Sulfur [S]

Characteristics

An: 16 N: 16
Am: 32.065 g/mol
Group No: 16
Group Name: Chalcogen
Block: p-block Period: 3
State: solid at 298 K
Colour: lemon yellow Classification: Non-metallic
Boiling Point: 717.87K (444.72oC)
Melting Point: 388.36K (115.21oC)
Critical temperature: 1314K (1041oC)
Density: (alpha) 2.08g/cm3
Density: (beta) 1.96g/cm3
Density: (gamma) 1.92g/cm3

Discovery Information

Who: Known to the ancients. Homer mentioned "pest-averting sulfur" in the 8th century BC.

Name Origin

Latin: sulfur (brimstone). "Sulfur" in different languages.

Sources

Found in pure form (near hot springs and in volcanic regions) and in ores like cinnabar (HgS), galena (PbS), alunite, barite (BaSO4), sphalerite (ZnS) and stibnite (Sb2S3).
sulfur is a component of many of the complex organic compounds that are found in crude oil. It also occurs as hydrogen sulfide in natural gas. This piece of sulfur came from a natural gas plant near Tioga, North Dakota (USA).
Primary producers are the USA and Spain. Around 54 million tons are produced annually.

Abundance

Universe: 500 ppm (by weight)
Sun: 400 ppm (by weight)
Carbonaceous meteorite: 41000 ppm
Earth’s Crust: 420 ppm
Seawater: 928 ppm
Human: 2 x 106 ppb by weight; 3.9 x 105 ppb by atoms

Uses

Used in matches, gunpowder, detergents, fireworks, batteries, fungicides, vulcanization of rubber, medicines, permanent wave lotion and pesticides. Its most important use is probably that of sulfuric acid (H2SO4). Sulfites are used to bleach paper and as a preservative in wine and dried fruit. Sodium or ammonium thiosulfate are used as photographic fixing agents. Magnesium sulfate, better known as Epsom salts, can be used as a laxative, a bath additive, an exfoliant, or a magnesium supplement for plants.

History

Homer mentioned "pest-averting sulfur" in the 8th century BC and in 424 BC, the tribe of Boeotia destroyed the walls of a city by burning a mixture of coal, sulfur, and tar under them.
sulfur was known in China since the 6th century BC, in a natural form that the Chinese had called ’brimstone’, or shiliuhuang that was found in Hanzhong. By the 3rd century, the Chinese discovered that sulfur could be extracted from pyrite. Chinese Daoists were interested in sulfur’s flammability and its reactivity with certain metals, yet its earliest practical uses were found in traditional Chinese medicine. A Song Dynasty military treatise of 1044 AD described different formulas for Chinese gun powder, which is a mixture of potassium nitrate (KNO3), carbon, and sulfur. Early alchemists gave sulfur its own alchemical symbol which was a triangle at the top of a cross.
In the late 1770s, Antoine Lavoisier helped convince the scientific community that sulfur was an element and not a compound. In 1867, sulfur was discovered in underground deposits in Louisiana and Texas. The overlying layer of earth was quicksand, prohibiting ordinary mining operations. Therefore the Frasch process was utilized.

Notes

The distinctive colours of Jupiter’s volcanic moon, Io, are from various forms of molten, solid and gaseous sulfur. sulfur has also been found in many types of meteorite.
Hydrogen Sulfide (H2S), is well known for its smell of rotten eggs!

Hazards

Carbon disulfide, carbon oxysulfide, hydrogen sulfide, and sulfur dioxide (SO2) should all be handled with care.
Although sulfur dioxide is sufficiently safe to be used as a food additive in small amounts, at high concentrations it reacts with moisture to form sulfurous acid which in sufficient quantities may harm the lungs, eyes or other tissues. In creatures without lungs such as insects or plants, it otherwise prevents respiration.
Hydrogen sulfide is quite toxic (more toxic than cyanide). 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.

Sulfur Compounds

Sulfamic acid H3NO3S [Carcinogenic & Toxic]
The most famous applicaton of sulfamic acid is in the synthesis of compounds that taste sweet. It is used as an acidic cleaning agent, typically for metals and ceramics. It is a replacement for hydrochloric acid for the removal of rust.
It’s other uses include; catalyst for esterification process, dye and pigment manufacturing, herbicide, coagulator for urea-formaldehyde resins, ingredient in fire extinguishing media, in the pulp and paper industry as a chloride stabilizer and for the synthesis of nitrous oxide by reaction with nitric acid.

Sulfur dioxide SO2 [Toxic]
It is produced by volcanoes and in various industrial processes. In particular, poor-quality coal and petroleum contain sulfur compounds, and generate sulfur dioxide when burned: the gas reacts with water and atmospheric oxygen to form sulfuric acid (H2SO4) and thus acid rain.
Sulfur dioxide is sometimes used as a preservative in alcoholic drinks, or dried apricots and other dried fruits. The preservative is used to maintain the appearance of the fruit rather than prevent rotting. This can give fruit a distinctive chemical taste. Prior to the development of Freons, sulfur dioxide was used as a refrigerant in home refrigerators.

Sulfuric Acid H2SO4 [Highly Corrosive & Toxic]
Sulfuric acid has many applications, and is produced in greater amounts than any other chemical besides water. World production in 2001 was 165 million tonnes. Principal uses include ore processing, fertilizer manufacturing, oil refining, wastewater processing, and chemical synthesis.
Sulfuric acid is produced in the upper atmosphere of Venus by the sun’s photochemical action on carbon dioxide, sulfur dioxide, and water vapour. Ultraviolet photons of wavelengths less than 169 nm can photodissociate carbon dioxide into carbon monoxide and atomic oxygen. Atomic oxygen is highly reactive; when it reacts with sulfur dioxide, a trace component of the Venusian atmosphere, the result is sulfur trioxide, which can combine with water vapour, another trace component of Venus’ atmosphere, to yield sulfuric acid.

Reactions of Sulfur

Reactions with water
Sulfur will not react with water under standard conditions.
Reactions with air
Sulfur burns in air to form sulfur(IV) dioxide.
S8(s) + 8O2(g) --> 8SO2(g)
 
Reactions with halogens
sulfur reacts with fluorine and burns to form sulfur(VI) hexafluoride.
S8(s) + 24F2(g) --> 8SF6(l)
Molten sulfur will react with molten sulfur to form disulfur dichloride.
S8(s) + 4Cl2(g) --> 4S2Cl2(l)
With excess chlorine and in the presence of a catalyst it is possible to make a mixture containing and equilibrium of sulfur(II) chloride and disulfur dichloride.
S2Cl2(l) + Cl2(g) --> 2SCl2(l)
Reactions with acids
Sulphur does not react with dilute non-oxidizing acids.
Reactions with bases
Sulfur reacts with hot aqueous potassium hydroxide to form sulfide and thiosulfate species.
S8(s) + 6KOH(aq) --> 2K2S3 + K2S2O3 + 3H2O(l)

Occurrence of Sulfur

Elemental sulfur can be found near hot springs and volcanic regions in many parts of the world, especially along the Pacific Ring of Fire. Such volcanic deposits are currently mined in Indonesia, Chile, and Japan. Sicily is also famous for its sulfur mines.
Significant deposits of elemental sulfur also exist in salt domes along the coast of the Gulf of Mexico, and in evaporites in eastern Europe and western Asia. The sulfur in these deposits is believed to come from the action of anaerobic bacteria on sulfate minerals, especially gypsum, although apparently native sulfur may be produced by geological processes alone, without the aid of living organisms. However, fossil-based sulfur deposits from salt domes are the basis for commercial production in the United States, Poland, Russia, Turkmenistan, and Ukraine.
Sulfur production through hydrodesulfurization of oil, gas, and the Athabasca Oil Sands has produced a surplus - huge stockpiles of sulfur now exist throughout Alberta, Canada.
Common naturally occurring sulfur compounds include the sulfide minerals, such as pyrite (iron sulfide), cinnabar (mercury sulfide), galena (lead sulfide), sphalerite (zinc sulfide) and stibnite (antimony sulfide); and the sulfates, such as gypsum (calcium sulfate), alunite (potassium aluminium sulfate), and barite (barium sulfate). It occurs naturally in volcanic emissions, such as from hydrothermal vents, and from bacterial action on decaying sulfur-containing organic matter.
The distinctive colours of Jupiter’s volcanic moon, Io, are from various forms of molten, solid and gaseous sulfur. There is also a dark area near the Lunar crater Aristarchus that may be a sulfur deposit. Sulfur is also present in many types of meteorites.

Allotropes of Sulfur

sulfur (S) is second only to carbon in the number of known allotropes formed. The existence of at least twenty-two sulfur allotropes has been demonstrated.
Sulfur forms an extensive series of generally yellow crystalline allotropes, Sn (where species with n up to 30 have been identified). 

Orthorhombic Sulfur [ S8 ]
The most common form, stable at room temperature and atmospheric pressure. Eight sulfur atoms bond covalently in crownlike rings. 

Cyclohexasulfur or Rhombohedral Sulfur [ S6 ]
Cyclohexasulfur was first reported in 1891. It is the densest of the sulfur allotropes and forms air-sensitive orange-red crystals containing chair-shaped, six-membered rings. 

Amorphous Sulfur [ S ]
The result of very rapid cooling of very hot sulfur. 

Disulfur [ S2 ]
The simplest allotrope of sulfur, it is a violet colour. It does not occur naturally at room temperature and pressure. It is commonly generated in the vapour generated from sulfur at temperatures above 700oC.
It has been detected by the Hubble Space Telescope in volcanic eruptions on Jupiter’s satellite, Io.

Isotopes of Sulfur

32S [16 neutrons]
Abundance: 95.02%
Stable with 16 neutrons
33S [17 neutrons]
Abundance: 0.75%
Stable with 17 neutrons 

34S [18 neutrons]
Abundance: 4.21%
Stable with 18 neutrons 

35S [19 neutrons]
Abundance: synthetic
Half life: 87,32 days [ beta- ]
Decay Energy: 0.167 MeV
Decays to 35Cl. 

36S [20 neutrons]
Abundance: 0.02%
Stable with 20 neutrons

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