.5 46.4
Total loss of sulphur (rag)
* All ~ 345 values are given here with respect to the starting sulphur. ** F is the fraction of isotopes exchanged (see text). t a is the fractionation factor (see text). t ? All 6 1 8 0 values are given here with respect to the starting water.
0.5 2.5 5
Time (days)
320 320 320
Temperature (°C)
Table 1 Chemical and isotopic data from the sulphur hydrolysis reactions
+0.2 +0.4 +0.3 0.0 +0.2
+0.4 +0.2
0.0 +0.3 +0.4 +0.1
+0.2 +0.1
-0.3 0.0 0.0
+1.8 +1.5 +0.4 +1.2 +1.6
+1.7 +2.1
+1.2 +2.1 +1.9 +2.2
+2.5 +1.6
+2.3 +2.0 +1.9
+21.1 +26.4 +29.0 +30.4 +30.9
+28.3 +28.8
+23.8 +26.6 +26.5 +27.1
+24.6 +24.2
+22.7 +22.9 +22.8
19.3 24.9 28.6 29.2 29.3529. 31
26.6 26.7526"7 ~
22.6 24.5 24.6-~ 4 l8 24.9 ~2
22.1 22.6522. 41
20.4 20.9 20.9 y20"9 a
(%o)
(%o)
(%o)
HSO4
A 34S H2S-HSO 4
6345
6345
Residual H2S sulphur (%0)
6 345*
0
0.342 0.150 0.024
0
0
0.088 0.011
0
0
0.024
(l-F)**
28.9
26.3
24.5
22.2
20.7
1000 In c~t
6 180
8.9 9.7 9.5 10.7 10.7
9.1 9.7
8.4 9.2
(%~)
HS04 tt
2.
z
4~
B. W. Robinson, Sulphur isotope equilibrium
447
1.0 0.5~N~4~
II/ \ It~
I1~
~
II/
",t,
l
200
T.c
1.7 \
T ('C) 260
320
-4.0
200
.s
26c
,
32°
\
tn
200
4-73x166
-5'0 ~.)
"01
-0031 I o
I 2
t 4
I I 6 8 Time ( d a y s )
I 10
-6"0
12
Fig. 1. First order rate plot of log ( l - F ) versus time for the sulphur isotope exchange reaction at 200, 260 and 320°C. The data points are shown with error bars and the t~- values are 1.7 -* 0.1, 0.3 +- 0.05 and 0.1 -+ 0.03 days for 200, 260 and 320°C respectively.
cess of ten times t~ to obtain equilibrium values. Fig. 2, an Arrhenius plot, gives the reaction rates and demonstrates that below 200°C reaction (8) is very slow. At all temperatures the runs were made in duplicate with run time in excess of ten times t~ to check the overall reproducibility. Normally the 1000 in ct values have an error of-+0.2 to 0.4%o and the difference between duplicates did not exceed 0.5%o. Fig. 3 gives all the 1000 lnc~ points and their correlation with temperature, together with error bars on each data point.
I 1.7
1"6
I I I 1"8 1-9 2"0 T -~ (" K )-t x 103
I 2.1
2.2
Fig. 2. An Arrhenius plot for the sulphur isotope exchange reaction with rate constants determined at 200, 260 and 320°C. The data points are shown with error bars and the experimental activation energy (A) is 13.3 kcal mole-1 .
500 I
4 0 0 350 300 I
I
I
I
T ('C) 250
200
150
I
I
I
40
30 _E o
~ 20
5. Discussion 10
The sulphur hydrolysis reaction (1) is relatively fast, compared to the sulphur isotope exchange reaction (8), and the products were found to reach a constant value after 5 hr at 250°C and 1½ hr at 300°C [5]. However, the reverse reaction is very slow at 200°C [4] and for this study equilibrium was only approached from one direction. The yield controlling parameters for the reaction are temperature and pH,
2
3
4 T-2 ( ' K ) -2 x 10 6
5
6
Fig. 3. A 1000 In a versus lIT 2 plot showing the theoretical data for SO4-H2S sulphur isotope exchange, Sakai [11] (as squares) and the experimental data points from this study for HSO~-H2S sulphur isotope exchange (as circles). The error on the latter points is the vertical diameter of the circles.
448
B. W. Robinson, Sulphur isotope equilibrium
and the concentration of the products increases with increasing temperature [4]. The precise mechanism for the sulphur hydrolysis reaction is not known. However, the initial effect of the reaction on the sulphur isotope exchange reaction is a kinetic one, which favours the light isotope in the sulphate and the heavy to enter the hydrogen sulphide. This effect is superimposed by the much larger exchange equilibrium effect, which proceeds at a much slower rate. The sulphur species formed by the hydrolysis reaction would be expected to have a mean/5 34S close to that of the elemental sulphur, i.e., at 200°C and 42 days the H2S and sulphate should have values about -7%o and +21%o respectively. However, the species show a strong depletion in 32S from this common mean, which is either due to the hydrolysis mechanism or the factors mentioned below. A variable excess of sulphur was used in the reactions and the concentrations of the products formed and the isotope exchange are independent of this. An indication that the overall reaction studied was reaction (1) is given by the H2S/H2SO4 ratio being close to 3:1 (see table 1). The slight loss of H2S may be caused during sampling or by solution in the liquid sulphur [4]. The solubility of H2S in liquid sulphur has been shown to be twice its aqueous solubility for temperatures between 200 and 320°C [23]. However, the relatively slow quench times of the reactions and the low weight ratios of sulphur to water would lead to little HzS being trapped in the sulphur. In addition to the reaction products, about 5 to 15 mg of sulphur is lost during the reaction. This phenomenon was noted in the silica tube experiments of Oana and Mizutani [5] and may be caused by diffusion into the glass, although there is no correlation with time or temperature. Since the total HzS is sampled after the runs, any isotopic fractionation between H2S(liqui d S), H2S(aq), and H2S(gas) during the reaction will have affected the results. Upon cooling to room temperature, most of the H2S will partition into the aqueous phase [24]. The isotopic fractionation factors for HzS in the phases mentioned have not been determined but, in analogy to the other systems, such as SO2(gas)-SOE(aq) [25], these can be expected to be small at high temperatures and their effect has been neglected in this study.
After a run, the sulphur cools as a large ovoid immiscible bleb, having a brown outer shell and a yellow interior, together with numerous small (0.5) mm yellow spherutes. These different phases are thought to represent the liquid and gaseous portions of sulphur present at high temperatures. Thus for the runs at 260°C and 8 days the phases were separated and analysed; the isotopic values did not differ outside of the errors of measurement. Fig. 1 demonstrates that equilibrium was achieved in the sulphur isotope exchange reactions, and the reaction rate is similar to that for the H2SO4-SO2 sulphur isotope equilibrium at 200°C [26]. From the Arrhenius plot (fig. 2) an experimental activation energy of 13.3 kcal/mole is calculated. No attempt has been made to investigate the influences of the sulphur hydrolysis pre-equilibrium and associated variations in reaction products with time and temperature on the isotopic fractionation rates. However, this is included here to facilitate future comparisons with other sulphur isotope exchange reactions. Fig. 3 (a 1000 lnc~ versus 1/T 2 plot) gives the sulphur isotope fractionation data from this study (HSO; H2S ) and the theoretical curve (SO~-H2S) from Sakai [ 11 ]. The experimental points probably form part of a curve of similar form to the theoretical curve. In fact, an inflection point at 260°C appears to be present. However, extrapolation is difficult, and the experimental points obtained in this study are defined by the least squares best fit straight line: 1000 lna(HSO]-H2S) = 5.07 (10 6 T -2) + 6.33 for temperatures between 200°C and 320°C. A direct comparison with Sakai's curve should not be made. More recent data on the partition function ratios of H2 34S and H2 32S [12] differ from Sakai's by 2%0 at 200°C, and better estimates for SO4 partition functions should soon be published..4dso, undoubtedly HSOT~--H2S and SO4-H2S will behave differently. How differently will not be known until calculations are done on HSOT, partition functions or the systems are compared experimentally. It has also been pointed out by Thode et al. [ 12] that, during experimental determinations, other sulphur species found could be quantitatively recovered as silver sulphide or barium sulphate. For the chemistry and conditions of the present study, this is extremely improbable.
B. IV. Robinson, Sulphur isotope equilibrium
6. Conclusions Isotope exchange reactions have been used to calculate the temperatures at depth in natural hydrothermal systems [22,27,28]. For the Wairakei geothermal area, New Zealand, the average A 34S(HsO4_HzS ) is 19%o [28] which, using the data obtained in this study, indicates a temperature o f about 350°C. The maximum measured temperature in this area is 270°C, and the isotope temperature probably represents that at greater depth. The exchange reaction studied can be applied to ore deposit studies: at the Echo Bay deposit, N.W.T., Canada, early formed acanthite has a 5 a4S value o f -20%o whilst the ore fluid is thought to have been in the HSO?, stability field at a temperature o f 150°C [29]. Using the H2S-Ag2S fractionation factor from [30] and the H S O 4 - H 2 S fractionation factor from an extrapolation of this study the ~ 34Szs value for the ore fluid is +20%o, thus suggesting trapped sea-water sulphate as the origin of the sulphur. However, when adapting the results of this study to sulphide-sulphate fractionation in hydrothermal ore solutions, several differences should be noted. For the average conditions o f an ore forming fluid, say 250°C, pH 3 . 5 - 7 . 5 and fs2 ~ 10-1°, sulphur isotope exchange in solution will probably not be as rapid as that during the reactions with pH 1 . 5 - 3 and fs2 ~ 10-3 - 10-6. In analogy to the oxygen isotope exchange rate between dissolved sulphate and water [21 ], the exchange rate was shown to be extremely sensitive to pH, and was very rapid at low pH. Also the sulphur isotope exchange rate between metallic sulphides is dependent on fS2, the rate increasing with increasing fs2 [31 ]. It thus appears that the reaction rates determined here could be slower in many natural systems.
Acknowledgements The high temperature part of this study was conducted in the laboratories of Chemistry Division, D.S.I.R., where the co-operation of my colleagues there is gratefully acknowledged. Discussions with Dr. T.A. Rafter, Dr. J.R. Hulston and Dr. W. Giggenbach have been o f value during the course o f this study.
449
Dr. C.B. Taylor kindly supplied the 8 180 analysis of the water starting material. Mr. D.W.N. Keith and Miss D.R. Pettis are thanked for their help with laboratory and mass spectrometry procedures.
References [ 1] T.A. Rafter, S.H. Wilson and B.W. Shilton, Sulphur isotopic variations in nature. Part 5, Sulphur isotopic variations in New Zealand geothermal bore waters, New Zealand J. Sci. 1 (1958) 103. [2] H. Sakai, Fractionation of sulphur isotopes in nature, Geochim. Cosmochim. Acta 12 (1957) 150. [3] S. Oana and H. lshikawa, Sulfur isotopic fractionation between sulfur and sulfuric acid in the hydrothermal solution of sulfur dioxide, Geochem. J. 1 (1966) 45. [4] A.J. Ellis and W. Giggenbach, Hydrogen sulphide ionization and sulphur hydrolysis in high temperature solution, Geochim. Cosmochim. Acta 35 (1971) 247. [5] S. Oana and Y. Mizutani, Sulphur isotopic fractionation between hydrogen sulphide and sulphuric acid formed by the 4S + 4H20 = 3H2S + H2SO4 reaction, Symp. Geochem. Soc. Japan, Nagoya (1966). [6] H. Ohmoto, Systematics of sulfur and carbon isotopes in hydrothermal ore deposits, Econ. Geol. 67 (1972) 551. [7] H.C. Urey, The thermodynamic properties of isotopic substances, J. Chem. Soc. (1947) 562. [8] J. Bigeleisen and M.G. Mayer, Calculation of equilibrium constants for isotopic exchange reactions, J. Chem. Phys. 15 (1947) 261. [9] J. Bigeleisen, Statistical mechanics of isotope effects on the thermodynamic properties of condensed systems, J. Chem. Phys. 34 (1961) 1485. [ 10] A.P. Tudge and H.G. Thode, Thermodynamic properties of isotopic compounds of sulphur, Can. J. Res. 28B (1950) 567. [11] H. Sakai, Isotopic properties of sulfur compounds in hydrothermal processes, Geochem. J. 2 (1968) 29. [12] H.G. Thode, C.B. Cragg, J.R. Hulston and C.E. Rees, Sulphur isotope exchange between sulphur dioxide and hydrogen sulphide, Geochim. Cosmochim. Acta 35 (1971) 35. [13] A.S. Quist, W.L. Marshall and H.R. Jolley, Electrical conductance of aqueous solutions at high temperature and pressure. Part II, The conductances and ionization constants of sulfuric acid-water solutions from 0° to 800 ° and pressures up to 4000 bars, J. Phys. Chem. 69 (1965) 2726. [14] A.C. Wahl and N.A. Bonner (eds.), Radioactivity applied to chemistry (Wiley and Sons, New York, 1951) p. 604.
450
B.W. Robinson, Sulphur isotope equilibrium
[15] T.A. Rafter, Oxygen isotopic composition of sulphates. Part I, A method for the extraction of oxygen and its quantitative conversion to carbon dioxide for isotope radiation measurements, New Zealand J. Sci. 10 (1967) 402. [16] H. Sakai and H.R. Krouse, Elimination of memory effects in 180/160 determinations in sulphates, Earth Planet. Sci. Letters 11 (1971) 369. [ 17] T.A. Rafter, Sulphur isotopic variations in nature. Part 2, A quantitative study of the reduction of barium sulphate by graphite for recovery of sulphide-sulphur for sulphur isotopic measurements, New Zealand J. Sci. and Tech. 38 (1957) 955. [ 18] J.R. Hulston and B.W. Shilton, Sulphur isotopic variations in nature. Part 4, Measurements of sulphur isotopic ratio by mass spectrometry, New Zealand J. Sci. 1 (1958) 91. [19] T.A. Rafter, Sulphur isotopic variations in nature. Part 1, The preparation of sulphur dioxide for mass spectrometer examination, New Zealand J. Sci. Tech. 38 (1957) 849. [20] J.R. Hulston and J. Beauchamp, Sulphur isotopic variations in nature. Part 9, A recomparison of the New Zealand sulphur standard with the meteoritic sulphur standard, New Zealand J. Sci. 6 (1963) 157. [21 ] R.M. Lloyd, Oxygen isotope behaviour in the sulfatewater system, J. Geophys. Res. 73 (1968) 6099. [22] Y. Mizutani and T.A. Rafter, Oxygen isotopic composition of sulphates. Part 3, Oxygen isotopic fractionation in the bisulphate ion-water system, New Zealand J. Sci. 12 (1969) 54.
[23] R. Fanelli, Solubility of hydrogen sulfide in sulphur, Ind. Eng. Chem. 41 (1949) 2031. [24] T.N. Kozintseva, Solubility of hydrogen sulfide in water and solutions under high temperature, in: Geochemical investigations in the field of high pressures and temperatures, N.I. Khitarov, ed. (Nauka. Pub. Office, Moscow, 1965) p. 121. [25] T.E. Eriksen, Sulphur isotope effects. II, The isotopic exchange coefficients for the sulfur isotopes 345-325 in the system SO2g-aqueous solutions of SO2, Acta Chem. Scand. 26 (1972) 581. [26] H.B. Dunford, A.G. Harrison and H.G. Thode, Equilibrium constants for the sulfur isotope exchange between SO2 and H2SO4, Can. J. Chem. 35 (1957) 817. [27] N. Nakai, Isotopic and chemical equilibrium of CH4-CO2 and pyrite-anhydrite in geothermal area in Japan, Abs. Int. Syrup. Hydrogeochem. and Biogeochem., Tokyo (1970). [28] A. Steiner and T.A. Rafter, Sulfur isotopes in pyrite, pyrrhotite, alunite and anhydrite from steam wells in the Taupo volcanic zone, New Zealand, Econ. Geol. 61 (1966) 1115. [29] B.W. Robinson, Studies of the Echo Bay silver deposit, N.W.T., Canada, Unpub. Ph.D. Thesis, University of Alberta (1971). [30] B.W. Robinson and H. Ohmoto, Mineralogy, fluid inclusions and stable isotopes of the Echo Bay U - A g - C u deposits, Northwest Territories, Canada, in preparation. [31] Y. Kajiwara and H.R. Krouse, Sulfur isotope partitioning in metallic sulfide systems, Can. J. Earth Sci. 8 (1971) 1397.
[]
SULPHUR SULPHUR
ISOTOPE EQUILIBRIUM
HYDROLYSIS
DURING
AT HIGH TEMPERATURES*
Brian W. ROBINSON Institute of Nuclear Sciences, Department of Scientific and Industrial Research, Lower Hutt, New Zealand
Revised version received 6 February 1973 Studies of the sulphur hydrolysis reaction, 4S + 4H20 ~ 3H2S + HSO4 + H~, were conducted between 200 and 320°C in sealed silica glass tubes. The isotope exchange reaction: H2 180 + HS 160~- # H2 160 + HS 180 4 is so rapid at the low pH (1.5-3) as to be unquenchable. However, the sulphur isotope exchange reaction: H2 34S + H 32SO4-# H2 32S + H 3450 4 gave t~ values of 0.1, 0.3 and 1.7 days at 320, 260 and 200°C respectively and equilibrium H2S - HSO4 sulphur isotope fractionation values of 20.9, 22.4, 24.8, 26.7 and 29.3%o at 320, 290, 260, 230 and 200°C respectively. This latter data is represented by: 1000 In a(HSO4_H2S) = 5.07 (10 6 T -2) + 6.33, and has valuable applications in geothermal and ore deposit studies.
1. Introduction Dissolved sulphate is one of the most abundant constituents of hot spring and geothermal waters. It is generally believed to have been produced from the oxidation of hydrogen sulphide and sulphur dioxide by atmospheric oxygen in the near surface environment [ 1,2]. However, concentrations of sulphate in the order o f several grams per litre often contained in hot spring water are difficult to reconcile with surface oxidation alone. Thus, both the disproportionation reaction of sulphurous acid and the reaction of sulphur with water (sulphur hydrolysis) have been pro: posed as important mechanisms for the occurrence of sulphate in volcanic thermal water [3,4]. The sulphur hydrolysis mechanism has often been regarded as improbable due to an insufficient excess of elemental sulphur. However, in the Tatun area, N. Taiwan and at Rotokaua, New Zealand, high temperature waters come intc contact with elemental sulphur and produce acidic, sulphate solutions [4]. In both areas sulphur containing beds occur to depths of a few hundred meters and produce acidic, sulphate rich waters, whereas deeper drillholes encounter neutral pH, hot waters low in sulphate. * I.N.S. contribution no. 563.
Furthermore, the above mentioned sulphur hydrolysis and disproportionation of sulphurous acid reactions, which are respectively: 4S + 4H:O ~ 3H2S + H2504
(1)
3HzSOs ~ 2HzSO4 + S + H20
(2)
also prompted sulphur isotope studies and Oana and Ishikawa [3] investigated sulphur isotope fractionation between sulphur and sulphuric acid whilst Oana and Mizutani [5] published data on the fractionation between hydrogen sulphide and sulphuric acid formed by reaction (1). Neither of these studies yielded a correlation between temperature and the isotope fractionatiop factors, probably because of slow reaction rates. The present study was therefore initiated to measure the sulphur isotope fractionation factor between hydrogen sulphide and sulphuric acid for reaction (1) using much longer reaction times. Also, the isotope exchange reaction between the species was studied at different time intervals to determine the rate of the reaction. The results are of relevance in geothermal studies where the sulphur isotope ratios of hydrogen sulphide and sulphuric acid from a surface discharge can indicate the water temperature at depth.
444
B. I4. Robinson, Sulphur isotope equilibrium
In addition, sulphur isotope exchange reactions are now of relevance in the study of hydrothermal ore deposits. Ohmoto [6] has shown that changes in the oxidation state of an ore fluid, in particular a change from oxidised to reduced sulphur solute species, or vice versa, can have a profound effect on the 8 34S values of precipitating sulphides, providing that equilibrium conditions prevail. If the temperature and chemistry of the ore solution and the 8 34S values of precipitated sulphides can be determined, then knowledge of the relevant sulphur isotope fractionation factors can be applied to calculate the 8 34S value of the total sulphur in solution (8 34Sxs ). This value, and not the 8 34S value of the sulphides, reflects the origin of the sulphur.
For simple isotope exchange reactions involving gaseous molecules, the equilibrium constants can be calculated by the methods given in Urey [7], Bigeleisen and Mayer [8] and Bigeleisen [9]. Tudge and Thode [ 10] calculated partition function ratios for sulphur compounds to determine the equilibrium constants for exchange reactions at 0°C and 25°C. Sakai [2,11], Melhuish in [1] and Thode et al. [12] extended this study to high temperatures. Theory predicts that the concentration of the heavy isotope increases with increase in the oxidation state [7] and [10]. Also considering the exchange equilibrium constant, K, l n K varies in proportion to l I T at low temperature, to l I T 2 at high temperature and approaches zero at very high temperatures. Fig. 3 gives a theoretical curve for sulphur isotope fractionation between SO4 and H2S on a 1000 l n a versus l I T z plot [11]. The fractionation factor, a, between two species A and B i s R a / R B, where R is the ratio of the heavy to the light isotope. Also the relative enrichment of one isotope in A with respect to B is given by the enrichment factor CAB = a -- 1 or AAB = 1000 (c~--l). In 6 * terminology it can be shown that: * The isotopic compositions of a sample (i) and the differences between samples are reported in permil deviation from a standard (s): Ri - Rs
R ~ - X 1000.
1 +
8A/1000
1 +
8B/lO00
and
8 A --8 B l O 0 0 ( a - 1 ) = 1 + 8B/1000 (4)
Hence AAB ~ 8 A - 8 B .
(3)
(5)
Fractionation factors obtained from theoretical calculations tend to be expressed in the form of 1000 in a, which can be expanded to give the following approximation: 1000
In a
~ 6n - 8 B "
(6)
For the sulphur hydrolysis reaction above 200°C the sulphate produced is present essentially as HSO~ions [ 13]. Thus the reaction can be expressed as 4S0) + 4H200) -~ 3H:S(aq) + I f + H S 0 4
2. Theoretical background
8i(%e) =
=
(7)
and the sulphur isotope exchange reaction being studied is H2 34S + H 32SO4 ~ H2 32S + H 34SO4
(8)
where all species are aqueous (aq). The oxygen isotope exchange reaction studied is HE x80 + HS 1604 ~- HE 160 + HS 1804 .
(9)
For these reactions a = K after isotopic equilibrium is attained, where K is the equilibrium constant. The half time, t~, for the reaction is defined as the time required for the value of 1 - F to fall to one half of the initial value, where F = A 34St/A 34Soo, the fraction of the isotope exchanged at time t [ 14]. A positive linear relationship on a In ( l - F ) versus time plot characterises a first order reaction where: k = in 2/t+(sec) and k is the rate constant for the reaction. The logarithm of the rate constants plotted against the reciprocal of the absolute temperature (Arrhenius plot) normally gives a straight line. For this line: Slope = - E / 2 . 3 0 R, where E is the activation energy and R the gas constant.
3. Experimental
techniques
3.1. Hydrolysis reactions
Silica glass tubes (35 ml) with a break seal at one end and loading-sealing facilities at the other were
B. W. Robinson, Sulphur isotope equilibrium
445
used for the reactions. Known amounts of pure sulphur (between 0.3 and 0.8 g weighed to -+0.1 mg) and boiled, distilled water (between 18 and 24 g weighed to -+0.01 g) were placed in the tubes. Remaining air was flushed out with nitrogen and the tubes were sealed under 1 mm nitrogen total pressure with the water frozen. 100 ml stainless steel reaction vessels carried the tubes as well as water to balance the steam pressure. Horizontally mounted furnaces, whose temperatures could be controlled to within -+l°C held the reaction vessels which had two chromel-alumel thermocouples strapped to the outside for a continuous record of temperatures and a monitor on thermal gradients. Internal thermocouples showed no difference between inside and outside temperatures. For quenching, the reaction vessels were cooled in air, then hot and finally cold water for 1 rain intervals each. More rapid cooling fractured the tubes.
the oxygen isotopes are given relative to the starting water and the overall error for the method is taken as +0.3%o. The sulphur isotope analyses of the silver sulphur and the sulphur were performed using the mass spectrometric techniques described in [ 18] on SO2 prepared by burning the silver sulphide and the sulphur in a stream of pure tank oxygen at 1100° and 800°C respectively [ 19]. Mixing and memory effects of the mass spectrometer were found to be negligible. Since A 34 S(HSO4_H2S) values are used, no further corrections to the data need be considered. All the analyses are given with respect to the sulphur starting material and the reproducibility for the whole procedure is normally -+0.1%o and never greater than -+0.2%o. The silver sulphide and elemental sulphur burnings have been shown to give equivalent 6 34S values within the quoted errors [20].
3. 2. Sulphur species separation
4. Results
dreakseals on the tubes were carefully opened under excess acidified silver nitrate solution, where the H2S(gas+aq) was precipitated as silver sulphide. At the dilutions used (total volume about 300 ml) silver sulphate is soluble and the excess silver ions were precipitated as silver chloride and filtered off. The sulphate was precipitated as barium sulphate and, after filtration, the remaining solution was tested for the presence of other sulphur species by oxidation with bromine water and the further addition of barium chloride solution. No other sulphur species were detected, confirming previous work [4,5]. After washing inside the tube the main body was broken and the sulphur filtered off. Between 10.9 and 29.3 mg of H2S and 11.9 to 29.4 mg of H2SO4 were produced in the reactions and the concentration of H2SO4 and the HzS/H2SO4 molar ratios are given in table 1.
All the experiments were run from the same water batch which had a 6 180 value of - 5 . 7 + 0.2%0 (SMOW). The 8 180 values of the bisulphate are given in table 1 with respect to the starting water, and are thus equivalent to A 180(HSO2~_H20) values. These values can be seen to increase with the run time and, for the maximum time at each temperature, to increase with decreasing temperature. However, under the conditions of high temperature and low pH of this study, reaction (9) is so rapid as to be unquenchable [21 ]. Thus the results are not equilibrium values for the run temperature, but equilibrium values frozen in at temperatures between 210 and 160°C [22]. Since no additional useful data on this system could be obtained without changing the chemistry of the runs, only half of the sulphate samples were analysed for the oxygen isotopic composition. For the sulphur isotope exchange reaction (8) runs at 320°C were conducted for 0.5, 2.5 and 5 days, at 260°C for 1, 2 and 8 days and at 200°C for 2.5, 5, 10, 21, and 42 days. The results of these runs are shown in table 1, and as data points in fig. 1 where the straight line plots demonstrate a first order reaction. Also the half times of the reactions can be determined as 1.7 + 0.1, 0.3 --- 0.05, 0.1 +- 0.03 days at 200, 260 and 320°C respectively. All reactions were run in ex-
3.3. Isotope techniques Oxygen isotope analysis of the sulphate was performed on CO2 evolved by the graphite reduction method [ 15], using an internally heated reaction vessel [ 16]. The barium sulphide formed is converted to silver sulphide for the analysis of the sulphur isotopes of the sulphate [ 17]. All the sulphate results for
4 4
1 2 8 8
10 10
2.5 5 10 21 42
290 290
260 260 260 260
230 230
200 200 200 200 200
0.0084 0.0056 0.0054 0.0046 0.0061
0.0088 0.0078
0.0092 0.0117 0.0107 0.0098
0.0135 0.0122
0.0121 0.0128 0.0129
ftzSO4 formed (moles/kg)
2.92 2.89 3.05 2.63 2.82
2.95 2.96
2.74 2.91 3.00 2.95
2.84 2.97
2.87 2.91 2.83
Molar ratio H2S/HzSO 4
38.6 30.2 23.8 20.7 27.5
38.5 37.6
28.1 33.6 39.9 32.6
49.6 40.1
31.8 49.5 46.4
Total loss of sulphur (rag)
* All ~ 345 values are given here with respect to the starting sulphur. ** F is the fraction of isotopes exchanged (see text). t a is the fractionation factor (see text). t ? All 6 1 8 0 values are given here with respect to the starting water.
0.5 2.5 5
Time (days)
320 320 320
Temperature (°C)
Table 1 Chemical and isotopic data from the sulphur hydrolysis reactions
+0.2 +0.4 +0.3 0.0 +0.2
+0.4 +0.2
0.0 +0.3 +0.4 +0.1
+0.2 +0.1
-0.3 0.0 0.0
+1.8 +1.5 +0.4 +1.2 +1.6
+1.7 +2.1
+1.2 +2.1 +1.9 +2.2
+2.5 +1.6
+2.3 +2.0 +1.9
+21.1 +26.4 +29.0 +30.4 +30.9
+28.3 +28.8
+23.8 +26.6 +26.5 +27.1
+24.6 +24.2
+22.7 +22.9 +22.8
19.3 24.9 28.6 29.2 29.3529. 31
26.6 26.7526"7 ~
22.6 24.5 24.6-~ 4 l8 24.9 ~2
22.1 22.6522. 41
20.4 20.9 20.9 y20"9 a
(%o)
(%o)
(%o)
HSO4
A 34S H2S-HSO 4
6345
6345
Residual H2S sulphur (%0)
6 345*
0
0.342 0.150 0.024
0
0
0.088 0.011
0
0
0.024
(l-F)**
28.9
26.3
24.5
22.2
20.7
1000 In c~t
6 180
8.9 9.7 9.5 10.7 10.7
9.1 9.7
8.4 9.2
(%~)
HS04 tt
2.
z
4~
B. W. Robinson, Sulphur isotope equilibrium
447
1.0 0.5~N~4~
II/ \ It~
I1~
~
II/
",t,
l
200
T.c
1.7 \
T ('C) 260
320
-4.0
200
.s
26c
,
32°
\
tn
200
4-73x166
-5'0 ~.)
"01
-0031 I o
I 2
t 4
I I 6 8 Time ( d a y s )
I 10
-6"0
12
Fig. 1. First order rate plot of log ( l - F ) versus time for the sulphur isotope exchange reaction at 200, 260 and 320°C. The data points are shown with error bars and the t~- values are 1.7 -* 0.1, 0.3 +- 0.05 and 0.1 -+ 0.03 days for 200, 260 and 320°C respectively.
cess of ten times t~ to obtain equilibrium values. Fig. 2, an Arrhenius plot, gives the reaction rates and demonstrates that below 200°C reaction (8) is very slow. At all temperatures the runs were made in duplicate with run time in excess of ten times t~ to check the overall reproducibility. Normally the 1000 in ct values have an error of-+0.2 to 0.4%o and the difference between duplicates did not exceed 0.5%o. Fig. 3 gives all the 1000 lnc~ points and their correlation with temperature, together with error bars on each data point.
I 1.7
1"6
I I I 1"8 1-9 2"0 T -~ (" K )-t x 103
I 2.1
2.2
Fig. 2. An Arrhenius plot for the sulphur isotope exchange reaction with rate constants determined at 200, 260 and 320°C. The data points are shown with error bars and the experimental activation energy (A) is 13.3 kcal mole-1 .
500 I
4 0 0 350 300 I
I
I
I
T ('C) 250
200
150
I
I
I
40
30 _E o
~ 20
5. Discussion 10
The sulphur hydrolysis reaction (1) is relatively fast, compared to the sulphur isotope exchange reaction (8), and the products were found to reach a constant value after 5 hr at 250°C and 1½ hr at 300°C [5]. However, the reverse reaction is very slow at 200°C [4] and for this study equilibrium was only approached from one direction. The yield controlling parameters for the reaction are temperature and pH,
2
3
4 T-2 ( ' K ) -2 x 10 6
5
6
Fig. 3. A 1000 In a versus lIT 2 plot showing the theoretical data for SO4-H2S sulphur isotope exchange, Sakai [11] (as squares) and the experimental data points from this study for HSO~-H2S sulphur isotope exchange (as circles). The error on the latter points is the vertical diameter of the circles.
448
B. W. Robinson, Sulphur isotope equilibrium
and the concentration of the products increases with increasing temperature [4]. The precise mechanism for the sulphur hydrolysis reaction is not known. However, the initial effect of the reaction on the sulphur isotope exchange reaction is a kinetic one, which favours the light isotope in the sulphate and the heavy to enter the hydrogen sulphide. This effect is superimposed by the much larger exchange equilibrium effect, which proceeds at a much slower rate. The sulphur species formed by the hydrolysis reaction would be expected to have a mean/5 34S close to that of the elemental sulphur, i.e., at 200°C and 42 days the H2S and sulphate should have values about -7%o and +21%o respectively. However, the species show a strong depletion in 32S from this common mean, which is either due to the hydrolysis mechanism or the factors mentioned below. A variable excess of sulphur was used in the reactions and the concentrations of the products formed and the isotope exchange are independent of this. An indication that the overall reaction studied was reaction (1) is given by the H2S/H2SO4 ratio being close to 3:1 (see table 1). The slight loss of H2S may be caused during sampling or by solution in the liquid sulphur [4]. The solubility of H2S in liquid sulphur has been shown to be twice its aqueous solubility for temperatures between 200 and 320°C [23]. However, the relatively slow quench times of the reactions and the low weight ratios of sulphur to water would lead to little HzS being trapped in the sulphur. In addition to the reaction products, about 5 to 15 mg of sulphur is lost during the reaction. This phenomenon was noted in the silica tube experiments of Oana and Mizutani [5] and may be caused by diffusion into the glass, although there is no correlation with time or temperature. Since the total HzS is sampled after the runs, any isotopic fractionation between H2S(liqui d S), H2S(aq), and H2S(gas) during the reaction will have affected the results. Upon cooling to room temperature, most of the H2S will partition into the aqueous phase [24]. The isotopic fractionation factors for HzS in the phases mentioned have not been determined but, in analogy to the other systems, such as SO2(gas)-SOE(aq) [25], these can be expected to be small at high temperatures and their effect has been neglected in this study.
After a run, the sulphur cools as a large ovoid immiscible bleb, having a brown outer shell and a yellow interior, together with numerous small (0.5) mm yellow spherutes. These different phases are thought to represent the liquid and gaseous portions of sulphur present at high temperatures. Thus for the runs at 260°C and 8 days the phases were separated and analysed; the isotopic values did not differ outside of the errors of measurement. Fig. 1 demonstrates that equilibrium was achieved in the sulphur isotope exchange reactions, and the reaction rate is similar to that for the H2SO4-SO2 sulphur isotope equilibrium at 200°C [26]. From the Arrhenius plot (fig. 2) an experimental activation energy of 13.3 kcal/mole is calculated. No attempt has been made to investigate the influences of the sulphur hydrolysis pre-equilibrium and associated variations in reaction products with time and temperature on the isotopic fractionation rates. However, this is included here to facilitate future comparisons with other sulphur isotope exchange reactions. Fig. 3 (a 1000 lnc~ versus 1/T 2 plot) gives the sulphur isotope fractionation data from this study (HSO; H2S ) and the theoretical curve (SO~-H2S) from Sakai [ 11 ]. The experimental points probably form part of a curve of similar form to the theoretical curve. In fact, an inflection point at 260°C appears to be present. However, extrapolation is difficult, and the experimental points obtained in this study are defined by the least squares best fit straight line: 1000 lna(HSO]-H2S) = 5.07 (10 6 T -2) + 6.33 for temperatures between 200°C and 320°C. A direct comparison with Sakai's curve should not be made. More recent data on the partition function ratios of H2 34S and H2 32S [12] differ from Sakai's by 2%0 at 200°C, and better estimates for SO4 partition functions should soon be published..4dso, undoubtedly HSOT~--H2S and SO4-H2S will behave differently. How differently will not be known until calculations are done on HSOT, partition functions or the systems are compared experimentally. It has also been pointed out by Thode et al. [ 12] that, during experimental determinations, other sulphur species found could be quantitatively recovered as silver sulphide or barium sulphate. For the chemistry and conditions of the present study, this is extremely improbable.
B. IV. Robinson, Sulphur isotope equilibrium
6. Conclusions Isotope exchange reactions have been used to calculate the temperatures at depth in natural hydrothermal systems [22,27,28]. For the Wairakei geothermal area, New Zealand, the average A 34S(HsO4_HzS ) is 19%o [28] which, using the data obtained in this study, indicates a temperature o f about 350°C. The maximum measured temperature in this area is 270°C, and the isotope temperature probably represents that at greater depth. The exchange reaction studied can be applied to ore deposit studies: at the Echo Bay deposit, N.W.T., Canada, early formed acanthite has a 5 a4S value o f -20%o whilst the ore fluid is thought to have been in the HSO?, stability field at a temperature o f 150°C [29]. Using the H2S-Ag2S fractionation factor from [30] and the H S O 4 - H 2 S fractionation factor from an extrapolation of this study the ~ 34Szs value for the ore fluid is +20%o, thus suggesting trapped sea-water sulphate as the origin of the sulphur. However, when adapting the results of this study to sulphide-sulphate fractionation in hydrothermal ore solutions, several differences should be noted. For the average conditions o f an ore forming fluid, say 250°C, pH 3 . 5 - 7 . 5 and fs2 ~ 10-1°, sulphur isotope exchange in solution will probably not be as rapid as that during the reactions with pH 1 . 5 - 3 and fs2 ~ 10-3 - 10-6. In analogy to the oxygen isotope exchange rate between dissolved sulphate and water [21 ], the exchange rate was shown to be extremely sensitive to pH, and was very rapid at low pH. Also the sulphur isotope exchange rate between metallic sulphides is dependent on fS2, the rate increasing with increasing fs2 [31 ]. It thus appears that the reaction rates determined here could be slower in many natural systems.
Acknowledgements The high temperature part of this study was conducted in the laboratories of Chemistry Division, D.S.I.R., where the co-operation of my colleagues there is gratefully acknowledged. Discussions with Dr. T.A. Rafter, Dr. J.R. Hulston and Dr. W. Giggenbach have been o f value during the course o f this study.
449
Dr. C.B. Taylor kindly supplied the 8 180 analysis of the water starting material. Mr. D.W.N. Keith and Miss D.R. Pettis are thanked for their help with laboratory and mass spectrometry procedures.
References [ 1] T.A. Rafter, S.H. Wilson and B.W. Shilton, Sulphur isotopic variations in nature. Part 5, Sulphur isotopic variations in New Zealand geothermal bore waters, New Zealand J. Sci. 1 (1958) 103. [2] H. Sakai, Fractionation of sulphur isotopes in nature, Geochim. Cosmochim. Acta 12 (1957) 150. [3] S. Oana and H. lshikawa, Sulfur isotopic fractionation between sulfur and sulfuric acid in the hydrothermal solution of sulfur dioxide, Geochem. J. 1 (1966) 45. [4] A.J. Ellis and W. Giggenbach, Hydrogen sulphide ionization and sulphur hydrolysis in high temperature solution, Geochim. Cosmochim. Acta 35 (1971) 247. [5] S. Oana and Y. Mizutani, Sulphur isotopic fractionation between hydrogen sulphide and sulphuric acid formed by the 4S + 4H20 = 3H2S + H2SO4 reaction, Symp. Geochem. Soc. Japan, Nagoya (1966). [6] H. Ohmoto, Systematics of sulfur and carbon isotopes in hydrothermal ore deposits, Econ. Geol. 67 (1972) 551. [7] H.C. Urey, The thermodynamic properties of isotopic substances, J. Chem. Soc. (1947) 562. [8] J. Bigeleisen and M.G. Mayer, Calculation of equilibrium constants for isotopic exchange reactions, J. Chem. Phys. 15 (1947) 261. [9] J. Bigeleisen, Statistical mechanics of isotope effects on the thermodynamic properties of condensed systems, J. Chem. Phys. 34 (1961) 1485. [ 10] A.P. Tudge and H.G. Thode, Thermodynamic properties of isotopic compounds of sulphur, Can. J. Res. 28B (1950) 567. [11] H. Sakai, Isotopic properties of sulfur compounds in hydrothermal processes, Geochem. J. 2 (1968) 29. [12] H.G. Thode, C.B. Cragg, J.R. Hulston and C.E. Rees, Sulphur isotope exchange between sulphur dioxide and hydrogen sulphide, Geochim. Cosmochim. Acta 35 (1971) 35. [13] A.S. Quist, W.L. Marshall and H.R. Jolley, Electrical conductance of aqueous solutions at high temperature and pressure. Part II, The conductances and ionization constants of sulfuric acid-water solutions from 0° to 800 ° and pressures up to 4000 bars, J. Phys. Chem. 69 (1965) 2726. [14] A.C. Wahl and N.A. Bonner (eds.), Radioactivity applied to chemistry (Wiley and Sons, New York, 1951) p. 604.
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B.W. Robinson, Sulphur isotope equilibrium
[15] T.A. Rafter, Oxygen isotopic composition of sulphates. Part I, A method for the extraction of oxygen and its quantitative conversion to carbon dioxide for isotope radiation measurements, New Zealand J. Sci. 10 (1967) 402. [16] H. Sakai and H.R. Krouse, Elimination of memory effects in 180/160 determinations in sulphates, Earth Planet. Sci. Letters 11 (1971) 369. [ 17] T.A. Rafter, Sulphur isotopic variations in nature. Part 2, A quantitative study of the reduction of barium sulphate by graphite for recovery of sulphide-sulphur for sulphur isotopic measurements, New Zealand J. Sci. and Tech. 38 (1957) 955. [ 18] J.R. Hulston and B.W. Shilton, Sulphur isotopic variations in nature. Part 4, Measurements of sulphur isotopic ratio by mass spectrometry, New Zealand J. Sci. 1 (1958) 91. [19] T.A. Rafter, Sulphur isotopic variations in nature. Part 1, The preparation of sulphur dioxide for mass spectrometer examination, New Zealand J. Sci. Tech. 38 (1957) 849. [20] J.R. Hulston and J. Beauchamp, Sulphur isotopic variations in nature. Part 9, A recomparison of the New Zealand sulphur standard with the meteoritic sulphur standard, New Zealand J. Sci. 6 (1963) 157. [21 ] R.M. Lloyd, Oxygen isotope behaviour in the sulfatewater system, J. Geophys. Res. 73 (1968) 6099. [22] Y. Mizutani and T.A. Rafter, Oxygen isotopic composition of sulphates. Part 3, Oxygen isotopic fractionation in the bisulphate ion-water system, New Zealand J. Sci. 12 (1969) 54.
[23] R. Fanelli, Solubility of hydrogen sulfide in sulphur, Ind. Eng. Chem. 41 (1949) 2031. [24] T.N. Kozintseva, Solubility of hydrogen sulfide in water and solutions under high temperature, in: Geochemical investigations in the field of high pressures and temperatures, N.I. Khitarov, ed. (Nauka. Pub. Office, Moscow, 1965) p. 121. [25] T.E. Eriksen, Sulphur isotope effects. II, The isotopic exchange coefficients for the sulfur isotopes 345-325 in the system SO2g-aqueous solutions of SO2, Acta Chem. Scand. 26 (1972) 581. [26] H.B. Dunford, A.G. Harrison and H.G. Thode, Equilibrium constants for the sulfur isotope exchange between SO2 and H2SO4, Can. J. Chem. 35 (1957) 817. [27] N. Nakai, Isotopic and chemical equilibrium of CH4-CO2 and pyrite-anhydrite in geothermal area in Japan, Abs. Int. Syrup. Hydrogeochem. and Biogeochem., Tokyo (1970). [28] A. Steiner and T.A. Rafter, Sulfur isotopes in pyrite, pyrrhotite, alunite and anhydrite from steam wells in the Taupo volcanic zone, New Zealand, Econ. Geol. 61 (1966) 1115. [29] B.W. Robinson, Studies of the Echo Bay silver deposit, N.W.T., Canada, Unpub. Ph.D. Thesis, University of Alberta (1971). [30] B.W. Robinson and H. Ohmoto, Mineralogy, fluid inclusions and stable isotopes of the Echo Bay U - A g - C u deposits, Northwest Territories, Canada, in preparation. [31] Y. Kajiwara and H.R. Krouse, Sulfur isotope partitioning in metallic sulfide systems, Can. J. Earth Sci. 8 (1971) 1397.