Oxidized sulfur species occurring

Oxidized sulfur species occurring in natural waters (sulfate, sulfite, thiosulfate] do not interact with the platinum electrode when in the presence of H2S and the pH-Eh-E52- relations found were similar to the above relations. Thus, the unambiguous relations found between pH, Eh and E22- in aqueous solutions of hydrogen sulfide can be employed to characterize solutions and water samples where hydrogen sulfide is the only reduced sulfur species present. 

This mechanism as a main cause for epithermal-type Au deposition is supported by sulfur isotopic data on sulfides. Shikazono and Shimazaki (1985) determined sulfur isotopic compositions of sulfide minerals from the Zn-Pb and Au-Ag veins of the Yatani deposits which occur in the Green tuff region. The values for Zn-Pb veins and Au-Ag veins are ca. +0.5%o to -f4.5%o and ca. -l-3%o to - -6%c, respectively (Fig. 1.126). This difference in of Zn-Pb veins and Au-Ag veins is difficult to explain by the equilibrium isotopic fractionation between aqueous reduced sulfur species and oxidized sulfur species at the site of ore deposition. The non-equilibrium rapid mixing of H2S-rich fluid (deep fluid) with SO -rich acid fluid (shallow fluid) is the most likely process for the cause of this difference (Fig. 1.127). This fluids mixing can also explain the higher oxidation state of Au-Ag ore fluid and lower oxidation state of Zn-Pb ore fluid. Deposition of gold occurs by this mechanism but not by oxidation of H2S-rich fluid. 

Table 2.1 lists atmospheric sulfur compounds. The principal sulfur compounds in the atmosphere are H2S, CH3SCH3, CS2, OCS, and SO2. Sulfur occurs in five oxidation states in the atmosphere. (See Box) Chemical reactivity of atmospheric sulfur compounds is inversely related to their sulfur oxidation state. Reduced sulfur compounds, those with oxidation state -2 or —1, are rapidly oxidized by the hydroxyl radical and, to a lesser extent, by other species, with resulting atmospheric lifetimes of a few days. The water solubility of sulfur species increases with oxidation state reduced sulfur species occur preferentially in the gas phase, whereas the (+6) compounds often tend to be found in particles or droplets. Once converted to compounds in the S(+6) state, sulfur species residence times are determined by removal by wet and dry deposition. 

Sulfur dioxide occurs in industrial and urban atmospheres at 1 ppb—1 ppm and in remote areas of the earth at 50—120 ppt (27). Plants and animals have a natural tolerance to low levels of sulfur dioxide. Natural sources include volcanoes and volcanic vents, decaying organic matter, and solar action on seawater (28,290,291). Sulfur dioxide is beHeved to be the main sulfur species produced by oxidation of dimethyl sulfide that is emitted from the ocean. 

A second complication is that we would like to decouple zero-valent sulfur from the element s other redox states, since Reaction 17.28 produces native sulfur, but the database does not include such a coupling reaction. Situations of this nature are not uncommon, occurring when an element in a certain oxidation state is stable as a solid, but no corresponding aqueous species occurs under geochemical conditions. To work the problem, we invent a ficticious zero-valent species S(aq) with an arbitrarily low stability. Setting log K for the reaction... 

Pyrite and arsenopyrite have similar oxidation and self-induced collectorless flotation behavior. It is generally suggested that anodic oxidation of pyrite occurs according to reactions (2-24) in acidic solutions (Lowson, 1982 Heyes and Trahar, 1984 Trahar, 1984 Stm et al., 1991 Chander et al., 1993). The oxidation of pyrite in basic solutions takes place according to reactions (2-25). Since pyrite is flotable only in strong acidic solutions, it seems reasonable to assume that reaction (2-24) is the dominant oxidation at acidic solutions. Whereas pyrite oxidizes to oxy-sulfur species with minor sulphur in basic solutions. 

Cycloaddition reactions of thioaldehydes and sulfines are most probably encountered in plants, as elegantly and soundly shown by the group of Eric Block during their investigation of sulfur products occurring in the Allium species (for a review see [91]). They were able [92, 93] to isolate bicyclic dithioacetal oxides, called zwiebelanes, and also to synthesise them from a thioxosulfine, already described in this review (Sect. 2.6, Scheme 18). An extremely rich stereochemical and analytical study has resulted. 

In addition, the report by Smock et al (1998) that when thiosulfate is used as the oxidant (instead of sulfate), the fractionation of sulfur is much lower (on the order of 15 per mil), reminds us of two critical items. First, there are many organisms now in culture that contribute to the sulfur cycle, and many of them (and their various chemistries) are not well characterized with regard to sulfur fractionation. Perhaps it would be wise to revisit the fractionation that occurs with the oxidation of various sulfur species, both photosynthetic and nonphotosynthetic. Second, we are reminded that many of the organisms that contribute to the environment today have not been grown in culture. 

Oxidation of reduced sulfur species. Oxidation of reduced sulfur species in the presence of oxygen can occur spontaneously, without bacterial mediation. Bacteria of the family Thiobacteriaceae are probably the most important bacteria involved in sulfur oxidation. Of these, bacteria of the genus Thiobacillus have been most studied (Goldhaber and Kaplan 1974 Cullimore 1991). The first product of sulfide oxidation abiotically or by Thiobaccillus is thought to be elemental sulfur according to... 

The true physiological role of these reductases in phytoplankton is not known and it is unclear whether electron transport out of the cell occurs in nature. Although the oxidation and reduction of the extracellular solutes may just be an adventitious reaction of these enzymes with no significance to the microorganism, under certain conditions, such as the presence of favorable redox couples, such electron export may occur. Because of the high half-saturation constants measured for metal reduction by Jones et al. (1987), it seems unlikely that these or similar trace metal complexes are the major electron acceptors in nature. We cannot rule out, however, the possibility of reduction of other (major) solutes such as sulfate or intermediate redox sulfur species. 

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