History & Discovery

Hydrogen sulfide is a colorless, flammable, and toxic gas known for its distinctive odor reminiscent of rotten eggs. Its natural occurrence results from the anaerobic bacterial decomposition of organic wastes, volcanic gases, hot springs, animal digestion, and various industrial processes. Hydrogen sulfide is also found as a natural component in natural gas and petroleum, constituting a minor fraction in oil (hundreds of ppm) but potentially forming a significant portion of natural gas, where concentrations can reach up to 5%. When the content of hydrogen sulfide exceeds 5.7 mg of H2S per cubic meter of natural gas, the gas is classified as sour. The process employed to remove hydrogen sulfide from sour gas is commonly known as sweetening the gas. Due to its association with anaerobic respiration in sewers and swamps, hydrogen sulfide is colloquially referred to as sewer gas, swamp gas, or stink damp.

Knowledge of hydrogen sulfide dates back to the 15th century when it was identified in water and referred to as sulfur water or sulfur vapors. Alchemists termed H2S as aer hepaticus (hepatic air), while early chemists adopted the name sulfuretted hydrogen, a term still in use today. The first chemist to prepare and describe hydrogen sulfide was Carl Wilhelm Scheele (1742–1786), who viewed it as a combination of sulfur, phlogiston, and heat. Further advancements in understanding the composition of H2S were made by Claude Louis Berthollet (1748–1822) in 1789, where he recognized its acidic nature.

Preparation & Production & Application

Hydrogen sulfide is generated as a byproduct of anaerobic respiration, particularly in the context of fermentation. Anaerobic respiration, employed primarily by bacteria and other microorganisms, allows them to fulfill their energy requirements using sulfate, elemental sulfur, and sulfur compounds as electron acceptors instead of oxygen.

A simplified representation of anaerobic respiration is articulated as follows:

The hydrogen sulfide produced can undergo oxidation back to sulfate or elemental sulfur when exposed to the atmosphere or an oxidizing environment. Moreover, hydrogen sulfide can react with other substances in sediments or soils, leading to the formation of sulfur minerals, such as pyrite (FeS_2), where hydrogen sulfide reacts with iron monosulfide:

The primary method for removing hydrogen sulfide from sour natural gas involves employing amine solutions. These solutions, containing monoethanolamine and diethanolamine as the principal amines, absorb hydrogen sulfide. The elemental sulfur can subsequently be recovered using the Claus process, wherein the amine-extractedH2Sis thermally oxidized to sulfur dioxide:

Sulfur is then produced by the combination of sulfur dioxide and hydrogen sulfide:

The Claus process achieves a 60–70% recovery of elemental sulfur, and to recover the majority of the remaining sulfur, hydrogen sulfide and sulfur dioxide are catalytically combined using activated alumina (Al_2O_3) or titanium catalysts.

Hydrogen sulfide finds limited commercial applications, primarily being employed in the production of elemental sulfur, sulfuric acid, and heavy water for nuclear reactors. The heavy water, D_2O, is produced through the water-hydrogen sulfide exchange process (Girdler-Sulfide or G-S process), which involves the exchange of hydrogen and deuterium in water and hydrogen sulfide at different temperatures. The resulting equilibrium favors the formation of semiheavy water (HDO) at colder temperatures and heavy isotope variety of hydrogen sulfide (HDS) at higher temperatures. This exchange occurs in a series of towers, enabling the production of reactor-grade heavy water enriched to approximately 30% in deuterium.

Despite its commercial applications, hydrogen sulfide poses a significant health risk due to its highly toxic nature. Precautionary measures, especially against inhalation, are crucial, given its toxicity. Workers in various industries, including oil and gas and agriculture, may face exposure, making hydrogen sulfide the leading cause of death from occupational inhalation of toxic chemicals. Even at low concentrations detectable by smell (0.02 ppm), caution is necessary, as disruptions to the olfactory sense occur at concentrations nearing 100 ppm. Concentrations higher than 100 ppm pose severe health risks, with 250 ppm potentially causing unconsciousness and 500–700 ppm resulting in fatality after prolonged exposure. While hydrogen sulfide exposure is typically not a household concern, faulty plumbing or groundwater sources may lead to unintentional exposure. The gas is also produced by intestinal bacteria in the digestive tract, contributing to the odor associated with flatulence and bad breath.

Reference

Richard L. Myers (2009). The 100 Most Important Chemical Compounds: A Reference Guide. Greenwood Publishing Group. October 1, 2009. https://doi.org/10.1021/ed086p1182