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Hydrogen Sulfide in the Environment


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This is a page about hydrogen sulfide spatial and temporal variation in the environment.

Hydrogen sulfide in sediments and water

Hydrogen sulfide develops in sediments (and thus in oil) under oxygen-free conditions where sulfur becomes the terminal electron acceptor. Natural sources include oceans, coastal wetlands, soils and plants, volcanoes, and biomass burning (Andreae, 1990), most of which have episodic rather than continual emissions. The “rotten egg” smell in wetlands and boggy areas are usually releases of hydrogen sulfide from sediments where the oxygen has all been consumed by reacting with organic matter, and sulfates become reduced sulfur and sulfides. In situations where oxygen is present, hydrogen sulfide readily oxidizes. For example, in seawater, hydrogen sulfide oxidizes to sulfate ions in hours to days (e.g. Andreae, 1990). Interestingly, in water with pH below 6.5, the hydrogen sulfide oxidation rate decreases (Morse et al, 1987), which could have implications for sulfide ore mines with acid mine drainage (where the acid would by sulfuric acid from the oxidation of sulfides, but hydrogen sulfide gas may still be present too).

Hydrogen sulfide in the atmosphere

In the atmosphere, there are generally very low ambient concentrations of hydrogen sulfide because it oxidizes into sulfur dioxide, or even sulfate aerosols. Ambient hydrogen sulfide concentrations have been reported ranging from 0.00071 to 0.066 ppm (AEGL vol 9, p 176). However, episodic emissions can create hydrogen sulfide concentrations that are orders of magnitude higher than these mean ambient concentrations. Anthropogenic sources of hydrogen sulfide include oil and gas drilling and production and refining, coke production, wood pulp production, liquid manure application, wastewater treatment, and most any processes that involve sulfur and organic matter, especially at high temperatures (WHO, 2000). Hydrogen sulfide has measured downwind of a pulp and paper mill in California, with peak concentration of 0.13 ppm and has been calculated near sulfate pulp mills in Finland to exceed 1000 ppm for short periods of time (WHO, 2000).

Hydrogen sulfide has a molecular weight of 34.08 grams per mole, so it is heavier than air. Thus, hydrogen sulfide gas will accumulate in low-lying areas. The spread and transport of hydrogen sulfide is influenced by wind and its photodissociation and oxidation rates. These rates at which hydrogen sulfide reacts to form other sulfur compounds are largely dependent upon oxidants in the atmosphere, namely hydroxyl radicals, ozone, and nitrogen dioxide (e.g. Wine et al, 1981; Glavas and Toby, 1975). The Agency for Toxic Substances and Disease Registry (ASTDR) reports that airborne hydrogen sulfide gas usually lasts between 1 and 42 days (ASTDR, 2016).

Emissions of hydrogen sulfide are usually episodic, such as a wetland disruption resulting in sediment gas release, unless the hydrogen sulfide emanates from a continual leak or is a processing bi-product. With emissions variability and active atmospheric chemistry, airborne hydrogen sulfide concentrations are likely to vary widely spatially and temporally.

Resources and references:

  1. Andreae, M. (1990). Ocean-atmosphere interactions in the global biogeochemical sulfur cycle, Marine Chemistry, 30:1-29. https://www.researchgate.net/profile/Meinrat_Andreae/publication/223295397_Ocean-Atmosphere_Interactions_in_the_Global_Biogeochemical_Sulfur_Cycle/links/02e7e52430e58ebb53000000/Ocean-Atmosphere-Interactions-in-the-Global-Biogeochemical-Sulfur-Cycle.pdf

  2. Agency for Toxic Substance and Disease Registry. 2016. Hydrogen Sulfide Fact Sheet. https://www.atsdr.cdc.gov/ToxProfiles/tp114-c1-b.pdf (note: it is unclear what original data was used to arrive at the conclusion that H2S usually lasts in environment 1-42 days).

  3. Glavas, S. and Toby, S. (1975). Reaction between ozone and hydrogen sulfide. Journal of Physical Chemistry, 19(8). http://pubs.acs.org/doi/abs/10.1021/j100575a004

  4. Morse, J.W., Millero, F.J., Cornwell, J.C. and Rickard, D. (1987). The chemistry of the hydrogen sulfide and iron sulfide systems in natural waters. Earth Science Reviews, 24: 1-42. (useful H2S chemistry): https://www.researchgate.net/profile/David_Rickard2/publication/222157377_The_Chemistry_of_the_Hydrogen_Sulfide_and_Iron_Sulfide_Systems_in_Natural_Waters/links/55705c9808ae7d0f5f900cb5/The-Chemistry-of-the-Hydrogen-Sulfide-and-Iron-Sulfide-Systems-in-Natural-Waters.pdf

  5. World Health Organization (2000). Air Quality Guidelines, Hydrogen Sulfide. http://www.euro.who.int/__data/assets/pdf_file/0019/123076/AQG2ndEd_6_6Hydrogensulfide.PDF

  6. Wine, P.H., Kreutter, N.M., Gump, C.A., Ravishankara, A.R. (1981). Kinetics of hydroxyl radical reactions with the atmospheric sulfur compounds hydrogen sulfide, methanethiol, ethanethiol, and dimethyl disulfide. Journal of Physical Chemistry, 85(18). http://pubs.acs.org/doi/abs/10.1021/j150618a019


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