Introduction to Sulfur Functional Groups
[Mercaptans (Thiols) - Sulfonic Acids (Sulfonates) - Carbon Disulfide - Hydrogen Sulfide]

James Richard Fromm

thiol.gif (1028 bytes) sulfide.gif (1092 bytes) disulfide.gif (1134 bytes)
Mercaptan (Thiol) Sulfide Disulfide

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Sulfenic Acid Sulfinic Acid Sulfonic Acid
rso3r.gif (1119 bytes) sulfoxide.gif (1000 bytes) sulfone.gif (1045 bytes)
Sulfonate Ester Sulfoxide Sulfone

Mercaptans (Thiols)

Sulfur is very similar to oxygen in its chemistry and many of the organic sulfur compounds behave in a way similar to that of their oxygen counterparts.


General Formula

Mercaptans (Thio-alcohols), organic chemical compounds of the type R-SH (R = an alkyl group).  The name is derived from mercurium captans, in reference to the fact that these compounds react readily with mercuric oxide to form crystalline mercury derivatives.  The mercaptans may be prepared by the action of the alkyl halides on an alcoholic solution of potassium hydrosulfide; by the reduction of the sulfo-chlorides, e.g. C2H5SO2Cl (chlorides of sulfonic acids), by heating the salts of esters of sulfuric acid with potassium hydrosulfide, and by heating the alcohols with phosphorus pentasulfide.  They are colorless liquids, which are insoluble in water and possess a characteristic offensive smell.  On oxidation by nitric acid they yield sulfonic acids.  They combine with aldehydes and ketones, with elimination of water and formation of mercaptals and mercaptols.  Methyl mercaptan, CH3-SH, is a liquid which boils at 5.8C, and forms a crystalline hydrate with water.  Ethyl mercaptan, C2H5-SH, is a colorless liquid which boils at 36.2C.  It is used commercially in the preparation of sulfonal.  The mercury salt, Hg(SC2H5)2, crystallizes from alcohol in plates.  When heated with alcohol to 190C it decomposes into mercury and ethyldisufide.

Mercaptans are malodorous substances otherwise similar to the corresponding alcohols, and parallel them in name.  For example: methyl mercaptan, CH3-SH,


Methyl Mercaptan,

is the counterpart to methyl alcohol (CH3OH).


When a thiol group is a substituent on an alkane, there are several ways of naming the resulting thiol:

Physical Properties


Many thiols are colorless liquids having an odor resembling that of garlic.  The odor of thiols is often strong and repulsive, particularly for those of low molecular weight.  Thiols bind strongly to skin proteins, and are responsible for the intolerable, persistent odor produced by feces, rotting flesh and the spraying of skunks.   Natural gas distributors began adding various forms of pungent thiols, usually ethanethiol, to natural gas, which is naturally odorless, after the deadly 1937 New London School explosion in New London, Texas.  Thiols are also responsible for a class of wine faults caused by an unintended reaction between sulfur and yeast.  However, not all thiols have unpleasant odors.  For example, grapefruit mercaptan, a monoterpenoid thiol, is responsible for the characteristic scent of grapefruit.

Boiling Points & Solubility

Due to the small electronegativity difference between sulfur and hydrogen, an S-H bond is practically nonpolar covalent.  Therefore, the S-H bond in the thiols have a lower dipole moment as compared to the alcohol's O-H bond.  Thiols show little association by hydrogen bonding, with both water molecules and among themselves.  Hence, they have lower boiling points and are less soluble in water and other polar solvents than alcohols of similar molecular weight.  Thiols are as soluble and have similar boiling points to isomeric sulfides.

Chemical Properties


The methods used in making thiols are analogous to those used to make alcohols and ethers.  The reactions are quicker and higher yielding because sulphur anions are better nucleophiles than oxygen atoms.

Thiols are formed when a halogenoalkane is heated with a solution of sodium hydrosulphide

CH3CH2Br + NaSH heated in ethanol (aq) arrow2.gif (113 bytes) CH3CH2SH + NaBr

In addition, disufides can be readily reduced by reducing agents like lithium aluminum hydride in dry ether to form two thiols.


The thiol group is the sulfur analog of the hydroxyl group (-OH) found in alcohols.   Since sulfur and oxygen belong to the same periodic table group, they share some similar chemical bonding properties.  Like alcohol, in general, the deprotonated form RS- (called a thiolate) is more chemically reactive than the protonated thiol form RSH

The chemistry of thiols is thus related to the chemistry of alcohols: thiols form thioethers, thioacetals and thioesters, which are analogous to ethers, acetals, and esters.  Furthermore, a thiol group can react with an alkene to create a thioether.   (In fact, biochemically, thiol groups may react with vinyl groups to form a thioether linkage.)


The sulfur atom of a thiol is quite nucleophilic, rather more so than the oxygen atom of an alcohol.  The thiol group is fairly acidic with a usual pKa around 10 to 11.   In the presence of a base, a thiolate anion is formed which is a very powerful nucleophile.  The group and its corresponding anion are readily oxidized by reagents such as bromine to give an organic disulfide (R-S-S-R).

2R-SH + Br2 arrow2.gif (113 bytes) R-S-S-R + 2HBr

Oxidation by more powerful reagents such as sodium hypochlorite or hydrogen peroxide yield sulfonic acids (RSO3H).

Biological Role

As the functional group of the amino acid cysteine, the thiol group plays an important role in biological systems. 

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When the thiol groups of two cysteine residues (as in monomers or constituent units) are brought near each other in the course of protein folding, an oxidation reaction can create a cystine unit with a disulfide bond (-S-S-). 

cystine.gif (1652 bytes)


Disulfide bonds can contribute to a protein's tertiary structure if the cysteines are part of the same peptide chain, or contribute to the quaternary structure of multi-unit proteins by forming fairly strong covalent bonds between different peptide chains.   The heavy and light chains of antibodies are held together by disulfide bridges.   Also, the kinks in curly hair are a product of cystine formation.  Permanents take advantage of the oxidizability of cysteine residues.  The chemicals used in hair straightening are reductants that reduce cystine disulfide bridges to free cysteine sulfhydryl groups, while chemicals used in hair curling are oxidants that oxidize cysteine sulfhydryl groups to form cystine disulfide bridges.  Sulfhydryl groups in the active site of an enzyme can form noncovalent bonds with the enzyme's substrate as well, contributing to catalytic activity.  Active site cysteine residues are the functional unit in cysteine proteases.

Other Thiols

Methanethiol Ethanethiol

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Coenzyme A

Coenzyme A (CoA, CoASH, or HSCoA) is a coenzyme, notable for its role in the synthesis and oxidation of fatty acids, and the oxidation of pyruvate in the citric acid cycle.  It is adapted from cysteamine, pantothenate and adenosine triphosphate.

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Lipoamide is a trivial name for 6,8-dithiooctanoic amide.  It is 5,8-dithiooctanoic acid's functional form where the carboxyl group is attached to protein (or any other amine) by an amide linkage (containing -NH2) to an amino group.

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Glutathione (GSH)

Glutathione (GSH), contains an unusual peptide linkage between the amine group of cysteine and the carboxyl group of the glutamate side chain.  Glutathione, an antioxidant, protects cells from toxins such as free radicals.

Thiol groups are kept in a reduced state within ~5 mM in animal cells.  In effect, glutathione reduces any disulfide bonds formed within cytoplasmic proteins to cysteines by acting as an electron donor.  Glutathione is found almost exclusively in its reduced form, since the enzyme which reverts it from its oxidized form (GSSG), glutathione reductase, is constitutively active and inducible upon oxidative stress.  In fact, the ratio of reduced to oxidized glutathione within cells is often used scientifically as a measure of cellular toxicity.

Sulfonic Acids - Sulfonates

The sulfur compounds analogous to the carboxylic acids are the sulfonic acids.  

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Sulfonic Acid

Sulfonic acids are produced by treatment of organic compounds with sulfuric acid. They have the structure -SO3H rather than the -COOH structure of carboxylic acids. The proton on a sulfonic acid is more acidic than the proton on the corresponding carboxylic acid.  Sulfonic acid groups are often used to make an organic compound more soluble in water, since the sulfonate ion is predominant at almost all values of pH.

Sulfonic acid is a hypothetical acid with formula H-SO2-OH.   This compound is a tautomer of sulfurous acid HO-S(=O)-OH, but less stable, and would likely convert to that very quickly if it were formed.  Although this compound is unimportant, there are many derived compounds, with formula R-S(=O)2-OH for various R.  These may then form salts or esters, called sulfonates.

Sulfonic acids are a class of organic acids with the general formula RSO2-OH, where R is usually a hydrocarbon side chain.  Sulfonic acids are typically much stronger acids than their carboxylic equivalents, and have the unique tendency to bind to proteins and carbohydrates tightly; most "washable" dyes are sulfonic acids (or have the functional sulfonyl group in them) for this reason.  They are also used as catalysts and intermediates for a number of different products. Sulfonic acid salts (sulfanates) are important as detergents, and the antibacterial sulfa drugs are also sulfonic acid derivatives.  The simplest example is methanesulfonic acid, CH3SO2OH, which is a reagent regularly used in organic chemistry.  p-Toluenesulfonic acid is also an important reagent.

Note that the sulfonic acids and sulfonates are analogous to carboxylic acids and carboxylates; in both cases, -C(=O)- is replaced by -S(=O)2-. Chemical properties are similar as well, although sulfonic acids are often even stronger acids than carboxylic acids, the hydrogen being easier to leave than in most compounds, and they readily form esters.

The sulfonic acid and sulfonate functional groups, -SO2OH and -SO2O-, are found in many chemical compounds, e.g. certain detergents and dyes as well as in strongly acidic cation exchange resins.

A sulfonate ion is an ion that contains the -S(=O)2-O- functional group. 

180-sulfonate.jpg (4012 bytes)

General Formula

The general formula is RSO2O-, where R is some organic group.  They are conjugate basis of sulfonic acids with formula RSO2OH.

Sulfonates, being weak acids, are good leaving groups in Sn1, Sn2, E1 and E2 reactions.

As is common, the same term is used for compounds containing this functional group, ionic salts, or similar covalent compounds, esters.


The preceding terms apply to salts and esters as well as anions.

Sulfonic Acid Chlorides

Sulfonic acid chlorides are a class of organic compounds with the general formula R-SO2-Cl.  These compounds react readily with alcohols to sulfonic esters.  Important acid chlorides are tosyl chloride, brosyl chloride, nosyl chloride and mesyl chloride.  One synthetic procedure to synthesize sulfonic acid chlorides is the Reed reaction.

Sulfonic Esters

Sulfonic esters are a class of organic compounds with the general formula R-SO2-OR.   Sulfonic esters are considered good leaving groups in nucleophilic aliphatic substitution.


A sulfoxide is a chemical compound containing a sulfinyl functional group attached to two carbon atoms.  Sulfoxides can be considered as oxidized sulfides.   (The use of the alternative name sulphoxide is discouraged by IUPAC.)   An example of an sulfoxide occurring in nature is alliin.

Nature of the Bond

The structural formula of a sulfoxide is generally represented as R-S(=O)-R' where R and R' are the organic groups.  The bond between the sulfur and oxygen atoms is not a normal double bond (with overlap of p-bonds between p-orbitals of adjacent atoms), and it is believed that electrostatic interaction between a negatively charged oxygen and positively charged sulfur accounts for most of the bonding.

300-sulfoxide-resonance.jpg (6987 bytes)

Not drawn is the electron lone pair on sulfur and the geometry around sulfur is actually tetrahedral as with carbon.  When the two organic residues are dissimilar, sulfoxides can be chiral.  Chiral sulfoxides find application in certain drugs such as esomeprazole and Armodafinil and they are also employed as chiral auxiliary.  Many chiral sulfoxides are prepared from asymmetric oxidation of achiral sulfides with a transition metal and a chiral ligand.


Sulfides are often the starting materials for sulfones by organic oxidation.  For example, dimethyl sulfide with oxidation state of -2 is oxidized to dimethyl sulfoxide with oxidation state 0.  Further oxidation takes the compound to dimethyl sulfone with oxidation state +2.

Methyl sulfoxides (like DMSO) have an acidic character, because the sulfoxide group stabilizes the anion that results from loss of a methyl hydrogen.

Dimethyl Sulfoxide

Dimethyl sulfoxide (DMSO) is the chemical compound with the formula (CH3)2SO.   This colorless liquid is an important polar aprotic solvent that dissolves both polar and nonpolar compounds and is miscible in a wide range of organic solvents as well as water.  It has a distinctive property of penetrating the skin very readily, allowing the handler to taste it.  Some describe it as an "oyster-like" taste, others claim it tastes like garlic.


Dimethyl sulfoxide is a by-product of wood pulping.  One of the leading suppliers of DMSO is the gaylord Chemical Corporation.



DMSO is an important polar aprotic solvent.  It is less toxic than other members of this class such as dimethylformamide, dimethylacetamide, N-methyl-2-pyrrolidone, HMPA.   Because of its excellent solvating power, DMSO is frequently used as solvent for chemical reactions involving the reactions of salts.  Because DMSO is only weakly acidic, it tolerates relatively strong bases.  The main problem with DMSO as a solvent is its high boiling point, thus its solutions are not typically evaporated but instead diluted to isolate the reaction product.  DMSO is an effective paint stripper, being much safer than many of the others such as nitromethane and dichloromethane.

In its deuterated form, i.e. DMSO-d6, it is a useful and inexpensive solvent for NMR spectroscopy, again due to its ability to dissolve a wide range of analytes and its own simple spectrum.  Disadvantages to the use of DMSO-d6 are its high viscosity, which broadens signals, and high boiling point, which interferes with sample recovery from the NMR solvent.  Often it is used with deuterochloroform, because the 1:1 mixture has a low viscosity.


The sulfur center in DMSO is nucleophilic toward soft electrophiles and the oxygen is basic toward hard nucleophiles.  The methyl groups of DMSO are somewhat acidic in character (pKa=35) due to the stabilization of the resultant carbanion by the S(O)R group.

DMSO reacts with methyl iodide to form a sulfoxonium salt [(CH3)3SO]I, which can be deprotonated with sodium hydride to form the sulfur ylide:

(CH3)2SO + CH3I arrow2.gif (113 bytes) [(CH3)3SO]I

[(CH3)3SO]I + NaH arrow2.gif (113 bytes) [(CH2)3CH2SO + NaI + H2

In organic synthesis, DMSO is used as an oxidant.  Products of ozonolysis, trioxolanes, are quenched with dimethyl sulfide to produce aldehydes and DMSO.


DMSO is used in the PCR reaction to inhibit secondary structures in the DNA template or the DNA primers.  It is added to the PCR mix before reacting, where it interferes with the self-complementarity of the DNA, allowing otherwise troublesome reactions to occur.  However, use of DMSO in PCR increases the mutation rate.


In cryobiology DMSO has been used as a cryoprotectant and is still an important constituent of cryoprotectant vitrification mixtures used to preserve organs, tissues, and cell suspensions.  It is particularly important in the freezing and long-term storage of embryonic stem cells and hematopoietic stem cells, which are often frozen in a mixture of 10% DMSO and 90% fetal calf serum.  As part of an autologous bone marrow transplant the DMSO is re-infused along with the patient's own hematopoietic stem cells.

Use of DMSO in medicine dates from around 1963, when a University of Oregon Medical School team, headed by Stanley Jacob, discovered it could penetrate the skin and other membranes without damaging them and could carry other compounds into a biological system.   Some people report an onion- or garlic-like taste after touching DMSO.  This effect is likely due to catabolic processes that reduce DMSO to dimethyl sulfide.  In the medical field DMSO is predominantly used as a topical analgesic, a vehicle for topical application of pharmaceuticals, as an anti-inflammatory and an antioxidant.  It has been examined for the treatment of numerous conditions and ailments.  The Food and Drug Administration (FDA) has approved DMSO usage only for the palliative treatment of interstitial cystitis.  Also, DMSO is commonly used in the veterinary field as a liniment for horses.

Because DMSO increases the rate of absorption of some compounds through organic tissues including skin, it can be used as a drug delivery system.

Dimethyl sulfoxide dissolves a variety of organic substances, including carbohydrates, polymers, peptides, as well as many inorganic salts and gases.  Loading levels of 50-60 wt. % are often observed vs 10-20 wt. % with typical solvents.  For this reason DMSO plays a role in sample management and High-throughput screening operations in drug design.


Because DMSO easily penetrates the skin, substances dissolved in DMSO may be quickly absorbed.  For instance, a solution of sodium cyanide in DMSO can cause cyanide poisoning through skin contact.  DMSO by itself has low toxicity.  Dimethyl sulfoxide can produce an explosive reaction when exposed to acid chlorides.   Recently, it was found that DMSO waste disposal into sewers can cause environmental odor problems in cities: Waste water bacteria transform DMSO under hypoxic (anoxic) conditions into dimethyl sulfide (DMS) that is slightly toxic and has a strong disagreeable odor, similar to rotten cabbage.

Carbon Disulfide (S=C=S)

Carbon disulfide is a colorless liquid with the formula CS2.  It has a pleasant odor that is like that of chloroform, although it is usually impure yellowish with an unpleasant odor like that of rotting radishes due to traces of other sulphurous species, such as carbonyl sulfide (COS).

Occurrence and Manufacture

Small amounts of carbon disulfide are found in gases released by volcanic eruptions or marshes.  CS2 once was manufactured by combining carbon (or coke) and sulfur at high temperatures.  A lower temperature reaction, requiring only 600C involves natural gas in the presence of kieselgel or alumina catalysts:

CH4 + 1/2 S8 arrow2.gif (113 bytes) CS2 + 2H2S

Chemical Properties

CS2 is structurally analogous to CO2, but more reactive toward nucleophiles and bases and more easily reduced. This difference in reactivity can be attributed to the weaker p donor-ability of the sulfido centers.

Addition of Nucleophiles

Nucleophiles such as amines afford dithiocarbamates:

2R2NH + CS2 arrow2.gif (113 bytes) [R2NH2+][R2NCS2-]

Xanthates form similarly from alkoxides:

RONa + CS2 arrow2.gif (113 bytes) [Na+][ROCS2-]

Sodium sulfide affords trithiocarbonate:

Na2S + CS2 arrow2.gif (113 bytes) [Na+]2[CS32-]


Sodium reduces CS2 to give the heterocycle "dmit2-"

3CS2 + 4Na arrow2.gif (113 bytes) Na2C3S5 + Na2S

Direct electrochemical reduction affords the tetrathiooxalate anion:

2CS2 + 2e arrow2.gif (113 bytes) C2S42-


Chlorination of CS2 is the route to Thiophosgene.

CS2 is a ligand for many metal complexes, forming pi complexes.  It is used to manufacture regenerated cellulose (the main ingredient of viscose rayon and cellophone), carbon tetrachloride and organic sulfur compounds including those mentioned above.

Health Effects

At very high levels, carbon disulfide may be life-threatening because it affects the nervous system.  People who breathed carbon disulfide near an accident involving a derailed railroad car showed changes in breathing and some chest pains, although these effects are probably due to the sulfur oxides formed in the ensuing fire.

Some workers who breathed high levels during working hours for at least 6 months had headaches, tiredness, and trouble sleeping, in some cases even serious central and peripheral nervous system disease such as toxic encephalitis and peripheral mixed neuropathy.  Most of these data come from the Viscose Rayon Industry, where small amounts of H2S may also have been present.  Among workers who breathed lower levels, some developed very slight changes in their nerves.  An increased risk of cardiovascular death has also been established in several countries.

Studies in animals indicate that carbon disulfide can affect the normal functions of the brain, liver, and heart.  After pregnant rats breathed carbon disulfide in the air, some of the newborn rats died or had birth defects.

Liquid carbon disulfide has caused skin burns.

Hydrogen Sulfide (H2S)

Hydrogen Sulfide, H2S, production is a natural by-product of yeast metabolism. Yeasts form H2S by reduction of sulfates, sulfites and elemental sulfur during synthesis of amino acids. Problems arise when yeast H2S production exceeds its utilization in forming amino acids and excess H2S "leaks" into the wine. Additional H2S can be formed by the natural breakdown of sulfur containing amino acids.

H2S is the most frequent volatile sulfur compound found in fermenting wine. Large amounts of H2S may be produced during fermentation. Under normal conditions, much of this H2S is volatilized from the wine along with CO2. The residual H2S may pose a serious problem due to its low sensory threshold and its potential reactivity. Several other volatile sulfur compounds arise from chemical modification of H2S.

Hydrogen sulfide, H2S, is a colorless, toxic, flammable gas that is responsible for the foul odor of rotton eggs and flatulence.  It often results when bacteria break down organic matter in the absence of oxygen, such as in swamps, and sewers (alongside the process of anaerobic digestion).  It also occurs in volcanic gases, natural gas and some well waters.  This is the odor that is commonly misattributed to elemental sulfur, which is in fact odorless.

Hydrogen sulfide is also known as sulfane, sulfur hydride, sour gas, sulfurated hydrogen, hydrosulfuric acid, sewer gas and stink damp.  IUPAC accepts the names "hydrogen sulfide" and "sulfane"; the latter one is used exclusively when naming more complicated compounds.


Hydrogen sulfide is a covalent hydride chemically 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 arrow2.gif (113 bytes) HS- + H+

Ka = 1.310-7 mol/L; pKa = 6.89.

The sulfide ion, S2-, is known in the solid state but not in aqueous solution (c.f. oxide).  The second dissociation contant 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 pKa2 is 192.

Hydrogen sulfide reacts with many metals 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.

Hydrogen sulfide is corrosive and renders some steels brittle, leading to sulfide stress cracking — a concern especially for handling acid gas and sour crude oil in the oil industry.

(Sulfides should not be confused with sulfites or sulfates, which contain the sulfite ion SO32- and the sulfate ion SO42-, respectively.)

Hydrogen sulfide burns to give the gas sulfur dioxide, which is more familiar to people as the odor of a burnt match.


Small amounts of hydrogen sulfide occur in crude petroleium but natural gas can contain up to 28%.  Volcanoes and hot springs emit some H2S, where it probably arises via the hydrolysis of sulfide minerals, i.e.

MS + H2O arrow2.gif (113 bytes) MO + H2S.

Normal average concentration in clean air is about 0.0001-0.0002 ppm.

Sulfate-reducing bacteria obtain their 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 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 fairly non-toxic sulfates.

It is also produced by Salmonella bacteria.

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

Manufacture and Use

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.

Industrial production focuses on separation of hydrogen sulfide from sour gas — natural gas with high content of H2S.

It is used in metallurgy for the preparation of metallic sulfides.  It also finds use in preparation of phosphors and oil additives, in separation of metals, removal of metallic impurities, and in organic chemical synthesis.  Hydrogen sulfide is also used in the separation of deuterium oxide, i.e. heavy water, from normal water via the Girdler Sulfide process.


Hydrogen sulfide is a highly toxic and flammable gas.  Because it is 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.


Hydrogen sulfide is considered a broad-spectrum poison, meaning that it can poison several different systems in the body, although the nervous sytem 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 caugh, 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.  Higher concentrations of 700-800 ppm tend to be fatal.

A practical test used in the oilfield industry to determine whether someone requires overnight observation for pullmonary edema is the knee test: if a worker that gets "gased" loses his balance and at least one knee touches the ground, the dose was high enough to cause pulmonary edema.  This is important as the worker may feel fine after some fresh air, and not think medical attention is needed, but the onset of pulmonary edema may occur many hours later when the worker is asleep: the worker's lungs could fill with fluid, and the sedative effects of the gas may prevent the worker from waking up.

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. (Only two other such gases are currently known: nitric oxide (NO) and carbon monoxide (CO).)  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.

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 37C to 2C 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 geologic time in the Earth's past.  The extinction 65 million years ago (at the Cretaceous — Tertiary boundary) was almost certainly caused by an asteroid impact.  However, evidence of asteroid impacts for many of the other mass extinctions has been weak.  In particular, the biggest extinction of them all (at the end of the Permian) has little evidence of a sudden die off because of an asteroid or comet impact.

The Permian mass extinction (sometimes known as the "Great Dying") killed 96% of ocean dwellers and 70% of plants, animals and even insects on land.  Evidence from the ratio of Carbon 12 — Carbon 13 isotopes suggests that the Permian extinction took place in spurts over around 100,000 years.  Studies of organic residues from these extinction boundaries show that the oceans were anoxic (oxygen depleted) and had species of shallow plankton that used H2S for oxidization.  Lastly these were periods when the Earth's carbon dioxide levels were high and therefore the water was warm.

The logic goes this way: during periods of great heat (caused by massive volcanic eruptions pumping CO2 and methane in to the atmosphere) the oceans can absorb less oxygen.  In areas on the ocean floor which were anoxic, the breakdown of organics created hydrogen sulfide.  Normally the oxygen-rich waters above would "cap" the water and oxidize any undissolved H2S gas before it rose to the surface, but the warmer waters lacked enough oxygen in many places and the H2S-rich water reached the surface in many areas.

This killed off oxygen-generating plants causing a runaway H2S effect.   Additionally, the H2S attacked the ozone layer causing further stress on other species, particularily oxygen-producing plankton and species in the Arctic areas (which would be farthest from the H2S blooms).

Small H2S blooms have been detected in modern times in the Dead sea and in the Atlantic ocean off the coast of Namibia.

Reported Sensory Thresholds for Sulfur Compounds

Compound Structure Sensory Description Range (ppb)
hydrogen sulfide H2S rotten egg, sewage-like 0.9 - 1.5
ethyl mercaptan CH3CH2SH burnt match, sulfidy, earthy 1.1 - 1.8
methyl mercaptan CH3SH rotten cabbage, burnt rubber 1.5
diethyl sulfide CH3CH2SCH2CH3 rubbery 0.9 - 1.3
dimethyl sulfide CH3SCH3 canned corn, cooked cabbage, asparagus 17- 25
diethyl disulfide CH3CH2SSCH2CH3 garlic, burnt rubber 3.6 - 4.3
dimethyl disulfide CH3SSCH3 vegetal, cabbage, onion-like at high levels 9.8 - 10.2
carbon disulfide CS2 sweet, ethereal, slightly green, sulfidy 5

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