Sulfur, ₁₆S

Sulfur
Alternative name	Sulphur (British spelling)
Allotropes	see Allotropes of sulfur
Appearance	Lemon yellow sintered microcrystals
Standard atomic weight Ar°(S)
	[1] 32.06±0.02 (abridged)[2]
Sulfur in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson O ↑ S ↓ Se phosphorus ← sulfur → chlorine
Atomic number (Z)	16
Group	group 16 (chalcogens)
Period	period 3
Block	p-block
Electron configuration	[Ne] 3s2 3p4
Electrons per shell	2, 8, 6
Physical properties
Phase at STP	solid
Melting point	alpha (α-S8): 388.36 K ​(115.21 °C, ​239.38 °F)
Boiling point	717.8 K ​(444.6 °C, ​832.3 °F)
Density (near r.t.)	alpha (α-S8): 2.07 g/cm3 beta (β-S8): 1.96 g/cm3 gamma (γ-S8): 1.92 g/cm3
when liquid (at m.p.)	1.819 g/cm3
Critical point	1314 K, 20.7 MPa
Heat of fusion	beta (β-S8): 1.727 kJ/mol
Heat of vaporization	beta (β-S8): 45 kJ/mol
Molar heat capacity	22.75 J/(mol·K)
Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 375 408 449 508 591 717
Atomic properties
Oxidation states	−2, −1, 0, +1, +2, +3, +4, +5, +6 (a strongly acidic oxide)
Electronegativity	Pauling scale: 2.58
Ionization energies	1st: 999.6 kJ/mol 2nd: 2252 kJ/mol 3rd: 3357 kJ/mol (more)
Covalent radius	105±3 pm
Van der Waals radius	180 pm
Spectral lines of sulfur
Other properties
Natural occurrence	primordial
Crystal structure	alpha (α-S8): ​orthorhombic (oF128)
Lattice constants	a = 1.0460 nm b = 1.2861 nm c = 2.4481 nm (at 20 °C)[3]
Crystal structure	beta (β-S8): ​monoclinic (mP48)
Lattice constants	a = 1.0923 nm b = 1.0851 nm c = 1.0787 nm β = 95.905° (at 20 °C)[3]
Thermal conductivity	0.205 W/(m⋅K) (amorphous)
Electrical resistivity	2×1015  Ω⋅m (at 20 °C) (amorphous)
Magnetic ordering	diamagnetic[4]
Molar magnetic susceptibility	alpha (α-S8): −15.5×10−6 cm3/mol (298 K)[5]
Bulk modulus	7.7 GPa
Mohs hardness	2.0
CAS Number	7704-34-9
History
Discovery	before 2000 BCE[6]
Recognized as an element by	Antoine Lavoisier (1777)
Isotopes of sulfur v e
Main isotopes Decay abun­dance half-life (t1/2) mode pro­duct 32S 94.8% stable 33S 0.760% stable 34S 4.37% stable 35S trace 87.37 d β− 35Cl 36S 0.02% stable 34S abundances vary greatly (between 3.96 and 4.77 percent) in natural samples.
Category: Sulfur view talk edit | references

Sulfur (also spelled sulphur in British English) is a chemical element; it has symbol S and atomic number 16. It is abundant, multivalent and nonmetallic. Under normal conditions, sulfur atoms form cyclic octatomic molecules with the chemical formula S₈. Elemental sulfur is a bright yellow, crystalline solid at room temperature.

Sulfur is the tenth most abundant element by mass in the universe and the fifth most abundant on Earth. Though sometimes found in pure, native form, sulfur on Earth usually occurs as sulfide and sulfate minerals. Being abundant in native form, sulfur was known in ancient times, being mentioned for its uses in ancient India, ancient Greece, China, and ancient Egypt. Historically and in literature sulfur is also called brimstone,^[7] which means "burning stone".^[8] Today, almost all elemental sulfur is produced as a byproduct of removing sulfur-containing contaminants from natural gas and petroleum.^[9][10] The greatest commercial use of the element is the production of sulfuric acid for sulfate and phosphate fertilizers, and other chemical processes. Sulfur is used in matches, insecticides, and fungicides. Many sulfur compounds are odoriferous, and the smells of odorized natural gas, skunk scent, bad breath, grapefruit, and garlic are due to organosulfur compounds. Hydrogen sulfide gives the characteristic odor to rotting eggs and other biological processes.

Sulfur is an essential element for all life, almost always in the form of organosulfur compounds or metal sulfides. Amino acids (two proteinogenic: cysteine and methionine, and many other non-coded: cystine, taurine, etc.) and two vitamins (biotin and thiamine) are organosulfur compounds crucial for life. Many cofactors also contain sulfur, including glutathione, and iron–sulfur proteins. Disulfides, S–S bonds, confer mechanical strength and insolubility of the (among others) protein keratin, found in outer skin, hair, and feathers. Sulfur is one of the core chemical elements needed for biochemical functioning and is an elemental macronutrient for all living organisms.

Characteristics

Physical properties

Sulfur forms several polyatomic molecules. The best-known allotrope is octasulfur, cyclo-S₈. The point group of cyclo-S₈ is D_4d and its dipole moment is 0 D.^[11] Octasulfur is a soft, bright-yellow solid that is odorless.^[a] It melts at 115.21 °C (239.38 °F),^[b] boils at 444.6 °C (832.3 °F).^[7] At 95.2 °C (203.4 °F), below its melting temperature, cyclo-octasulfur begins slow changing from α-octasulfur to the β-polymorph.^[13] The structure of the S₈ ring is virtually unchanged by this phase change, which affects the intermolecular interactions. Cooling of molten sulfur gives freezing point in 119.6 °C (247.3 °F),^[14] as it predominantly consists of the β-S₈ molecules.^[c] Between its melting and boiling temperatures, octasulfur changes its allotrope again, turning from β-octasulfur to γ-sulfur, again accompanied by a lower density but increased viscosity due to the formation of polymers.^[13] At higher temperatures, the viscosity decreases as depolymerization occurs. Molten sulfur assumes a dark red color above 200 °C (392 °F). The density of sulfur is about 2 g/cm³, depending on the allotrope; all of the stable allotropes are excellent electrical insulators.

Sulfur sublimes more or less between 20 °C (68 °F) and 50 °C (122 °F).^[18]

Sulfur is insoluble in water but soluble in carbon disulfide and, to a lesser extent, in other nonpolar organic solvents, such as benzene and toluene.

Chemical properties

Under normal conditions, sulfur hydrolyzes very slowly to mainly form hydrogen sulfide and sulfuric acid:

The reaction involves adsorption of protons onto S
₈ clusters, followed by disproportionation into the reaction products.^[19]

The second, fourth and sixth ionization energies of sulfur are 2252 kJ/mol, 4556 kJ/mol and 8495.8 kJ/mol, respectively. A composition of products of sulfur's reactions with oxidants (and its oxidation state) depends on that whether releasing out of a reaction energy overcomes these thresholds. Applying catalysts and / or supply of outer energy may vary sulfur's oxidation state and a composition of reaction products. While reaction between sulfur and oxygen at normal conditions gives sulfur dioxide (oxidation state +4), formation of sulfur trioxide (oxidation state +6) requires temperature 400–600 °C (750–1,100 °F) and presence of a catalyst.

In reactions with elements of lesser electronegativity, it reacts as an oxidant and forms sulfides, where it has oxidation state −2.

Sulfur reacts with nearly all other elements with the exception of the noble gases, even with the notoriously unreactive metal iridium (yielding iridium disulfide).^[20] Some of those reactions need elevated temperatures.^[21]

Allotropes

Sulfur forms over 30 solid allotropes, more than any other element.^[22] Besides S₈, several other rings are known.^[23] Removing one atom from the crown gives S₇, which is of a deeper yellow than S₈. HPLC analysis of "elemental sulfur" reveals an equilibrium mixture of mainly S₈, but with S₇ and small amounts of S₆.^[24] Larger rings have been prepared, including S₁₂ and S₁₈.^[25][26]

Amorphous or "plastic" sulfur is produced by rapid cooling of molten sulfur—for example, by pouring it into cold water. X-ray crystallography studies show that the amorphous form may have a helical structure with eight atoms per turn. The long coiled polymeric molecules make the brownish substance elastic, and in bulk this form has the feel of crude rubber. This form is metastable at room temperature and gradually reverts to the crystalline molecular allotrope, which is no longer elastic. This process happens within a matter of hours to days, but can be rapidly catalyzed.

Isotopes

Sulfur has 23 known isotopes, four of which are stable: ³²S (94.99%±0.26%), ³³S (0.75%±0.02%), ³⁴S (4.25%±0.24%), and ³⁶S (0.01%±0.01%).^[27][28] Other than ³⁵S, with a half-life of 87 days, the radioactive isotopes of sulfur have half-lives less than 3 hours.

The preponderance of ³²S is explained by its production in the so-called alpha-process (one of the main classes of nuclear fusion reactions) in exploding stars. Other stable sulfur isotopes are produced in the bypass processes related with ³⁴Ar, and their composition depends on a type of a stellar explosion. For example, proportionally more ³³S comes from novae than from supernovae.^[29]

On the planet Earth the sulfur isotopic composition was determined by the Sun. Though it is assumed that the distribution of different sulfur isotopes should be more or less equal, it has been found that proportions of two most abundant sulfur isotopes ³²S and ³⁴S varies in different samples. Assaying of these isotopes ratio (δ³⁴S) in the samples allows to make suggestions about their chemical history, and with support of other methods, it allows to age-date the samples, estimate temperature of equilibrium between ore and water, determine pH and oxygen fugacity, identify the activity of sulfate-reducing bacteria in the time of formation of the sample, or suggest the main sources of sulfur in ecosystems.^[30] However, there are ongoing discussions about what is the real reason of the δ³⁴S shifts, biological activity or postdeposital alteration.^[31]

For example, when sulfide minerals are precipitated, isotopic equilibration among solids and liquid may cause small differences in the δ³⁴S values of co-genetic minerals. The differences between minerals can be used to estimate the temperature of equilibration. The δ¹³C and δ³⁴S of coexisting carbonate minerals and sulfides can be used to determine the pH and oxygen fugacity of the ore-bearing fluid during ore formation.

Scientists measure the sulfur isotopes of minerals in rocks and sediments to study the redox conditions in the oceans in the past. Sulfate-reducing bacteria in marine sediment fractionate sulfur isotopes as they take in sulfate and produce sulfide. Prior to 2010s, it was thought that sulfate reduction could fractionate sulfur isotopes up to 46 permil^[32] and fractionation larger than 46 permil recorded in sediments must be due to disproportionation of sulfur compounds in the sediment. This view has changed since the 2010s as experiments show that sulfate-reducing bacteria can fractionate to 66 permil.^[33] As substrates for disproportionation are limited by the product of sulfate reduction, the isotopic effect of disproportionation should be less than 16 permil in most sedimentary settings.^[34]

In most forest ecosystems, sulfate is derived mostly from the atmosphere; weathering of ore minerals and evaporites contribute some sulfur. Sulfur with a distinctive isotopic composition has been used to identify pollution sources, and enriched sulfur has been added as a tracer in hydrologic studies. Differences in the natural abundances can be used in systems where there is sufficient variation in the ³⁴S of ecosystem components. Rocky Mountain lakes thought to be dominated by atmospheric sources of sulfate have been found to have measurably different ³⁴S values than lakes believed to be dominated by watershed sources of sulfate.

The radioactive ³⁵S is formed in cosmic ray spallation of the atmospheric ⁴⁰Ar. This fact may be used for proving the presence of recent (not more than 1 year) atmospheric sediments in various things. This isotope may be obtained artificially by different ways. In practice, the reaction ³⁵Cl + n → ³⁵S + p is used by irradiating potassium chloride with neutrons.^[35] The isotope ³⁵S is used in various sulfur-containing compounds as a radioactive tracer for many biological studies, for example, the Hershey-Chase experiment.

Because of the weak beta activity of ³⁵S, its compounds are relatively safe as long as they are not ingested or absorbed by the body.^[36]

Natural occurrence

³²S is created inside massive stars, at a depth where the temperature exceeds 2.5×10⁹ K, by the fusion of one nucleus of silicon plus one nucleus of helium.^[37] As this nuclear reaction is part of the alpha process that produces elements in abundance, sulfur is the 10th most common element in the universe.

Sulfur, usually as sulfide, is present in many types of meteorites. Ordinary chondrites contain on average 2.1% sulfur, and carbonaceous chondrites may contain as much as 6.6%. It is normally present as troilite (FeS), but there are exceptions, with carbonaceous chondrites containing free sulfur, sulfates and other sulfur compounds.^[38] The distinctive colors of Jupiter's volcanic moon Io are attributed to various forms of molten, solid, and gaseous sulfur.^[39]

It is the fifth most common element by mass in the Earth. 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. These deposits are polycrystalline, with the largest documented single crystal measuring 22×16×11 cm.^[40] Historically, Sicily was a major source of sulfur in the Industrial Revolution.^[41] Lakes of molten sulfur up to about 200 m (660 ft) in diameter have been found on the sea floor, associated with submarine volcanoes, at depths where the boiling point of water is higher than the melting point of sulfur.^[42]

Native sulfur is synthesised by anaerobic bacteria acting on sulfate minerals such as gypsum in salt domes.^[43][44] Significant deposits in salt domes occur along the coast of the Gulf of Mexico, and in evaporites in eastern Europe and western Asia. Native sulfur may be produced by geological processes alone. Fossil-based sulfur deposits from salt domes were once the basis for commercial production in the United States, Russia, Turkmenistan, and Ukraine.^[45] Currently, commercial production is still carried out in the Osiek mine in Poland. Such sources are now of secondary commercial importance, and most are no longer worked.

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 sulfate minerals, such as gypsum (calcium sulfate), alunite (potassium aluminium sulfate), and barite (barium sulfate). On Earth, just as upon Jupiter's moon Io, elemental sulfur occurs naturally in volcanic emissions, including emissions from hydrothermal vents.

The main industrial source of sulfur is now petroleum and natural gas.^[9]

Compounds

Common oxidation states of sulfur range from −2 to +6. Sulfur forms stable compounds with all elements except the noble gases.

Electron transfer reactions

Sulfur polycations, S₈²⁺, S₄²⁺ and S₁₆²⁺ are produced when sulfur is reacted with oxidising agents in a strongly acidic solution.^[46] The colored solutions produced by dissolving sulfur in oleum were first reported as early as 1804 by C.F. Bucholz, but the cause of the color and the structure of the polycations involved was only determined in the late 1960s. S₈²⁺ is deep blue, S₄²⁺ is yellow and S₁₆²⁺ is red.^[13]

Reduction of sulfur gives various polysulfides with the formula S_x²⁻, many of which have been obtained in crystalline form. Illustrative is the production of sodium tetrasulfide:

Some of these dianions dissociate to give radical anions, such as S₃⁻ gives the blue color of the rock lapis lazuli.

This reaction highlights a distinctive property of sulfur: its ability to catenate (bind to itself by formation of chains). Protonation of these polysulfide anions produces the polysulfanes, H₂S_x, where x = 2, 3, and 4.^[48] Ultimately, reduction of sulfur produces sulfide salts:

The interconversion of these species is exploited in the sodium–sulfur battery.

Hydrogenation

Treatment of sulfur with hydrogen gives hydrogen sulfide. When dissolved in water, hydrogen sulfide is mildly acidic:^[7]

Hydrogen sulfide gas and the hydrosulfide anion are extremely toxic to mammals, due to their inhibition of the oxygen-carrying capacity of hemoglobin and certain cytochromes in a manner analogous to cyanide and azide (see below, under precautions).

Combustion

The two principal sulfur oxides are obtained by burning sulfur:

Many other sulfur oxides are observed including the sulfur-rich oxides include sulfur monoxide, disulfur monoxide, disulfur dioxides, and higher oxides containing peroxo groups.

Halogenation

Sulfur reacts with fluorine to give the highly reactive sulfur tetrafluoride and the highly inert sulfur hexafluoride.^[49] Whereas fluorine gives S(IV) and S(VI) compounds, chlorine gives S(II) and S(I) derivatives. Thus, sulfur dichloride, disulfur dichloride, and higher chlorosulfanes arise from the chlorination of sulfur. Sulfuryl chloride and chlorosulfuric acid are derivatives of sulfuric acid; thionyl chloride (SOCl₂) is a common reagent in organic synthesis.^[50] Bromine also oxidizes sulfur to form sulfur dibromide and disulfur dibromide.^[50]

Pseudohalides

Sulfur oxidizes cyanide and sulfite to give thiocyanate and thiosulfate, respectively. Zdroj:https://en.wikipedia.org?pojem=Sulfur
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