Part II The Sulfur Cycle Downloaded from http://iwaponline.com/ebooks/book/chapter-pdf/772252/9781789060966_0011.pdf by guest on 27 September 2021 Downloaded from http://iwaponline.com/ebooks/book/chapter-pdf/772252/9781789060966_0011.pdf by guest on 27 September 2021 Chapter 2 The chemical sulfur cycle Ralf Steudel 2.1 INTRODUCTION Sulfur is one of the most important elements for life as well as for the chemical and pharmaceutical industries. Even in extraterrestrial space, sulfur compounds are abundant albeit in low concentrations. Sulfur contributes to only 0.07 wt% of the crust of the Earth but elemental sulfur and numerous sulfur-containing minerals occur in substantial deposits. Important sulfidic minerals are, for example, pyrite FeS2, galena PbS, zinc-blende (sphalerite) ZnS, cinnabar HgS, chalcopyrite CuFeS2, and chalcocite Cu2S. Weathering and oxidation of the sulfides has resulted in large deposits of water-insoluble or poorly soluble sulfate minerals such as gypsum Ca[SO4]·2H2O, bassanite Ca[SO4] · 0.5H2O, anhydrite Ca[SO4] and baryte Ba[SO4]. Gypsum is, by volume, the most abundant sulfate mineral on Earth. Ocean water contains 2.7 g L−1 sulfate, river waters only ca. 0.01 g L−1. Sulfur compounds are constituents of all organisms and consequently of all biomass and materials which originated from these sources such as wood, peat, coal and crude oil as well as their derivatives. Combustion of such materials releases not only sulfur dioxide (SO2) but also traces of carbonyl sulfide (COS). The latter is assimilated by plants as part of their sulfur metabolism. Dimethyl sulfide (DMS) is released to the atmosphere in enormous quantities by phytoplankton in the oceans, and hydrogen sulfide (H2S) as well as SO2 and carbon dioxide (CO2) are emitted by volcanoes. Sulfate reduction by the © 2020 The Authors. This is an Open Access book chapter distributed under the terms of the Creative Commons Attribution Licence (CC BY-NC-ND 4.0), which permits copying and redistribution for noncommercial purposes with no derivatives, provided the original work is properly cited (https:// creativecommons.org/licenses/by-nc-nd/4.0/). This does not affect the rights licensed or assigned from any third party in this book. The chapter is from the book Environmental Technologies to Treat Sulfur Pollution: Principles and Engineering, 2nd Edition, Piet N.L. Lens (Ed.). DOI: 10.2166/9781789060966_0011 Downloaded from http://iwaponline.com/ebooks/book/chapter-pdf/772252/9781789060966_0011.pdf by guest on 27 September 2021 12 Environmental Technologies to Treat Sulfur Pollution ubiquitous sulfur bacteria in anoxic environments such as ponds, lakes, swamps and coastal waters also produces H2S. The human body contains 2 g S per kg, in other words 140 g S for a person of 70 kg. Large natural underground deposits of elemental sulfur exist in the USA, Mexico and Poland, but today elemental sulfur is mainly produced by the desulfurization of crude oil, of sour natural gas and of coal. Only on a very small scale is sulfur still mined in volcanic areas such as Indonesia. Historically, Southern Italy (Sicily) was the main origin of elemental sulfur during the industrialization of Western Europe in the early 19th century. In 1900, Sicily produced 500,000 t of elemental sulfur. Another important source of sulfur for production of sulfuric acid is pyrite with large deposits in many countries. The average atomic weight of sulfur is 32.066 representing the natural mixture of the isotopes 32S (95.0 mol%), 33S (0.76%), 34S (4.22%) and 36S (0.02%). The relative atomic weights of most elements, however, vary slightly owing to natural variations in the abundances of their isotopes. This variation is used to determine the origin of a particular sample (a mineral or biological material). In the case of sulfur, the variation may be +0.01 units, while individual samples can be determined at an accuracy of +0.00015 (mainly by mass spectrometry). Due to the historic developments in analytical methods, the published atomic weights of the chemical elements have changed over the years. The artificial radioactive nuclide 35S is used for labeling experiments; it decomposes with a half-life of 87.2 d by β-emission to 35Cl. More than 200 years of scientific research on sulfur and its compounds has resulted in a vast body of literature which cannot easily be searched for reliable information. Moreover, this literature contains errors and contradictions since earlier workers, not having the methods available that are standard today, often made claims that have not always been subsequently confirmed. However, reliable reviews written by experts in the field are available, above all the many volumes of Gmelin Handbook of Inorganic Chemistry in which the chemical literature is critically and exhaustively evaluated. On sulfur and its compounds 22 volumes have appeared, dating from 1939 to 1996 and covering the literature up to 1991. Unfortunately, no further volumes have been produced. Other reliable reviews on inorganic and analytical sulfur chemistry have been published (in alphabetical order) by Devillanova (2006), Holleman-Wiberg (2017), Karchmer (1970), Müller and Krebs (1984), Nickless (1968), Schmidt and Siebert (1973), Steudel and Eckert (2003), Steudel (2003a, b, c, d and Steudel, 2020), Steudel and Chivers (2019) as well as Szekeres (1974). 2.1.1 Oxidation states and redox potentials The complexity of sulfur chemistry originates from the many oxidation states and coordination numbers sulfur atoms can assume, as well as from the tendency of Downloaded from http://iwaponline.com/ebooks/book/chapter-pdf/772252/9781789060966_0011.pdf by guest on 27 September 2021 The chemical sulfur cycle 13 sulfur in the oxidation state zero to catenate, forming chains and rings of an astonishing variety. Sulfur atoms and ions can adopt any coordination number between 1 (e.g. CS2) and 8 (e.g. in solid Na2S with antifluorite structure); sulfur oxidation states range from −2to+6(Table 2.1). In Table 2.1, the nine oxidation states of sulfur are illustrated by typical examples. Most of them play a role in aqueous systems in which redox reactions occur either as a result of microbiological activity or simply following the thermodynamics of the system in non-enzymatic reactions. However, chemical systems are not always composed according to the requirements of thermodynamics. High activation enthalpies may keep exergonic reactions from proceeding at ambient temperatures, resulting in a chemical composition far from equilibrium (Licht & Davis 1997). The equilibrium composition of an aqueous system containing just sulfur and oxygen is shown in the Pourbaix diagram in Figure 2.1. Depending on the redox potential, the pH value, the temperature and the overall concentration of sulfur, the relative stability areas of sulfide HS−, elemental sulfur (S), as well as sulfate 2− − [SO4] and hydrogen sulfate [HSO4] are shown (Garrels & Naeser, 1958; Williamson & Rimstidt, 1992). The different areas of this diagram indicate which species will predominate at a given potential and pH value. As the overall sulfur concentration decreases, the smaller the existence area of elemental sulfur becomes. Sulfite, thiosulfate and other sulfur oxoanions (e.g. polythionates) never predominate, regardless of pH and potential. In other words, these species exist in water only under non-equilibrium conditions or as minority species. The thiosulfate, polythionate and disulfite ions are typical examples of anions with mixed oxidation states of sulfur (see below). Table 2.1 The oxidation states of sulfur atoms in common compounds. Oxidation Examples State – 2− –2 dihydrogen sulfide H2S, hydrogen sulfide ion HS , sulfide ion S as in FeS 2− –1 disulfane H2S2, disulfide ion [S2] as in pyrite FeS2 0 elemental sulfur Sn, organic polysulfanes RZSnZR +1 dichlorodisulfane ClZSZSZCl 2− +2 sulfur dichloride SCl2, sulfoxylate ion [SO2] 2− +3 dithionite ion [S2O4] 2− +4 sulfur dioxide SO2, sulfite ion [SO3] 2− − +5 dithionate ion [S2O6] , organic sulfonates [RZSO3] 2− 2− +6 sulfur trioxide SO3, sulfate ion [SO4] , peroxosulfate ion [SO5] Downloaded from http://iwaponline.com/ebooks/book/chapter-pdf/772252/9781789060966_0011.pdf by guest on 27 September 2021 14 Environmental Technologies to Treat Sulfur Pollution Figure 2.1 Pourbaix diagram for the binary system sulfur/oxygen in water at 25°C and 1.013 bar with the sum of the activities of all sulfur-containing ions equal to 0.1 mM. 2.1.2 Catenation of sulfur atoms As in hydrogen sulfide H–S–H and dimethyl sulfide CH3–S–CH3, sulfur atoms can form two covalent bonds with other atoms or with itself to form chain-like units –S–S–S– of practically unlimited length (‘catenation’). These chains may be terminated by single atoms such as H or Cl, by groups like CH3 or SO3H, by ions such as S− or may ‘bite their own tail’ forming rings of various sizes. Corresponding examples are listed in Table 2.2. A special case are the 2− polysulfide anions [Sn] in which the chains are terminated by negatively charged sulfur atoms which are iso-electronic with Cl atoms. Therefore, 2− polysulfide anions [S–Sn–S] are iso-electronic with dichlorosulfanes Cl–Sn–Cl. Table 2.2 gives those values of n which have been determined in compounds isolated in pure form (column 2) or which have been detected in mixtures by high-performance liquid chromatography (HPLC), proton nuclear magnetic resonance (1H-NMR) spectroscopy or ion chromatography (column 3). From these data it is obvious that there is seemingly no limitation to the values of n. It is just that the preparation of the higher-molecular species becomes increasingly difficult since the solubility and thermal stability decrease with increasing values of n. Polymeric sulfur (Sμ) is insoluble in all solvents (excepting liquid sulfur) and therefore is considered to consist of very long chains and/or very large rings.