Hydrogen sulfide

Hydrogen sulfide
IUPAC name	Hydrogen sulfide, sulfane
Other names	Sulfuretted hydrogen; sulfane; sulfur hydride; sour gas; sulfurated hydrogen; hydrosulfuric acid; sewer gas; stink damp
Identifiers
CAS number	[7783-06-4]
RTECS number	MX1225000
Properties
Molecular formula	H2S
Molar mass	34.082 g/mol
Appearance	Colorless gas.
Density	1.363 g/L, gas.
Melting point	-82.30°C (190.85 K)
Boiling point	-60.28°C (212.87 K)
Solubility in water	0.25 g/100 mL (40°C)
Acidity (pKa)	6.89 19±2 (See Text)
Structure
Molecular shape	Bent
Dipole moment	0.97 D
Hazards
Main hazards	Toxic, flammable.
R-phrases	R12, R26, R50
S-phrases	S1/2, S9, S16 S36, S38, S45, S61
Flash point	-82.4 °C
Related Compounds
Related hydrogen compounds	water; hydrogen selenide; hydrogen telluride
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)

Hydrogen sulfide (hydrogen sulphide in British English) is a chemical compound with the formula H₂S. It is a colorless, toxic, and flammable gas, and is responsible for the foul odor of rotten eggs and flatulence. It often results from the bacterial breakdown of organic matter in the absence of oxygen, such as in swamps and sewers (anaerobic digestion). It also occurs in volcanic gases, natural gas and some well waters. The odor of H₂S is commonly misattributed to elemental sulfur, which is odorless.

Occurrence

Deposit of sulfur on a rock, caused by volcanic gases

Small amounts of hydrogen sulfide occur in crude petroleum, but natural gas can contain up to 90 percent. Volcanoes and hot springs emit some H₂S, where it probably arises via the hydrolysis of sulfide minerals. That is, a metal sulfide may react with water to produce the metal oxide and hydrogen sulfide, as follows:

MS + H₂O to give MO + H₂S.

Sulfate-reducing bacteria obtain energy by oxidizing organic matter or hydrogen with sulfates, producing H₂S. These microorganisms are prevalent in low-oxygen environments, such as in swamps and standing waters. Sulfur-reducing bacteria (such as Salmonella) and some archaea obtain their energy by oxidizing organic matter or hydrogen with elemental sulfur, also producing H₂S. Other anaerobic bacteria liberate hydrogen sulfide when they digest sulfur-containing amino acids, for instance during the decay of organic matter. H₂S-producing bacteria also operate in the human colon, and the odor of flatulence is largely due to trace amounts of the gas. Such bacterial action in the mouth may contribute to bad breath. Evidence exists that hydrogen sulfide produced by sulfate-reducing bacteria in the colon may cause or contribute to ulcerative colitis.

About 10% of total global emissions of H₂S are due to human activity. By far the largest industrial route to H₂S occurs in petroleum refineries: the hydrodesulfurization process liberates sulfur from petroleum by the action of hydrogen. The resulting H₂S is converted to elemental sulfur by partial combustion via the Claus process, which is a major source of elemental sulfur. Other anthropogenic sources of hydrogen sulfide include coke ovens, paper mills (using the sulphate method), and tanneries. H₂S arises from virtually anywhere where elemental sulfur comes into contact with organic material, especially at high temperatures.

Hydrogen sulfide can be present naturally in well water. In such cases, ozone is often used for its removal. An alternative method uses a filter with manganese dioxide. Both methods oxidize sulfides to less toxic sulfates.

A buildup of hydrogen sulfide in the atmosphere could have caused the Permian-Triassic extinction event 252 million years ago.

General properties

Hydrogen sulfide is a covalent hydride structurally related to water (H₂O) since oxygen and sulfur occur in the same periodic table group.

Hydrogen sulfide is weakly acidic, dissociating in aqueous solution into hydrogen cations H⁺ and the hydrosulfide anion HS⁻:

H₂S → HS⁻ + H⁺

K_a = 1.3×10⁻⁷ mol/L; pK_a = 6.89.

The sulfide ion, S²⁻, is known in the solid state but not in aqueous solution (c.f. oxide). The second dissociation constant of hydrogen sulfide is often stated to be around 10⁻¹³, but it is now clear that this is an error caused by oxidation of the sulfur in alkaline solution. The current best estimate for pK_a2 is 19±2.

Hydrogen sulfide reacts with many metals cations to produce the corresponding metal sulfides. Well-known examples are silver sulfide (Ag₂S), the tarnish that forms on silver when exposed to the hydrogen sulfide of the atmosphere, and cadmium sulfide (CdS), a pigment also known as cadmium yellow. Transition metal sulfides are characteristically insoluble, thus H₂S is commonly used to separate metal ions from aqueous solutions. (Sulfides should not be confused with sulfites or sulfates, which contain the sulfite ion SO₃²⁻ and the sulfate ion SO₄²⁻, respectively.)

Hydrogen sulfide is corrosive and renders some steels brittle, leading to sulfide stress cracking—a concern especially for handling "sour gas" and sour crude oil in the oil industry. Hydrogen sulfide burns to give the gas sulfur dioxide, which is more familiar as the odor of a burnt match.

Production

Hydrogen sulfide is obtained by its separation from sour gas—that is, natural gas with a high content of H₂S. It can be produced by reacting hydrogen gas with molten elemental sulfur at about 450 °C. Hydrocarbons can replace hydrogen in this process. Sulfate-reducing bacteria produce hydrogen sulfide under ambient conditions by the reduction of sulfate or from elemental sulfur. The standard lab preparation is to gently heat iron sulfide with a strong acid (see below). A less well-known but more convenient alternative is to react aluminum sulfide with water:

H₂O + Al₂S₃ → H₂S + Al₂O₃.

Al₂S₃ is readily prepared by ignition of the elements.

Uses

Production of thioorganic compounds

Several organosulfur compounds are produced using hydrogen sulfide. These include methanethiol, ethanethiol, and thioglycolic acid.

Alkali metal sulfides

Upon combining with alkali metal bases, hydrogen sulfide converts to alkali hydrosulfides such as sodium hydrosulfide and sodium sulfide, which are used in the degradation of biopolymers. The depilation of hides and the delignification of pulp by the Kraft process are both effected by alkali sulfides.

In analytical chemistry

Hydrogen sulfide used to have importance in analytical chemistry for well over a century, in the qualitative inorganic analysis of metal ions. For such small-scale laboratory use, H₂S was made as needed in a Kipp generator by reaction of sulfuric acid (H₂SO₄) with ferrous sulfide FeS. Kipp generators were superseded by the use of thioacetamide, an organic solid that converts in water to H₂S. In these analyses, heavy metal (and nonmetal) ions (e.g. Pb(II), Cu(II), Hg(II), As(III)) are precipitated from solution upon exposure to H₂S. The components of the resulting precipitate redissolve with some selectivity.

A precursor to metal sulfides

As indicated above, many metal ions react with hydrogen sulfide to give the corresponding metal sulfides. This conversion is widely exploited. In the purification of metal ores by flotation, mineral powders are often treated with hydrogen sulfide to enhance the separation. Metal parts are sometimes passivated with hydrogen sulfide. Catalysts used in hydrodesulfurization are routinely activated with hydrogen sulfide, and the behavior of metallic catalysts used in other parts of a refinery is also modified using hydrogen sulfide.

Miscellaneous applications

Hydrogen sulfide is also used in the separation of deuterium oxide, i.e. heavy water, from normal water via the Girdler Sulfide process.

Safety

Hydrogen sulfide is a highly toxic and flammable gas. Being heavier than air, it tends to accumulate at the bottom of poorly ventilated spaces. Although very pungent at first, it quickly deadens the sense of smell, so potential victims may be unaware of its presence until it is too late. For more information see an MSDS.

Toxicity

Hydrogen sulfide is considered a broad-spectrum poison, meaning that it can poison several different systems in the body, although the nervous system is most affected. The toxicity of H₂S is comparable with that of hydrogen cyanide. It forms a complex bond with iron in the mitochondrial cytochrome enzymes, thereby blocking oxygen from binding and stopping cellular respiration. Since hydrogen sulfide occurs naturally in the environment and the gut, enzymes exist in the body capable of detoxifying it by oxidation to (harmless) sulfate. Hence low levels of sulfide may be tolerated indefinitely. However, at some threshold level, the oxidative enzymes will be overwhelmed. This threshold level is believed to average around 300-350 ppm. Many personal safety gas detectors are set to alarm at 10 ppm and to go into high alarm at 15 ppm (Utility, sewage & petrochemical workers).

An interesting diagnostic clue of extreme poisoning by H₂S is the discoloration of copper coins in the pockets of the victim. Treatment involves immediate inhalation of amyl nitrite, injections of sodium nitrite, inhalation of pure oxygen, administration of bronchodilators to overcome eventual bronchospasm, and in some cases hyperbaric oxygen therapy.

Exposure to lower concentrations can result in eye irritation, a sore throat and cough, nausea, shortness of breath, and fluid in the lungs. These symptoms usually go away in a few weeks. Long-term, low-level exposure may result in fatigue, loss of appetite, headaches, irritability, poor memory, and dizziness. Chronic exposures to low level H₂S (around 2 ppm) has been implicated in increased miscarriage and reproductive health issues amongst Russian and Finnish wood pulp workers, but the reports hadn't (as of circa 1995) been replicated. Higher concentrations of 700-800 ppm tend to be fatal.

• 0.0047 ppm is the recognition threshold, the concentration at which 50% of humans can detect the characteristic odor of hydrogen sulfide [1], normally described as resembling "a rotten egg."
• 10-20 ppm is the borderline concentration for eye irritation.
• 50-100 ppm leads to eye damage.
• At 150-250 ppm the olfactory nerve is paralyzed after a few inhalations, and the sense of smell disappears, often together with awareness of danger,
• 320-530 ppm leads to pulmonary edema with the possibility of death.
• 530-1000 ppm causes strong stimulation of the central nervous system and rapid breathing, leading to loss of breathing;
  • 800 ppm is the lethal concentration for 50% of humans for 5 minutes exposure(LC50).
• Concentrations over 1000 ppm cause immediate collapse with loss of breathing, even after inhalation of a single breath.

A practical test used in the oilfield industry to determine whether someone requires overnight observation for pulmonary edema is the knee test: if a worker that gets "gassed" loses his balance and at least one knee touches the ground, the dose was high enough to cause pulmonary edema.

Function in the body

Hydrogen sulfide is produced in small amounts by some cells of the mammalian body and has a number of biological functions. (nitric oxide (NO) and carbon monoxide (CO) are also implicated as gaseous signalling agents.) It is produced from cysteine by various enzymes. It acts as a vasodilator and is also active in the brain, where it increases the response of the NMDA receptor and facilitates long term potentiation, which is involved in the formation of memory. Eventually the gas is converted to sulfites and further oxidized to thiosulfate and sulfate.^{[citation needed]} Due to its effects similar to NO (without it's potential to form peroxides by interacting with superoxide), hydrogen sulfide is now recognized as a potential cardioprotective agent. Vasoactivity of garlic is caused by catabolism of its polysulfides to H₂S, a reaction which could depend on reduction mediated by glutathione. In trisomy 21 (the most common form of Down syndrome) the body produces an excess of hydrogen sulfide.

Induced hibernation

In 2005 it was shown that mice can be put into a state of suspended animation by applying a low dosage of hydrogen sulfide (80 ppm H₂S) in the air. The breathing rate of the animals sank from 120 to 10 breaths per minute and their temperature fell from 37 °C to 2 °C above ambient temperature (in effect, they had become cold-blooded). The mice survived this procedure for 6 hours and afterwards showed no negative health consequences.

Such a hibernation occurs naturally in many mammals and also in toads, but not in mice. (Mice can fall into a state called clinical torpor when food shortage occurs). If the H₂S-induced hibernation can be made to work in humans, it could be useful in the emergency management of severely injured patients, and in the conservation of donated organs.

As mentioned above, hydrogen sulfide binds to cytochrome oxidase and thereby prevents oxygen from binding, which leads to the dramatic slowdown of metabolism. Animals and humans naturally produce some hydrogen sulfide in their body; researchers have proposed that the gas is used to regulate metabolic activity and body temperature, which would explain the above findings.

In 2006, it was shown that the blood pressure of mice treated in this fashion with hydrogen sulfide did not significantly decrease.

Participant in the sulfur cycle

Hydrogen sulfide is a central participant in the sulfur cycle, the biogeochemical cycle of sulfur on Earth. As mentioned above, sulfur-reducing and sulfate-reducing bacteria derive energy from oxidizing hydrogen or organic molecules in the absence of oxygen by reducing sulfur or sulfate to hydrogen sulfide. Other bacteria liberate hydrogen sulfide from sulfur-containing amino acids. Several groups of bacteria can use hydrogen sulfide as fuel, oxidizing it to elemental sulfur or to sulfate by using dissolved oxygen, metal oxides (e.g. Fe oxyhyroxides and Mn oxides) or nitrate as oxidant. The purple sulfur bacteria and the green sulfur bacteria use hydrogen sulfide as electron donor in photosynthesis, thereby producing elemental sulfur. (In fact, this mode of photosynthesis is older than the mode of cyanobacteria, algae and plants which uses water as electron donor and liberates oxygen.)

H₂S implicated in mass extinctions

Hydrogen sulfide has been implicated in some of the five mass extinctions that have occurred in the Earth's past. Although asteroid impacts are thought to have caused some extinctions, the Permian mass extinction (sometimes known as the "Great Dying") may have been caused by hydrogen sulfide. Organic residues from these extinction boundaries indicate that the oceans were anoxic (oxygen depleted) and had species of shallow plankton that metabolized H₂S. The formation of H₂S may have been initiated by massive volcanic eruptions, which emitted CO₂ and methane into the atmosphere which warmed the oceans, lowering their capacity to absorb oxygen which would otherwise oxidize H₂S. The increased levels of hydrogen sulfide could have killed oxygen-generating plants as well as depleted the ozone layer causing further stress. Small H₂S blooms have been detected in modern times in the Dead sea and in the Atlantic ocean off the coast of Namibia.

See also

• Sulfide
• Sulfur
• Water

References

ISBN links support NWE through referral fees

• Brown Jr., Theodore L., H. Eugene LeMay, Bruce Edward Bursten, and Julia R. Burdge. 2002. Chemistry: The Central Science. 9th ed. Upper Saddle River, NJ: Prentice Hall. ISBN 0130669970.

• Chang, Raymond. 2006. Chemistry. 9th ed. New York, NY: McGraw-Hill Science/Engineering/Math. ISBN 0073221031.

• Cotton, F. Albert, and Geoffrey Wilkinson. 1980. Advanced Inorganic Chemistry. 4th ed. New York, NY: Wiley. ISBN 0-471-02775-8.

• "Hydrogen Sulfide," Committee on Medical and Biological Effects of Environmental Pollutants, University Park Press, 1979, Baltimore. ISBN 0-8391-0127-9

External links

• International Chemical Safety Card 0165. Retrieved December 8, 2007.
• Concise International Chemical Assessment Document 53. Retrieved December 8, 2007.
• National Pollutant Inventory - Hydrogen sulfide fact sheet. Retrieved December 8, 2007.
• NIOSH Pocket Guide to Chemical Hazards. Retrieved December 8, 2007.
• MSDS safety data sheet. Retrieved December 8, 2007.
• Kamoun, P. Abstract of survey article on H₂S as used by the body. Retrieved December 8, 2007.

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