Geochemical Journal, Vol. l,'-.pp. 45 to 50, 1966.
Sulfur isotopic fractionation between sulfur and sulfuric acid in the hydrothermal solution of sulfur dioxide
SHINYA OANA and HITOMI ISHIKAWA Department of Earth Sciences, Nagoya University, Chikusa, Nagoya, Japan
(Received June 18, 1966)
Abstract-Sulfur and sulfur compounds involved in the chemical reaction, 3H2S03= 2H2SO4-{-S+H20 in the temperature range from 150 to 300°C have been quantitatively determined. After the disproportionation reaction of sulfurous acid is completed, sulfur decreases according to the chemical reaction, 4 S+4 H2O = 3 H2S-I-H2S04. These two reactions may be responsible for the production of sulfuric acid in hot spring water. 343 is enriched in sulfuric acid and depleted in sulfur revealing 16.1 to 21.1©00 fractionation between them. This would account for the enrichment of 34S in sulfuric acid encountered in hot spring water.
INTRODUCTION Sulfate is one of the most abundant constituents dissolved in hot spring water. It has long been believed that all of it is produced from hydrogen sulfide and sulfur dioxide by oxidation with atmospheric oxygen. Sulfate is ubiquitous in hot spring and volcanic areas associated with hydrogen sulfide and sulfur dioxide, which, after being exposed to atmospheric air or encountered by dissolved oxygen in groundwater, form sulfuric acid and/or sulfur. The production of sulfuric acid by this process, however, is controlled by the supply of hydrogen sulfide, sulfur dioxide and oxygen, and so is restricted to the near surface regions where atmos pheric oxygen is supplied directly, or dissolved in groundwater. The several grams of sulfate per liter often contained in hot spring water cannot be explained to have been supplied solely by atmospheric oxidation of hydrogen sulfide and sulfur dioxide. IWASAKI and OZAWA (1960) introduced the disproportionation reaction of sul furous acid, 3 H2SO3=2 H2SO4+S+H20, to explain the occurrence of sulfate in hot spring water. This seems to be pertinent from the standpoint of oxygen-free pro duction of sulfuric acid. DEINES (1934) used the same reaction to account for the accumulation of native sulfur in the vicinity of Vulkano. This may sound reason able for some specific instances. The production and accumulation of native sulfur half as much as the production of sulfuric acid, however, cannot be understood to be ubiquitous. And OANA postulates that the native sulfur thus produced might react with water to form hydrogen sulfide and sulfuric acid. 46 S. OANA and H. ISHIKAWA
SAKAI (1957) determined O34S for hydrogen sulfide, sulfur dioxide, native sulfur and sulfate in hot spring water. He observed a definite enrichment of 3'S in sulfur dioxide with respect to hydrogen sulfide according to the sulfur istopic exchange equilibrium. Sulfate samples which were found near fumaroles and believed to have been derived from hydrogen sulfide, however, showed no enrichment, or even slight depletion of 34S, if any. Sulfate in hot spring water is marked with a high S34S value and cannot be supposed to have been derived from hydrogen sulfide. But he postulated that sulfate in hot spring water was produced by the oxidation of sulfur dioxide. The disproportionation reaction of sulfurous acid appears to be unidirectional from the form of the reaction equation. Therefore, no sulfur isotopic fractionation is likely to occur, or if any, a depletion of 3'S into sulfuric acid should be expected. And this cannot explain high S34S values of some sulfuric acid samples. We thus came to run a series of laboratory experiments to study the disproportionation re action in terms of stoichiometry and isotopic fractionation.
EXPERIMENTAL
1. Disproportionation reaction of sulfurous acid. Tank sulfur dioxide is bubbled through 50 ml of distilled water in a glass test tube to purge dissolved oxygen out of it and to saturate it with sulfur dioxide.* Three test tubes are sealed and placed in an autoclave with an appropriate amount of water, which nearly equalizes the pressures inside and outside the test tubes. The autoclave is heated to a desired temperature, let stand for 3.5 hr and cooled to the room temperature . The sulfur occurs as a button in each test tube and is weighed directly. Residual un reacted sulfur dioxide is determined by iodometry. Sulfuric acid that has formed is precipitated with barium chloride as barium sulfate, which weighed , reduced to hydrogen sulfide in a reducing mixture of hydriodic acid, hypophosphorous acid and hydrochloric acid. This hydrogen sulfide is fixed as cadmium sulfide and then converted to silver sulfide (THODE et al., 1961). 2. Reaction of sulfur with water. Native sulfur is put in a test tube with dis tilled water. Dissolved oxygen is purged by nitrogen . Three test tubes are sealed and treated in an autoclave charged with water . Sulfur and water react to form hydrogen sulfide and sulfuric acid. Hydrogen is absent , as confirmed by Nakai with a gas chromatograph.
3. Preparation of sulfur dioxide for the isotope analysis . A setup to prepare sulfur dioxide is shown in Fig. 1. Silver sulfide or sulfur is burned at 1 atmos pheric pressure and 1,250 °C in a slow stream of carbon dioxide-free atmospheric air in a silica tubing. Sulfur dioxide is collected in the U tube , D, which is im mersed in a n-pentane bath cooled with liquid nitrogen . C3 is closed, C, opened, * The production of sulfuric acid by this process is negligibly s mall, cf. 0-2 in Table 1, Sulfur isotopic fractionation between sulfur and sulfuric acid 47
vacuum line air pirani gauge C1
C4 C3 C2
manometer
D C B A
Fig. 1. S02 preparation line A, cooled with liquid nitrogen. B, cooled with dry ice in acetone. C, cooled with liquid nitrogen. D, cooled with frozen n-pentane. and liquid nitrogen is removed from D. The temperature of D rises from the boiling point of nitrogen, -195.6 °C to the melting point of n-pentane, -1308 °C. At this temperature carbon dioxide with a pressure of 0.4 mm Hg is pumped off, leaving sulfur dioxide with a pressure of less than 10-3 mm Hg without any practical loss. The removal of carbon dioxide is checked up by a pirani gauge. An evacuation for 3 min after the pirani gauge reading dropped to zero eliminates carbon dioxide completely as it cannot be detected by the mass spectrometer over the background. Sulfur dioxide thus purified is stored in a sample flask and is subjected to the isotopic ratio determination. 4. Sulfur isotopic ratio determination. 34S/32S has been measured with a mass spectrometer installed at the Institute for Thermal Spring Research, Okayama University, assisted by Dr. H. Sakai.
RESULT AND DISCUSSION
S34S values are indicated with respect to the tank sulfur dioxide (assumed 0.0%0) used throughout this experiment and shown in Table 1, Fig. 2 and 3. The experi mental error has not been checked specifically for this experiment. However, we believe that it does not exceed ±0.3%0 according to our experience dealing with sample preparations. The error in the ratio determination itself is ±0.1%or less and the overall error is involved mostly in the sample preparation. The ratio of sulfuric acid to sulfur both formed from sulfur dioxide is 2 and this suggests that they are formed by the disproportionation reaction of sulfurous acid. 3 H2SO3=2 H2SO4+S+H20 (1)
This reaction goes on fast at elevated temperatures (200 'C and over). With our 48 S. DANA and H. ISHIKAW A
Table 1. Sulfur and sulfur compounds in the hydrothermal solution of sulfur dioxide
Sample Temp. Time H2SO3, residual H2SO4, found S, found Molar a34STIPRO,a34S ratio -a34Ss total mg S °o of mg S a34S*, mg a34S* No. °C hr H2SO4/S %o 50 ml total S /50 ml %o /50 ml %9 %e 0-2 room 1 614 6.1 temp. week 10-1 150 3.5 537 45.9 438 + 6.1 196 9.9 2.23 +16.0 +0,8 6-B 150 3.5 369 + 6.6 149 -13 .5 +20.1 -0.1 5-A 200 3.5 53.7 3.1 1,022 +10.4 647 8.1 1.58 +18.5 +4.2 11-B 200 3.5 34.1 4.1 535.2 + 7.5 269 9.8 1.99 +17.3 +1.8 11-C 200 3.5 52.9 4.9 684.4 + 7.5 343 -11 .1 1.97 +18.6 +1.3 1-A 250 3.5 22.5 1.8 798.6 + 8.3 402 8.6 1.99 +16.9 +2.7 2-A 300 3.5 73.9 4.8 973.0 + 9.6 495 -11 .3 1.97 +20.9 +2.6 2-B 300 3.5 108 6.7 996.9 +10.2 511 -10 .9 1.95 +21.1 +3.2 8-1 300 1.0 11.0 2.0 360.6 + 8.6 179 -10 .7 2.02 +19.3 +2.1
* 3346 with respect to the tank S02 which is assumed to be 0.0%8.
0 rrom sulfurous acid SH;S04-cSS 0 From sulfurous acid SS3', %D From sulfur and water %D From sulfur and water 18 16 22 14 cP 20 0 12 0 CD 0 0 10 0 8 0 8 0 0 18 CZ) 6 8 4 HIS04 150 200 250 300 2 Temp. °C
0 Fig. 2. 31S fractionation to sulfuric acid -2 and sulfur from sulfur dioxide and sulfur S -4 -6
-8 0 0 -10 0 0 -12 0 8)
0 , Fig. 3. 39S fractionation between sulfuric 150 200 250 300 acid and sulfur Temp. °C Sulfur isotopic fractionation between sulfur and sulfuric acid 49 present autoclave we cannot quench the reaction at any desired time and temper ature and the test tubes are let stand in the autoclave to cool down to the room temperature after usually 3.5 hr of reaction. From the data in Table 1, however, it is sure that the chemical reaction has approached an equilibrium. Sulfuric acid is enriched in 34S and sulfur is depleted in 34S (Fig. 2). 834S values for total sulfur are the weighted means calculated from those for sulfuric acid and sulfur, remaining sulfurous acid being small and hence neglected, It is not appreci ably deviated from the O31S value of the starting material, sulfur dioxide, 0.0%0. Fractionation of 3'S between sulfuric acid and sulfur is about 20%0 at the studied temperatures, which is the same extent of fractionation between sulfuric acid and sulfur encountered in some hot springs. A small increase in 834S up to 4.2 in one case from the starting sulfur dioxide to the calculated total is not explained yet. The disproportionation reaction of sulfurous acid is unidirectional, and sulfurous acid decreases exponentially. Sulfur may react with water to produce sulfuric acid and hydrogen sulfide, which then reacts with sulfurous acid to yield sulfur and water, establishing a cycle of chemical reactions (2) and (3),
8S+8H2O=6H2S+2H2SO4 (2) 6 H2S+3 H2SO3=9 S+9H20 (3)
3 H2S03=2 H2S04+S+H20 (4) eventually resulting in the disproportionation of sulfurous acid. This cycle may account for the fractionation of 34S. Its possibility has been confirmed by the ex periment 2. The result is given in Table 2 and plotted in Fig. 2 and 3. Sulfur
Table 2. Sulfuric acid and hydrogen sulfide produced from sulfur and water
Sample Temp. Time S, used recovered H2SO4, found H2S, found Molar ratio a34S n,so, No. °C hr mg 434'S*,%0 /5g mlS 631S*%a mg S/50 ml H2S/H2SO4-a34Ss, %®
10 150 3.5 530.8 trace trace 3-A 300 3.5 1,030 -6 .0 16.2 +18.3 24.9 1.54 +24.3 3-B 300 3.5 935.7 16.4 50.3 3.07 3-C 300 3.5 991.6 15.8 43.5 2.75
* PS , with respect to the tank SO2 which is assumed to be 0.0%. reacts with water to produce sulfuric acid and hydrogen sulfide. The ratio of hydrogen sulfide to sulfuric acid in Table 2 is not uniform. In cases of 3-C and especially 3-A hydrogen sulfide is likely to have escaped before the analysis. The ratio 3.07 for 3-B sounds most reasonable and is qualified to suggest the chemical reaction (2). As far as there is sulfurous acid, the hydrogen sulfide produced reacts with this and yields sulfur still to maintain the overall reaction (4). As soon as the 50 S. OANA and H. ISHIKAWA sulfurous acid disappears, however, the sulfur starts decreasing, the reaction (2) persisting. This would be the case most common in natural environment of hot springs. Fractionation shown in Fig. 3 seems to increase slightly with increasing temper ature. This is the reverse of what is expected from theoretical calculations (SAKAI, 1957). An explanation may be as follows. The reaction (2) is so sluggish that it takes time for the reaction (1) to establish an isotopic exchange equilibrium. The lower the temperature, the more the deviation from the equilibrium. The time required for a hydrothermal solution in our present autoclave to cool down to the room temperature must also be considered, and chemical and isotopic equilibria would probably be shifted. Further experiments with sulfur and water as well as those on the disproportionation reaction of sulfurous acid with a quenching device has been started and will be reported in the next paper.
ACKNOWLEDGEMENTS
We are indebted to Dr. HITOSHI SAKAI of the Institute for Thermal Spring Research, Okayama University, who assisted us with his mass spectrometer. Our thanks are due to Dr. YOSHIHIKO MIZUTANI who gave us suggestions in the sample preparation and chemical analysis, and to Dr. NoBUYuKI NAKAI who analyzed a sample with the gas chromatograph.
References
DEINES, O. VON (1934) Die Entstehung der vulkanischen Schwefelablagerungen. Naturwiss. 22, 129-134. IWASAKI, I. and OZAWA, T. (1960) Genesis of sulfate in acid hot spring. Bull. Chem. Soc. Japan 33, 1018-1019. SAKAI, H. (1957) Fractionation of sulphur isotopes in nature. Geochim. Cosmochim. Acta 12, 150-169. THODE, H. G., MONSTER, J. and DUNFORD, H. B. (1961) Sulphur isotope geochemistry. ibid. 25, 159-175.