Sulfur

ស្ពាន់ធ័រគឺជាធាតុគីមី។ វាមាននិមិត្តសញ្ញា S និងលេខអាតូមិក 16។ វាមានច្រើនក្រៃលែង ពហុវ៉ាឡង់ និងមិនមែនលោហធាតុ។ នៅក្រោមលក្ខខណ្ឌធម្មតា អាតូមស្ពាន់ធ័របង្កើតជាម៉ូលេគុល octatomic cyclic ជាមួយនឹងរូបមន្តគីមី S8 ។ សារធាតុស្ពាន់ធ័រគឺជាសារធាតុរឹងគ្រីស្តាល់ពណ៌លឿងភ្លឺនៅសីតុណ្ហភាពបន្ទប់។

ស្ពាន់ធ័រគឺជាធាតុដ៏សម្បូរបែបបំផុតទី 10 ដោយម៉ាស់នៅក្នុងសកលលោក និងជាធាតុទី 5 ទូទៅបំផុតនៅលើផែនដី។ ទោះបីជាពេលខ្លះត្រូវបានរកឃើញនៅក្នុងទម្រង់ដើមសុទ្ធក៏ដោយ ស្ពាន់ធ័រនៅលើផែនដីជាធម្មតាកើតឡើងជាសារធាតុរ៉ែស៊ុលហ្វីត និងស៊ុលហ្វាត។ ដោយមានភាពសម្បូរបែបនៅក្នុងទម្រង់ដើម ស្ពាន់ធ័រត្រូវបានគេស្គាល់នៅសម័យបុរាណ ដែលត្រូវបានលើកឡើងសម្រាប់ការប្រើប្រាស់នៅក្នុងប្រទេសឥណ្ឌាបុរាណ ក្រិកបុរាណ ចិន និងអេហ្ស៊ីបបុរាណ។ តាមប្រវត្តិសាស្ត្រ និងក្នុងអក្សរសិល្ប៍ ស្ពាន់ធ័រត្រូវបានគេហៅថា brimstone [8] ដែលមានន័យថា "ថ្មឆេះ"។ ស្ពាន់ធ័រធាតុស្ទើរតែទាំងអស់ត្រូវបានផលិតជាអនុផលនៃការដកសារធាតុកខ្វក់ដែលមានផ្ទុកស្ពាន់ធ័រចេញពីឧស្ម័នធម្មជាតិ និងប្រេង។[10][11] ការប្រើប្រាស់ពាណិជ្ជកម្មដ៏អស្ចារ្យបំផុតនៃធាតុគឺការផលិតអាស៊ីតស៊ុលហ្វួរិកសម្រាប់ជីស៊ុលហ្វាត និងផូស្វាត និងដំណើរការគីមីផ្សេងទៀត។ ស្ពាន់ធ័រត្រូវបានគេប្រើនៅក្នុងការប្រកួត ថ្នាំសំលាប់សត្វល្អិត និងថ្នាំសម្លាប់ផ្សិត។ សមាសធាតុស្ពាន់ធ័រជាច្រើនមានក្លិនស្អុយ ហើយក្លិនឧស្ម័នធម្មជាតិ ក្លិនស្អុយ ក្លិនមាត់មិនល្អ ក្រូចថ្លុង និងខ្ទឹមសគឺដោយសារសមាសធាតុសរីរាង្គ។ អ៊ីដ្រូសែនស៊ុលហ្វីតផ្តល់ក្លិនលក្ខណៈដល់ស៊ុតរលួយ និងដំណើរការជីវសាស្រ្តផ្សេងទៀត។

ស្ពាន់ធ័រគឺជាធាតុសំខាន់សម្រាប់ជីវិតទាំងអស់ស្ទើរតែតែងតែមាននៅក្នុងទម្រង់នៃសមាសធាតុសរីរាង្គឬស៊ុលហ្វីតដែក។ អាស៊ីតអាមីណូ (សារធាតុប្រូតេអ៊ីនពីរគឺ cysteine ​​និង methionine និងជាច្រើនផ្សេងទៀតដែលមិនមានកូដ៖ cystine, taurine ។ cofactors ជាច្រើនក៏មានផ្ទុកស្ពាន់ធ័រ រួមទាំង glutathione និងប្រូតេអ៊ីនដែក-ស្ពាន់ធ័រ។ Disulfides, S-S bonds ផ្តល់កម្លាំងមេកានិច និងភាពមិនរលាយនៃប្រូតេអ៊ីន keratin (ក្នុងចំនោមផ្សេងទៀត) ដែលមាននៅក្នុងស្បែកខាងក្រៅ សក់ និងរោម។ ស្ពាន់ធ័រគឺជាធាតុគីមីសំខាន់មួយដែលត្រូវការសម្រាប់ដំណើរការជីវគីមី និងជាសារធាតុម៉ាក្រូសារធាតុចិញ្ចឹមសម្រាប់សារពាង្គកាយមានជីវិតទាំងអស់។

 Characteristics 

 Physical properties 
ធ្វេង|រូបភាពតូច|As a solid, sulfur is a characteristic lemon yellow; when burned, sulfur melts into a blood-red liquid and emits a blue flame.
Sulfur forms several polyatomic molecules. The best-known allotrope is octasulfur, cyclo-S8. The point group of cyclo-S8 is D4d and its dipole moment is 0 D. Octasulfur is a soft, bright-yellow solid that is odorless. It melts at , and boils at . At , below its melting temperature, cyclo-octasulfur begins slowly changing from α-octasulfur to the β-polymorph. The structure of the S8 ring is virtually unchanged by this phase transition, which affects the intermolecular interactions. Cooling molten sulfur freezes at , as it predominantly consists of the β-S8 molecules. 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. At higher temperatures, the viscosity decreases as depolymerization occurs. Molten sulfur assumes a dark red color above . 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  and .

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  clusters, followed by disproportionation into the reaction products.

The second, fourth and sixth ionization energies of sulfur are 2252 kJ/mol, 4556 kJ/mol and 8495.8 kJ/mol, respectively. The composition of reaction products of sulfur with oxidants (and its oxidation state) depends on whether releasing of reaction energy overcomes these thresholds. Applying catalysts and/or supply of external energy may vary sulfur's oxidation state and the composition of reaction products. While reaction between sulfur and oxygen under normal conditions gives sulfur dioxide (oxidation state +4), formation of sulfur trioxide (oxidation state +6) requires a temperature of  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 except noble gases, even with the notoriously unreactive metal iridium (yielding iridium disulfide). Some of those reactions require elevated temperatures.

 Allotropes 

ធ្វេង|រូបភាពតូច|The structure of the cyclooctasulfur molecule, S8
Sulfur forms over 30 solid allotropes, more than any other element. Besides S8, several other rings are known. 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. Larger rings have been prepared, including S12 and S18.

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 it 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 over a matter of hours to days, but can be rapidly catalyzed. 

 Isotopes 

Sulfur has 23 known isotopes, four of which are stable: 32S (), 33S (), 34S (), and 36S ().Sulfur. Commission on Isotopic Abundances and Atomic Weights 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.

On the planet Earth the sulfur isotopic composition was determined by the Sun. Though it was assumed that the distribution of different sulfur isotopes would be more or less equal, it has been found that proportions of the two most abundant sulfur isotopes 32S and 34S varies in different samples. Assaying of the isotope ratio (δ34S) in the samples suggests 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. However, there are ongoing discussions over the real reason for the δ34S shifts, biological activity or postdeposit alteration.

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 past oceans. Sulfate-reducing bacteria in marine sediment fractionate sulfur isotopes as they take in sulfate and produce sulfide. Prior to the 2010s, it was thought that sulfate reduction could fractionate sulfur isotopes up to 46 permil 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 showed that sulfate-reducing bacteria can fractionate to 66 permil. 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.

In 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 to verify the presence of recent (up to 1 year) atmospheric sediments in various materials. 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. 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.

 Natural occurrence 
រូបភាពតូច|Sulfur vat from which railroad cars are loaded, Freeport Sulphur Co., Hoskins Mound, Texas (1943)
រូបភាពតូច|Most of the yellow and orange hues of Io are due to elemental sulfur and sulfur compounds deposited by active volcanoes.
រូបភាពតូច|Sulfur extraction, East Java
រូបភាពតូច|A man carrying sulfur blocks from Kawah Ijen, a volcano in East Java, Indonesia, 2009
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. 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. The distinctive colors of Jupiter's volcanic moon Io are attributed to various forms of molten, solid, and gaseous sulfur. In July 2024, elemental sulfur was discovered to exist on Mars by surprise, after the Curiosity rover ran over and crushed a rock revealing sulfur crystals inside it.

Sulfur 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 mined in Indonesia, Chile, and Japan. These deposits are polycrystalline, with the largest documented single crystal measuring . Historically, Sicily was a major source of sulfur in the Industrial Revolution. Lakes of molten sulfur up to about  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.

Native sulfur is synthesized by anaerobic bacteria acting on sulfate minerals such as gypsum in salt domes. 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. Such sources have become of secondary commercial importance, and most are no longer worked but commercial production is still carried out in the Osiek mine in Poland.

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 has become petroleum and natural gas.