Sulfur is a relatively old element that has been on mankind's radar for thousands of years. It was referenced in ancient literature, such as Homer's Odyssey, and was even mentioned in the Bible. In fact, during ancient times, sulfur and carbon were the only nonmetals known to man. Recently, sulfur has been noted as a key element on the planet of Venus and in its atmosphere. So, how can sulfur truly be 'defined?'

Some might define sulfur as the second member of the oxygen family. Sulfur has six valence electrons and enjoys having a two minus formal charge. Sulfur appears yellow in color, and because it is a member of Group 6, sulfur serves as a relatively good oxidizing agent. On the periodic table, sulfur is characterized simply by the letter S as its elemental abbreviation.

All of this provides a good description of what sulfur is; however, to truly define the element, we might turn to a more traditional dictionary definition. There, sulfur can be defined as 'a nonmetallic that exists in several forms, which burns with a blue flame and suffocating odor.'

Natural Occurrence And Distribution

Many important metal ores are compounds of sulfur, either sulfides or sulfates. Some important examples are galena (lead sulfide, PbS), blende(zinc sulfide, ZnS), pyrite (iron disulfide, FeS2), chalcopyrite (copper ironsulfide, CuFeS2), gypsum (calcium sulfate dihydrate, CaSO4∙2H2O) and barite (barium sulfate, BaSO4). The sulfide ores are valued chiefly for their metal content, although a process developed in the 18th century for making sulfuric acid utilized sulfur dioxide obtained by burning pyrite. Coal, petroleum, and natural gas contain sulfur compounds.


In sulfur, allotropy arises from two sources: (1) the different modes of bonding atoms into a single molecule and (2) packing of polyatomic sulfur molecules into different crystalline and amorphous forms. Some 30 allotropic forms of sulfur have been reported, but some of these probably represent mixtures. Only eight of the 30 seem to be unique; five contain rings of sulfur atoms and the others contain chains.

In the rhombohedral allotrope, designated ρ-sulfur, the molecules are composed of rings of six sulfur atoms. This form is prepared by treating sodium thiosulfate with cold, concentrated hydrochloric acid, extracting the residue with toluene, and evaporating the solution to give hexagonal crystals. ρ-sulfur is unstable, eventually reverting to orthorhombic sulfur (α-sulfur).

A second general allotropic class of sulfur is that of the eight-membered ring molecules, three crystalline forms of which have been well characterized. One is the orthorhombic (often improperly called rhombic) form, α-sulfur. It is stable at temperatures below 96 °C. Another of the crystalline S8 ring allotropes is the monoclinic or β-form, in which two of the axes of the crystal are perpendicular, but the third forms an oblique angle with the first two. There are still some uncertainties concerning its structure; this modification is stable from 96 °C to the melting point, 118.9 °C. A second monoclinic cyclooctasulfur allotrope is the γ-form, unstable at all temperatures, quickly transforming to α-sulfur.

An orthorhombic modification, S12 ring molecules, and still another unstable S10 ring allotrope are reported. The latter reverts to polymeric sulfur and S8. At temperatures above 96 °C, the α-allotrope changes into the β-allotrope. If enough time is allowed for this transition to occur completely, further heating causes melting to occur at 118.9 °C; but if the α-form is heated so rapidly that the transformation to β-form does not have time to occur, the α-form melts at 112.8 °C.

Just above its melting point, sulfur is a yellow, transparent, mobile liquid. Upon further heating, the viscosity of the liquid decreases gradually to a minimum at about 157 °C, but then rapidly increases, reaching a maximum value at about 187 °C; between this temperature and the boiling point of 444.6 °C, the viscosity decreases. The colour also changes, deepening from yellow through dark red, and, finally, to black at about 250 °C. The variations in both colour and viscosity are considered to result from changes in the molecular structure. A decrease in viscosity as temperature increases is typical of liquids, but the increase in the viscosity of sulfur above 157 °C probably is caused by rupturing of the eight-membered rings of sulfur atoms to form reactive S8 units that join together in long chains containing many thousands of atoms. The liquid then assumes the high viscosity characteristic of such structures. At a sufficiently high temperature, all of the cyclic molecules are broken, and the length of the chains reaches a maximum. Beyond that temperature, the chains break down into small fragments. Upon vaporization, cyclic molecules (S8 and S6) are formed again; at about 900 °C, S2 is the predominant form; finally, monatomic sulfur is formed at temperatures above 1,800 °C.

Commercial Production

Elemental sulfur is found in volcanic regions as a deposit formed by the emission of hydrogen sulfide, followed by aerial oxidation to the element. Underground deposits of sulfur associated with salt domes in limestone rock provide a substantial portion of the world’s supply of the element. These domes are located in the Louisiana swamplands of the United States and offshore in the Gulf of Mexico.

Where deposits of sulfur are located in salt domes, as they are along the coast of the Gulf of Mexico, the element was recovered by the Frasch process, named after German-born U.S. chemist Herman Frasch. Ordinary underground mining procedures were inapplicable since highly poisonous hydrogen sulfide gas accompanies the element in the domes. Beginning in 1894, the Frasch process, which takes advantage of the low melting point of sulfur (112 °C), made sulfur of a high purity (up to 99.9 percent pure) available in large quantities and helped establish sulfur as an important basic chemical commodity. Wells were drilled from 60 to 600 m (200 to 2,000 feet) into the sulfur formation and then lined with a 15-cm (6-inch) pipe in which an air pipe and a water pipe of smaller diameter were concentrically placed. Superheated water, injected into the circular space between the three- and six-inch pipes, penetrated the cap rock through holes on the bottom of the pipe. As the sulfur melted, it settled to the bottom of the deposit. From there it was pumped to the surface by applying air pressure through the central pipe. Several such wells operated under the ocean floor in the Gulf of Mexico. The sulfur was collected in reservoirs, or sumps, and from there transferred to vats or bins to solidify for storage and stockpiling. Vats contained as much as 300,000 tons of sulfur. Frasch-process sulfur produced at the Gulf Coast salt domes constituted the major source of U.S. sulfur production and dominated the world market until approximately 1970. Thereafter, non-Frasch sources such as the purification of sour (high sulfur-content) petroleum, the refining of natural gas, and improved methods for obtaining sulfur from metal sulfides gained a greater share of the market. The Frasch process is still used today in Poland and Russia.


About 9,000,000 tons of sulfur are recovered in the United States each year from natural gas, petroleum refinery gases, pyrites, and smelter gases from the processing of copper, zinc, and lead ores. In most cases sulfur is separated from other gases as hydrogen sulfide and then converted to elemental sulfur by the Claus process, which involves the partial burning of hydrogen sulfide to sulfur dioxide, with subsequent reaction between the two to yield sulfur. Another important source is the sulfur dioxide emitted into the atmosphere by coal-fired steam power plants. In the early 1970s techniques to collect this sulfur dioxide and convert it into usable sulfur were developed.

A few of the non-Frasch processes for sulfur production may be mentioned.

  1. Sulfur-bearing rock is piled into mounds. Shafts are bored vertically and fires set at the top of the shafts. The burning sulfur provides sufficient heat to melt the elemental sulfur in the rock layers below, and it flows out at the bottom of the pile. This is an old process, still used to some extent in Sicily. The product is of low purity and must be refined by distillation. The air pollution in the area of the process is so great that its operation is limited to certain times of the year when prevailing winds will carry the fumes away from populated areas.
  2. Rock bearing sulfur is treated with superheated water in retorts, melting the sulfur, which flows out. This process is a modification of the Frasch method.
  3. Sulfates (such as gypsum or barite) may be treated with carbon at high temperatures, forming the metal sulfides CaS or BaS (the Chance-Claus process). The metal sulfides can be treated with acid, generating hydrogen sulfide, which in turn can be burned to give elemental sulfur.
  4. Tremendous tonnages of sulfur are available from smelter operations and from power production by combustion of fossil and sour petroleum fuels, some of which contain as much as 4 percent sulfur. Thus, generation of electrical power and heat represent a major source of atmospheric pollution by sulfur dioxide. Unfortunately, recovery and purification of sulfur dioxide from stack gases are expensive operations.


Wherever such metals as lead, zinc, copper, cadmium, or nickel (among others) are processed, much of the sulfuric acid needed in the metallurgical operations may be obtained on the site by converting sulfur dioxide, produced by roasting the ores, to sulfur trioxide, SO3, and thence to sulfuric acid.

Sulfur available in bulk from commercial production usually is more than 99 percent pure, and some grades contain 99.9 percent sulfur. For research purposes, the proportion of impurities has been reduced to as little as one part in 10,000,000 by the application of procedures such as zone melting, column chromatography, electrolysis, or fractional distillation. China, Canada, Germany and Japan led the world in sulfur production in the early 21st century.

Uses Of Sulfur

Sulfur is so widely used in industrial processes that its consumptionoften is regarded as a reliable indicator of industrial activity and the state of the national economy. Approximately six-sevenths of all the sulfur produced is converted into sulfuric acid, for which the largest single use is in the manufacture of fertilizers (phosphates and ammonium sulfate). Other important uses include the production of pigments, detergents, fibres, petroleum products, sheet metal, explosives, and storage batteries; hundreds of other applications are known. Sulfur not converted to sulfuric acid is used in making paper, insecticides, fungicides, dyestuffs, and numerous other products.


Sulfur forms compounds in oxidation states −2 (sulfide, S2−), +4 (sulfite, SO32−), and +6 (sulfate, SO42−). It combines with nearly all elements. An unusual feature of some sulfur compounds results from the fact that sulfur is second only to carbon in exhibiting catenation—i.e., the bonding of an atom to another identical atom. This allows sulfur atoms to form ring systems and chain structures. The more significant sulfur compounds and compound groups are as follows.

One of the most familiar sulfur compounds is hydrogen sulfide, also known as sulfureted hydrogen, or stinkdamp, H2S, the colourless, extremely poisonous gas responsible for the characteristic odour of rotten eggs. It is produced naturally by the decay of organic substances containing sulfur and is often present in vapours from volcanoes and mineral waters. Large amounts of hydrogen sulfide are obtained in the removal of sulfur from petroleum. It was formerly used extensively in chemical laboratories as an analytical reagent.

All the metals except gold and platinum combine with sulfur to form inorganic sulfides. Such sulfides are ionic compounds containing the negatively charged sulfide ion S2−; these compounds may be considered as salts of hydrogen sulfide. Some inorganic sulfides are important ores of such metals as iron, nickel, copper, cobalt, zinc, and lead.

Several oxides are formed by sulfur and oxygen; the most important is the heavy, colourless, poisonous gas sulfur dioxide, SO2. It is used primarily as a precursor of sulfur trioxide, SO3, and thence sulfuric acid, H2SO4. It is also utilized as a bleach and an industrial reducing agent. Other noteworthy applications include its use in food preservation and for fruit ripening. (See also sulfur oxide.)

Sulfur forms a wide variety of compounds with halogen elements. In combination with chlorine it yields sulfur chlorides such as disulfur dichloride, S2Cl2, a corrosive, golden-yellow liquid used in the manufacture of chemical products. It reacts with ethylene to produce mustard gas, and with unsaturated acids derived from fats it forms oily products that are basic components of lubricants. With fluorine, sulfur forms sulfur fluorides, the most useful of which is sulfur hexafluoride, SF6, a gas employed as an insulator in various electrical devices. Sulfur also forms oxyhalides, in which the sulfur atom is bonded to both oxygen and halogen atoms. When such compounds are named, the term thionyl is used to designate those containing the SO unit and the term sulfuryl for those with SO2. Thionyl chloride, SOCl2, is a dense, toxic, volatile liquid used in organic chemistry to convert carboxylic acids and alcohols into chlorine-containing compounds. Sulfuryl chloride, SO2Cl2, is a liquid of similar physical properties utilized in the preparation of certain compounds that contain sulfur, chlorine, or both.

Sulfur forms some 16 oxygen-bearing acids. Only four or five of them, however, have been prepared in the pure state. These acids, particularly sulfurous acid and sulfuric acid, are of considerable importance to the chemical industry. Sulfurous acid, H2SO3, is produced when sulfur dioxide is added to water. Its most important salt is sodium sulfite, Na2SO3, a reducing agent employed in the manufacture of paper pulp, in photography, and in the removal of oxygen from boiler feedwater. Sulfuric acid is one of the most valuable of all chemicals. Prepared commercially by the reaction of water with sulfur trioxide, the compound is used in manufacturing fertilizers, pigments, dyes, drugs, explosives, detergents, and inorganic salts and esters.

The organic compounds of sulfur constitute a diverse and important subdivision of organic substances. Some examples include the sulfur-containing amino acids (e.g., cysteine, methionine, and taurine), which are key components of hormones, enzymes, and coenzymes. Significant, too, are the synthetic organic sulfur compounds, among them numerous pharmaceuticals (sulfa drugs, dermatological agents), insecticides, solvents, and agents such as those used in preparing rubber and rayon.

Acidified KMnO4 can be decolorized by SO2.Potassium permanganate(KMnO4) has a purple colour. Sulphur dioxide is a reducing agent. Potassium permanganate is an oxidizing agent. When sulphur dioxide reacts with potassium permanganate the solution decolourizes. colour changes from Purple to transparent(Colourless).

What is sulfur dioxide?

Sulfur dioxide (SO2) is a colorless, reactive air pollutant with a strong odor. This gas can be a threat to human health, animal health, and plant life.

The main sources of sulfur dioxide emissions are from fossil fuel combustion and natural volcanic activity. Hawai'i Volcanoes National Park (NP) is unique in the national park system because it sometimes has extremely high concentrations of sulfur dioxide — far higher than any other national park, or even most urban areas.

How can sulfur dioxide affect your health?

Sulfur dioxide irritates the skin and mucous membranes of the eyes, nose, throat, and lungs. High concentrations of SO2 can cause inflammation and irritation of the respiratory system, especially during heavy physical activity. The resulting symptoms can include pain when taking a deep breath, coughing, throat irritation, and breathing difficulties. High concentrations of SO2 can affect lung function, worsen asthma attacks, and worsen existing heart disease in sensitive groups. This gas can also react with other chemicals in the air and change to a small particle that can get into the lungs and cause similar health effects.

Who is at risk?

People sensitive to sulfur dioxide include:

  • People with lung diseases, such as asthma, chronic bronchitis, and emphysema will generally have more serious health effects at higher SO2 levels.
  • Children are at higher risk from SO2 exposure because their lungs are still developing. They are also more likely to have asthma, which can get worse with SO2 exposure.
  • Older adults may be more affected by SO2 exposure, possibly because they are more likely to have pre-existing lung or cardiovascular disease.
  • Active people of all ages who exercise or work outdoors have higher exposure to sulfur dioxide than people who are less active.

Hawai'i Volcanoes NP visitors, residents, and park staff downwind of the volcanic SO2 emissions can be exposed to unhealthy levels of pollution. Since it is not possible to control volcanic activity, the National Park Service created a sulfur dioxide advisory program, which gives out warnings to let people know when unhealthy levels of this pollutant are present. Advisories encourage people to limit their exposure when necessary.

How can I avoid unhealthy exposure?

You can take simple steps to reduce your exposure to unhealthy air. First, visit the Current Conditions Website to find out about current sulfur dioxide conditions and the health advisory level.

When possibly unhealthy sulfur dioxide pollution happens, your chances of being affected increase with high levels of activity and the length of time you are active outdoors. If your planned activity has long or heavy physical exertion and the sulfur dioxide levels are high, you may want to limit or stop your activity. For recommended ways to protect yourself at high levels of sulfur dioxide, consult the Health Advisory Table.

What are the NPS sulfur dioxide health advisories?

A SO2 air pollution advisory program was created at Hawai'i Volcanoes NP to deliver timely information about possible unhealthy air pollution conditions that could affect the health of visitors, island residents, and park personnel. Using the Environmental Protection Agency (EPA) air quality index, the NPS SO2 health advisories for Hawai'i Volcanoes NP help you understand what local air quality means to your health. The air quality index is divided into six levels of health concern: 

Scale describing the health advisory levels for sulfur dioxide (SO2), ranging from good (green) to hazardous (maroon)

Understanding Sulfur Dioxide Health Advisory Levels

  • Good (0–0.1 ppm)
  • No cautionary statement.
  • Moderate (0.1–0.2 ppm)
  • Unusually sensitive people should consider reducing prolonged or heavy exertion outdoors.
  • Unhealthy for Sensitive Groups (0.2–1.0 ppm)
  • Active children and adults, and people with lung disease, such as asthma, should reduce prolonged or heavy exertion outdoors.
  • Unhealthy (1.0–3.0 ppm)
  • Active children and adults, and people with lung disease, such as asthma, should avoid prolonged or heavy exertion outdoors. Everyone else, especially children, should reduce prolonged or heavy exertion outdoors.
  • Very Unhealthy (3.0–5.0 ppm)
  • Active children and adults, and people with lung disease, such as asthma, should avoid all outdoor exertion. Everyone else, especially children, should avoid prolonged or heavy exertion outdoors.
  • Hazardous ( > 5.0 ppm) 
  • Triggers health warnings of emergency conditions. Entire population is more likely to be affected. Avoid outdoor activities & remain indoors. Leave the area if directed by Civil Defense.

The SO2 and weather data used in this program are collected by the National Park Service at the Jaggar Museum and Kilauea Visitor Center monitoring sites. The SO2 concentrations measured at the monitoring sites are reviewed every 15 minutes and one of six advisory levels of health concern are assigned for that 15-minute period for each site.

How does sulfur dioxide affect national parks?

Hawai'i Volcanoes NP is significantly impacted by sulfur dioxide because the high levels create a human health concern. Sulfate particles can also create haze and reduce visibility at Hawai'i Volcanoes NP and other national parks. Sulfur dioxide can convert to acids in the atmosphere and come down from the atmosphere in rain, snow, or fog, or as dry particles. This atmospheric depositioncan damage vegetation, affect soils, acidify lakes and streams, and ruin memorials, buildings, and statues at our national cultural monuments.


• How to define a sutli bomb so the ban stands legal scrutiny?

• How to stop crackers from entering the metropolis?

• What should be done with existing stocks?

• If caught with sutli bomb, who to punish — child or parent?

• Would it be all right to visit houses or carry out arrests during the festival?

• Will a police force stretched with security duties be able to attend to all the complaints?

Sulfur dioxide air quality index hazardous condition icon (maroon)
Sulfur dioxide air quality index very-unhealthy condition icon (purple)
Sulfur dioxide air quality index unhealthy condition icon (red)
Sulfur dioxide air quality index unhealthy-for-sensitive-groups condition icon (orange)
Sulfur dioxide air quality index moderate condition icon (yellow)