سلفر (گندرف)

**گندرف** (Sulphur) ھڪ ڪيميائي عنصر آھي. ھن جو نشان "S" ۽ ايٽمي نمبر "16" آھي. اهو گهڻو (abundant)، ملٽي ويلنٽ ۽ غير ڌاتو آهي. عام حالتن ۾، گندرف جا ايٽم ڪيميائي فارمولا S8 سان چڪري آڪٽا ائٽامڪ (cyclic octoatomic) ماليڪيول ٺاهيندا آهن. عنصري گندرف ڪمري جي حرارت تي هڪ روشن پيلو ۽ ڪرسٽل وارو ٺوس آهي.

گندرف ڪائنات ۾ وزن جي لحاظ کان ڏهين نمبر تي سڀ کان وڌيڪ ۽ زمين تي پنجين نمبر تي سڀ کان وڌيڪ عام عنصر آهي. جيتوڻيڪ ڪڏهن ڪڏهن خالص، اصلي شڪل ۾ ملي ٿو، پر زمين تي گندرف عام طور تي سلفائيڊ ۽ سلفيٽ معدنيات طور ملي ٿو. اصلي شڪل ۾ وافر هجڻ ڪري، سلفر قديم زماني ۾ سڃاتو ويندو هو، قديم هندستان، يونان، چين ۽ مصر ۾ ان جي استعمال لاءِ ذڪر ڪيو ويندو هو. تاريخي طور تي سلفر کي برم اسٽون (brimestone) به سڏيو ويندو آهي، جنهن جو مطلب آهي "جلندڙ پٿر". تقريبن سڀئي عنصري سلفر قدرتي گئس ۽ پيٽروليم مان سلفر تي مشتمل آلودگي کي هٽائڻ جي ضمني پيداوار جي طور تي پيدا ڪيا ويندا آهن. عنصر جو سڀ کان وڏو تجارتي استعمال سلفيٽ ۽ فاسفيٽ ڀاڻ ۽ ٻين ڪيميائي عملن لاءِ سلفرڪ ايسڊ جي پيداوار آهي. سلفر کي ماچس، حشرات مار دوائن ۽ فنگسائيڊ ۾ استعمال ڪيو ويندو آهي. ڪيترائي سلفر مرڪب بدبودار آهن ۽ قدرتي گئس، اسڪنڪ  بدبوءِ، انگور ۽ لہسن جي بوءِ آرگنو سلفر مرڪب جي ڪري آهي. هائيڊروجن سلفائيڊ سڙيل انڊن ۽ ٻين حياتياتي عملن کي خاصيت واري بو ڏئي ٿو.

سلفر سڀني زندگين لاءِ هڪ ضروري عنصر آهي، تقريبن هميشه آرگنو سلفر مرڪب يا ڌاتو سلفائيڊ جي صورت ۾. امينو ايسڊ (ٻه پروٽينوجينڪ: سيسٽين ۽ ميٿيونين، ۽ ٻيا ڪيترائي غير ڪوڊ ٿيل: سيسٽين، تورين، وغيره) ۽ ٻه وٽامن (بايوٽين ۽ ٿاامين) زندگي لاءِ اهم آرگنوسلفر مرڪب آهن. ڪيترن ئي ڪوفيڪٽرز ۾ سلفر پڻ شامل آهي، جنهن ۾ گلوٽاٿيون، ۽ آئرن-سلفر پروٽين شامل آهن. ڊِسلفائيڊس، ايس-ايس بانڊ، ٻاهرين چمڙي ۽ وارن ۾ موجود (ٻين جي وچ ۾) پروٽين ڪيريٽين جي ميڪانياتي طاقت ۽ حل نه ٿيڻ جي صلاحيت ڏين ٿا.

## خاصيتون

### طبعي خاصيتون

سلفر ڪيترائي پولي ايٽامڪ ماليڪيول ٺاهيندو آهي. سڀ کان وڌيڪ مشهور ايلوٽروپ آڪٽاسلفر، سائڪلو-S8 آهي. سائڪلو-S8 جو پوائنٽ گروپ D4d آهي ۽ ان جو ڊائپول لمحو 0 - D آهي.   آڪٽاسلفر هڪ نرم، روشن پيلو ٺوس آهي جيڪو بدبودار آهي. اهو (115.21°C)  (239.38 °F) تي ڳري ٿو، ۽ (444.6°C) (832.3 °F) تي اُبلندو آهي. (95.2°C) (203.4 °F) تي، ان جي پگھلڻ واري درجه حرارت کان هيٺ، سائڪلو-آڪٽاسلفر آهستي آهستي α-  آڪٽاسلفر کان β-پوليمورف ۾ تبديل ٿيڻ شروع ڪري ٿو. S8 رنگ جي جوڙجڪ هن مرحلي جي منتقلي سان عملي طور تي تبديل نه ٿيندي آهي، جيڪا بين-ماليڪيولر رابطي کي متاثر ڪري ٿي. پگھريل سلفر کي ٿڌو ڪرڻ سان (119.6°C) (247.3 °F) تي جمي ويندو آهي،  ڇاڪاڻ ته اهو گهڻو ڪري β-S8 ماليڪيولن تي مشتمل هوندو آهي. پنهنجي پگھلڻ ۽ اُبلندڙ گرمي پد جي وچ ۾، آڪٽاسلفر پنهنجي ايلوٽروپ کي ٻيهر تبديل ڪري ٿو، β-آڪٽاسلفر کان γ-سلفر ڏانهن رخ ڪري ٿو، ٻيهر گهٽ کثافت سان گڏ پر پوليمر جي ٺهڻ جي ڪري ويسڪوسيٽي وڌي ٿي. وڌيڪ گرمي پد تي، ڊيپوليمرائيزيشن ٿيڻ سان ويسڪوسيٽي گهٽجي ويندي آهي. پگھريل سلفر 200 °C (392 °F) کان مٿي هڪ ڳاڙهو ڳاڙهو رنگ اختيار ڪري ٿو. سلفر جي کثافت تقريباً 2 g/cm3 آهي، جيڪو ايلوٽروپ تي منحصر آهي؛ سڀئي مستحڪم ايلوٽروپ بهترين برقي انسولٽر آهن.   

سلفر 20 °C (68 °F) ۽ 50 °C (122 °F) جي وچ ۾ تصعید جو عمل (حالت کي سڌو سنئون ٺوس کان بخارات (گيس) ۾ تبديل ڪرڻ) (sublimation) ڪندو آهي. سلفر پاڻيءَ ۾ حل نه ٿيندڙ آهي پر ڪاربان ڊائي سلفائيڊ (CS2) ۽ گهٽ حد تائين، ٻين غير قطبي نامياتي محلولن ۾، جهڙوڪ بينزين ۽ ٽولوئين، ۾ حل ٿيندڙ آهي.

### ڪيميائي خاصيتون
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&nbsp;kJ/mol, 4556&nbsp;kJ/mol and 8495.8&nbsp;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

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 (). Other than 35S, with a half-life of 87 days, the radioactive isotopes of sulfur have half-lives less than 3&nbsp;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 (δ<sup>34</sup>S) 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 δ<sup>13</sup>C 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 <sup>40</sup>Ar. 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 <sup>35</sup>Cl + 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

32S is created inside massive stars, at a depth where the temperature exceeds 2.5×109&nbsp;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 accidentally discovered to exist on Mars after the Curiosity rover drove 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.

## حوالا

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