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Hydrogen Sulfide
(redirected from Hydrosulphuric acid)

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hydrogen sulfide, chemical compound, H2S, a colorless, extremely poisonous gas that has a very disagreeable odor, much like that of rotten eggs. It is slightly soluble in water and is soluble in carbon disulfide. Dissolved in water, it forms a very weak dibasic acid that is sometimes called hydrosulfuric acid. Hydrogen sulfide is flammable; in an excess of air it burns to form sulfur dioxide and water, but if not enough oxygen is present, it forms elemental sulfur and water. Hydrogen sulfide is found naturally in volcanic gases and in some mineral waters. It is often formed during decay of animal matter. It is a part of many unrefined carbonaceous fuels, e.g., natural gas, crude oil, and coal; it is obtained as a byproduct of refining such fuels. It may be made by reacting hydrogen gas with molten sulfur or with sulfur vapors, or by treating a metal sulfide (e.g., ferrous sulfide, FeS) with an acid. Hydrogen sulfide reacts with most metal ions to form sulfides; the sulfides of some metals are insoluble in water and have characteristic colors that help to identify the metal during chemical analysis. Hydrogen sulfide also reacts directly with silver metal, forming a dull, gray-black tarnish of silver sulfide (Ag2S).
hydrogen sulfide [′hī·drə·jən ′səl‚fīd]
(inorganic chemistry)
H2S Flammable, toxic, colorless gas with offensive odor, boiling at -60°C; soluble in water and alcohol; used as an analytical reagent, as a sulfur source, and for purification of hydrochloric and sulfuric acids. Also known as hydrogen disulfide.

Hydrogen Sulfide 

(also sulfuretted hydrogen), H2S, the simplest compound of sulfur with hydrogen. A colorless gas, hydrogen sulfide has the odor of rotten eggs when in an extremely rarefied state.

Hydrogen sulfide was first studied in detail by K. Scheele in 1777. It is a component of volcanic gases and certain mineral waters (at Kemeri, Piatigorsk, and Matsesta in the USSR). In the Black Sea, the compound is found at depths exceeding 150 m. Hydrogen sulfide is constantly being formed in the putrefaction of organic material of animal origin.

At – 60.38°C, hydrogen sulfide is transformed into a colorless liquid, and at – 85.6°C, it crystallizes. Solid hydrogen sulfide exists in three phases, with transition points at – 170°C and – 147°C. The hydrogen sulfide molecule is polar and has an ionization potential of 10.5 volts. Under standard conditions, one volume of water dissolves approximately three volumes of hydrogen sulfide, with the formation of weak hydrosul-furic acid. The solubility of hydrogen sulfide decreases upon heating. The crystal hydrate H2S.6H20 may be obtained by cooling a saturated aqueous solution of hydrogen sulfide. The ignition temperature of hydrogen sulfide in air is approximately 300°C; hydrogen sulfide burns with a light blue flame. The reaction here is 2H2S + 3O2 = 2H2O + 2SO2 when there is an excess of oxygen and 2H2S + O2 = 2H2O + 2S when the oxygen is insufficient.

Mixtures of hydrogen sulfide and air are explosive when the former constitutes between 4 and 45 percent by volume. An aqueous solution of hydrogen sulfide (hydrosulfuric acid) gradually becomes turbid upon exposure to air as a consequence of the separation of sulfur. Hydrogen sulfide reacts with most metals and metal oxides in the presence of moisture or upon heating, forming the corresponding sulfides. It is a strong reducing agent. Halogens are reduced by H2S to their corresponding hydrogen compounds, and sulfuric acid is reduced to sulfur dioxide and sulfur:

H2SO4 + H2S = 2H2O + SO2 + S

Hydrogen sulfide is obtained by heating sulfur in a stream of hydrogen: H2 + S ⇄ H2S. The equilibrium of this reaction below 350°C is shifted to the right; at higher temperatures it is shifted to the left. Thermal dissociation of hydrogen sulfide begins at 400°C and is practically complete at approximately 1700°C.

In the laboratory, hydrogen sulfide is obtained by the action of dilute acids on ferrous sulfide:

FeS + 2HCl = FeCl2 + H2S

On an industrial scale, it is obtained through the purification of natural, petroleum, and coke-oven gases. Hydrogen sulfide is one of the most important reagents in chemical analysis. It is used in industry mainly in the production of sulfur. On a smaller scale, it finds use in the production of sulfuric acid and in organic synthesis. In balneotherapy, hydrogen sulfide is used as a therapeutic agent.

Hydrogen sulfide is very poisonous. Its maximum permissible concentration in the air at industrial facilities is 0.01 milligram per liter (mg/l).

Poisoning by hydrogen sulfide can occur in the extraction and refining of petroleum having a high sulfur content, in the production of viscose fibers and sulfur dyes, and in the cleaning and repair of sewer systems. Plants where sugar and leather goods are produced present similar hazards. Acute poisoning results from concentrations of 0.2–0.3 mg/l, and chronic poisoning from concentrations of 0.02 mg/l. Concentrations above 1 mg/l are fatal. The toxicity of hydrogen sulfide manifests itself in irritation of the mucous membranes of the eyes and upper respiratory passages and in an inhibitory effect on the respiratory enzymes in respiratory tissues.

In cases of acute poisoning that are not severe, conjunctivitis, corneal edema, and catarrh of the upper respiratory passages develop. In cases of moderate severity, there are also symptoms of damage to the central nervous system. In severe cases, toxic pulmonary edema and coma are possible, and in fulminant forms, paralysis of respiratory and cardiac activity can occur. With chronic poisoning, functional disorders of the nervous system develop, as well as malnutrition, anemia, bronchitis, tremor of the fingers and eyelids, and pain in the muscles and along the nerve trunks.

Preventive measures against hydrogen sulfide poisoning include the reduction of hydrogen sulfide pollution of the air in working areas, preliminary and periodic medical examinations, and the use of devices, such as respirators, to protect the respiratory organs.


Professional’nye bolezni, 3rd ed. Moscow, 1973.


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